Download GENERAL AND INORGANIC CHEMISTRY STUDY GUIDE

DANYLO HALYTSKY LVIV NATIONAL MEDICAL UNIVERSITY
DEPARTMENT of GENERAL, BIOINORGANIC, PHYSICAL and
COLLOIDAL CHEMISTRY
GENERAL AND INORGANIC CHEMISTRY
STUDY GUIDE
for the 1st year students of pharmaceutical faculty
(Part 2. Inorganic chemistry)
L’VIV – 2015
General and inorganic chemistry study guide for the 1st year students
of pharmaceutical faculty (Part 2. Inorganic chemistry)
Методичні вказівки з загальної та неорганічної хімії для студентів І
курсу фармацевтичного факультету (Частина 2. Неорганічна хімія)
(англійською мовою)
Методичні вказівки уклали: доценти Роман О.М., Кленіна О.В.,
Огурцов В.В., асистент Маршалок О.І.
За загальною редакцією: доцента Роман О.М.
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Safety Rules
The chemistry laboratory is not a dangerous place to work as long as all
necessary precautions are taken seriously. In the following paragraphs, those
important precautions are described. Everyone who works and performs
experiments in a laboratory must follow these safety rules at all times. Students
who do not obey the safety rules will not be allowed to enter and do any type of
work in the laboratory and they will be counted as absent. It is the student’s
responsibility to read carefully all the safety rules before the first meeting of the
lab.
Eye Protection: Because the eyes are particularly susceptible to permanent
damage by corrosive chemicals as well as flying objects, safety goggles must be
worn at all times in the laboratory. Prescription glasses are not recommended since
they do not provide a proper side protection. No sunglasses are allowed in the
laboratory. Contact lenses have potential hazard because the chemical vapors
dissolve in the liquids covering the eye and concentrate behind the lenses. If you
have to wear contact lenses consult with your instructor. If possible try to wear a
prescription glasses under your safety goggles. In case of any accident that a
chemical splashes near your eyes, immediately wash your eyes with lots of water
and inform your instructor. Especially, when heating a test tube do not point its
mouth to anyone.
Always assume that you are the only safe worker in the lab. Work defensively.
Never assume that everyone else as safe as you are. Be alert for other’s mistakes.
Cuts and Burns: Remember you will be working in a chemistry laboratory and
many of the equipment you will be using are made of glass and it is breakable.
When inserting glass tubing or thermometers into stoppers, lubricate both the
tubing and the hole in the stopper with water. Handle tubing with a piece of towel
and push it with a twisting motion. Be very careful when using mercury
thermometer. It can be broken easily and may result with a mercury contamination.
Mercury vapor is an extremely toxic chemical.
When you heat a piece of glass it gets hot very quickly and unfortunately hot
glass look just like a cold one. Handle them with a tong. Do not use any cracked or
broken glass equipment. It may ruin an experiment and worse, it may cause serious
injury. Place it in a waste glass container. Do not throw them into the wastepaper
container or regular waste container.
Poisonous Chemicals: All of the chemicals have some degree of health
hazard. Never taste any chemicals in the laboratory unless specifically directed to
do so. Avoid breathing toxic vapors. When working with volatile chemicals and
strong acids and bases use ventilating hoods. If you are asked to taste the odor of a
substance does it by wafting a bit of the vapor toward your nose. Do not stick your
nose in and inhale vapor directly from the test tube. Always wash your hands
before leaving the laboratory.
Eating and drinking any type of food are prohibited in the laboratory at all
times. Smoking is not allowed. Anyone who refuses to do so will be forced to leave
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the laboratory.
Clothing and Footwear: Everyone must wear a lab coat during the lab and no
shorts and sandals are allowed. Students who come to lab without proper clotting
and shoes will be asked to go back for change. If they do not come on time it will
be counted as an absence. Long hair should be securely tied back to avoid the risk
of setting it on fire. If large amounts of chemicals are spilled on your body,
immediately remove the contaminated clothing and use the safety shower if
available. Make sure to inform your instructor about the problem. Do not leave
your coats and back packs on the bench. No headphones and Walkman are allowed
in the lab because they interfere with your ability to hear what is going on in the
lab.
Fire: In case of fire or an accident, inform your instructor at once. Note the
location of fire extinguishers and, if available, safety showers and safety blankets as
soon asyou enter the laboratory so that you may use them if needed. Never perform
an unauthorized experiment in the laboratory. Never assume that it is not necessary
to inform your instructor for small accidents. Notify him/her no matter how slight it
is.
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Thematic schedule of practice and laboratory studies in inorganic chemistry
for the 1st year students of pharmaceutical faculty
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The topics
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Chemical elements and their classification. A human and biosphere
General characteristic of s-elements. Hydrogen and its compounds
s-Elements of the ІA group of the Periodic Table. Alkali metals
s-Elements of the ІІA group of the Periodic Table. Beryllium,
Magnesium, and Alkaline earth metals.
General characteristic of p-elements. Boron, Alluminium and
properties of their compounds
р-Elements of IVА group. Carbon, Silicon
р-Elements of IVА group. Germanium family elements
(Germanium, Tin, and Lead)
р-Elements of VА group. Nitrogen and its compounds in the
negative oxidation states
р-Elements of VА group. Nitrogen and its compounds in the
positive oxidation states
р-Elements of VА group. Chemical properties of Phosphorus and
its compounds
р-Elements of VА group. Arsenic family elements (Arsenic,
Antimony, and Bismuth)
р-Elements of VІА group. Oxygen, Sulfur, Selenium, Tellurium
р-Elements of VІІА group. Halogens
General characteristic of d-elements. d-Elements of IВ group.
Copper, Silver, Gold
d-Elements of IІВ group. Chemical properties of Zinc, Cadmium,
Mercury
d-Elements of VІВ group. Chromium elements family
d-Elements of VІІВ group. Manganese elements family
d-Elements of VІІІВ group. Iron and its compounds
d-Elements of VІІІВ group. Cobalt and Nickel compounds.
Platinum metals
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Topic 1
Chemical elements and their classification. A human and
biosphere
1. Objectives
Nowadays we know 110 chemical elements that make up more than 10 million
of organic and hundreds thousands of inorganic substances. There are different
classifications of elements, depending on the criteria by which chemical elements
are divided. These classifications helps to orient on modern chemical and
pharmaceutical information.
2. Learning Targets
To be able to classify the chemical elements according to different criteria. To
understand the relationship between the biological role of nutrients and their form
in human body.
To interpret on the basis of Vernadsky’s doctrine such concepts as biosphere,
noosphere. To explain the regularities of migration of chemical elements in the
biosphere.
3. Self Study Section
3.1. Syllabus Content
The concept of the chemical elements; their classification by origin, chemical
properties, the structure of the outer energy level, spreading in nature and
importance for living organism. Classification of bioelements; their content in
human body. Connection between physico-chemical parameters of the elements and
their position in the periodic system and the content in the body.
V. Vernadsky’s doctrine about biosphere and biogeochemistry. The concept of
migration of chemical elements. Relationship between endemic diseases and
features of biogeochemical provinces.
A human and biosphere. Noosphere. Technological progress and ecology.
3.2. Overview
About 16 elements are used in formation of chemical compounds from
which living organisms are made. These 16 elements and a few others, which occur
in a particular organism, are called bioelements. Bioelement is any chemical
element that is found in the molecules and compounds that make up a living
organism. Some of the more prominent representatives are called macronutrients,
whereas those appearing only at the level of parts per million or less are referred to
as micronutrients. These nutrients perform various functions, including the
building of bones and cell structures, regulating the body’s pH, carrying charge,
and driving chemical reactions.
The main six elements are: C, H, O, N, P, and S, and they’re called primary
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p-elements
s-elements
Sodium Na
Potassium K
Magnesium Mg
Calcium Ca
biometals
Oxygen O
Carbon C
Nitrogen N
Phosphorus P
Sulfur S
Hydrogen H
organogens
bioelements. These elements are present as constituents of biomolecules, in
inorganic matrix substances, and in water. Minerals are rarely present in large
amounts. The above 6 bioelements plus Ca, K, Na, Cl, Mg and Fe make up 99.9%
of the biomass. The remaining elements occur mainly as trace elements, which are
needed only in catalytic quantities. While the light metals are usually present as
mobile cations, the heavy metals are generally fixed as stable components of
organic complexes.
Main macroelements:
Are the part of albumens (proteins), fats, nucleic
acids, and also hormones and enzymes. Their mass
in an organism is about 100 and more grammes per
70 kg of living mass.
Are found in biological liquids (thus, К in
intracellular and Na - in extracellular). Besides,
Mg and Ca are included in the composition of
bone tissue. They participate in the processes of
excitation and inhibition of the central nervous
system, and also stimulate some metabolic
processes.
d-elements
Iron Fe
Copper Cu
Zink Zn
Manganese Mn
Cobalt Co
Nickel Ni
Chromium Cr
Molybdenum Mo
Biogenic microelements
Main microelements:
Are contained in an organism within the limits of
10-2-10-6 % mass. They are included in a significant
number of enzymes (metaloenzymes), some
vitamins (B12) and hormones (insulin). They are
involved in the processes of hematopoiesis,
reproduction, growth, and metabolism. Their
biological functions are closely related to the
processes of complexing between bioligands and
metal ion due to free atomic orbitals
A chemical element is a pure chemical substance consisting of one type of atom
distinguished by its atomic number, which is the number of protons in its nucleus.
Elements are divided into metals, metalloids, and non-metals. Familiar examples of
elements include carbon, oxygen (non-metals), silicon, arsenic (metalloids),
aluminium, iron, copper, gold, mercury, and lead (metals).
The lightest chemical elements, including hydrogen, helium (and smaller
amounts of lithium, beryllium and boron), are thought to have been produced by
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various cosmic processes and cosmic-ray spallation. While most elements are
generally viewed as stable, a small amount of natural transformation of one element
to another also occurs at the present time through decay of radioactive elements as
well as other natural nuclear processes.
Of the 98 naturally occurring elements, those with atomic numbers 1 through 40
are all considered stable. Elements with atomic numbers 41 through 82 are
apparently stable (except technetium, element 43 and promethium, element 61,
which are unstable) but theoretically unstable, and thus possibly mildly radioactive.
The half-lives of elements 41 through 82 are so long, however, that their
radioactive decay remains undetected by experiment. These "theoretical
radionuclides" have half-lives at least 100 million times longer than the estimated
age of the universe. Elements with atomic numbers 83 through 98 are unstable to
the point that their radioactive decay can be detected. Some of these elements,
notably thorium (atomic number 90) and uranium (atomic number 92), have one or
more isotopes with half-lives long enough to survive as remnants of the explosive
stellar nucleosynthesis that produced the heavy elements before the formation of
our solar system. The very heaviest elements (those beyond californium, atomic
number 98) undergo radioactive decay with half-lives so short that they do not
occur in nature and must be synthesized.
Chemical elements may also be categorized by their origin on Earth, with the
first 98 considered naturally occurring, while those with atomic numbers beyond 98
have only been produced artificially as the synthetic products of man-made nuclear
reactions.
Of the 98 naturally occurring elements, 84 are considered primordial and either
stable or metastable (apparently stable but theoretically unstable or radioactive).
The remaining 14 naturally occurring elements possess half lives too short for them
to have been present at the beginning of the Solar System, and are therefore
considered transient elements. Of these 14 transient elements, 7 (polonium,
astatine, radon, francium, radium, actinium, and protactinium) are relatively
common decay products of thorium, uranium, and plutonium. The remaining 7
transient elements (technetium, promethium, neptunium, americium, curium,
berkelium, and californium) occur only rarely, as products of rare nuclear reaction
processes involving uranium or other heavy elements.
Elements with atomic numbers 1 through 40 are all stable, while those with
atomic numbers 41 through 82 (except technetium and promethium) are metastable.
The half-lives of these metastable "theoretical radionuclides" are so long (at least
100 million times longer than the estimated age of the universe) that their
radioactive decay has yet to be detected by experiment. Elements with atomic
numbers 83 through 98 are unstable to the point that their radioactive decay can be
detected. Some of these elements, notably thorium (atomic number 90) and
uranium (atomic number 92), have one or more isotopes with half-lives long
enough to survive as remnants of the explosive stellar nucleosynthesis that
produced the heavy elements before the formation of our solar system. For
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example, at over 1.9·1019 years, over a billion times longer than the current
estimated age of the universe, bismuth-209 (atomic number 83) has the longest
known alpha decay half-life of any naturally occurring element. The very heaviest
elements (those beyond californium, atomic number 98) undergo radioactive decay
with half-lives so short that they do not occur in nature and must be synthesized.
A human and biosphere
In Vernadsky’s theory of the Earth’s development, the noosphere is the third
stage in the Earth’s development, after the geosphere (inanimate matter) and the
biosphere (biological life). Just as the emergence of life fundamentally transformed
the geosphere, the emergence of human cognition fundamentally transformed the
biosphere. In this theory, the principles of both life and cognition are essential
features of the Earth’s evolution, and must have been implicit in the Earth all along.
This systemic and geological analysis of living systems complements Charles
Darwin’s theory of natural selection, which looks at each individual species, rather
than at its relationship to a subsuming principle. Vernadsky defined the future
evolutionary state of the biosphere as the Noosphere, the sphere of reason. The
term "Noosphere" was first coined by the French mathematician and philosopher,
Edouard Le Roy (1927). "Le Roy, building on Vernadsky’s ideas and on
discussions with Teilhard de Chardin [they both attended Vernadsky’s lectures on
biogeochemistry at the Sorbonne in 1922-1923], came up with the term
"noosphere", which he introduced in his lectures at the College de France in 1927.
Vernadsky saw the concept as a natural extension of his own ideas predating Le
Roy’s choice of the term". Le Roy understood the noosphere as a shell of the Earth
or a "thinking stratum", including various components, such as industry, language,
and other forms of rational human activity. Le Roy’s concept was developed by De
Chardin, who considered the noosphere as something external to the biosphere - a
progression from biological to psychological and spiritual evolution. Teilhard
based his conception based on philosophical writings, and was completely ignorant
of Vernadsky’s biogeochemical approach. Vernadsky developed his concept of the
noosphere out of his theory of the biosphere, combining his biogeochemical works
with his own work in philosophy of science.
According to Vernadsky, the biosphere became a real geological force that is
changing the face of the Earth, and the biosphere is changing into the noosphere. In
Vernadsky’s interpretation (1945), the noosphere, is a new evolutionary stage of
the biosphere, when human reason will provide further sustainable development
both of humanity and the global environment.
Vernadsky made an important contribution to science in general, and in ecology
in particular. It is essentially Vernadsky’s theory of the biosphere, expounded in his
work "Biosfera" (1926) that is embodied in the global approach to ecological
problems today. To solve global ecological problems that may endanger even the
very existence of humanity in the future, a cultivation of a new worldview among
people, and especially young generations, is absolutely needed.
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Vernadsky’s views are stated in the Foreword to the English-language edition as
follows:
1. Life occurs on a spherical planet. Vernadsky is the first person in history to
come [to] grips with the real implications of the fact that Earth is a self-contained
sphere.
2. Life makes geology. Life is not merely a geological force, it is the geological
force. Virtually all geological features at Earth’s surface are bio-influenced, and are
thus part of Vernadsky’s biosphere.
3. The planetary influence of living matter becomes more extensive with time.
The number and rate of chemical elements transformed and the spectrum of
chemical reactions engendered by living matter are increasing, so that more parts of
Earth are incorporated into the biosphere.
Vernadsky could not make the connection between thought as a non-energetic
"something" and the transformations involving human activity, such as thought
leading to creation of new compounds or thought causing the significant effects in
all the other spheres of the planet. However, what Vernadsky did in creating the
idea of the Noosphere was to bring into scientific discussion the idea of Mind, the
idea of mental operations within the Biosphere. From a scientific perspective, this
was an enormous jump. Proceeding from where Vernadsky left off, we can define
the Noosphere as an envelope within the Biosphere that is comprised of the sum
total of all human mental operations. Using a characteristic Vernadskian naming
convention, we could say that these mental operations are composed of "Mental
Matter" either within or alongside the Living Matter of the Biosphere. Unlike the
transformations of Living Mattter within the Biosphere, the only way we have of
detecting the operation of Mental Matter is by its effects on human behavior. Mind,
then, is a mechanism for transforming Mental Matter into behavior which then
carries out the transformations of energy that we can measure with conventional
methods. The idea that human behavior is our measuring tool of the state of the
Noosphere is crucial to our understanding.
Glossary
Atmosphere - The air envelope surrounding the Earth. Earth’s atmosphere is the
layer of gases surrounding the planet Earth retained by the Earth’s gravity
(Wikipedia Encyclopedia, 2004).
Biosphere - The totality of living organisms with their environment, i.e. those
layers of the Earth and the Earth’s atmosphere in which living organisms are
located (VanDeVeer and Pierce, 2003). Vernadsky defined ecology (originally
intended as the "economy of nature") as the science of the biosphere.
Ecology - The branch of science that studies the distribution and abundance of
living organisms, their habitats, and the interactions between them and their
environment - which includes both biotic (non-living) elements like climate and
geology, and biotic ones like other species (Wikipedia Encyclopedia, 2004).
Hydrosphere - The water envelope surrounding the Earth. Hydrosphere
describes collective mass of water that is found under, on and over the surface of
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the Earth (Wikipedia Encyclopedia, 2004).
Lithosphere - The outer solid shell of the Earth (Wikipedia Encyclopedia,
2004), i.e. the outer layers of the earth made up of the crust and the outer portion of
the mantle (the layer between the core and the crust) (VanDeVeer and Pierce,
2003).
Noosphere - Literally, "the envelope of mind" (Smil, 2002, p. 12) or the "sphere
of human thought" (Wikipedia Encyclopedia, 2004). "The Noosphere is the last of
many stages in the evolution of the biosphere in geological history" (Vernadsky,
1945, p. 10).
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. –
1991.
3.4. Test Yourself
а) Review Questions
1. What is a chemical element?
2. What is the most abundant element?
3. What are the main elements in the human body?
4. Which chemical elements are called bioelements? Give their classifications
according to the biological significance and content in the human body.
5. What biometals are part of chlorophyll and hemoglobin?
6. Why the balance between macro- and micronutrients is an important factor in the
proper functioning of the organism? How is it maintained?
b) Problems to Solve
1. Animal bones containe 2.12% of phosphorus, 7.56% of calcium and 1.51% of
magnesium. Find the mass percentage of these elements in the ash of the bones,
which is 27% of their mass.
Answer: Р – 7.85 %; Са – 28 %; Mg – 5.6 %.
2. The content of magnesium in plasma and cellular elements of the blood are
respectively 1.33 and 2.125 mmol/kg. This blood consists of 58% of plasma and
42% of the cells. Find magnesium content in a blood (mmol/kg).
Answer: 1.66 mmol/kg.
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3. Human blood contains 60% plasma and 40% blood cells. Calculate the mass
percentage of water in the blood, if its mass percentage in plasma is 92%, mass
percentage in cells - 64%.
Answer: Cp(Н2О) blood = 80.2 %.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) Discussion and explanation of main questions of the topic:
– Classification of chemical elements based on different criteria.
– Relationship between the biological role of nutrients and their form in human
body.
– Main concepts of Vernadsky’s doctrine about the biosphere and noosphere.
– Regularities of chemical elements migration in the biosphere.
b) Solving of typical numerical problems.
5. Conclusions and Interpretations. Lesson Summary
Topic 2
General characteristic of s-elements. Hydrogen and its
compounds
1. Objectives
Hydrogen forms 0.15 % of the Earth’s crust; it is the major constituent of water.
0.5 ppm of hydrogen H2 and varial proportions as water vapour are present in the
atmosphere. Hydrogen is also a major component of biomass, consituing the 14%
by weight. Hydrogen occurs naturally in the atmosphere.
Hydrogen can form compounds with most elements and is present in water and
all organic compounds. It plays a particularly important role in acid-base chemistry,
in which many reactions involve the exchange of protons between soluble
molecules. As the only element for which the Schrödinger equation can be solved
analytically, study of the energetics and bonding of the hydrogen atom has played a
key role in the development of quantum mechanics.
2. Learning Targets:
To carry out chemical experiments and to write equations of chemical reactions
those characterize properties of hydrogen and its compounds.
To know its biological role and uses of hydrogen compounds in medical
practice.
3. Self Study Section
3.1. Syllabus Content
General characteristics of Hydrogen. Position in the periodic table of elements.
Reactions with oxygen, halogens, metal oxides. Characteristics and reactivity of
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hydrogen compounds with other common elements: oxygen, nitrogen, carbon,
sulfur. Ions of hydrogen, hydronium and ammonium.
Water as an important compound of hydrogen. Its physical and chemical
properties. Aquacomplexes and crystall hydrates. Distilled and non-pyrogenic
water - preparation and use in pharmacy. Natural water, pollution of water, mineral
water.
Hydrogen peroxide. The structure of the molecule. Methods of obtaining. Acidbase and redox properties of hydrogen peroxide, use in medicine and pharmacy.
3.2. Theoretical Backgrounds
The most common naturally occurring isotope of hydrogen contains one
electron and an atomic nucleus of one proton. In ionic compounds it can take on
either a positive charge (becoming a cation, a bare proton) or a negative charge
(becoming an anion known as a hydride). Hydrogen can form compounds with
most elements and is present in water and all organic compounds. It plays a
particularly important role in acid-base chemistry, in which many reactions involve
the exchange of protons between soluble molecules. H2 reacts directly with other
oxidizing elements. A violent and spontaneous reaction can occur at room
temperature with chlorine and fluorine, forming the corresponding hydrogen
halides, HCl and HF. While H2 is not very reactive under standard conditions, it
does form compounds with most elements. Millions of hydrocarbons are known,
but they are not formed by the direct reaction of elementary hydrogen and carbon.
Hydrogen can form compounds with elements that are more electronegative, such
as halogens (e.g., F, Cl, Br, I) and chalcogens (O, S, Se); in these compounds
hydrogen takes on a partial positive charge. When bonded to fluorine, oxygen, or
nitrogen, hydrogen can participate in a form of strong noncovalent bonding called
hydrogen bonding, which is critical to the stability of many biological molecules.
Hydrogen also forms compounds with less electronegative elements, such as metals
and metalloids, in which it takes on a partial negative charge. These compounds are
often known as hydrides. In compounds of hydrogen, the most common oxidation
states of hydrogen are: 1, and -1.
Н20 + 2е– → 2Н–
KH, NaH, BaH2
The human body contains about 10% of hydrogen. The most important
compounds of hydrogen are water and hydrogen peroxide. Oxidizing properties of
hydrogen peroxide are used in the bleaching of substances, such as hair, ivory,
feathers, and delicate fabrics, which would be destroyed by other agents. It is also
used medicinally, in the form of a 3% aqueous solution, as an antiseptic and throat
wash.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
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1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. –
1991.
3.4. Self Assessment Exercises
а) Review Questions
1. How to explain the position of hydrogen in the periodic table of elements.
Specify the oxidation states for hydrogen. Give the formulas of corresponding
compounds.
2. What is the most common hydrogen compound?
3. Give the examples of reactions which characterize the chemical properties of
hydrogen.
4. Give the examples of acidic and basic hydrides. Write down equation of the
preparation of litium tetrahydroaluminate.
5. What is a hydration reaction? Give the examples of basic and acidic oxides
hydration reactions.
6. Hydrolysis. Write the molecular and ionic equations of the hydrolysis of the
following salts: KCN, Li3PO4, Cr2(SO4)3, CuCl2, CH3COONH4. Indicate the
pH.
7. Give the examples of protolytic reactions.
8. Complete and balance the equations of the following reactions. Define the
properties of hydrogen peroxide (reducing or oxidizing agent):
H2O2 + FeSO4 + H2SO4 → …
H2O2 + KMnO4 + H2SO4 → …
H2O2 + Cr2(SO4)3 + NaOH → …
H2O2 + KI + H2SO4 → …
H2O2 + PbS → …
H2O2 + H2S → …
H2O2 + I2 → …
H2O2 + K2Cr2O7 + H2SO4 → …
9. Which compound is used in cosmetic preparations to bleach facial and body
hair?
10. List the major types of mineral waters and specify their use.
b) Types of Numerical Problems and Their Solving Strategies
1. What volume (in cm3) of 4 N HCl solution is required for dissolving of 10 grams
of zinc?
Answer: 76.5 cm3.
2. Calculate the mass percentage of hydrogen peroxide if 25 cm3 of solution reacts
completely with 22 cm3 of 2 N solution of potassium in acidic medium.
Answer: 2.99 %.
3. What is the density of “oxyhydrogen” gas (2 volumes of Н2 and 1 volume of О2)
by hydrogen, oxygen and an air?
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Answer: а) 6.0; b) 0.375; c) 0.414.
4. How much heat does the organism loss if it losses 650 g of water through skin?
Answer: 1589 kJ.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
− position of hydrogen in the periodic table, atomic structure, valence;
− properties of atoms (ionic radius, covalent radius, ionization energy,
electronegativity);
− natural occurrence, biological role of hydrogen;
− chemical properties: acid-base, redox, complex formation. The most
important compounds that are used in medical practice and national
economy.
4.2. Experimental Guidelines:”Chemical properties of Hydrogen and its
compounds”
4.2.1. Comparing reduction properties of molecular and atomic Hydrogen.
Place 2-3 cm3 of acidified KMnO4 solution into a test-tube and pass through the
solution hydrogen gas from Kipp’s apparatus. Observe the change of color of the
solution. Add zinc granules into the same test-tube. What happened with the
solution of potassium permanganate? Write equations of the reactions. Make a
conclusion.
4.2.2. Hydrogen peroxide as an acid.
To 2-3 cm3 of aluminium salt solution drop by drop add sodium hydroxide
solution to dissolve the precipitate (do not give the excess of NaOH). Then, add
hydrogen peroxide drop by drop. What is observed? Explain this phenomenon and
write the corresponding equations of the reactions.
4.2.3. Decomposition of hydrogen peroxide.
Pour into a test-tube 3.2 cm3 of 3% hydrogen peroxide solution. What is
observed? Add a few grains of manganese (II) oxide. How will change the rate of
hydrogen peroxide decomposition? Write equation of the reaction. What gas is
released? Make a conclusion.
4.2.4. Oxidizing properties of hydrogen peroxide.
а) To 2-3 ml of potassium iodide add 2-3 ml of sulfuric acid and a solution of
hydrogen peroxide dropwise. Observe the change of color of the solution. Write
equation of the reaction.
b) To 1-2 cm3 of soluble salt of lead add 1 cm3 of sodium sulfide or hydrogen
sulphide solution. Drain off the liquid above the precipitate and add 1.2 cm3 of 3%
solution of H2O2. What is observed? Write equations of the reactions.
4.2.5. Reducing properties of hydrogen peroxide.
16
To 1-2 ml of potassium permanganate add 1 ml of sulfuric acid and 3%
hydrogen peroxide solution dropwise. Observe the change of color of the solution.
Write equation of the reaction.
Make a conclusion about the redox properties of hydrogen peroxide. Give the
values of redox potentials of hydrogen peroxide; indicate which properties are
more typical for this compound.
5. Conclusions and Interpretations. Lesson Summary
Topic 3
s-Elements of the ІA group of the Periodic Table.
Alkali metals
1. Objectives
Chemistry of life involves many chemical elements. There are about 27
elements which have been found essential in the biochemical reactions.
Interestingly most of these elements are of low atomic numbers. Sodium and
potassium are quite abundant, ranking sixth and seventh among all elements in the
Earth’s crust, but the other alkali metals are rare. Sodium and potassium ions are
components of numerous silicate crystal lattices seen in the Earth’s crust, but since
most of their compounds are water soluble, they are also important constituents of
seawater and underground deposits of brine. Sodium chloride obtained from such
brines is the chief commercial source of sodium, while potassium can be obtained
from such ores as sylvite (KCl) or carnallite (KCl·MgCl2·6H2O). Both sodium
(Na+) and potassium (K+) ions are essential to living systems. Na+ is the main
cation in fluids surrounding the cells, while K+ is most important inside the cells.
Na+ plays a role in muscle contraction, and both K+ and Na+ play a role in
transmitting of nerve impulses. Sodium and potassium are essential to all
organisms. Their mono-positive ions are structure promoters for both poly- nucleic
acids and proteins. Potassium is an important enzyme activator and plays an
important role in nerve action and cardiac function. Potassium is required in the
cell glucose metabolism and protein synthesis. Sodium is relatively harmless except
in excessive amounts, whereas potassium is moderately toxic to mammals when
injected intravenously, otherwise it is harmless.
2. Learning Targets
To carry out chemical reactions and write equations which can be used to
characterize the properties of IA group s-elements.
To know biological role of Lithium, Potassium, Sodium and their use.
3. Self Study Section
3.1. Syllabus Content
General characteristics of IA group elements. Occurrence in nature. Biological
17
role of s-elements in mineral balance of a human body. Macroelements.
The difference between lithium and other alkali metals. Binary compounds of
alkali metals: hydrides, oxides, peroxides, superoxide.
Alkali metals hydroxides, salts, their properties and use. Use of lithium, sodium
and potassium compounds in medicine.
3.2. Overwiev
The alkali metals exhibit many of the physical properties common to metals,
although their densities are lower than those of other metals. Alkali metals have one
electron in their outer shell, which is loosely bound. This gives them the largest
atomic radii of the elements in their respective periods. Their low ionization
energies result in their metallic properties and high reactivities. An alkali metal can
easily lose its valence electron to form the univalent cation. Electronegativity and
ionization energy increase from left to right and from bottom to top. Alkali metals
have the lowest electronegativity and ionization energy. Francium is the least
electronegative element. Atomic radius increases from right to left and from top to
bottom. Francium is the largest element. Boiling points and melting points increase
going from bottom to top: lithium has the highest boiling point and francium has
the lowest boiling point in Group 1. Alkali metals have low electronegativities.
They react readily with nonmetals, particularly halogens.
Electron configuration
ns1
Valence
I
Oxidation state
+1
All alkali metals are strong reducing agents. The alkali metals react directly
with the halogens, and with sulfur, bromine, hydrogen at the heating:
2Na + 2F2 → 2NaF;
2K + S → 2K2S;
2Li + 2Br2 → 2LiBr;
2Na + H2 → 2NaH.
Only lithium reacts with nitrogen at room temperature and its nitride is the only
stable alkali metal nitride:
6Li + N2 → 2Li3N.
The sodium, potassium, rubidium, and cesium elements also combine violently
with water to form hydroxides:
2Na + 2H2O → 2NaОН + Н2↑;
2Na + 2H2O → 2NaOH+ H2↑;
2K + 2H2O → 2KOH+ H2↑;
2Rb + 2H2O → 2RbOH + H2↑;
2Cs + 2H2O→ 2CsOH + H2↑.
The alkali metals react directly with many elements. All combine swiftly with
oxygen in air to form white oxides Me2O, MeO, peroxides Ме2О2, МеО2, and
superoxides МеО2. Only lithium reacts with oxygen to form Li2O oxide. All except
lithium react further to form yellow peroxides, Me2O2. Potassium, rubidium, and
18
cesium are sufficiently reactive so that yellow superoxides (whose general formula
is MeO2) can be formed.
Peroxide compounds decompose by action of water and dilute acids to form
H2O2 and O2:
Na2O2 + H2SO4 = Na2SO4 + H2O2;
2KO2 + H2SO4 = K2SO4 + H2O2 + O2↑;
2KO3 + 2 H2O = 2KOH + H2O2 + 2O2↑.
Other binary compounds also react vigorously with water:
NaH + H2O = NaOH + H2↑;
Na4C + 4H2O = 4NaOH + CH4↑.
With acids alkali metals form compounds with lower oxidation states of
nitrogen and sulfur:
8K + 10HNO3 (dil.) = 8KNO3 + NH4NO3 + 3H2O;
8Na + 10HNO3 (conc.) = 8NaNO3 + N2O↑ + 5H2O.
These oxides, hydrides, hydroxides, and sulfides all dissolve in water to give
basic solutions, and they are among the strong bases.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. –
1991.
3.4. Self Assessment Exercises
а) Review Questions
1. What are the common physical features and chemical properties of alkali
metals?
2. What alkali metals can form superoxides?
3. Which element is the most electronegative: Francium, Potassium, or Lithium?
4. Compounds that generally look like Me2O2 are formed with the metal and
oxygen ion. Define the kind of oxygen ion.
5. Complete and balance the following equation: Li2O2 + H2O →?
6. Which of the alkali metals has a higher melting point: Sodium (Na) or Francium
(Fr)? Explain.
7. What is the main difference in the chemical properties of lithium and other
alkali metals?
8. Write equations of the following reactions:
19
Na + O2 → …
K + C2H2 → …
K2O4 + CO2 → …
KOH + O3 → …
Li + N2 → …
Na2O2 + CO2 → … .
9. Write the molecular and ionic equations of the hydrolysis of the following salts:
LiCN, K2S, K2CO3, NaClO.
10. How to store alkali metals and how to recycle their residues?
11. Explain the significance of sodium, potassium, magnesium and calcium in
biological fluids.
12. What alkali metal compounds are used in medicine?
b) Problems to Solve
1. Whan is the number of moles (in mmol) of 0.9% sodium chloride NaCl solution
(density = 1 g/сm3). Volume of solution is 100 сm3.
Answer: 15.4 mmol.
2. What mass of Na2CO3·10Н2О is needed to prepare 250 cm3 of 0.1 N sodium
carbonate solution?
Answer: 3.575 g
3. Calculate рН, constant of hydrolysis and degree of hydrolysis of 0.1 М sodium
carbonate solution.
Answer: рН = 11.7, Кh = 2.2·10–4, h = 4.7.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of IA group s-elements and
their compounds»
4.2.1. Flame test.
Heat salts of IA group elements strongly in a flame. What is observed? Mark the
color of metals ions flame.
Flame Colors
All alkali metals have their own specific flame color. The colors are caused by the
difference in energy among the valence shell of s and p orbitals, which corresponds
to wavelengths of visible light. When the element is introduced into the flame, its
outer electrons are excited and jump to a higher electron orbital. The electrons then
fall and emit energy in the form of light. The different colors of light depend on
how much energy or how far the electron falls back to a lower energy level. For this
reason, they are often used in fireworks. Each alkali metal has a unique color and is
easily identifiable.
Group 1 Element
Flame Color
Lithium
Carmine
20
Sodium
Golden Yellow
Potassium
Red/Violet
Rubidium
Blue/Violet
Cesium
Blue/Violet
4.2.2. Reactions of identification of K+ and Na+ ions.
а) To 1–2 ml of sodium chloride (NaCl) solution in a test-tube add identical
volume of potassium hexa-hydroxostibate(V) (K[Sb(OH)6]) solution. What is
observed? Write equation of the reaction.
b) To 1–2 ml of potassium chloride (KCl) solution in a test-tube add identical
volume of sodium hydrotartrate (NaHC4H4O6) solution. Cool the test-tube in a
stream of water and rub it by glass stick. What is observed? Write equation of the
reaction.
4.2.3. Hydrolysis of alkali metals salts.
Place 2-3 cm3 of solutions of NaHCO3, K2S, Na3PO4, Na2HPO4, NaH2PO4 salts
into test-tubes and determine their pH using universal indicator paper, methyl
orange and phenolphthalein indicators. Write equations of the hydrolysis of listed
salts.
5. Conclusions and Interpretations. Lesson Summary
Topic 4
s-Elements of the ІІA group of the Periodic Table.
Beryllium, Magnesium, and Alkaline earth metals.
1. Objectives
Elements of Group IIA of the Periodic Table are known as the Alkaline Earth
Metals. The name "Alkaline" results from their slight solubility in water, and
"Earth" is derived from their inability to decompose when exposed to heat. The IIA
group of elements comprises beryllium, magnesium, calcium, strontium, barium
and radium. They follow alkali metals in the periodic table. These (except
beryllium) are known as alkaline earth metals. Magnesium is essential to all
organisms. It is an integral part of chlorophyll, the green pigment in plants
responsible for photosynthetic reaction. Its deficiency in plants is characterized by
yellowing of leaves (chlorosis). A typical human adult requires about 200-300 mg
of magnesium daily. If the diet is very rich in phosphates, magnesium may
precipitate out as magnesium phosphate. This most commonly occurs in infections
of the urinary tract. Calcium is also an essential element for all organisms. It forms
solid skeletal materials such as bones, and acts as trigger for muscular contraction
and the release of hormones. Its deficiency is caused due to its actual absorption.
21
One major difficulty is the tendency for calcium to be precipitated by a large
number of anions present in food. In this regard, phosphate ions interfere to the
greatest extent. Therefore, a high protein diet which is rich in phosphates will be
unfavorable to calcium absorption. Bones are building up from a precipitate of
octacalcium phosphate, laid down on a framework of the protein collagen by cells
known as osteoblasts. The salt is then slowly converted to the normal form of bone,
calcium hydroxyapatite. In cases of hypocalcaemia, deposits of calcium
triphosphate are formed as stones in the bladder and kidney.
2. Learning Targets
To carry out chemical reactions and write equations which can be used to
characterize the properties of s-elements of IIA group.
To know biological role of Calcium, Magnesium and their use in medicine.
Toxic effects of Strontium.
3. Self Study Section
3.1. Syllabus Content
General characteristics of s-elements of IIA group. Reducing properties of
elements. Comparison of beryllium, magnesium and calcium properties. Reactions
of simple substances with water, acids and bases solutions.
Beryllium. Chemical properties. sр-hybridization of atomic orbitals of
beryllium. Beryllium oxide and hydroxide, their amphoteric properties. Aqua- and
hydroxocomplexes of beryllium. Solubility and hydrolysis of beryllium salts.
Magnesium. Magnesium oxide and hydroxide. Solubility and hydrolysis of
magnesium salts. Mg2+ ion as a complex formation agent. Chlorophyll.
Alkaline earth metals. General characteristics. Physical properties and
occurrence. Chemical properties. Basic oxides and hydroxides of the alkaline earth
metals. Solubility in water. Reactions of identification of Mg2+, Ca2+, Sr2+, Ba2+
ions. Hardness of water. Methods of softening.
Calcium compounds in the bone tissue. The toxic action of beryllium and
barium. The biological role of calcium and magnesium. Uses of magnesium,
calcium and barium compounds in medicine and pharmacy.
3.2. Overwiev
Electron configuration
ns2
Valence
II
Oxidation state
+2
ІІА group metals are strong reducing agents. This is indicated by large negative
values of their reduction potentials. However, their reducing power is less than
those of their corresponding alkali metals. Beryllium has less negative value
compared to other alkaline earth metals. However, it has reducing nature due to
large hydration energy associated with the small size of Be2+ ion and relatively
large value of the atomization enthalpy of the metal.
ІІА group metals react with nonmetals. Calcium, strontium and barium are
22
readily attacked by air to form the oxides ЕО. They also react with water with
increasing vigour even in cold to form hydroxides:
Е + 2Н2О = Е(ОН)2 + Н2↑.
Other binary compounds of these elements also react vigorously with water:
CaO + H2O = Ca(OH)2;
Ca3N2 + 6H2O = 3Ca(OH)2 + 2NH3↑;
CaC2 + 2H2O = Ca(OH)2 + C2H2↑.
With acids alkaline earth metals form compounds with lower oxidation state of
nitrogen and sulfur:
4Са + 10HNO3 (dil) = 4Са(NO3)2 + NH4NO3 + 3H2O;
4Са + 10HNO3 (conc) = 4Са(NO3)2 + N2O↑ + 5H2O;
4Mg + 5H2SO4 (conc) = 4MgSO4 + H2S↑ + 4H2O.
The solubility, thermal stability and the basic character of hydroxides of these
elements increase with increasing of atomic number from Mg(OH)2 to Ba(OH)2.
The alkaline earth metal hydroxides are, however, less basic and less stable than
alkali metal hydroxides. Beryllium hydroxide is amphoteric in nature as it reacts
with both acid and alkali.
Beryllium, the first member of the Group II metals, shows anomalous behaviour
as compared to magnesium and the rest of the members. Further, it shows diagonal
relationship to aluminium. The ionic radius of Be2+ is estimated to be 31 pm; the
charge/radius ratio is nearly the same as that of the Al3+ ion. Hence beryllium
resembles aluminium in some ways. Some of the similarities are: (a) like
aluminium, beryllium is not readily attacked by acids because of the presence of an
oxide film on the surface of the metal; (b) beryllium hydroxide dissolves in excess
of alkali to give a beryllate ion, [Be(OH)4]2- just as aluminium hydroxide gives
aluminate ion, [Al(OH)4]-; (c) the chlorides of both beryllium and aluminium have
Cl- bridged chloride structure in vapour phase. Both chlorides are soluble in
organic solvents and are strong Lewis acids. They are used as Friedel Craft
catalysts. Beryllium and aluminium ions have strong tendency to form complexes:
BeF42-, AlF63-. Beryllium reacts with strong acids and bases:
Be + 2HCl + 4H2O = [Be(H2O)4]Cl2 + H2↑;
Be + 2NaOH + 2H2O = Na2[Be(OH)4] + H2↑;
BeO + H2SO4 + 4H2O = [Be(H2O)4]SO4 + H2O;
BeO + 2NaOH = Na2BeO2 + H2O;
t

→
Na2BeS2;
t
BeS + SiS2 
→ BeSiS3;
BeS + Na2S
BeF4 + 2KF = K2[SiF6];
BeF2 + SiF4 = Be[SiF6];
Be(OH)2 + 2HCl = BeCl2 + 2H2O;
Be(OH)2 + 2NaOH = Na2[Be(OH)4].
Reactions of identification:
Ca2+ + (NH4)2C2O4 → CaC2O4↓ + 2NH4+
23
white crystalline
2+
Sr + K2CrO4 → SrCrO4↓ + 2K+
yellow
Ba2+ + H2SO4 → BaSO4↓ + 2H+
white
Hardness of water is determined by the concentration of multivalent cations in
water. Common cations found in hard water include Ca2+ and Mg2+. There are two
types of water hardness, temporary and permanent.
Temporary hardness is caused by the bicarbonate ion, HCO3-, being present in the
water. This type of hardness can be removed by boiling the water to expel the CO2,
as indicated by the following equation:
Ca(HCO3)2 → CaCO3 + CO2 + H2O
Permanent hardness is caused by calcium and magnesium nitrates, sulphates, and
chlorides etc. This type of hardness cannot be eliminated by boiling.
The hardness of water is referred to by three types of measurements: grains per
gallon, milligrams per liter (mg/L), or parts per million (ppm). Typically, the water
produced by Fairfax Water is considered "moderately hard" to "hard." The table
below is provided as a reference.
Grains Per
Gallon
Milligrams Per Liter
(mg/L) or Parts Per
Million (ppm)
Millimolesequvivalents per
liter (mmol-eq/L)
Classification
less than 1.0
less than 17.1
less than 4
Soft
1.0 - 3.5
17.1 - 60
4-8
Slightly Hard
3.5 - 7.0
60 - 120
-
Moderately Hard
7.0 - 10.5
120 - 180
8-12
Hard
over 10.5
over 180
over 12
Very Hard
The hardness of water is expressed also in mole-equivalents of Ca2+ and Mg2+
ions contained in 1 L of water (mmol-eq/L).
Hardness of drinking water of Lviv is within 7-8mmol-eq/L.
Formula that is used for calculation of water hardness:
Т =
m1
m2
mn
+
+ ⋅⋅⋅ +
E1 ⋅ V
E2 ⋅V
En ⋅V
,
where
Т – hardness of water;
Е1, Е2, …, Еn – equivalent masses of cations or salts that cause hardness;
V – volume of water, L;
m1, m2, …, mn – masses of cations or salts, mg.
24
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. –
1991.
3.4. Self Assessment Exercises
а) Review Questions
1. Write down chemical formulas of main natural compounds of IIA group metals.
How can be pure metals obtained from them?
2. True or False: alkaline earth metals do not react vigorously with water.
3. Which of the following is not an alkaline earth metal: Ba, K, Mg, Be, Ra?
4. Compare the alkali metals and alkaline earth metals with respect to: (a)
ionisation enthalpy, (b) basicity of oxides and (c) solubility of hydroxides.
5. Compare the solubility and thermal stability of the following compounds of the
alkali metals with those of the alkaline earth metals: (a) nitrates, (b) carbonates,
(c) sulphates.
6. In what ways does lithium show similarities to magnesium in its chemical
behavior?
7. Explain why beryllium and magnesium do not give color to flame whereas other
alkaline earth metals do so.
8. Describe the change in acid-base properties of oxides:
ВеО–MgО–СаО–SrO–ВаО.
9. Write the equations of the reactions to make the following transformations:
Na→ Na2O2 → O2 → MgО → MgCl2 → MgОНCl;
Na2CO3 → NaНCO3 → Na2CO3 → ВaCO3 → CO2 → СaCO3 → Сa(НCO3)2.
10. Write the molecular and ionic equations of the hydrolysis of the following salts:
CaS, CaCO3, Ca(CN)2, CaH2, Ba(NO2)2, BaSO3, MgClO, SrCl2.
11. Complete and balance the equations of the following reactions:
BeS + Na2S → …
Be(OH)2 + NaOH → …
[Be(H2O)4]Cl2 + H2O → …
BeS + SiS2 → … .
12. What alkaline metal is a main component in our bones?
13. One of the alkaline earth elements has only radioactive isotopes. What is the
name of this element?
25
b) Problems to Solve
1. Calculate the temporary hardness of water if for softening of 100 cm3 of water it
is necessary to add 6.02 g of sodium hydroxide?
Answer: 1.5 mmol-eq/l.
2. What is the general hardness of water if 1L of water contains 48.6 mg of
CaHCO3 and 24.6 mg of MgSO4.
Answer: 1 mmol-eq/l.
3. Calculate the solubility of СаС2О4 in water and in 0.1 М НCl solution if
Ksp(СаС2О4) = 4·10–9, аnd Кgen. = К1·К2 = 6.5·10–2·6.1·10–5 = 4·10–6.
Answer: 6.3·10–15 mol/l, 3.2·10–3 mol/l.
4. Equal volumes of 0.02 M solutions of calcium chloride and sodium carbonate
are mixed. Will a precipitate of calcium carbonate form?
Answer: yes.
5. What mass of MgSO4·7Н2О must be dissolved in 150 ml of water to prepare 10
% MgSO4 solution?
Answer: 38.7 g.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of s-elements of IIA group
and their compounds»
4.2.1. Flame test.
To heat strongly salts of group IIA elements in a flame. What is observed? Mark
the color of metals ions flame.
4.2.2. Preparation and properties of beryllium and calcium hydroxides.
а) To 1–2 ml of beryllium chloride (BeCl2) solution in a test-tube add sodium
hydroxide solution (NaOH) drop by drop. What is observed? Divide precipitate in
two test-tubes. Add solution of sulfuric acid into the first test-tube, and solution of
sodium hydroxide into the second. What is observed? Write equations of the
reactions in molecular and ionic forms.
b) To few pieces of calcium oxide in a test-tube add 3-2 cm3 of water. Observe
heating and increasing the amount of substance. Add more water, shake and divide
into two test-tubes. Into the first add few drops of phenolphthalein, into the second
– sodium hydrogencarbonate solution. Explain the changes that occur. Write
equations of the reactions.
4.2.3. Preparation and properties of magnesium hydroxide.
To 1–2 ml of magnesium chloride (MgCl2) solution in a test-tube add solution
of sodium hydroxide (NaOH) by drops. What is observed? Divide precipitate in
26
two test-tubes. Into the first test-tube add 2M HCl solution dropwise till its
complete dissolution and count the number of drops. Repeat the same experiment
with a solution in the second test-tube, but previously add 2 M solution of
ammonium chloride to it. In which case were spent more drops of HCl to dissolve
the precipitate? Explain this phenomenon. Write equations of the reactions.
4.2.4. Reactions of identification of Mg2+, Ca2+, Вa2+, Sr2+ ions.
а) Into the test-tube pour 1-2 cm3 solution of magnesium salt, add 1 cm3 of 2 M
hydrochloric acid and 1 cm3 of sodium hydrogenphosphate. To the prepared
mixture add drop by drop 1 cm3 of ammonia solution.
b) To 1–2 ml of calcium chloride (CaCl2) solution in a test-tube add identical
volume of solution of ammonium oxalate ((NH4)2C2O4). What is observed? Write
equation of reaction.
c) To 1-2 cm3 of barium salt solution pour the same amount of sodium sulfate.
Check solubility of obtained precipitate in hydrochloric acid solution. Write
equation of the reaction.
d) To 1-2 cm3 of strontium salt solution pour the same amount of potassium
chromate solution and heat. Write equation of the reaction.
4.2.5. Determination of temporary hardness of water.
For analysis, into a conical flask (250 ml by volume) pour 100 cm3 of water
investigated and add 2-3 drops of methyl orange. The content of the flask is titrated
with the 0.1 M HCl solution to change the color of the indicator. Titration is
complete when the sample solution will change yellow color on orange. Repeat
experiment 3-4 times. Write equation of the reaction:
Ca(HCO3)2 + 2HCl = CaCl2 + 2CO2 + 2H2O
The total number of Ca2+ and Mg2+ ions (general hardness of water) is
calculated by the formula:
Hhardness = С N ⋅V ⋅1000 (mmol-eq/l);
V1
where:
CN – normality of HCl solution;
V – volume of HCl solution, used for titration;
V1 – volume of water taken for the analysis.
5. Conclusions and Interpretations. Lesson Summary
Topic 5
General characteristic of p-elements. Boron, Alluminium
and properties of their compounds
1. Objectives
Aluminium is the most abundant metal in the Earth, making up about 8% of the
Earth’s crust and occurring in igneous rocks such as feldspars and micas. Boron is a
27
trace element in humans and is essential for some plants. Lack of boron can lead to
stunted plant growth, while an excess can also cause harm by inhibiting growth.
The ubiquity of aluminium in nature would suggest a biological function yet, until
recently, no specific function had been found. However, aluminium compounds are
toxic to most plants and animals - in animals they act as neurotoxins. Due to its
high reactivity, Al 3+ is able to interfere with several biological functions, including
enzymatic activities in key metabolic pathways, including Krebs cycle enzymes
such as succinic dehydrogenase. For some years now there has been concern about
the possible role of aluminium in a number of neurological disorders such as
Alzheimer’s disease. Scientists have observed increased levels of aluminium in the
brain tissues of some patients suffering from Alzheimer’s disease, amyotrophic
lateral sclerosis and Parkinson’s disease. Although various hypotheses have been
put forward to explain this, there is insufficient evidence to say that aluminium is
causative.
2. Learning Targets
To write equations of the reactions which can be used to characterize the
properties of p-elements of IIIA group. To get practical skills in Al3+ and BO33–
determining.
To know biological role of Aluminum, Boron and their uses.
3. Self Study Section
3.1. Syllabus Content
General characteristics of IIІА group elements. Electron deficiency and its
influence on the properties of elements and their compounds. General
characteristics of Boron. Simple substance and its chemical activity. Borides.
Compounds with hydrogen (boranes). Boron halogenides, hydrolysis and complex
formation. Boron oxide and boric acids. Equilibrium in aqueous solution. Sodium
tetraborate. Boric acid esthers. Organoaluminium compounds of boron. The
biological role of boron. Antiseptic properties of boric acid and its salts.
Aluminium. General characteristics. Simple substance and its chemical activity.
Amphoteric properties of aluminum and its oxide and hydroxide. Aluminate.
Aluminum ion as a complexing agent. Anhydrous aluminum salts and crystalline
hydrates. Halides. Aluminum hydride. Uses of aluminum and its compounds in
medicine and pharmacy.
3.2. Overview
Group IIIA of the Periodic Table contains five elements: boron, aluminium,
gallium, indium and thallium. The 'Header' element of each Group in the Periodic
Table often displays properties anomalous to the rest of the Group. Group IIIA is
no exception in that. This is the first group of the Periodic Table containing a nonmetal boron. The remaining four elements of group IIIA - aluminium, gallium,
indium and thallium, sometimes known as the 'Poor Metals' - have markedly
different physical and chemical properties from boron. For this reason, boron and
aluminium appear to be the most dissimilar elements in the Group. The elements in
28
the boron group are characterized by having three electrons in their outer energy
levels (valence layers). These elements have also been referred to as earth metals
and as triels.
Electronic configuration – ns2np1. Common oxidation state +3, but they can
have also an oxidation state +1. This ability increases from boron to thallium.
With the exception of hydrogen and helium, boron is the only non-metal with
less than four valence electrons, and thus forms only covalent compounds - the
boron ion does not exist. The behavior of boron to air depends upon the
crystallinity of the sample, temperature, particle size, and purity. Boron does not
react with air at room temperature. At higher temperatures, boron burns to form
boron (III) oxide, B2O3:
4B + 3O2 → 2B2O3.
Boron does not react with water under normal conditions.
Crystalline boron does not react with boiling hydrochloric acid, HCl, or boiling
hydrofluoric acid, HF. Powdered boron oxidizes slowly when treated with
concentrated nitric acid, HNO3. At the heating it reacts with concentrated H2SO4
and HNO3, and “royal water”:
B + 3HNO3(conc) =H3BO3 + 3NO2↑;
B + 4HCl + HNO3 = H[BCl4] + NO↑ + 2H2O.
Crystalline boron reacts with melts of alkalies in the presence of oxidizing
agent:
2B + 2NaOH + KClO3 → 2NaBO2 + KCl + H2O.
Powdered boron reacts with concentrated alkali solutions:
2B + 2KOH + 2H2O → 2KBO2 + 3H2↑.
Boron reacts with many metals at the heating (MnB, CaB6, Cr4B, Cr2B, Cr5B3
та ін.). Boron forms oxide B2O3 when heated in an atmosphere of oxygen at high
temperature:
T = 700 0 C
→ 2B2O3 (Boron oxide or Boric anhydride).
4B + 3O2  
Boron also forms nitride BN when heated in the atmosphere of nitrogen or
ammonia, and sulfide B2S3 when heated with sulfur:
∆
→
2BN;
∆
2B + 2NH3 
→ 2BN + 3H2;
∆
→
B2S3.
2B + 3S 
2B + N2
Boron oxide B2O3 forms ortoboric acid with water. Ortoboric acid reacts with
sodium hydroxide and sodium tetraborate is formed:
4H3BO3 + 2NaOH → Na2B4O7 + 7H2O.
Boron can form trichloride either by passing chlorine over heated boron or by
passing chlorine over heated mixture of its oxide and charcoal:
2B + 3Cl2
∆
→
2BCl3;
29
∆
B2O3 + 3C + 3Cl2 
→ 2BCl3 + 3CO
The boron halides react with water to give boric acid and the hydrogen halides.
For example,
BCl3 + 3H2O → H3BO3 + 3HCl.
Boron fluoride forms stable complex with F– ion:
4BF3 + 3HOH → H3BO3 + 3H[BF4].
Boron compounds have low toxicity to humans and other mammals, but are
very toxic to many insects, especially ants and cockroaches. A number of
commercial insecticides contain boric acid or other boron compounds. Boric acid is
also widely used as antiseptic, eyewash and as a treatment for some yeast
infections. It is a very weak acid and does not cause irritation of the skin or eyes.
Another of the uses of boron in medicine is in a cancer treatment known as Boron
Neutron Capture Therapy (BNCT). The treatment involves introducing the stable
isotope boron-10 into cancer cells. This boron isotope can absorb neutrons that
cause it to produce an alpha particle and a lithium ion. These particles carry a lot of
energy, but do not travel far; all their energy is released within the target cell,
damaging it, but leaving adjacent cells unharmed. It is thus possible to destroy
tumors by adding boron-10 to a compound that is absorbed more readily by
cancerous cells, then bombarding the tumor with neutrons.
Aluminium is a silvery white metal. The surface of aluminium metal is covered
with a thin layer of oxide that helps to protect the metal from attack by air. So,
normally, aulumium metal does not react with air. Aluminium burns in oxygen with
a brilliant white flame to form the trioxide aluminium(III) oxide, Al2O3. It is an
active metal:
4Al + 3О2 → 2Al2O3 + Q.
If the oxide layer is damaged, the aluminium metal is exposed to attack:
2Al + 6H2O → 2Al(OH)3 + 3H2↑.
Aluminium metal reacts vigorously with all the halogens to form aluminium
halides. So, it reacts with chlorine, Cl2, bromine, Br2, and iodine, I2, to form
respectively aluminium(III) chloride, AlCl3, aluminium(III) bromide, AlBr3, and
aluminium(III) iodide, AlI3:
2Al + 3Cl2→ 2AlCl3.
Aluminium dissolves in sodium hydroxide with the evolution of hydrogen gas,
H2, and the formation of aluminates:
2Al + 2NaOH + 6H2O → 2Na[Al(OH)4] + 3H2↑.
Concentrated nitric and sulfuric acids passivate aluminium metal. It reacts with
the acids at the heating:
2Al + 6H2SO4(conc) → Al2(SO4)3 + 3SO2↑ + 6H2O;
Al + 6HNO3(conc) → Al(NO3)3 + 3NO2↑ + 3H2O.
Oxide Al2O3 and hydroxide Al(OH)3 both have amphoteric properties:
Al2O3 + 6HCl → 2AlCl3 + 3H2O;
30
Al2O3 + 2NaOH + 3H2O → 2Na[Al(OH)4];
Al(OH)3 + 3HCl → AlCl3 + 3H2O;
Al(OH)3 + NaOH → Na[Al(OH)4].
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. –
1991.
3.4. Self Assessment Exercises
а) Review Questions
1. What is the electron configuration for Boron and how much unpaired electrons
does it have?
2. Is boric acid a protic acid? Explain.
3. Explain what happens when boric acid is heated.
4. Write reactions to justify amphoteric nature of aluminium.
5. Write the molecular and ionic equations of the hydrolysis of the following salts:
Al2(CO3)3, Al2(SO4)3, Al2S3, Na2B4O7·10H2O. Determine рН of medium.
6. Write equations of the transformation reactions: ortoboric acid → metaboric
acid → tetraboric acid → boron oxide.
7. How to make the next cycle of transformations:
Al→AlCl3→Al(NO3)3→KAlO2→K[Al(OH)4]→Al(OH)3→AlCl3→Al.
8. Name the following coordination compounds: BF3⋅NH3, K[BF4], K3[Al(OH)6],
Na3[AlF6], [Al(H2O)6]Cl3.
9. What is a boron deficiency?
10. What are the diffirent uses of boron?
b) Problems to Solve
1. What volume of hydrogen (at standart conditions) and mass of sodium
metaluminate can be obtained at the reaction of 27.2 kg (ω=25 %) of sodium
hydroxide solution with an excess of aluminum?
Answer: 615.4 L; 1394 kg.
2. At the burning of some amount of diborane 510 kJ of heat were evolved.
Prepared boron oxide melts with 42 g of baking soda NaHCO3. What is the mass
of obtained salt? Heat effect of diborane burning is 2040 kJ/mol.
Answer: 33 g.
31
3. Calculate [H+], рН and mass percentage of 0.001 М ortoboric acid H3BO3
solution (Кa = 5.7⋅10–10).
Answer: 7.55⋅10–7; 6.12; 0.0062 %.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: “Chemical properties of р-elements of IIIА
group”
4.2.1. Aluminum hydroxide amphoteric properties.
To 3-5 ml of an aluminum salt solution add sodium hydroxide solution till a
considerable amount of aluminum hydroxide will form. Divide the precipitate into
two test-tubes. Add some drops of hydrochloric acid solution to the first part of the
precipitate and the excess of sodium hydroxide solution to the second part. Observe
the dissolving of the precipitate in both cases. Write equations of the reactions of
formation of aluminum hydroxide precipitate and its dissolving in alkalis and acids
(the evidence of its amphoteric properties).
4.2.2. Hydrolysis of aluminium salts.
а) Write the equations of hydrolysis of aluminum sulfate and aluminum chloride
in molecular and ionic forms. Point out the medium of the solutions (pH). Confirm
the correctness of the conclusion by adding 2-3 drops of methyl orange to the
solutions of the salts (what is the color of the indicator?)
b) To 2–3 ml of aluminum sulfate solution add the same volume of sodium
carbonate solution. What is observed? Write the equation of the reaction between
these salts taking into account their hydrolysis and dissolving of a precipitate
formed in acids and bases.
4.2.3. Preparation of ortoboric acid.
Carefully (!) add 1 ml of concentrated sulfuric acid H2SO4 to 3 ml of hot 30%
solution of borax Na2B4O7 · 10H2O. Cool the test-tube with cold water steam till
the white crystals of boric acid will form. Write equation of reaction.
4.2.4 Sodium tetraborate hydrolysis.
Add 2-3 drops of phenolphthalein to 2–3 ml of sodium tetraborate solution.
What is observed? Write equation of the hydrolysis reaction. Point out the pH of
the medium of the solution.
4.2.5. The test to borate-ion (BO33- - ion).
To 3–4 ml of ortoboric acid solution or 0.5 g of boric acid powder in a vapor
bowl add 3-4 drops of concentrated sulfuric acid H2SO4 solution and pour 2-3 ml of
ethyl alcohol. Stir the mixture well. Put the bowl into exhaust-hood and set fire to
the alcohol. What color of the flame is observed? Write down equation of the
32
reaction of formation of the complex ether of ortoboric acid and ethyl alcohol at the
presence of sulfuric acid.
5. Conclusions and Interpretations. Lesson Summary
Topic 6
р-Elements of IVА group. Carbon, Silicon
1. Objectives
Carbon is one of the most common elements on Earth, and indeed it is common
is our everyday life. It is all around us. The most common molecules containing
carbon are carbon dioxide (CO2) and methane (CH4). All kinds of scientists study
carbon - biologists investigate the origins of life, oceanographers measure the
acidification of the oceans, and engineers develop diamond films for their tools, to
name just a few.
Silicon, the second most abundant element on Earth, is an essential part of the
mineral world. Its stable tetrahedral configuration makes it incredibly versatile and
is used in various ways in our everyday life. Found in everything from spaceships
to synthetic body parts, silicon can be found all around us, and sometimes even in
us. 27.6% of the Earth’s crust is made up of silicon. Although it is so abundant, it is
not usually found in its pure state, but rather its dioxide and hydrates. SiO2 is
silicon’s only stable oxide, and is found in many crystalline varieties. Its purest
form is quartz, but also as jasper and opal. Silicon can also be found in feldspar,
micas, olivines, pyroxenes and even in water.
2. Learning Targets
To write equations of the reactions which can be used to characterize the
properties of Carbon and its compounds. To get practical skills in reactions of ІVА
group elements identification.
To know biological role of Carbon and uses of its compounds in medical
practice.
3. Self Study Section
3.1. Syllabus Content
General characteristic of ІVА group elements. Carbon allotropes. Hybridisation.
Carbon as the basis of all organic molecules. Biological role of carbon. Physical
and chemical properties of its inorganic compounds. Activated charcoal.
Compounds of carbon with negative value of the oxidation state. Carbides, their
properties and use.
Compounds of carbon(ІІ). Carbon oxide(ІІ), its acid-base and redox properties.
Carbon oxide(ІІ) as a ligand.
Hydrogen cyanide. Toxic action.
Carbon dioxide(IV). Equlibrium in water solution. Carbonic acid, carbonates
and hydrogencarbonates. Hydrolysis and thermolysis of carbonic acid salts.
33
Compounds of carbon with halogens and sulfur. Carbon chloride(IV). Carbon
disulfide and tiocarbonates. Thiocyanates and cyanates. Physical and chemical
properties.
Silicon. General characteristic. The biological role. Silicides. Compounds with
hydrogen (silane), hydrolysis of silane. Silicon tetrafluoride and tetrachloride, their
hydrolysis. Hexafluorosilicates.
Compounds of silicon with oxygen, silicon dioxide(IV) (silica). Glass, its
properties and stability. Silicic acids. Silicates, their solubility and hydrolysis.
Silicone polymers. The use of silicon compounds in medicine.
3.2. Theoretical Backgrounds
Group IV of the Periodic Table of the Elements contains carbon (C), silicon
(Si) and several heavy metals. Members of this group conform well to general
periodic trends. The atomic radii increase as you move down the group, and
ionization energies decrease. Metallic properties increase as you move down the
group. Carbon is a non-metal; silicon and germanium are metalloids; and tin and
lead are poor metals. Electronic configuration - ns2np2. Short electronic
configuration of carbon 2s22p2. Valence ІІ and IV, oxidation states: –4; 0; +2; +4.
Carbon has two important qualities: small size and a unique electron configuration.
Since it is small, the p-orbital electrons overlap considerably and enable pi bonds to
form. Carbon is often referenced for its allotropes. It is second next to sulfur as the
element with the most allotropes. Carbon has three main solid state allotropes:
graphite, diamond and fullerenes (or fullerenes’ more memorable name:
Buckyballs). These allotropes differ greatly in form but are widely used in modern
production.
Carbon has both oxidizing and reducing properties:
1) С0 + nе– → С–n (oxidizing agent), reacts with some metals to give the metals
carbides under the high temperatures:
2K + 2C → K2C2;
4Al + 3C → Al4C3.
2) С0 – ne– → C+n, where n = 2, 4 (reducing agent). For example, it reacts with
oxygen to form oxides:
2C + O2 → 2CO;
2CO + O2 → 2CO2.
Carbon reacts also with sulfur and forms CS2, which reacts with some salts and
alkalis:
C + 2S → CS2;
Na2S + CS2 → Na2CS3;
KOH + CS2 → K2COS2 + H2O;
K2CS3 + 2HCl → 2KCl + H2CS3.
Solution of hydrogen cyanide in water is called hydrocyanic acid, or prussic
acid:
NH3 + CO → HCN + H2O.
34
Cyanide ion is active complex formation ion, it can form stable complexes with
many metals. When alkali metal cyanide is fused with sulfur, a thiocyanate is
formed, for example:
KCN + S → KSCN.
Hydrogen cyanide and its compounds are used for many chemical processes,
including fumigation, the case hardening of iron and steel, electroplating, and the
concentration of ores. It also is employed in the preparation of acrylonitrile, which
is used in the production of acrylic fibres, synthetic rubber, and plastics.
Carbon forms compounds with halogens (СНal4). CCl4 is liquid, that does not
mix with water. Compounds of carbon with fluorine – CF4 and CCl2F2 – freones are
used in refrigerators.
3) The most important inorganic carbon compounds are carbon monoxide and
carbon dioxide. Both are produced by combustion of any fuel containing carbon:
2C + O2 → 2CO;
2CO + O2 →2CO2.
Carbon monoxide CO is about 200 times better than O2 at bonding to
hemoglobin, the protein which transports O2 through the bloodstream from the
lungs to the tissues. Consequently a small concentration of CO in the air you
breathe can inhibit transport of O2 to the brain, causing drowsiness, loss of
consciousness, and death (After a few minutes of breathing undiluted auto exhaust,
more than half your hemoglobin will be incapable of transporting O2, and you will
faint). Because CO is colorless and odorless, your senses cannot detect it, and
people must constantly be cautioned not to run cars in garages or other enclosed
spaces. It doesn’t react with water:
CuO + CO → Cu + CO2↑;
Fe2O3 + 3CO → 2Fe + 3CO2↑.
Carbon dioxide reacts with sodium hydroxide solution in the cold to give either
sodium carbonate or sodium hydrogencarbonate solution - depending on the
reacting proportions:
CO2 + H2O → H2CO3;
2NaOH + CO2 → Na2CO3 + H2O;
NaOH + CO2 → NaHCO3.
Carbonic acid is the inorganic compound with the formula H2CO3. It is also a
name sometimes given to solutions of carbon dioxide in water, which contain small
amounts of H2CO3. It is an acid. When dissolved in water carbon dioxide exists in
equilibrium with carbonic acid. Carbonic acid plays an important role in keeping
the body’s pH stable. The normal pH of bodily fluids is around 7.4 and must be
kept close to this value in order for the body to function properly. If the pH
changes, whether up or down, enzymes can stop functioning, muscles and nerves
can start weakening, and metabolic activities becomes impaired. The bicarbonate
ion released from carbonic acid serves as a buffer that helps resist changes in pH.
This means it can act as an acid or a base as the need arises.
35
The salts of carbonic acids are called bicarbonates (or hydrogen carbonates)
and carbonates.
O C
O C
O
O
O H
O H
O C
O
H
O
Thermal stability: carbonates are decomposed to carbon dioxide and oxide
upon heating. Where as bicarbonates give carbonate, water and carbon dioxide.
Thermal stability of IA and IIA group carbonates (also of bicarbonates) increases
down the group as the polarizing power of the metal ion decreases.
CaCO3 → CaO + CO2↑,
Na2CO3 → NaHCO3 + CO2↑ + H2O.
Solubility in water: except Li2CO3, the IA group carbonates are fairly soluble in
water. The solubility increases down the group as the ionic nature increases. IIA
group carbonates are sparingly soluble in water as their lattice energies are higher
(it is due to increase in covalent nature). There is no clear solubility trend observed
down this group. But IIA group carbonates are soluble in a solution of CO2 due to
formation of HCO3-.
Reaction of СО32– identification:
CaCO3 + HCl → CaCl2 + CO2↑ + H2O;
Ca(OH)2 + CO2 → CaCO3↓ + H2O.
CaCO3 dissolves when excess of carbon dioxide СО2 is passed into the solution:
CaCO3 + CO2 + H2O → Ca(HCO3)2.
Although silicon plays a much smaller role in biology, it still plays an important
role in our world. It is the second most common element in the Earth’s crust (after
oxygen) and is the backbone of the mineral world. It is neither a metal nor
nonmetal, but a metalloid. Silicon is an inert metal, mainly reacting with halogens.
It may have acted as a catalyst in the formation of the earliest organic molecules.
Plants depend on silicates (such as [SiO4]4-) to hold nutrients in the soil, where their
roots can absorb them. People around the world have been using silicon (primarily
in the silica SiO2 molecule) for millennia in the creation of ceramics and glass. In
more recent history, the name "Silicon Valley" attests to its importance in the
computing industry - if carbon is the backbone of human intelligence, silicon is the
backbone of artificial intelligence. Silicon is found in beach sand, and is useful in
making concrete and brick.
Silicon forms two allotropic modifications: crystalline silicon, semiconductor
and amorphous silicon – brown powder, more reactive than the first one. Oxidation
states of silicon in its compounds: –4, 0, +4.
Silicon is rather uncreative element and is not attacked by acids:
Si + 2F2 = SiF4↑;
Si + O2 = SiO2;
Si + 3H2O(g) = H2SiO3 + 2H2↑;
3Si + 2N2 = Si3N4.
36
Silicon shows oxidizing properties when reacts with metals:
Si + 2Mg = Mg2Si.
Silicon reacts with many metals at high temperature forming silicides:
Ca2Si + 4H2O → 2Ca(OH)2 + SiH4↑;
Ca2Si + 4HCl → 2CaCl2 + SiH4↑.
Like the organic compounds of carbon, the oxygen compounds of silicon which
make up most of the Earth’s crust have already been described. These substances
illustrate a major contrast between the chemistry of carbon and silicon. The latter
element does form a few compounds, called silanes, which are analogous to the
alkanes, but the Si—Si bonds in silanes are much weaker than Si-O bonds.
Consequently the silanes combine readily with oxygen from air, forming Si-O-Si
linkages. Unlike the alkanes, which must be ignited with a spark or a match before
they will burn, silanes catch fire of their own accord in air:
2Si4H10 + 13O2 → 4SiO2 + 5H2O.
Silanes are very reactive compounds and strong reducing agents:
SiH4 + (x + 2)H2O → SiO2⋅xH2O + 4H2↑;
SiH4 + 2NaOH + H2O → Na2SiO3 + 4H2↑.
All four of the silicon tetrahalides are known. Unlike the carbon tetrahalides,
the silicon tetrahalides are completely hydrolized in water. The only exception is
silicon tetrafluoride, which produces the hexafluorosilicate ion. Silicon
tetrachloride as a fuming liquid is used in the manufacture of elemental silicon.
Silicon tetrahalides complitly hydrolyze in water solution:
SiCl4 + 3H2O → H2SiO3 + 4HCl;
SiF4 + 3H2O → H2SiO3 + 4HF;
SiF4 + 2HF → H2[SiF6].
Silicon doesn’t double bond with oxygen. Silicon atoms are bigger than carbon.
That means that silicon-oxygen bonds will be longer than carbon-oxygen bonds.
The most common compound of silicon is SiO2 - silicon(IV) oxide (silica).
Silicon dioxide reacts with sodium hydroxide solution, but only if it is hot and
concentrated. Sodium silicate solution is formed:
SiO2 + 2NaOH → Na2SiO3 + H2O;
SiO2 + Na2CO3 → Na2SiO3 + CO2↑.
Silicon dioxide doesn’t react with water, because of the difficulty of breaking
up the giant covalent structure. It is very hard and stable substance, it does not react
with any acids except hydrofluoric acid:
SiO2 + 4HF → SiF4↑ + 2H2O.
Silicon dioxide is weak oxidizing agent (it reacts with strong reducing agents at
the high temperature):
2Mg + SiO2 → 2MgO + Si;
3C + SiO2 → SiC + 2CO.
Silicon forms some acids with the general formula [SiOx(OH)4-2x]n. Some
simple silicic acids have been identified, but only in very dilute aqueous solution,
37
such as metasilicic acid (H2SiO3), orthosilicic acid (H4SiO4, pKa1=9.84, pKa2=13.2
at 25 °C), disilicic acid (H2Si2O5), and pyrosilicic acid (H6Si2O7); however in the
solid state these probably condense to form polymeric silicic acids of complex
structure. All silicic acids are very weak acids:
Na2SiO3 + 2H2O + CO2 → 2NaHCO3 + H2SiO3.
Silicates are the minerals containing silicon and oxygen in tetrahedral SiO44units which are linked together in several patterns. Potassium and sodium silicates
are soluble salts; their aqueous solutions are called liquid glass (silicate glue).
Na2SiO3 and K2SiO3 hydrolyze in water solution:
2Na2SiO3 + H2O ⇄ Na2Si2O5 + 2NaOH.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. –
1991.
3.4 Test Yourself
а) Review Questions
1. What is the hybridization state of carbon in (a) CO32-, (b) diamond, (c)
graphite?
2. Write the ground-state electron configuration and orbital notation for the atom
of carbon.
3. What is an electron configuration for an atom of silicon in excited state?
4. Why CO2 is a gas and SiO2 is a solid?
5. Give one method for industrial preparation and one for laboratory preparation
of CO and CO2 each.
6. Write equations of hydrolysis of CaC2, Al4C3, K2CO3, KHCO3. Specify the pH
of the medium.
7. How can be CO2 gas detected?
8. Write the reactions of thermal decomposition of calcium carbonate, calcium
hydrogencarbonate, ammonium hydrogencarbonate and carbonate, magnesium
carbonate, sodium carbonate. Explain, in which case there is no thermal
decomposition of the compounds.
9. Write the equations of the following transformation:
C →CH4→CO→CO2→CaCO3→Ca(HCO3)2→CaCO3→CaO→Ca(OH)2.
10. Write the reactions of the preparation and hydrolysis of sodium silicate and
38
aluminum silicide.
11. Write the equations of the following transformation:
Si → Mg2Si → SiH4 → Si → Na2SiO3 → H2SiO3 → H2Si2O5
↓↑
↑
SiO2 SiH4
b) Problems to Solve
1. Calcium and aluminum carbides were hydrolyzed and formed a mixture of gases
which is in 1.6 times lighter than oxygen. Calculate the mass percentage of
carbides in initial mixture.
Answer: 47.06 % СаС2, 52.94 % Al4C3.
2. Calculate degree of dissociation, hydrogen concentration and pH of 0.05 moleq/L H2CO3 solution. Ka = 4⋅10–10.
Answer: 8.94⋅10–5; 4.47⋅10–6 mol/L; рН = 5.35
3. Calculate the general hardness of water if for softening of 50 cm3 of water it is
necessary to add 10.6 g of sodium carbonate.
Answer: 4 mol-eq/L.
4. Define the configuration of glass if it contains 13.8 % Na2O; 12.7 % CaO and
73.5 % SiO2.
Answer: Na2O⋅CaO⋅6SiO2.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: “Chemical properties of Carbon and Silicon and
their compounds”
4.2.1. Hydrolysis of calcium carbide.
Into a test tube with a vent tube pour 5 ml of water and add 0.1 g of calcium
carbide. Pass the released gas through benzene solution of bromine (in fume
cupboard). Write equations of the reactions of calcium carbide hydrolysis and
interaction of gas with bromine.
4.2.2. Hydrolysis of carbonates.
а) To 1–2 ml of solution of sodium or potassium carbonate (Na2CO3 or K2CO3)
in a test-tube add 2–3 drops of the phenolphthalein indicator solution. What is
observed? Write equation of the reaction.
b) To 1–2 ml of sodium carbonate (Na2CO3) solution in a test-tube add solution
of iron (III) chloride (FeCl3) and heat the mixture. What is observed? Write
equation of the reaction.
4.2.3. Thermal decomposition of ammonium, sodium, calcium carbonate and
sodium hydrogencarbonate.
39
Into four test-tubes add 0.5 g of ammonium, sodium and calcium carbonate and
sodium hydrogencarbonate. Close tubes with stoppers and vent tubes which are
immersed into a solution of lime water. Heat the content of each test-tube in the
flame of a gas burner. Mark whether the turbidity of the lime water in all cases is
observed. Write equations of the reactions.
4.2.4. Reactions of carbonates with acid.
Into four test tubes that contain 0.5 g of ammonium carbonate, sodium, calcium
and sodium hydrogencarbonate add 5 ml of HCl solution. What is observed? Write
equations of the reactions.
4.2.5. Hydrolysis of silicates.
а) To 1–2 ml of sodium silicate (Na2SiO3) solution in a test-tube add 2–3 drops
of the phenolphthalein indicator solution. What is observed? Write equation of the
reaction.
b) Mix 2–3 ml of sodium silicate (Na2SiO3) solution with double volume of
ammonium chloride solution and heat the mixture. Observe formation of silica gel
acid precipitate. Identify the smell of the released gas. Write equation of the
reaction.
4.2.6. Preparation of silicic acid.
To 1–2 ml of sodium silicate Na2SiO3 solution in a test-tube add solution of
concentrated sulfuric acid H2SO4 by drops (carefully). What is observed? Write
equation of the reaction.
5. Conclusions and Interpretations. Lesson Summary
Topic 7
р-Elements of IVА group. Germanium family elements
(Germanium, Tin, and Lead)
1. Objectives
Lead rarely occurs as a pure element in the Earth. Its most common ore is
galena, or lead sulfide (PbS). Other ores of Lead are anglesite, or lead sulfate
(PbSO4); cerussite, or lead carbonate (PbCO3); and mimetite (PbCl2·Pb3(AsO4)2).
Tin is a highly workable metal that was once as valuable as silver for jewelry,
coins, and special dishware. Today it is used as sheets in the construction of
buildings and roofs, for soldering or joining metal parts, for storage containers, and
in alloys like bronze. Germanium is widely used in semiconductors, infrared
prisms, reflectors in projectors, wide angle lenses and dentistry. The biological
function of these metals is not entirely clear. It is known that lead and its
compounds are toxic. Germanium is not thought to be essential to the health of
plants or animals. Some of its compounds present a hazard to human health,
40
however. For example, germanium chloride and germanium fluoride (GeF4) are a
liquid and a gas, respectively, that can be very irritating to the eyes, skin, lungs, and
throat. The substance may cause effects on the blood, resulting in lesions of blood
cells. Exposure may result in death. Excess of lead in the body is the cause of
serious violations of the central nervous system and mechanisms of synthesis of
hemoglobin. These metals are used in medical practice: tin in dentistry and some
compounds of lead (lead acetate and aluminum acetate) are used for treatment of
some diseases.
2. Learning Targets
To study chemical properties of germanium, tin, lead and their compounds, be
able to write the reactions that characterize chemical properties of p-elements of
group IV. To study the methods of Sn2+ and Pb2+ ions determining in the
environment. To know the basic compounds of germanium, tin and lead; their
biological role and applications in chemistry and medicine.
3. Self Study Section
3.1. Syllabus Content
Genaral characteristics of Germanium, Lead and Tin. Compounds with
hydrogen. Compounds with halogens EF2 and EF4, their behavior in aqueous
solutions. Oxides. Amphoteric properties of oxides. Stannic acid. Stannites
(Na2SnO2) and stannates (Na2SnO3). in and lead hydroxocomplexes. Reducing
properties of tin (II) compounds. Lead (IV) oxide as a strong oxidizing agent.
Soluble and insoluble salts of tin and lead. Redox reactions in solutions. The toxic
effects of Pb compounds.
Uses of lead compounds (lead (ІІ) oxide and lead acetate) in medicine and
pharmacy. Uses of tin and lead compounds in the pharmaceuticals analysis. Toxic
effect of leadorganic compounds.
3.2. Theoretical Backgrounds
All of these elements have four electrons in their outermost energy level.
Germanium is metalloid; it can form +4 ions. Tin and lead both are metals while
flerovium is a synthetic, radioactive (its half-life is very short), element that may
have a few noble-gas-like properties, though it is still most likely a post-transition
metal. These elements become more metallic in character with increasing atomic
weight, and while the chemical properties of lead bear some resemblance to those
of the other members of the group, it is chemically most similar to the metal, tin. In
its compounds, lead usually has an oxidation state of +2, which means that it
donates two electrons to other atoms or molecules. Less commonly, it can have an
oxidation state of +4. At temperature of 250 °C, germanium slowly oxidizes to
GeO2. Germanium dissolves slowly in concentrated sulfuric acid, and is insoluble
in diluted acids and alkalis. It reacts violently with molten alkalis to produce
[GeO3]2-. The common oxidation states that Germanium occurs in are +4 and +2.
Under rare conditions, Germanium also occurs in oxidation states of +3, +1, and -4.
There are two forms of oxides of germanium, germanium dioxide (GeO2) and
41
germanium monoxide (GeO). The two most important compounds of germanium
are the dioxide (GeO2) and the tetrachloride (GeCl4). Germanates, formed by
heating the dioxide with basic oxides, include zinc germanate (Zn2GeO4), used as a
phosphor (a substance that emits light when energized by radiation). The
tetrachloride, that is an intermediate in germanium obtaining from its natural
sources, is a volatile, colorless liquid that freezes at about -50° C and boils at 84°
C.
Lead combines with oxygen to form several oxides. “Red lead,” formed by
heating of lead in air, has the formula Pb3O4, but is thought to be a compound of
lead oxide (PbO) and lead dioxide (PbO2). Lead oxide, also known as litharge, is
formed when the metal is heated strongly in air and can take the form of a yellow
powder or a red crystalline material. “White lead” is basic lead carbonate
(2PbCO3·Pb(OH)2). It was formerly widely used in paints due to its strong white
color before being largely replaced by non-toxic titanium dioxide.
Metallic properties of elements in a row Ge–Sn–Pb increase. When hot and
concentrated HNO3 reacts with the metals, the metals are oxidized to germanium
oxide GeO2, stannic acid (H2SnO3) and lead nitrate while the acid gets reduced to
NO2:
Ge + 4 HNO3 = GeO2 + 4 NO2↑ + 2 H2O;
germanium(IV) oxide
Sn + 4 HNO3 = H2SnO3 + 4 NO2↑ + H2O;
β -stannic acid
Pb + 4 HNO3 = Pb (NO3)2 + 2 NO2↑ + 2 H2O.
lead(II) nitrate
Germanium, lead and tin show metallic properties but with active metals they
react as silicon. Germanides, stannides and plumbides are formed:
Ge + 2 NaOH + 2H2O2 = NaGeO3 + 3 H2O;
Sn + 2 NaOH = Na2SnO2 + H2↑;
Pb(OH)2 + 2 NaOH = Na2PbO2 + 2 H2O.
Lead is resistant to corrosion by most acids, due to the fact that the majorities of
lead salts have little or no solubility in water and form a layer that protects the lead
from further action. It will, however, react with acetic and nitric acids, as the salts
formed by these reactions - lead acetate and lead nitrate, respectively - are very
soluble. Basic character of their oxides and hydroxides increases with increasing of
ions radii: GeO2 shows acidic properties, PbO – basic. Stability of covalent
hydrides ЕН4 decreases in a row Ge – Sn – Pb.
Tin and lead, although with very low abundances in the crust, are nevertheless
common in everyday life. They occur in highly concentrated mineral deposits, can
be obtained easily in the metallic state from those minerals, and are useful as metals
and as alloys in many applications. Germanium, on the other hand, forms few
characteristic minerals and is most commonly found only in small concentrations in
association with the mineral zinc blende and in coals. Although germanium is
indeed one of the rarer elements, it assumed importance upon recognition of its
42
properties as a semiconductor.
Probably the best known of the properties of lead is its toxicity. Cases of acute
lead poisoning are rare, but it is a cumulative poison, and chronic exposure to low
levels of lead can lead to a variety of serious symptoms. It deactivates the enzymes
that manufacture hemoglobin, leading to a build-up of the precursor chemical - this
can paralyze the gut, resulting in constipation and abdominal pain, and cause a
build-up of fluid in the brain, causing headaches. Over a longer period, it causes
anemia and neurological problems.
Chronic lead poisoning has been a significant problem due to the widespread
use of lead in applications that have allowed it to enter the environment. For
example, metallic lead was formerly used in water pipes and lead compounds have
been used in paints. These uses have been discontinued in most countries, and lead
piping replaced by non-toxic alternatives. The biggest source of lead in the
environment has been the compound tetraethyl lead, which was added to gasoline
to achieve smoother combustion. Due to concerns about the health effects of lead in
the environment, particularly on children in urban areas, leaded gasoline has also
been phased out in many countries.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. –
Kharkiv: Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–
131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company.
– 1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry.- Prentice Hall. –
1991.
3.4. Test Yourself
а) Review Questions
1. Write equations of the dissolution reactions of alloy of tin and lead in:
а) НNO3 diluted; b) НNO3 concentrated at heating;
c) Н2SO4 diluted and concentrated.
2. Write the molecular and ionic equations of the hydrolysis of the following salts:
а) tin (II) chloride; b) tin (IV) chloride; c) lead (II) nitrate.
3. Write equations of the reactions that show: а) reducing properties of tin (II)
chloride with Hg (II) and Bi (III) salts; b) oxidizing properties of PbO2.
4. Write equations of the transformation reactions:
а) Sn → SnO → Na2SnO2 → SnS → (NH4)2SnS3 → H2SnS3 → SnS2;
b) Pb → Pb(NO3)2 → PbI2 → K2[PbI4].
43
5. Write equations of the following reactions:
b) PbS + H2O2 → …;
а) PbS + HNO3 → …;
c) Na2SnO3 + HCl(excess) → …; d) PbO2 + HCl(excess) → …;
f) SnCl2(excess) + HgCl2 → …
e) SnS2 + Na2S →…;
b) Problems to Solve
1. What volume of 2 M sodium hydroxide solution should be added to 200 g of 5%
aqueous solution of stannous (II) chloride to get sodium stannite?
Answer: 105.25 ml.
2. Calculate normality and mass percentage of stannous (II) chloride solution
which was prepared at mixing 2.5 L of 22 % solution (ρ = 1.19 g/ml) and 1.5 L
of 4 % solution (ρ = 1.03 g/ml).
Answer: 15.9 %; 1.9 mol-eq/L.
3. Considering that natural lead consists of four isotopes: 204Pb – 1.37 %, 206Pb –
25.15 %, 207Pb – 21.11 %, 208Pb – 52.38 %, determine the atomic mass of Pb.
Answer: 207.25 а.m.u.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: “Chemical properties of the p-elements of
Germanium subgroup and their compounds”
4.2.1. Amphoteric properties of stannous (II) hydroxide and lead (II) hydroxide.
To 1-2 ml of Sn (II) and Pb (II) salt solutions add drop by drop a solution of
sodium hydroxide to form stannous (II) and lead (II) hydroxides. Show their
amphoteric properties by adding the excess of sodium hydroxide solution and acid
solution to prepared precipitates. What acid should be taken to dissolve the
precipitate of lead (II) hydroxide and why? Write equations of the reactions of
stannous (II) and lead (II) hydroxide obtaining.
4.2.2. Hydrolysis of Tin (II), Tin (IV) and Lead (II) salts.
а) To 1-2 ml of stannous (II) chloride solution and lead (II) nitrate solution add
2-3 drops of methyl orange. What is observed? Write equations of the hydrolysis.
Specify the pH.
b) Open for 1-2 minutes a glass with anhydrous stannous (IV) chloride
(experiment is carried out only in a ventilating hood or fume cupboard). What is
observed? Write equation of hydrolysis of this salt. How to explain the formation
of smoke?
4.2.3. Tin (II)-ion as a reducing agent.
а) Reaction with mercury (II) chloride.
To 1–2 ml of mercury (II) chloride HgCl2 solution in a test-tube add solution of
44
stannous chloride SnCl2. What is observed? Write the equation of the reaction. To
obtained precipitate add SnCl2 solution. Write equation of the reaction of white
precipitate of mercury (I) chloride obtaining and its transformation into gray
precipitate of metallic mercury.
b) Reaction with bismuth (III) salts.
To 1–2 ml of stannous chloride SnCl2 solution in a test-tube add solution of
sodium hydroxide NaOH drop by drop till precipitate forms. Dissolve obtained
precipitate in the excess of NaOH and add 1 ml of bismuth (III) nitrate Bi(NO3)3
solution. What is observed? Write equations of the reactions.
4.2.4. Oxidizing properties of lead dioxide.
Into the test tube containing 6.5 ml of sulfuric acid and 2 drops of manganese
(II) sulfate, add 0.1 g of lead dioxide powder. Boil the content of the tube and
check the color of the solution. Write equation of the reaction.
4.2.5. Reactions of identification of the lead (II)-ion.
а) Reaction with potassium iodide.
To 1–2 ml of lead (II) nitrate (Pb(NO3)2) solution in a test-tube add a solution
of potassium iodide (KI) by drops. What is observed? Write the equation of the
reaction. Carry out dissolution of the precipitate by heating in the presence of 5.3
ml of acetic acid. Observe formation of golden crystals of lead (II) iodide after
cooling.
b) Reaction of Pb (II) salt with potassium chromate.
Into the test-tube add 4-5 drops of Pb(II) salt solution and 5-6 drops of
potassium chromate. What is the color of precipitate? Dissolve the precipitate in
hydrochloric acid, acetic acid and alkali solution. Write equations of the reactions
of lead (II) chromate obtaining and its dissolution in acids and alkali solution.
c) Reaction of Pb(II) salt with hydrogen sulfide and sodium sulfide.
To 1-2 ml of Pb(II) salt solution add drop by drop an aqueous solution of
hydrogen sulfide H2S or sodium sulfide Na2S. What is the color of precipitate?
Write equation of the reaction of black precipitate of lead (II) sulfide obtaining.
Write equation of the reaction of lead (II) sulfide with nitric acid and hydrogen
peroxide.
4.2.6. Preparation and properties of thiostannic acid (тіостанатної кислоти).
To 2-3 ml of stannous chloride SnCl2 solution add hydrogen sulphide solution
till precipitate forms. Decant the upper layer of the solution. To obtained
precipitate add ammonium polysulfide solution. What is observed? Write equations
of the reactions. To obtained solution add few drops of hydrochloric acid and heat
it (only in a ventilating hood). What is the color of precipitate? Write equation of
the decomposition reaction of thiostannic acid.
5. Conclusions and Interpretations. Lesson Summary
45
Topic 8
р-Elements of VА group. Nitrogen and its compounds in
the negative oxidation states
1. Objectives
Nitrogen is a gaseous element that is abundant in the atmosphere as the
molecule of nitrogen (N2). Nitrogen is the most abundant terrestrial element in an
uncombined state, as it makes up 78 percent of Earth’s atmosphere as N2, but it is a
minor component (19 parts per million) of Earth’s crust. Nitrogen exists as two
isotopes: 14N (99.63% relative abundance) and 15N (0.4% abundance). In its
reduced state nitrogen is essential for life because it is a constituent of the
nucleotides of deoxyribonucleic acid (DNA) and ribonucleic acid (RNA) molecules
that encode genetic information and of the amino acids of proteins.
2. Learning Targets
To study the chemical properties of nitrogen and its compounds with negative
values of the oxidation state.
To know the biological role and use of these compounds in medicine and the
national economy.
3. Self Study Section
3.1. Syllabus Content
General characteristics of the elements of VA group. Nitrogen, phosphorus,
arsenic. Their biological role in the nature and human body.
Nitrogen. General characteristics. Compounds with different oxidation states.
Nitrogen as a simple substance. The reasons for its low chemical activity. Nitrogen
molecule as a ligand. Compounds with negative oxidation states. Nitrides. Acidbase and redox properties of ammonia. Amides. Ammonia ion and its salts, acidbase properties, thermal decomposition. Acid-base and redox properties of
hydrazine and hydroxylamine.
3.2. Theoretical Backgrounds
Nitrogen is one of the most interesting of all chemical elements. It is not a very
active element. It combines with relatively few other elements at room temperature.
Yet, the compounds of nitrogen are enormously important both in living organisms
and in industrial applications. Five of the top fifteen chemicals that are produced
synthetically by chemical producers are compounds of nitrogen or the element
itself. Nitrogen makes up more than three-quarters of the Earth’s atmosphere. It is
also found in a number of rocks and minerals in the Earth’s surface. It ranks about
number 32 among the elements in terms of abundance in the Earth’s crust.
A nitrogen atom has the electronic structure represented by 1s22s22p3. The five
outer shell electrons screen the nuclear charge quite poorly, with the result that the
effective nuclear charge felt at the covalent radius distance is relatively high. Thus
nitrogen atoms are relatively small in size and high in electronegativity, being
46
intermediate between carbon and oxygen in both of these properties. The electronic
configuration includes three half-filled outer orbitals, which give the atom the
capacity to form three covalent bonds. The nitrogen atom should therefore be a
very reactive species, combining with most other elements to form stable binary
compounds, especially when the other element is sufficiently different in
electronegativity to impart substantial polarity to the bonds.
Valence of nitrogen III and IV in [NH4]+ ion.
Oxidation states of nitrogen vary from –3 to +5.
oxidation
state
–3
–2
examples
NH3
N2H4
–1
NH2OH
0
+1
+2
+3
+4
+5
N2
N2O
NO
N2O3
HNO2
NO2
N2O5
HNO3
Under normal conditions nitrogen reacts only with lithium and when heated
interacts with other metals (K, Na, Ca, Mg), and nonmetals (O2, H2, Si, halogens).
Among compounds with negative oxidation states important are ammonia NH3,
ammonium hydroxide NH4OH, ammonium salts NH4Cl, NH4NO3, hydrazine N2H4,
hydroxylamine NH2OH, and metal nitrides Na3N, Mg3N2.
These compounds are readily soluble in water, forming a weak base:
NH3 + H2O ⇄ NH3 ⋅ H2O ⇄ NH4+ + OH–, Кb = 1.8⋅10-5;
NH2OH + НОН ⇄ NH3OH+ + ОН–, Кb = 2⋅10-8.
Nitrides of s-elements of I and II groups are easily decomposed by water,
forming alkali and ammonia:
K3N + 3HOH = KOH + NH3;
Ca3N2 + 6HOH = 3Ca(OH)2 + 2NH3.
Nitrogen can act as a ligand in complexation reactions with ions of d-elements:
CuCl2 + 4NH3 = [Cu(NH3)4]Cl2;
AgBr(s) + 2NH3 = [Ag(NH3)2]Br.
Oxidation state of nitrogen in ammonia and ammonia salts has the lowest value
(−3). It has strong reducing properties in the redox reactions:
4NH3 + 3О2 = 2N2 + 6H2O;
NH3 + 3Cl2 = NCl3 + 3HCl;
2NH3 + 3NiО = N2 + 3Ni + 3H2O;
2NH4Cl + 3CuO = 3Cu + N2 + 2HCl + 3H2O.
Hydrazine shows reducing properties; hydroxylamine may be both oxidizing
agent and reducing agent:
N2H4 + 2І2 + 4NaOH = N2 + 4NaI + 4H2O;
2NH2OH + 4FeSO4 + 3H2SO4 = 4Fe2(SO4)3 + (NH4)2SO4 + 2H2O;
2NH2OH + I2 + 2KOH = N2 + 2KI + 4H2O.
Ammonia is a colorless, highly irritating gas with a pungent, suffocating odor at
room temperature. In pure form, it is known as anhydrous ammonia and is
hygroscopic (readily absorbs moisture). Ammonia has alkaline properties and is
corrosive. Ammonia gas dissolves easily in water to form ammonium hydroxide, a
47
caustic solution and weak base. Ammonia gas is easily compressed and forms a
clear liquid under pressure. Ammonia is usually shipped as a compressed liquid in
steel containers. Ammonia is not highly flammable, but containers of ammonia may
explode when exposed to high heat. Hydrazine is a product of oxidation of
ammonia by NaClO:
2NH3 + NaClO = N2H4 + NaCl + H2O.
About 80% of the ammonia produced by industry is used in agriculture as
fertilizer. Ammonia is also used as a refrigerant gas, for purification of water
supplies, and in the manufacture of plastics, explosives, textiles, pesticides, dyes
and other chemicals. It is found in many household and industrial-strength cleaning
solutions. Household ammonia cleaning solutions are manufactured by adding
ammonia gas to water and can be between 5 and 10% ammonia. Ammonia
solutions for industrial use may be concentrations of 25% or higher and are
corrosive. Hydrazine and hydroxylamine are intermediates in the cycle of
atmospheric nitrogen fixation by the enzyme nitrogenase.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. –
1991.
3.4. Test Yourself
а) Review Questions.
1. Give the comparative characteristics of the subgroups of nitrogen atoms by
specifying: a) electronic configuration; b) the valence opportunities; c)
oxidation state.
2. Describe the chemical properties of nitrogen. What substances react with
nitrogen and under what conditions? Give examples of the reactions.
3. Specify the most important methods of obtaining of ammonia in the laboratory
and industry.
4. Write the scheme of ammonium cation formation and specify the type of the
chemical bond in the ion.
5. Give examples of addition and substitution reactions that are typical for
ammonia.
6. What acid-base properties are typical for nitrogen compounds with negative
48
values of oxidation state? Give examples of the reactions that confirm these
properties.
7. Write equations of the following reactions and balance them:
a) NH3 + KMnO4 → MnO2 + ⋅⋅⋅; b) N2H4 + KMnO4 + H2SO4 → MnSO4 + ⋅⋅⋅;
c) NH3 + СаOCl2 → CaCl2 + ⋅⋅⋅; d) N2H4 + H2O2 → N2 + ⋅⋅⋅.
8. Write equations of the following reactions and balance them:
a) NH2ОН + Н2SO3 → ;
b) NH2ОН + KMnO4 + H2SO4 → MnSO4 + ⋅⋅⋅;
t
c) NH2ОН 
.
→
9. Write equations of the thermal decomposition reactions of the following salts:
ammonium carbonate, ammonium dichromate, ammonium dihydrogenphosphate.
10. Perform a simple test how to detect the ammonia gas using the litmus paper.
b) Problems to Solve
1. Write reactions of the following transformations:
N2 → NO →NO2 → HNO3 → NH4NO3 → NH3 →N2H4
↓
N2O → N2 → Ca3N2.
2. Calculate the heat effect of combustion reaction of hydrazine if
∆Hoformation(N2H4(l))=50.5 kJ/mol, ∆Hoformation(H2O(l))=-285.8 kJ/mol.
Answer: 622.1 kJ.
3. Calculate the density by air of gas mixture obtained at the thermal
decomposition of ammonium hydrogencarbonate.
Answer: 0.9.
4. 2.5 L of ammonia were dissolved in 5 L of water (at standart conditions). What
is the molar concentration of the prepared solution?
Answer: 0.02 mol/L.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of Nitrogen».
4.2.1. Preparaton of nitrogen.
Into a test tube with the vent tube add 4-5 ml of saturated sodium nitrite and
ammonium chloride solutions, close the tube and heat it. What is observed? What
gas is released? Write equation of the reaction.
4.2.2. Preparation of ammonia.
In a porcelain mortar mix 2-3 g of ammonium chloride with the same amount of
calcium hydroxide. What is the odor of a mixture? What gas is released? Write
equation of the reaction.
49
4.2.3. Reaction of identification of ammonium-ion.
а) To 1–2 ml of ammonium hydroxide (NH4OH) solution in a test-tube add
Nessler's reagent (a solution of potassium tetraiodomercurate(II), K2[HgI4]). What
is observed? Write equation of the reaction:
I
Hg
I
Hg
NH2
2 K 2 [H gI 4 ] + 2 K O H + N H 4 C l
I
+ 5K I + + K C l + 2 H 2 O
b) Into a test-tube add a few drops of any ammonium salt and 3-4 drops of 2 M
alkali solution. Heat the test tube and put to the top of the test-tube red litmus
paper. Explain why litmus paper changes color into the blue and write equations of
the reactions.
4.2.4. The shift in the equilibrium of ammonia solution.
To a solution of ammonia, add a few drops of phenolphthalein. Pour colored
solution into four test tubes. Into the first tube add some crystalline ammonium
acetate, into the second - diluted solution of HCl, heat the third tube to boiling, and
leave the fourth tube for comparison. How does adding of CH3COONH4, HCl and
heating shift the balance? Write equation of equilibrium in the ammonia - water
system.
4.2.5. Properties of an aqueous ammonia solution.
а) On a glass put 2-3 drops of an aqueous ammonia solution, add 1-2 drops of
phenolphthalein. What is the color of the indicator?
b) Into a test tube add 2-3 drops of an aqueous ammonia solution and pour the
same amount of iron (II) sulfate solution. What is observed? Write equation of the
reaction and explain the change in color of the precipitate.
4.2.6. Preparation of nickel ammoniacate.
To a solution of nickel (II) sulfate add excess of an aqueous ammonia solution.
What is observed? Explain the change in color of the solution and write equation of
the reaction.
4.2.7. Sublimation of ammonium chloride.
Into dry long test-tube place about 1 g of ammonium chloride, fix it with the
tripod clamps in an inclined position and heat. Observe the formation of white gas
on the walls of the upper, cooler part of the tube. Explain this phenomenon.
4.2.8. Thermal decomposition of ammonium salts.
а) Place into a dry test-tube 1-2 g of ammonium dihydrogenphosphate
NH4H2PO4 or ammonium hydrogencarbonate NH4NCO3 and heat. Lift the red
litmus paper to upper hole of the test-tube. What is observed? Write equation of the
reaction.
b) Place into a test-tube about 1 g of ammonium dichromate, fix it vertically in
a tripod and heat to the beginning of the reaction. What is observed? What gas is
50
released? Similarly to the previous experiment, check whether it is ammonia. Make
a conclusion about the thermal stability of ammonium salts.
4.2.9. Ammonium salts hydrolysis.
Into two test-tubes add a few crystals of NH4Cl ammonium chloride and
(NH4)2CO3 ammonium carbonate and dissolve them in 2-3 ml of distilled water.
Lift the red litmus paper to upper hole of the test-tube. What is observed? Write
equation of the hydrolysis reaction.
4.2.10. Reducing properties of hydrazine.
а) Dissolve a few crystals of hydrazine sulfate [N2H6]SO4 in 2-3 ml of water in
a test-tube. Add a few drops of alkali to the solution and slowly add the iodine
water. What is observed? Write equation of the reaction.
b) To potassium permanganate solution acidified with sulfuric acid, add drop by
drop a solution of hydrazine sulfate. Explain the cause of discoloration of KMnO4
solution and write equation of the redox reaction, whereas hydrazine is oxidized to
nitrogen.
4.2.11. Oxidation-reduction properties of hydroxylamine.
To acidified solutions of potassium permanganate and iron (II) sulfate drop by
drop add NH2OH hydroxylamine solution. Explain the change in a color of the
solutions. Write equations of the reactions and make the conclusion about redox
properties of hydroxylamine.
5. Conclusions and Interpretations. Lesson Summary
Topic 9
р-Elements of VА group. Nitrogen and its compounds in
the positive oxidation states
1. Objiectives
Among inorganic compounds of nitrogen oxides, nitrous and nitric acids and
salts of these acids are important. Nitrogen (II) oxide is important for medicine as it
serves as bioregulator of blood pressure. Nitrogen (I) oxide is used for anesthesia.
The two most common compounds of nitrogen are potassium nitrate (KNO3)
and sodium nitrate (NaNO3). These two compounds are formed by decomposing
organic matter that has potassium, or sodium present. These compounds are often
found in fertilizers and biproducts of industrial waste. Most nitogen compunds have
a positive Gibbs free energy (reactions are not spontanous). As a result of the
decomposition of nitrogen-containing compounds under the action of bacteria
nitrogen is oxidised to nitrates, which are absorbed by plants. Nitrates are - source
of nitrogen for green plants and mushrooms. Ammonium and calcium nitrates are
used in agriculture as mineral fertilizer.
51
2. Learning Targets
To study the chemical properties of nitrogen and nitrogen compounds with
positive oxidation states.
To know biological role and use of the nitrogen compounds in the national
economy and medicine.
3. Self Study Section
3.1. Syllabus Content
Compounds of nitrogen with a positive oxidation state. Nitrogen oxides.
Methods of preparation. Acid-base and redox properties. Nitrous acid and nitrites.
Nitric acid and nitrates, acid-base and redox properties. Thermal decomposition.
"Royal water". Toxic action of nitrogen oxides and nitrates.
3.2. Theoretical Backgrounds
Nitrogen has 5 electrons in its valence shell. It has a valance III with respect to
hydrogen and a valance up to V with respect to oxygen. So, it can combine with
various elements to form many compounds.
There are five major oxides of nitrogen (N2O, NO, N2O3, NO2, N2O5), two
acids: HNO2 nitrous, and HNO3 nitric, and salts of these acids - nitrites and
nitrates.
Nitrous oxide, commonly known as laughing gas, is a chemical compound with
the formula N2O. N2O3 and N2O5 are anhydrides of the nitrous and nitric acids. All
nitrogen oxides except N2O are toxic. Nitrogen (II) oxide plays an important role in
the regulation of cardiovascular activity, because it supports tonus of blood vessels
walls.
Nitrous acid, HNO2 is considerably less stable than HNO3 and tends to
disproportionate into NO and NO2. Thermal decomposition of nitrous acid:
2HNO2 
→ NO2 + NO + H2O.
It is normally made by action of a strong acid, such as H2SO4, on a cold solution
of a nitrite salt, such as NaNO2. Nitrous acid is a weak acid (Ka= 4.0⋅10–4, рКа =
3.4).
Nitrogen shows both oxidizing and reducing properties in chemical reactions:
t
+2
+3
2H N O2 + H2S–2 = S0 + 2 N O + 2H2O;
oxidizing agent
+3
+7
+5
+2
5H N O2+2K Mn O4+3H2SO4 = 5H N O3+2 Mn SO4+K2SO4+3H2O;
reducing agent
Salts of nitrous acid are toxic compounds because of their effect on the Fe (II)
ions, which are part of hemoglobin.
Nitric acid belongs to the strong monobasic acids (рКа = 1.6). It is decomposed
by light:
hν
4HNO3 → 4NO2 + 2H2O + O2.
As a strong acid HNO3 reacts with metal oxides, alkalis, salts. Nitric acid is a
strong oxidizing agent. In its interaction with metals hydrogen never releases:
52
4Zn + 10HNO3(dil.) = 4Zn(NO3)2 + NH4NO3 + 3H2O;
8Al + 30HNO3(dil.) = 8Al(NO3)3 + 3N2O + 15H2O;
Cu + 4HNO3(conc.) = Cu(NO3)2 + 2NO2 + 2H2O.
Nitric acid reacts with nonmetals such as S, P, C, B. Corresponding acids are
formed:
0
+5
+5
+2
3 Р + 5H N O3 + 2H2O = 3H3 P O4 + 5 N O.
↓5е– ↑3е–
HNO3 reacts with some substances which show reducing properties (Bі2S3, ZnS,
As2S3 etc.):
ZnS + 8HNO3 = ZnSO4 + 8NO2 + 4H2O.
Salts of nitric acid also show strong oxidizing properties in the red-ox reactions,
for example in interactions with active metals or salts:
2KNO3 + 6FeSO4 + 4H2SO4 = 3Fe2(SO4)3 + 2NO↑ + K2SO4 + 5H2O;
NaNO3 + 4Zn + 7NaOH + 6H2O = 4Na2[Zn(OH)4] + NH3.
Salts of nitric acid melt and decompose at the heating with releasing of oxygen.
Other decomposition products depend on the activity of the metal cation, which is
part of salt :
2KNO3 
→ KNO2 + O2↑;
t
2Cu(NO3)2 
→ 2CuO + 4NO2↑ + O2↑;
t
AgNO3 
→ 2Ag + 2NO2↑ + O2↑.
t
The oxides play a large role in living organisms. They can be useful, yet
dangerous.
• Dinitrogen monoxide (N2O) is an anesthetic used at the dentist as a
laughing gas.
• Nitrogen dioxide (NO2) is harmful. It binds to hemoglobin molecules not
allowing the molecule to release oxygen throughout the body. It is
released from cars and is very harmful.
• Nitrate (NO3-) is a polyatomic ion.
• The more unstable nitogen oxides allow for space travel.
Nitrates and nitrites are known to cause several health effects. These are the
most common effects:
- Reactions with haemoglobin in blood, causing the oxygen carrying capacity of the
blood to decrease (nitrite);
- Decreased functioning of the thyroid gland (nitrate);
- Vitamin A shortages (nitrate);
- Fashioning of nitro amines, which are known as one of the most common causes
of cancer (nitrates and nitrites).
But from a metabolic point of view, nitric oxide (NO) is much more important
than nitrogen alone. This is a vital body messenger for relaxing muscles, and today
we know that it is involved in the cardiovascular system, the immune system, the
53
central nervous system and the peripheral nervous system. The enzyme that
produces nitric oxide, called nitric oxide synthesis, is abundant in the brain.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Describe the electronic structure of NO molecules by the method of valence
bonds. Characterize the bonds in the molecule.
2. Write molecular formulas of all nitrogen oxides. Which of them react with a
solution of Ca(OH)2? Write equations of the reactions.
3. Give examples of the reactions that characterize the acid-base and redox
properties of nitric acid.
4. Complete and balance the equations of the reactions:
t
a) Pb(NO3)2 
b) NO + KMnO4 + H2SO4 → MnSO4 + …;
→ …;
d) NO2 + Ca(OH)2 → … .
c) NO2 + NH3 → N2 + …;
5. Complete and balance the equations of the reactions:
a) KNO2 + K2Cr2O7 + Н2SO4 → …;
b) KNO2 + Br2 + KOH → …;
c) As2S5 + HNO3(dil.) → …;
d) I2 + HNO3 → HIO3 →… .
6. How does nitric acid react with metals and non-metals? Give some examples of
the reactions. What does occur at the reaction of HNO3 with different metals?
7. Write equations of the following reactions:
t
a) LiNO3 + Al2O3 
b) NaNO3 + Al + NaOH → NaAlO2 + …;
→ …;
c) KNO3 + FeCl2 + HCl → KNO2 + … .
8. What are the features of thermal decomposition of nitrates? Write equations of
thermal decomposition of: a) nitrous and nitric acids; b) ammonium carbonate
and ammonium dichromate; c) lithium nitrate and aurum nitrate.
9. Write formulas nitrogen compounds with positive values of oxidation state
which are used in medical practice.
b) Problems to Solve
1. Write equations of the transformation of the following cycle:
Сu(NO3)2 → NO2 → HNO3 → NO → KNO3 → KNO2.
54
2. As a result of unknown metal nitrate thermal decomposition, metal nitrite was
obtained. Mass of metal nitrate was 4.54 g and mass of metal nitrite is 3.82 g.
Specify the atomic mass of the metal.
Answer: 39.1.
3. To the mixture of copper and copper (II) oxide excess of concentrated nitric acid
was added. As a result of the reaction 26.88 L (at standart conditions) of brown
gas were obtained. Calculate the mass percentage of CuO in the initial mixture.
Answer: 48.8 %.
4. Calculate the mass of 60% nitric acid HNO3 solution which can be obtained
from the 1L of ammonia (at standart conditions), if its yield is 95%.
Answer: 4.5 kg.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental guidelines: «Chemical properties of nitrogen compounds».
4.2.1. Oxidation-reduction properties of the nitrous acid and its salts.
а) To a solution of KMnO4 (acidified with sulfuric acid) pour 2-3 ml of
potassium nitrite. What is observed? Write equation of the reaction.
b) Into a test-tube add 2-3 ml of potassium nitrite (acidified with sulfuric acid)
and solution of potassium dichromate. Observe the change of the color. Write
equation of the reaction.
c) To potassium iodide (acidified with sulfuric acid) add potassium nitrite
solution. What is the color of obtained solution? Write equation of the reaction.
4.2.2. Oxidizing properties of the nitric acid.
а) In two test tubes put a copper wire, into the first one pour diluted nitric acid
HNO3, into the second - concentrated HNO3. What is observed? Write equations of
the reactions.
b) To 0.5 g of zinc dust in a test-tube add 2–3 ml of diluted nitric acid HNO3.
What is observed? Write equation of the reaction. Drain the solution above the
precipitate into another test-tube. Add Nessler's reagent (a solution of potassium
tetraiodomercurate(II), K2[HgI4]) to obtained solution to determine the ammonium
cation. Write equations of the reactions.
c) To FeSO4 solution add concentrated nitric acid HNO3. Mix the prepared
solution and add solution of ammonium tiocyanate to determine Fe(III) ions. Write
equations of the reactions.
4.2.3. Thermal decomposition of nitrates.
Into a dry test-tube add a small amount of sodium nitrate and heat it to melt the
salt. What gas is released? How to confirm it?
55
Make the same experiment with lead and magnesium nitrate. Write equations of
the reactions and make the conclusions about the thermal decomposition of nitrates.
5. Conclusions and Interpretations. Lesson Summary
Topic 10
р-Elements of VА group. Chemical properties of Phosphorus
and its compounds
1. Objectives
In the natural world phosphorous is never encountered in its pure form, but only
as phosphates, which consist of a phosphorous atom bonded to four oxygen atoms.
This can exists as the negatively charged phosphate ion (PO43-), which is how it
occurs in minerals, or as organophosphates in which there are organic molecules
attached to one, two or three of the oxygen atoms.
The amount of phosphorous that is naturally present in food varies considerably
but can be as high as 370 mg/100 g in liver, or can be low, as in vegetable oils.
Foods rich in phosphorous include tuna, salmon, sardines, liver, turkey, chicken,
eggs and cheese (200 g/100 g).
There are many phosphate minerals, the most abundant being forms of apatite.
Fluoroapatite provides the most extensively mined deposits. The chief mining areas
are Russia, USA, Morocco, Tunisia, Togo and Nauru. World production is 153
million tones per year. In the oceans, the concentration of phosphates is very low,
particularly at the surface. The reason lies partly within the insolubility of
aluminum and calcium phosphates, but in any case in the oceans phosphate is
quickly used up and falls into the deep as organic debris. There can be more
phosphate in rivers and lakes, resulting in excessive algae growth.
2. Learning Targets
To study chemical properties of phosphorus and its compounds. To be able to
characterize them by using chemical reactions. To know biological value of
phosphorus and its compounds and their use in the national economy and medicine.
3. Self Study Section
3.1. Syllabus Content
Phosphorus. General characteristics. Allotropic modifications of phosphorus.
Chemical activity of phosphorous compounds. Phosphides and phosphine. The
comparison of the phosphides and phosphine with the corresponding compounds of
nitrogen.
Phosphorus compounds with positive oxidation states. Hydrolysis of the
halides. Oxides of phosphorous.
Ortophosphorous and hypophosphorous acids, structure of molecules, acid-base
and redox properties. Phosphoric acid and its ions. Dihydrogenphosphates,
56
hydrogenphosphates and phosphates. Pyrophosphoric acid. Metaphosphoric acid.
Reaction of phosphate ion identification. The biological role of phosphorus and
its compounds.
3.2. Theoretical Backgrounds
Phosphorus was first discovered in 1669 by German physician Hennig Brand.
Phosphorous is a multivalent nonmetal of the nitrogen group. It is found in
nature in several allotropic forms, and is an essential element for the life of
organisms. White phosphorus exhibits a and b modifications, with a transition
temperature between the two forms at -3.8°C. Ordinary phosphorus is a waxy white
solid. It is colorless and transparent in its pure form. Phosphorus is insoluble in
water, but soluble in carbon disulfide. Phosphorus burns spontaneously in air to its
pentoxide. It is highly poisonous, with a lethal dose of ~50 mg. White phosphorus
should be stored under water and handled with forceps. It causes severe burns when
in contact with skin. White phosphorus is converted to red phosphorus when
exposed to sunlight or heated in its own vapor to 250°C. Unlike white phosphorus,
red phosphorus does not phosphoresce in air, although it still requires careful
handling. Red phosphorous can vary in colour from orange to purple, due to slight
variations in its chemical structure. The third form, black phosphorous, is made
under high pressure, looks like graphite and has the ability to conduct electricity.
The electron configuration of the outer orbitals of the phosphorus atom is
3s23p3; oxidation states of +5, + 3, and –3 are most characteristic for phosphorus in
its compounds.
Oxidation state
Examples of
compounds
–3
–1
0
PH3
H3PO2
P4
+3
P2O3
H3PO3
+5
P2O5
H3PO4
Like nitrogen, phosphorus forms mainly covalent bonds in its compounds.
Compounds with ionic bonds, such as the phosphides Na3P and Ca3P2, are very
few. Unlike nitrogen, phosphorus has free 3d orbitals with relatively low energies,
which makes an increase in the coordination number possible and leads to the
formation of donor-acceptor bonds.
The burning of phosphorus in an excess of oxygen yields the pentoxide P4O10
(or P2O5); an insufficiency of oxygen results in the formation of mainly the trioxide
P4O6 (or P2O3). Phosphorus pentoxide is produced commercially by burning
elemental phosphorus in an excess of dry air. Subsequent hydration of P4O10 yields
orthophosphoric acid (H3PO4) and polyphosphoric acids (Hn+2PnO3n+1). In addition,
phosphorus forms phosphorous acid (H3PO3), hypophosphoric acid (H4P2O6), and
hypophosphorous acid (H3PO2), as well as two peracids, namely perphosphoric
acid (H4P2O8) and monoperphosphoric acid (H3PO5).
Phosphorus combines directly with all halogens, liberating a large amount of
heat and forming trihalides (PX3, X being a halogen), pentahalides with the general
formula PX5, and oxyhalides, for example, POX3. The fusion of phosphorus and
57
sulfur at temperatures below 100 °C yields solid solutions based on the two
elements; temperatures above 100 °C bring about the exothermic reaction for the
formation of the crystalline sulfides P4S3, P4S5, P4S7, and P4S10. Of these, only P4S5
decomposes into P4S3 and P4S7 when heated above 200 °C; the others melt without
decomposition. The known oxysulfides of phosphorus are P2O3S2, P2O2S3, P4O4S3,
P6O10S5, and P4O4S3. Compared with nitrogen, phosphorus is less capable of
forming compounds with hydrogen. Hydrogen phosphide, or phosphine (PH3), and
diphosphine (P2H4) can be obtained only by indirect means. The known compounds
with nitrogen include the nitrides PN, P2N3, and P3N5 - solid, chemically stable
substances obtained by passing phosphorus vapor and nitrogen through an electric
arc.
At temperatures above 2000 °C, phosphorus reacts with carbon to form the
carbide PC3, a substance that is not soluble in ordinary solvents and that reacts with
neither acids nor alkalies. When heated with metals, phosphorus forms phosphides.
White phosphorus is the form that occurs most commonly at room temperatures.
It is very reactive. It combines with oxygen so easily that it catches fire
spontaneously (automatically). As a safety precaution, white phosphorus is stored
under water in chemical laboratories.
Phosphorus combines easily with the halogens. The halogens are the elements
that make up Group 17 (VIIA) of the periodic table. They include fluorine,
chlorine, bromine, iodine, and astatine. Phosphorus also combines with metals to
form compounds known as phosphides.
Important compounds of phosphorus are phosphine PH3, oxides, acids and
corresponding salts of phosphorus. Phosphine, PH3, is a very toxic gas with garlic
like odor. Phosphine can be readily prepared by reacting calcium or aluminum
phosphide with dilute acid. Pure phosphine is not spontaneously flammable. It is
readily oxidized by air when ignited and explosive mixtures may be formed.
Phosphine is sparingly soluble in water, and an aqueous solution of PH3 is neither
acidic nor basic. It is soluble in very strong acids. Phosphine can react with some
acids to yield phosphonium (PH+4) salts. The best known is colorless phosphonium
chlorate, which is produced according to:
PH3 + HClO4 = [PH4]ClO4.
Phosphine shows reducing properties in redox reactions:
2PH3 + 2O2 = H3PO4;
3PH3 + 8HNO3 = 3H3PO4 + 8NO + 4H2O.
Phosphorus acids are oxoacids of phosphorus. There are a large number of
these and some cannot be isolated and are only known through their salts.
Ortophosphorous acid, Н3РО3 or Н2[РО3Н], is dibasic acid, that contains
phosphorous in oxidation state +3. It forms hydrogenphosphates when reacts with
alkalies:
H2[PO3H] + 2NaOH = Na2HPO3 + 2H2O.
Ortophosphorous acid acts both as a reducing and as an oxidizing agent in the
redox reactions:
58
H3PO3 + AgNO3 + H2O = H3PO4 + 2Ag + 2HNO3;
H3PO3 + 3Zn + 3H2SO4 = PH3 + 3ZnSO4 + 3H2O.
Hypophosphorous acid or phosphinic acid, H3PO2 or H2PO(OH), is monobasic
acid, that contains phosphorous in oxidation state +1.
Pyrophosphoric acid, H4P2O7 or (OH)2(O)P-O-P(O)(OH)2, is tetrabasic acid,
containing phosphorous in formal oxidation state +5.
Phosphoric acid, H3PO4 or PO(OH)3, is tribasic acid, containing phosphorous in
oxidation state +5. Phosphoric acid is one of the most widely known and used acid.
The salts of phosphoric acids - phosphates - have found wide application, whereas
phosphites and hypophosphites are less widely used:
+NaOH
→
NaH2PO4 + H2O
sodium dihydrogenphosphate
Н3РО4
2NaOH
 +
 →
Na2HPO4 + H2O
sodium hydrogenphosphate
3NaOH
+

→
Na3PO4 + H2O
sodium phosphate
Phosphates of ammonium and alkali metals are soluble in water, they hydrolyze
in aqueous solutions. Phosphates insoluble in water form acidic salts when react
with acid:
Ca3(PO4)2 + 4HNO3 → Ca(H2PO4)2 + 2Ca(NO3)2.
Concentrated phosphoric acids are used in fertilizers for agriculture and farm
production. Phosphates are used for special glasses, sodium lamps, in steel
production, in military applications (incendiary bombs, smoke screenings etc.), and
in other applications as: pyrotechnics, pesticides, toothpaste, detergents.
Phosphorus can be found in the environment most commonly as phosphates.
Phosphates are important substances in the human body, because they are a part of
DNA materials and they take part in energy distribution. Phosphates can also be
found commonly in plants. Phosphate is a dietary requirement, the recommended
intake is 800 mg/day, a normal diet provides between 1000 and 2000 mg/day,
depending on the extent to which phosphate rich foods are consumed.
Phosphorus is essential to the health of plants and animals. Many essential
chemicals in living cells contain phosphorus. One of the most important of these
chemicals is adenosine triphosphate (ATP). ATP provides the energy to cells they
need to stay alive and carry out all the tasks they have to perform. Phosphorus is
critical to the development of bones and teeth. Nucleic acids also contain
phosphorus. Nucleic acids are chemicals that perform many functions in living
organisms. For example, they carry the genetic information in a cell. They tell the
cell what chemicals it must make. It also acts as the "director" in the formation of
those chemicals.
The daily recommended amount of phosphorus for humans is one gram. It is
fairly easy to get that much phosphorus every day through meat, milk, beans, and
grains.
59
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Write the ground-state electron configuration and orbital notation for the atom
of phosphorous.
2. Write equation od the reaction of phosphorus preparation.
3. Give examples of reactions of oxidation, reduction, disproportionation and
polymorphic transformations of phosphorus.
4. What kinds of compounds does phosphorus form with hydrogen and active
metals? How are they received and how do they hydrolyze?
5. Compose molecular and graphic formulas of oxides, halides and oxo-halides of
phosphorus and corresponding hydrated compounds.
6. Write formulas and names of all acids that contain phosphorus. What is the
basicity of acids?
7. Write the molecular and ionic equations of the hydrolysis of the following salts:
sodium phosphate, sodium hydrogenphosphate and dihydrogenphosphate. What
is the pH of these salts solutions? Which phosphates have the better solubility in
water?
8. Complete and balance the equations of the following reactions:
a) P4 + Ba(OH)2 + H2O → Ba(H2PO2)2 + …;
b) P4 + СuSO4 + KOH → Cu + …;
c) P + HClO3 + Н2O → …;
d) PH3 + KMnO4 + H2SO4 = MnSO4 + … .
9. What are phosphates, and what is the difference between phosphorus,
phosphates and phosphoric acid.
10. Are phosphates essential to life?
11. Write the equations of the reactions to make the following transformation:
Ca3(PO4)2 → P → P4O10 → HPO3 → H3PO4 → Ca(H2PO4)2?
b) Problems to Solve
1. Density of phosphorus by air is 4.27 at 800, and it decreases in 2 times at
1500 0C. What is the nubmer of phosphorus atoms both in the first and in the
second case?
60
Answer: two atoms, four atoms.
2. 49 kg of anhydrous phosphoric acid reacts with 13.44 m3 of ammonia at standart
conditions. Determine the composition the obtained salts.
Answer: 46 kg of NH4H2PO4 and 13.2 kg of (NH4)2HPO4.
0
3. Calculate ∆ H 298
of calcium phosphate formation, proceeding from the heat
effect of the reaction:
3CaO(s) + P2O5(s) = Ca3(PO4)2; ∆ H 0 = –739 kJ
Answer: –4137.5 kJ/mol.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of Phosphorous and its
compounds».
4.2.1. Oxidation of phosphorus by oxygen.
Ignite red phosphorus in a porcelain cup and cover it with a glass funnel. Leave
the products for the next experiment. What is observed? Write equation of the
combustion of phosphorus.
4.2.2. Oxidation of phosphorus by nitric acid.
Into a test-tube pour 2 ml of 30% nitric acid HNO3 and add a small amount of
red phosphorus. Fix the tube in the tripod and gently boil until phosphorus
dissolution. Write equation of the reaction.
4.2.3. Phosphate anhydride hydration.
Pay attention to the product obtained in the experiment 4.2.1. Explain this
phenomenon. Pour the obtained product with distilled water into a clean test-tube
and divide it on two parts.
To the first test-tube add silver nitrate AgNO3 solution to determine
metaphosphate-ions. To the other part of the solution add few drops of nitric acid
HNO3 solution and boil it. Neutralize the solution with amonia water and determine
the presence of a phosphate-ion (formation of yellow precipitate Ag3PO4). Write
equations of the reactions and electrolitical dissociation of the obtained acids.
4.2.4. Hydrolysis of phosphoric acid salts.
Determine the pH of solutions of the following salts K3PO4, K2HPO4 and
KH2PO4 using the universal indicator paper. Write equations of the hydrolysis
reactions. Explain the different pH values of solutions.
4.2.5. Reaction of identification of PO42- ions.
а) To 1-2 drops of phosphate salt add 2 ml of mixture of (NH4)2MoO4 and
NH4NO3 in a solution of nitric acid. Heat the prepared solution. Observe formation
61
of the precipitate. Write equation of the reaction in ionic form; specify the color of
the precipitate.
HPO42– + 3NH4+ +12MoO42– + 23H+ = (NH4)3H4[P(Mo2O7)6]↓ + 10H2O.
b) To 3-4 drops of sodium phosphate add the same amount of silver nitrate
solution. Observe the formation of a yellow precipitate. Test its solubility in water,
in nitric acid solution and ammonium hydroxide solution. Write equations of the
corresponding reactions.
4.2.6. Hydrolysis of phosphorous halides.
Gently add 2-3 drops of phosphorous trichloride PCl3 into a test-tube with water
and a few crystals of phosphorous pentachloride PCl5 into a second test-tube. Add
to each test-tube 2-3 drops of methyl orange indicator and determine the pH of the
medium. Write equations of the hydrolysis.
5. Conclusions and Interpretations. Lesson Summary
Topic 11
р-Elements of VА group. Arsenic family elements (Arsenic,
Antimony, Bismuth)
1. Objectives
Arsenic, antimony and bismuth, three related elements of group 15, are all
found in trace quantities in nature and have interesting biological properties and
uses. While arsenic is most well known as a poison - and indeed the contamination
of groundwater by arsenic is becoming a major health problem in Asia - it also has
uses for the treatment of blood cancer and has long been used in traditional chinese
medicine. Antimony and bismuth compounds are used in the clinic for the treatment
of parasitic and bacterial infections.
The historical uses of arsenic were pharmaceutical and medicinal. Arsenic was
also commonly used in pigments, poisons and in the manufacturing of glass. A
major modern use for arsenic was as pesticides in agriculture.
Antimony is used in the metallurgy industry, especially in alloys. When it is
added to other metals such as lead, it hardens them. It is employed for the
manufacture of battery plates and in type metal as well as solders ammunition and
electric cable coverings.
The principle uses for bismuth are in low melting alloys in metallurgical
additives for aluminum, carbon steel and malleable iron in pearlescent cosmetic
pigments in medicine and in a variety of other smaller specialized applications. The
largest single use of bismuth continues to be in the pharmaceutical field.
2. Learning Targets
To study the chemical properties of Arsenic, Antimony, Bismuth and their
compounds. To be able to write reactions that characterize properties of these
elements.
62
To know the methods of Arsenic, Antimony and Bismuth compounds obtaining
and their use in medical practice.
3. Self Study Section
3.1. Syllabus Content
The elements of Arsenic subgroup. General characteristics. Compounds of
arsenic, antimony and bismuth with hydrogen in comparison with ammonia and
phosphine.
Detection of arsenic and antimony by the Marsh test.
Compounds with positive oxidation states. Oxides and hydroxides of elements
and their acid-base and redox properties. Arsenites and arsenates. Their acid-base
and redox properties. Salts of antimony and bismuth. Oxosalts formation.
Bismuthates and their stability.
Application in medicine and pharmacy of oxides and salts of arsenic, antimony
and bismuth and compounds p-elements of VA group.
3.2. Theoretical Backgrounds
Arsenic is a highely poisonous metalloid. Since it is a metalloid, arsenic has a
high density, moderate thermal conductivity, and limited ability to conduct
electricity. The oxidation states of arsenic are +5, +3, +2, +1 and -3. The electronic
configurations of atoms are 3d104s24p3, 4d105s25p3, 4f145d106s26p3. The three
allotropic forms of arsenic are yellow, black and gray, with gray being the most
common. The oxide of arsenic is amphoteric which means it can act as both an acid
and a base. Arsenic is mainly obtained by the heating of arsenic containing sulfides:
→ FeS + As↑
FeAsS 
The As(g) deposits as As(s) which can then further be used to make other
compounds. Arsenic can also be obtained by the reduction of arsenic(III) oxide:
As2О3 + 3С = 2As + 3CO↑.
Antimony is also a metalloid. The oxidation states of antimony are +3, -3, and
+5. Atimony exhibits allotropy with the most stable being the metallic form which
has the same properies as arsenic of high density, moderate thermal conductivity
and limited ability to conduct electricity. The oxide of antimony is antimony (III)
oxide which is amphoteric, meaning it can act as both an acid and base. Antimony
is obtained mainly from its sulfide ores. At low temperatures, antimony vaporizes.
Along with arsenic, antimony is commonly used in making alloys of other metals.
Bismuth is a metallic element. The oxidation states of bismuth are +3 and +5.
Bismuth is a poor metal that is similar to both arsenic and antimony. Bismuth is
commonly used in cosmetic products and medicine. Out of the group, bismuth has
the lowest electronegativity and ionization energy which means that it is more
likely to lose an electron than the rest of the Group 5 elements. This is why
Bismuth is the most metallic of Group 5. Bismuth is also a poor electrical
conductor. The oxide of bismuth is bismuth(III) oxide which acts as a base, an
expected property of metal oxides.
When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this
t0
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reaction have an odor resembling garlic. This odor can be detected on striking
arsenide minerals such as arsenopyrite with a hammer. Arsenic (and some arsenic
compounds) sublimes upon heating at atmospheric pressure, converting directly to
a gaseous form without an intervening liquid state at 887 K (614 °C). Arsenic
makes arsenic acid with concentrated nitric acid, arsenious acid with dilute nitric
acid, and arsenic trioxide with concentrated sulfuric acid.
Arsenic, antiomony and bismuth react with nitric acid:
3As + 5HNO3 + 2H2O = 3H3AsO4 + 5NO↑;
3Sb + 5 HNO3 + 8H2O = 3H[Sb(OH)6] + 5NO↑;
Bi + 4HNO3 = Bi(NO3)3 + NO↑ + 2H2O.
Bismuth dissolves in concentrated sulfuric acid to make bismuth(III) sulfate and
sulfur dioxide:
6H2SO4 + 2Bi → Bi2(SO4)3 + 3SO2 + 6H2O.
It reacts with nitric acid to make bismuth(III) nitrate:
Bi + 6HNO3(conc.) → Bi(NO3)3 + 3NO2 + 3H2O.
It also dissolves in hydrochloric acid, but only with oxygen present:
4Bi + 3O2 + 12HCl → 4BiCl3 + 6H2O.
Arsenic forms colorless, odorless, crystalline oxides As2O3 ("white arsenic")
and As2O5, which are hygroscopic and readily soluble in water to form acidic
solutions. Arsenic(V) acid is a weak acid with the formula H3AsO4. More
descriptively written as AsO(OH)3. Its salts are called arsenates.
Arsenic acid is prepared by treating arsenic trioxide with concentrated nitric
acid:
As2O3 + 2 HNO3 + 3 H2O → 2 H3AsO4+ N2O3.
The resulting solution is cooled to give colourless crystals of H3AsO4.
Arsenous acid, also known as arsenious acid, is the inorganic compound with
the formula H3AsO3 or As(OH)3. It is known to occur in aqueous solutions, but it
has not been isolated as a pure material.
The preparation of As(OH)3 involves a slow hydrolysis of arsenic trioxide in
water. Addition of base converts arsenous acid to the arsenite ions.
Arsenic-containing compounds are highly toxic and carcinogenic. The
anhydride form of arsenous acid, arsenic trioxide, is used as a herbicide, pesticide,
and rodenticide.
Arsenic, antiomony and bismuth when heated with active metals form arsenides
antimonides and bismuthides:
2E + 3Ca = Ca3E2.
They are easy to interact with halogens to form halides ЕHal3, ЕHal5. However,
not all EHal5 halides are prepared. Only AsF5, SbCl5, SbF5 and BiF5 are known.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
64
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Write the symbols and charges for atoms and ions of the following elements:
Arsenic, Antimony and Bismuth. How many electrons each of these atoms and
ions contain?
2. What is the product of Bismuth(III) nitrate hydrolysis reaction?
3. How will arsenic, antimony and bismuth react with chlorine and oxygen? Write
equations of the reactions.
4. Describe the relationship of arsenic, antimony and bismuth to hydrochloric,
sulfuric and nitric acid. Write equations of the reactions.
5. How and why do the acid-base properties of oxides (III) and hydroxides (III) of
the elements change from Arsenic to Bismuth? Write examples of the reactions.
6. Write equation of the following redox reactions:
а) Sb2S3 + HNO3 + H2O → H[Sb(OH)6] + … ;
b) NaBiO3 + Mn(NO3)2 + HNO3 → … ;
c) AsH3 + KMnO4 + H2SO4 → H3AsO4 + … ;
d) Sb + KClO3 + H2SO4 → Sb2(SO4)3 + … .
7. How are tioarsenites and tioarsenates obtained? What happens with the
acidification of solutions of these salts? Write the corresponding equations of
the reactions.
b) Problems to Solve
1. What mass of sodium bismuthate can be obtained at oxidation 4.66 g of bismuth
(III) oxide with chlorine in basic medium?
Answer: 5.6 g.
2. What volume of 0.5 M iodine solution should be taken for oxidation of 450 ml
of 0.1 N sodium arsenite solution?
Answer: 45 ml.
3. Calculate the enthalpy change in the reaction of Sb2S3 burning if the enthalpy of
formation of antimony (III) sulfide, antimony (III) oxide and sulfur dioxide are
–149.2 kJ/mol, –693 and –296,8 kJ/mol respectively.
Answer: 1434.2 kJ/mol.
65
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of Arsenic subgroup
elements and their compounds».
• All compounds of Arsenic and Antimony are poisonous, so experiments
should be carried out carefully.
• After laboratory work wash hands well!
4.2.1. Amphoteric properties of antimony (III) and Bismuth (III) compounds.
а) To 2.3 ml of antimony (III) chloride solution add dropwise a solution of
sodium hydroxide to precipitate formation. Divide the obtained precipitate into two
test-tubes equally. Into the first tube add the excess of an alkali, and to the second concentrated hydrochloric acid. Write equations of the corresponding reactions.
b) To 2.3 ml of bismuth (III) nitrate solution add sodium hydroxide solution to
precipitate formation. Divide the content of the tube into two test-tubes. Into the
first tube add the excess of an alkali, and to the second - concentrated hydrochloric
acid. Write equations of the corresponding reactions.
4.2.2. Reducing properties of Arsenic (III) compounds.
Into two test tubes pour 1 ml of sodium arsenite acidified with sulfuric acid.
Into the first tube add a few drops of iodine water, and to the second - potassium
permanganate. Observe discoloration of solutions in test tubes. Write equations of
the corresponding reactions.
4.2.3. Oxidizing properties of Arsenic and Antimonic acids salts.
а) To 1 ml of sodium arsenate add the same amount of hydrochloric acid and
0.5 ml of potassium iodide. Observe the coloring of a solution. Write equation of
the reaction.
b) To 1 ml of sodium antimonate add the same amount of hydrochloric acid and
0.5 ml of potassium iodide. Observe coloring of a solution. Write equation of the
reaction.
4.2.4. Oxidizing properties of Bismuth (III) salts.
To 1 ml of stannous (II) chloride solution add sodium hydroxide solution to
dissolving of precipitate. To the resulting solution add a few drops of bismuth (III)
nitrate. Observe the appearance of the precipitate and a quick change in its color as
a result of the formation of bismuth.Write equations of the reactions.
4.2.5. Preparation of Arsenic subgroups sulfides.
а) Through heated sodium arsenite solution (strongly acidified with the
hydrochloric acid) pass hydrogen sulfide gas. Observe the precipitate formation.
What is the color of precipitate? Write equation of the reaction.
b) Into one test-tube add 0.5 ml of antimony (III) chloride, and into the second 66
0.5 ml of antimony (V) chloride. To both tubes add 3-5 drops of hydrogen sulphide
water. What is the color of precipitates? Write equations of the reactions.
c) Carry out experiment similarly to 4.2.5.b and obtain precipitate of bismuth
(III) sulfide. What is the color of the precipitate? Write equation of the reaction.
Do not pour out precipitates from test-tubes, leave them for the next
experiment!
4.2.6. Obtaining of Arsenic and Antimony tiosalts.
а) To obtained in the previous experiment precipitates add 3-4 ml of ammonium
sulfide solution and mix well. Write equations of the reactions. Why bismuth
sulfide does not dissolve?
b) Into two test-tubes pour 1-2 ml of silver nitrate solution. Into the first testtube add arsenic acid salt solution, into the second - arsenous acid salt solution.
What is the color of precipitate? Write equations of the reactions.
5. Conclusions and Interpretations. Lesson Summary
Topic 12
р-Elements of VІА group. Oxygen, Sulfur, Selenium, Tellurium
1. Objectives
Oxygen is necessary for the survival of all animal and human life on Earth.
Animals and humans breathe in oxygen and breathe out carbon dioxide.
One important use of oxygen is in medicine. People who have trouble breathing
are given extra doses of oxygen. In many cases, this "extra oxygen" keeps people
alive after they would otherwise have died.
Oxygen and ozone are examples of allotropes (from the Greek meaning "in
another manner"). By definition, allotropes are different forms of an element.
Because they have different structures, allotropes have different chemical and
physical properties. Ozone has three atoms in each molecule. The chemical formula
is O3. Like nascent oxygen, ozone does not exist for very long under normal
conditions. It tends to break down and form dioxygen.
Ozone does occur in fairly large amounts under special conditions. For
example, there is an unusually large amount of ozone in the Earth’s upper
atmosphere. That ozone layer is important to life on Earth. It shields out harmful
radiation that comes from the Sun. Ozone is also sometimes found closer to the
Earth’s surface. It is produced when gasoline is burned in cars and trucks. It is part
of the condition known as air pollution. Ozone at ground level is not helpful to life,
and may cause health problems for plants, humans, and other animals.
Sulfur is an essential element for all life, and is widely used in biochemical
processes. In metabolic reactions, sulfur compounds serve as both fuels (electron
donors) and respiratory (oxygen-alternative) materials (electron acceptors). Sulfur
in organic form is present in the vitamins biotin and thiamine, the latter being
67
named for the Greek word for sulfur. Sulfur is an important part of many enzymes
and in antioxidant molecules like glutathione and thioredoxin. Organically bonded
sulfur is a component of all proteins, as the amino acids cysteine and methionine.
Disulfide bonds are largely responsible for the mechanical strength and insolubility
of the protein keratin, found in outer skin, hair, and feathers, and the element
contributes to their pungent odor when burned.
The biological role of selenium and tellurium still not been studied enough. It is
known that selenium affects the function of sexual glands and tellurium accelerates
the normal secretion of the liver and removes cholesterol from the body.
2. Learning Targets
To study the structural features of the p-elements and VIA group compounds.
To write equations of the chemical reactions that characterize the properties of
these elements and their compounds.
To study the biological role and use in medical practice compounds of VIA
group p-elements.
3. Self Study Section
3.1. Syllabus Content
General characteristics of the elements of VІА group. Oxygen. General
characteristics, occurence in nature. Features of the electronic structure of oxygen
molecules. Stereochemistry and nature of bonds in molecule of Ozone. Binary
compounds: oxides, peroxides, superoxides, ozonides. Compound of oxygen with
fluorine. The biological role of oxygen. Use of oxygen and ozone in medicine and
pharmacy.
General characteristics and biological role of Sulphur. Compounds of sulfur
with negative oxidation states. Acid-base and redox properties of hydrogen sulfide.
Metal and non-metal sulphides, their water solubility and hydrolysis. Identification
reaction of sulfide-ion.
Sulfur (IV) compounds - oxide, sulfurous acid, sulfites and hydrogensulfites,
their acid-base and redox properties. The interaction of sulfites with sulfur.
Identification reaction of sulfite-ion. Properties of thiosulfate: reactions with acids,
oxidizing agents (chlorine, iodine), metal cations, complexation reactions.
Identification reaction of thiosulfate-ion.
Sulfur (VI) compounds – oxide, hexafluoride, dioxochloride, sulfuric acid,
sulfates. Theis acid-base and redox properties. Oleum. Disulfuric acid.
Chlorosulfonic acid.
The use of sulfur compounds in medicine, pharmacy and pharmaceutical
analysis. Selenium and tellurium. General characteristics. Acid-base and redox
properties of the compounds. The biological role of selenium.
3.2. Theoretical Backgrounds
The Oxygen Family, also called the chalcogens, consists of the elements found
in Group 6 of the periodic table and is considered part of the Main Group elements.
It consists of the elements oxygen, sulfur, selenium, tellurium and polonium. These
68
can be found in nature in both free and combined states.
All elements of the oxygen family have 6 electrons in their outermost shell. The
electron configurations for each element are shown below:
Oxygen: 1s2 2s2 2p4
Sulfur: 1s2 2s2p6 3s2p4
Selenium: 1s2 2s2p6 3s2p6d10 4s2p4
Tellurium: 1s2 2s2p6 3s2p6d10 4s2p6d10 5s2p4
Polonium: 1s2 2s2p6 3s2p6d10 4s2p6d10f14 5s2p6d10 6s2p4
As one moves down the group, metallic character increases, with tellurium
being a metalloid and polonium a metal. Melting point, boiling point, density,
atomic radius, and ionic radius all increase going down the group. Ionization
energy decreases going down the group. The most common oxidation state is -2,
however sulfur can also exist at a +4 and +6 state and +2, +4, and +6 oxidation
states are possible for Se, Te, and Po.
Oxygen is the first element in Group VIA of the periodic table. Typically,
compounds that have oxygen in the oxidation state of two are referred to as oxides.
When oxygen reacts with metals, it forms oxides that are mostly ionic in nature.
These can dissolve in water and react to form hydroxides, which is why they can be
called basic anhydrides or basic oxides. Nonmetal oxides, which form covalent
bonds, are simple molecules with low melting and boiling points.
Compounds that contain oxygen with an oxidation state of -1 are referred to as
peroxides. Examples of this type of compound are Na2O2 and BaO2. Because
oxygen has an oxidation state of -1/2 in O2–, it is called a superoxide ion.
Oxygen is rarely featured as the central atom in a structure and can never have
more than 4 elements bonded to it due to its small size and its inability to create an
expanded valence shell. When it reacts with hydrogen, it forms water, which is
extensively hydrogen-bonded, has a large dipole moment and is considered an
universal solvent.
There are a wide variety of oxygen-containing compounds, both organic and
inorganic, including oxides, peroxides and superoxides, alcohols, phenols, ethers,
and carbonyl-containing compounds such as aldehydes, ketones, esters, amides,
carbonates, carbamates, carboxylic acids and anhydrides.
Sulfur is very unique in its ability to form a wide range of allotropes, more than
any other element in the periodic table. The most common state for sulfur to be in
is the solid S8 ring, as this is the most thermodynamically stable form at room
temperature. Sulfur exists in the gaseous form in five different forms (S, S2, S4, S6,
and S8). In order for sulfur to get to these states one must apply a sufficient amount
of heat.
Two very common oxides of sulfur are sulfur dioxide (SO2) and sulfur trioxide
(SO3). Sulfur dioxide is formed when sulfur is combusted in the air, which
produces a toxic gas that has a strong odor. These two compounds are used in the
production of sulfuric acid, which can be used in a variety of reactions. Sulfuric
69
acid is one of the top manufactured chemicals in the US, and is primarily used in
the manufacture of fertilizers.
Sulfur also exhibits a wide range of oxidation states, with values ranging from
-2 to +6. It is often the central ion in a compound and can easily hold up to 6 atoms
around itself. When in the presence of hydrogen it forms the compound hydrogen
sulfide, H2S, which is a poisonous gas, without hydrogen bonds and a very small
dipole moment. Hydrogen sulfide can easily be recognized by its strong odor that is
similar to that of rotten eggs, but this smell can be detected and low, nontoxic
concentrations. This reaction with hydrogen epitomizes how different oxygen and
sulfur act despite their common valence electron configuration and common
nonmetallic properties.
A very large variety of sulfur-containing compounds exist, many of them being
organic. The prefix thio- in from of the name of an oxygen-containing compound
means that the oxygen atom has been substituted with a sulfur atom. General
categories of sulfur-containing compounds include thiols (mercaptans),
thiophenols, organic sulfides (thioethers), disulfides, thiocarbonyls, thioesters,
sulfoxides, sulfonyls, sulfamides, sulfonic acids, sulfonates, and sulfates.
Selenium can be seen as a red or black amorphous, or a red or grey crystaline
structure, which is its most stable structure. Selenium has properties very similar to
those of sulfur; however, it is more metallic even though it is still classified as a
nonmetal. It acts as a semiconductor and therefore is often used in the manufacture
of rectifiers, which are devices that convert alternating currents to direct currents.
Selenium also has photoconductivity, which means that in the presence of light the
electrical conductivity of selenium increases. It is also used in the drums of laser
printers and copiers. In addition, it has found increased use now that lead has been
removed from plumbing brasses.
It is rare to find selenium in its elemental form in nature, and so typically it must
be removed through a refining process, usually involving copper. It can often be
found in soils and in plant tissues that have bioaccumulated the element. In large
doses, the element is toxic, however many animals require it as an essential
micronutrient. Selenium atoms are found in the enzyme glutathione peroxidase,
which destroys lipid-damaging peroxides. For humans, it is an essential cofactor in
maintaining the function of one’s thyroid gland. In addition, some research has
shown there to be a correlation between selenium-deficient soils and an increased
risk of contracting the HIV/AIDS virus.
Tellurium is the metalloid of the Oxygen family, with a silvery white color and
a metallic luster similar to tin at room temperature. Like selenium, it also displays
photoconductivity. It is an extremely rare element, and is most commonly found as
a telluride of gold. It is often used in metallurgy in combination with copper, lead,
and iron. In addition, it is used in solar panels and memory chips for computers. It
is not toxic or carcinogenic, however when humans get exposed to too much of it
they develop a garlic-like breath.
70
Polonium is a very rare, radioactive metal. There are 33 different isotopes of the
element and all of the isotopes are radioactive. It exists in a variety of states, and
has two metallic allotropes. It dissolves easily into dilute acids. It does not exist in
nature in compounds; however it can be manipulated to form ones synthetically in
the lab. It is used as an alloy with beryllium to act as a neutron source for nuclear
weapons. It is a highly toxic element. The radiation it emits makes it very
dangerous to handle. It can be immediately lethal when applied at the correct
dosage, or cause cancer if long-term exposure to the radiation occurs. Methods to
treat humans who have been contaminated with polonium are still being researched,
and it has been shown that chelation agents could possible help to decontaminate
humans.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. What properties increase going down the Oxygen family?
2. Write an electronic formula of oxygen, sulfur and selenium atoms and specify
their possible valences.
3. What compounds belong to polysulfides? How can they be obtained?
4. Describe the most important properties of hydrogen sulfide H2S:
а) reducing properties as an example of reactions with iodine, potassium
permanganate and oxygen in acidic medium;
b) acidic properties of hydrosulfuric acid; compare its strength with the strength
of carbonic acid;
c) how to detect hydrogen sulfide in the air?
5. What is the most common oxidation state for elements in the Oxygen Family?
6. How many elements in the Oxygen Family are metals, and which one(s)?
7. Why sulfur oxide (IV) and sulfurous acid are characterized by red-ox duality?
Show this property on the examples of reactions with hydrogen sulfide and
potassium permanganate.
8. Write equations of the reactions between sodium thiosulfate and chlorine,
iodine, sulfuric acid and silver chloride.
71
9. What is the most abundant element by mass in the Earth’s crust and in the
human body?
10. Write equations of the red-ox reactions:
a) РbS + О3 → ... ;
b) FeS2 + НNО3(conc.) → ... ;
c) Н2Sе + К2Сr2O7 + Н2SO4 →... ; d) Na2Te + NаОН + Сl2 → ... ;
e) SO2 + SеO2 + Н2O → ... .
b) Problems to Solve
1. What volume of hydrogen sulfide (at normal conditions or STP) must be passed
through 200 g of 16% lead acetate to complete sedimentation of Pb2+ ions in the
form of lead sulfide?
Answer: 2.2 L.
2. What volume of sulfur oxide (at normal condition or STP) must be dissolved in
1 L of water to obtain 5% H2SO3 solution?
Answer: 14.2 L.
3. Calculate the pH of 0.01 N solution of Nа2SО3.
Answer: 9.45.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of VIA group elements and
their compounds».
4.2.1. Properties of sulfur.
а) Reaction of sulfur with concentrated nitric acid.
Into a test-tube add a few crystals of sulfur and 4-5 drops of concentrated nitric
acid and boil the mixture carefully. What gas is released? Demonstrate the presence
of sulphate ions in a solution. Write equation of the reaction in molecular and ionic
form.
b) Reaction of sulfur with concentrated sodium hydroxide.
Into a test-tube add a few crystals of sulfur and 4-5 drops of concentrated
sodium hydroxide and boil the mixture carefully. Demonstrate the presence of
sulphide ions in a solution. Write equation of the reaction.
4.2.2. Sulfide-ion as a reducing agent.
a) To 1–2 ml of solution of bromine (Br2) in a test-tube add identical volume
of sodium sulfide (Na2S) solution. What is observed? Write equation of the
reaction.
72
b) To 1–2 ml of potassium permanganate (KMnO4) solution in a test-tube add
identical volume of sulfuric acid (H2SO4) solution and a solution of sodium sulfide
(Ns2S). What is observed? Write equation of the reaction.
c) To 1-2 ml of potassium dichromate (K2Cr2O7) solution in a test-tube add
identical volume of sulfuric acid (H2SO4) solution and a solution of sodium sulfide
(Na2S). What is observed? Write equation of the reaction.
d) To 1-2 ml of iron sulfate (FeSO4) solution in a test-tube add identical
volume of sulfuric acid (H2SO4) solution and a solution of sodium sulfide (Na2S).
What is observed? Write equation of the reaction.
4.2.3. Preparation of sulfides.
Into test-tubes with 1-2 ml of corresponding salts of potassium, barium, zinc,
manganese, copper, lead and iron (III) add 2-3 drops of sodium sulfide solution.
What is observed? Analyze their solubility in a hydrochloric acid solution. Write
equations of the reactions.
4.2.4. Sulfite-ion as a reducing and oxidizing agent.
а) To 1–2 ml of potassium permanganate (KMnO4) solution in a test-tube add
identical volume of sulfuric acid (H2SO4) solution and a solution of sodium sulfite
(Na2SO3). What is observed? Write equation of the reaction.
b) To 1–2 ml of sodium sulfide (Na2S) solution in a test-tube add sulfuric acid
(H2SO4) solution and a solution of sodium sulfite. What is observed? Write
equation of the reaction.
4.2.5. Oxidizing properties of sulfuric acid (experiment is carried out in a fume
cupboard!).
а) Reaction of sulfuric acid with metals.
Into three test-tubes pour 1-2 ml of sulfuric acid diluted. Add to each of the testtubes a small piece of metal - zinc, iron and copper (one metal to one test-tube).
Repeate the experiment with concentrated sulfuric acid. What is observed? Write
equations of the reactions.
b) Reaction of sulfuric acid on the paper.
Write something on a piece of filter paper with a glass stick wetted with diluted
sulfuric acid. Let it to dry. What is observed? Write equation of the reactions.
4.2.6. Properties of tiosulfates.
а) To 1–2 ml of sodium tiosulfate (Na2S2O3) solution in a test-tube add identical
volume of hydrochloric acid (HCl) solution. What is observed? Write equation of
the reaction.
b) To 1–2 ml of sodium tiosulfate (Na2S2O3) solution in a test-tube add a
solution of iodine (I2). What is observed? Write equation of the reaction.
c) To 1–2 ml of sodium tiosulfate (Na2S2O3) solution in a test-tube add a
solution of chlorine (Cl2). What is observed? Identify the sulphate ions in the
received solution. Write equation of the reaction.
5. Conclusions and Interpretations. Lesson Summary
73
Topic 13
р-Elements of VІІА group. Halogens
1. Objectives
Because of their great reactivity, the free halogen elements are not found in
nature. In combined form, fluorine is the most abundant of the halogens in Earth’s
crust. The percentages of the halogens in the igneous rocks of Earth’s crust are 0.06
fluorine, 0.031 chlorine, 0.00016 bromine, and 0.00003 iodine. Astatine and
element 117 do not occur in nature, because they consist of only short-lived
radioactive isotopes.
The halogen elements show great resemblances to one another in their general
chemical behaviour and in the properties of their compounds with other elements.
There is, however, a progressive change in properties from fluorine through
chlorine, bromine, and iodine to astatine — the difference between two successive
elements being most pronounced with fluorine and chlorine. Fluorine is the most
reactive of the halogens and, in fact, of all elements, and it has certain other
properties that set it apart from the other halogens.
Chlorine is the best known of the halogen elements. The free element is widely
used as a water-purification agent, and it is employed in a number of chemical
processes. Sodium chloride, of course, is one of the most familiar chemical
compounds. Fluorides are known chiefly for their addition to public water supplies
to prevent tooth decay, but organic fluorides are also used as refrigerants and
lubricants. Iodine has been used for many years as a disinfectant in "tincture of
iodine". Iodine compounds are used as catalysts, drugs, and dyes. Iodine is most
familiar as an antiseptic, and bromine is used chiefly to prepare bromine
compounds that are used in flame retardants and as general pesticides. In the past
ethylene dibromide was extensively used as an additive in leaded gasoline.
2. Learning Targets
To study the structural features of the p-elements of VIA group and their
compounds. To write equations of the chemical reactions that characterize the
properties of the VIA group elements and their compounds.
To study the biological role of halogens and their compounds. Use in medical
practice.
3. Self Study Section
3.1. Syllabus Content
General characteristics of the halogens. Properties of fluorine as the most
electronegative element. Simple substances, their chemical activity.
Compounds of halogens with hydrogen. Solubility in water. Acid-base and redox properties. Ionic and covalent halides. Halide ions as ligands in complex
compounds. Reactions of identification of halide ions.
Halogens with positive oxidation states. Compounds with oxygen. Reactions of
halogens with water and aqueous solutions of alkalis. Oxoacids of halogen and their
74
salts. Chlorate, bromates and iodates. The biological role of chlorine, fluorine,
bromine and iodine.
The bactericidal action of chlorine and iodine. The use of bleach, iodine and
fluoride, chloride, bromide, iodide for disinfecting and sterilizing.
3.2. Theoretical Backgrounds
All halogens possess the oxidation state 0 in their diatomic forms. Fluorine
exhibits the oxidation states of −1 (F− ion) and +1 (hypofluorous acid). The
principal oxidation states of chlorine, bromine, and iodine are −1, +1, +3, +5, and
+7.
Halogens display physical and chemical properties typical for nonmetals. They
have relatively low melting and boiling points that increase steadily down the
group. Near room temperature, the halogens span all of the physical states: fluorine
and chlorine are gases, bromine is a liquid, and iodine is a solid. All of the elements
are colored, with the color becoming more intense moving down the group.
Fluorine gas is pale yellow, and chlorine gas is a yellowish green. Liquid bromine
and its vapors are brownish red. Solid iodine appears as shiny, dark gray crystals,
and the vapors are a deep purple. The halogens are poor thermal and electrical
conductors in all phases, and as solids they are brittle and crumbly. The halogens
have distinctive, unpleasant odors, will burn exposed flesh, and are toxic.
Halogens readily combine with hydrogen to form the hydrogen halides which
are colourless gaseous covalent molecules:
H2(gas) + Cl2(gas) → 2HCl(gas).
The hydrogen halides dissolve in water to form very strong acids with solutions
of pH 1 e.g. hydrogen chloride forms hydrochloric acid in water HCl(aq) or H+Cl–
(aq) because they are fully ionised in aqueous solution. An acid is a substance that
forms H+ ions in water.
Aqueous solutions of the hydrogen halides are often called mineral acids
because they are literally acids prepared from minerals. Hydrochloric acid is
prepared by reacting table salt with sulfuric acid, for example, and hydrofluoric
acid is prepared from fluorite and sulfuric acid:
2NaCl(s) + H2SO4(aq) → 2HCl(aq) + Na2SO4(aq);
2CaF2(s) + H2SO4(aq) → 2HF(aq) + CaSO4(aq).
Under certain conditions, it is possible to isolate neutral oxides of the halogens,
such as Cl2O, Cl2O3, ClO2, Cl2O4, Cl2O6, and Cl2O7. Cl2O7, for example, can be
obtained by dehydrating perchloric acid, HClO4. These oxides are notoriously
unstable compounds that explode when subjected to either thermal or physical
shock. Some are so unstable they detonate when warmed to temperatures above
-40oC.
Chlorine reacts with the OH- ion to form chloride ions and hypochlorite (OCl-)
ions:
Cl2(aq) + 2 OH-(aq) → Cl-(aq) + OCl-(aq) + H2O(l).
75
This is a disproportionation reaction in which one-half of the chlorine atoms are
oxidized to hypochlorite ions and the other half are reduced to chloride ions.
When the solution is hot, this reaction gives a mixture of the chloride and chlorate
(ClO3-) ions:
3Cl2(aq) + 6 OH-(aq) → 5Cl-(aq) + ClO3-(aq) + 3H2O(l).
Under carefully controlled conditions, it is possible to convert a mixture of the
chlorate and hypochlorite ions into a solution that contains the chlorite (ClO2-) ion:
ClO3-(aq) + ClO-(aq) → 2ClO2-(aq).
The last member of this class of compounds, the perchlorate ion (ClO4-), is
made by electrolyzing solutions of the chlorate ion.
Oxyanions and Oxyacids of Chlorine
Oxyanions
Oxidation
State
of the Chlorine
Compound
+1
+3
+5
+7
ClOClO2ClO3ClO4-
Oxyacids
Name
hypochlorite
chlorite
chlorate
perchlorate
Compound
HClO
HOClO
HOClO2
HOClO3
Name
hypochlorous acid
chlorous acid
chloric acid
perchloric acid
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
76
3.4. Test Yourself
а) Review Questions.
1. Write the electronic configurations of halogen atoms and specify their valency
and oxidation state.
2. Give examples of the reactions of halogens preparation in the industry and in
the laboratory.
3. Write equations of the reactions of halogens with metals (Fe, Au, Cu) and
nonmetals (S, F, C).
4. How does the strength of acids change in a row HF-HI?
5. Write and give names of oxoacids of chlorine. How does the strength and redox properties of acids change in a row НС1О–НС1О4?
6. Write equations of the following transformations:
а) КС1 → Сl2 → КС1O3 → КС1O4 → С12O7 → C1O2;
b) І2 → КІ → І2 → НIO3 → І2О5 → І2 → НIO3 → НІ.
7. What compounds of halogens are used in medical practice for disinfecting and
sterilizing?
b) Problems to Solve
1. What mass of iodine will release at the reaction of the excess of potassium
iodide with 300 ml of 6% KMnO4 solution (ρ =1.04 g/ml) in an acidic medium?
Answer: 15.05 g.
2. Calculate рН of 0.01 М sodium hypochlorite.
Answer: 9.5.
3. To a solution containing C1-, Br-, I- ions a solution of silver nitrate was added.
Determine in what sequence will silver halides precipitate. Calculate the mass of
AgCl in 300 ml of saturated solution.
Answer: 0.06 g.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of VIIА group elements and
their compounds».
4.2.1. Halides preparation.
Fill three dry test-tubes: the first – with chloride, the second – with bromide and
the third – with potassium iodide. Into each test-tube add the same amount of dry
potassium permanganate. Wet the contents of the tubes with a solution of sulfuric
acid and heat. Observe the formation of halogens on the background of white
paper. Write equations of the reactions.
4.2.2. Properties of halogens.
а) Solubility of bromine and iodine in organic solvents. Into one test-tube pour
77
1-2 ml of bromine water, into second – iodine water. Pour the same amount of CCl4
into each test-tube. Mix the obtained solutions, check their color.
b) Oxidizing properties of chlorine, bromine and iodine water.
ba) Pour 2-3 ml of hydrogen sulphide water into three test-tubes and add: into
the first - chlorine, into the second – bromine and into the third - iodine water.
Observe the changes and write equations of the reactions.
bb) To solution of potassium bromide and potassium iodide add chlorine water.
Observe the releasing of bromine and iodine. Write equations of the reactions.
bc) To solution of potassium iodide add chlorine water. Write equation of the
reaction.
4.2.3.Preparation of hydrogen chloride and research of its properties.
Into a test-tube with a gas tube place 0.5-1 g of sodium chloride, pour 3-4 ml of
concentrated sulfuric acid and heat it gently. Check formation of a gas with the blue
litmus paper, and on a piece of paper moistened with a solution of potassium
permanganate. Write equation of the reaction.
4.2.4. Reducing properties of halide- ions.
Pour into three separate test-tubes 1-2 ml of potassium chloride, potassium
bromide and potassium iodide. Add the same amount of potassium dichromate,
acidified with sulfuric acid. Write equations of the reactions.
4.2.5. Preparation of silver halides.
Obtain insoluble and slightly soluble salts of СаF2, АgСl, АgВr, АgI, Cu2I2,
РbI2, РbСl2 using solutions of salts of calcium, silver, copper and lead.
4.2.6. Preparation and properties of chlorine compounds with oxygen.
а) Oxidizing properties of the hypochlorites.
Obtain solution of sodium chromite: into a test-tube pour 1-2 ml of
chromium(III) chloride and excess of sodium hydroxide. Then add NaClO solution
to change the color of the solution. Write equations of the transformation reactions
of sodium chromite to sodium chromate.
b) Preparation of potassium iodate.
Add a few crystals of iodine to the heated solution of potassium hydroxide (heat
gently!) with mass percentage of KOH 50%. A white precipitate of potassium
iodate will form. Write equation of the reaction.
5. Conclusions and Interpretations. Lesson Summary
Topic 14
General characteristics of d-elements. d-Elements of IB group.
Copper, Silver, Gold
1. Objectives
Gold, silver and copper were the first metals that early humans came across.
78
They would have been found in streambeds and rivers, washed out of the rocks.
They are called “native” metals, because they are sufficiently unreactive to be
found in the ground in their elemental forms.
Copper is essential in all higher plants and animals. Copper is found in a variety
of enzymes, including the copper centers of cytochrome c oxidase, the Cu-Zn
containing enzyme superoxide dismutase, and is the central metal in the oxygen
carrying pigment hemocyanin. The RDA for copper in normal healthy adults is 0.9
mg/day.
Copper is carried mostly in the bloodstream on a plasma protein called
ceruloplasmin. Though when copper is first absorbed in the gut it is transported to
the liver bound to albumin. An inherited condition called Wilson’s disease causes
the body to retain copper, as it is not excreted by the liver into the bile. This
disease, if untreated, can lead to brain and liver damage.
Elemental gold and silver have no known toxic effects or biological use,
although gold salts can be toxic to liver and kidney tissue. Like copper, silver also
has antimicrobial properties. The prolonged use of preparations containing gold or
silver can also lead to the accumulation of these metals in body tissue; the results
are the irreversible but apparently harmless pigmentation conditions known as
chrysiasis and argyria respectively.
These metals are much less active than members of their respective periods that
lie farther to the left in the periodic table. They do not displace H2 from acidic
solutions. Their uses are based mostly on their resistance to corrosion and their
exceptional abilities to conduct heat and electricity.
2. Learning Targets
To know equations of the chemical reactions that characterize properties of the
d-elements of IB group and their compounds.
To study the biological role of Copper, Silver, Gold and their compounds.
Antimicrobial properties of copper (II) and silver (I) ions.
3. Self Study Section
3.1. Syllabus Content
General characteristics of the group IB elements. Physical and chemical
properties of simple substances. Reactions with acids, oxygen, halogens.
Compounds of copper (I) and copper (II), their acid-base and red-ox properties,
ability to form complexes. Complex compounds of copper (II) with ammonia,
amino acids.
Oxide and halides of copper (I). Complex compounds of copper (I) with
chlorides and ammonia. The use of copper compounds in medicine and pharmacy.
Silver compounds, their acid-base and red-ox properties. The ability to form
complexes with halide-ions, ammonia, thiosulfate ions. The antimicrobial
properties of Ag+ ions. The use of silver compounds in medicine and
pharmaceutical analysis.
Gold. Oxidation of gold by oxygen in the presence of cyanide ions. Attitude of
79
gold to "aqua regia". Compounds of gold (I), gold (III) and their acid-base and redox properties, ability to form complexes. The use of gold and its compounds in
medicine and pharmacy.
3.2. Theoretical Backgrounds
In the Periodic Table of Elements, the chemical elements are organised
according to their chemical and physical properties. In the case of copper, silver
and gold, these form a vertical triad, called Group IB.
Like other groups, the members of this family show patterns in electron
configuration, especially in the outermost shells, resulting in trends in chemical
behavior:
Cu 3d104s1,
Ag 4d105s1,
Au 5d106s1.
Oxidation states of Cu, Ag, Au: +1, +2, +3, +5.
Copper is the most reactive of the three metals, and forms two series of
compounds, those with an oxidation number of one and those with an oxidation
number of two.
Silver is a very inactive metal. It does not react with oxygen in the air under
normal circumstances. It does react slowly with sulfur compounds in the air,
however. The product of this reaction is silver sulfide (Ag2S), a black compound.
Silver does not react readily with water, acids, or many other compounds. It does
not burn except as silver powder.
Cu and Ag are able to react with H2SO4(aq) or HNO3(aq) Thus shifting the
oxidation numbers of Cu and Ag to Cu2+ and Ag+:
Cu + 2H2SO4(conc.) = CuSO4 + SO2↑ + 2H2O;
Ag + 2HNO3 = AgNO3 + NO2↑ + H2O.
However, Au will not react with either H2SO4(dil.) or HNO3(dil.), rather it will
react with what is called “aqua regia”, which is one part HNO3 and three parts HCl:
Au + HNO3 + 4HCl = H[AuCl4] + NO↑ + 2H2O;
2Au + 6H2SeO4 = Au2(SeO4)3 + 2SeO2 + 6H2O.
The group 1 metals do not react with diluted hydrochloric and sulfuric acid. All
metals react with cyanide solutions, for example as gold:
Au + O2 + 8KСN + 2H2O = 4K[Au(CN)2] + 4KOH.
Oxides. CuO can be obtained at the termal decomposition of Cu(NO)2 and
Cu(OH)2:
t
2Cu(NO3)2 
→
2CuO + 4NO2 + O2;
t
Cu(OH)2 
→
CuO + H2O.
Silver oxide Ag2O can be obtained as follows:
t
2AgNO3 + 2KOH 
→
2KNO3 + Ag2O + H2O.
Decomposition of Au2O3 at heating:
t
2Au2O3 
→
4Au + 3O2.
80
Hydroxides. Preparation of hydroxides:
CuCl2 + 2NaOH = Cu(OH)2 + 2NaCl.
AgOH does not form:
2AgNO3 + 2NaOH = Ag2O + 2NaNO3 + H2O.
Au(OH)3 exists, it is called "acid gold".
CuO and Cu(OH)2 have amphoteric properties, in which basic properties
dominate. Au2O3 and Au(OH)3 are amphoteric compounds, in which acidic
properties dominate. With concentrated alkalis they form hydroxocomplexes:
Cu(OH)2 + 2KOH → K2[Cu(OH)4]
potassium tetrahydroxocuprate(II);
Au(OH)3 + KOH → K[Au(OH)4]
potassium tetrahydroxoaurate (II).
Salts are soluble in water: CuCl2⋅2H2O, Cu(NO3)2⋅6H2O and CuSO4⋅5H2O;
AgNO3; AgClO3; AgClO4; AuCl3; AuBr3. They can form complex compounds:
2CuSO4 + 2NH4OH = (CuOH)2SO4 + (NH4)2SO4;
(CuOH)2SO4 + NH4SO4 + 6NH4OH = 2[Cu(NH3)4]SO4 + 8H2O;
tetraamminecuprum (II) sulfate
AgCl + 2NH3 = [Ag(NH3)2]Cl
diammineargentum(I)chloride;
AgBr + 2Na2S2O3 = Na3[Ag(S2O3)2] + NaBr.
sodium ditiosulfatoargentate(I)
Ag+ ion shows red-ox properties (silver mirror reaction):
HCOH + 2[Ag(NH3)2]OH → 2Ag↓ + 4NH3↑ + H2O + HCOOH.
Copper occurs in its native form in Chile, China, Mexico, Russia and the USA.
Various natural ores of copper are: copper pyrites (CuFeS2), cuprite or ruby copper
(Cu2O), copper glance (Cu2S), malachite, (Cu(OH)2CuCO3), and azurite
(Cu(OH)22CuCO3).
Copper, although potentially toxic in excessive amounts, is essential for life.
Copper is shown to have antimicrobial properties which make it useful for hospital
doorknobs to keep diseases from being spread. Eating food in copper containers is
known to increase the risk of copper toxicity.
Elemental gold and silver have no known toxic effects or biological use,
although gold salts can be toxic to liver and kidney tissue. Like copper, silver also
has antimicrobial properties. The prolonged use of preparations containing gold or
silver can also lead to the accumulation of these metals in body tissue; the results
are the irreversible but apparently harmless pigmentation conditions known as
chrysiasis and argyria respectively.
The medical uses of silver include its incorporation into wound dressings, and
its use as an antibiotic coating in medical devices. Wound dressings containing
silver sulfadiazine or silver nanomaterials may be used to treat external infections.
Silver is also used in some medical applications, such as urinary catheters and
endotracheal breathing tubes, where there is tentative evidence that it is effective in
reducing catheter-related urinary tract infections and ventilator-associated
81
pneumonia respectively. The silver ion (Ag+) is bioactive and in sufficient
concentration readily kills bacteria in vitro. Silver and silver nanoparticles are used
as an antimicrobial in a variety of industrial, healthcare and domestic applications.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Write an electron configuration of Cu, Ag and Au elements. Indicate their
possible valences.
2. Write equations of the reactions: а) copper and silver with air (greening and
blackening of copper and silver products); b) copper and silver with diluted and
concentrated sulfuric and nitric acids solutions; c) dissolution of gold in the
"aqua regia"; d) dissolution of silver chloride in ammonium hydroxide and
sodium thiosulfate solutions; e) copper (II) hydroxide with an excess of
concentrated sodium and ammonia hydroxide.
3. The extraction of copper and silver from copper and silver ores. Write
equations of the corresponding reactions.
b) Problems to Solve
1. What mass of silver nitrate is needed to prepare 10 g of 2% eye drops solution?
Answer: 0.2 g
2. Calculate the solubility of AgI in water (in g/l and mol/l) if Ksp(AgI) = 8.3.10–17.
Answer: 2.88⋅10–9 mol/l; 6.76⋅10–7 g/l.
3. Determine the mass of the precipitate that is formed by mixing 26.8 g of silver
nitrate solution and 9.3 g of aluminum chloride solution.
Answer: 22.62 g.
4. Write the formula of complex ions Cu2+, Au3+ with coordination number 4 and
ligands: NH3, S2O32–, H2O, C1–, CN–. Add the external coordination sphere and
give the name of complexes.
5. Write molecular and ionic equation of copper sulfate and aurum (III) chloride
hydrolysis.
82
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: “General characteristics of d-elements of the IB
group. Cu, Ag, Au”.
4.2.1. Reducing properties of copper, silver and their compounds.
а) Dip the surface of metallic zinc and iron into a solution of copper salt and
leave for 5 minutes. What is observed? Write equation of the reaction.
b) Pour 2 ml of silver nitrate solution into a clean dry test-tube and add
ammonia solution to dissolve the formed precipitate. Then add 1-2 ml of 10%
glucose solution. Immerse prepared solution into a glass with hot water. What is
observed? Write equations of the reactions.
4.2.2. Preparation and properties of silver and copper hydroxide.
а) To 2-3 ml of copper sulfate solution add solution of sodium hydroxide.
Divide prepared precipitate into three test-tubes. To the first one add 2-3 ml of
hydrochloric acid, to the second - the same amount of sodium hydroxide, and heat
the third test-tube till boiling. Compare the color of solutions and precipitates and
write equations of the reactions.
b) To a solution of sodium hydroxide add 1-2 drops of silver nitrate. Observe
the formation of brown precipitate. Write equation of the reaction. Divide the
obtained precipitate into two test-tubes and add to the first test-tube ammonia
solution and to the second hydrochloric acid solution. Write equations of the
reactions.
4.2.3. Preparation of copper and silver complex compounds.
а) To obtained copper(II) hydroxide solution add an excess of ammonium
hydroxide. Observe the change of color of the solution. Write equation of the
reaction.
b) To obtained precipitate of silver bromide add a solution of sodium
thiosulfate. What is observed? Write equation of the reaction.
4.2.4. Preparation of copper and silver sulfides.
To solutions of copper (II) nitrate and silver nitrate add hydrogen sulfide
solution. What is observed? Write equations of the reactions. The work is carried
out in the fume hood.
4.2.5. Hydrolysis of copper (II) salts.
Into two test-tubes add 2-3 drops of copper (II) sulfate (chloride) and use litmus
paper to determine the medium of the reaction. Into one of the test-tubes add 2-3
drops of hot sodium carbonate solution. Observe the formation of a green
precipitate of copper (II) hydroxocarbonate and releasing of a gas. Write equation
of the hydrolysis reaction of copper (II) salt.
83
4.2.6. Oxidizing properties of silver compounds.
Into test-tube add 2-3 drops of 0.5 M stannous (II) chloride solution and add
drop by drop a solution of potassium (sodium) hydroxide to dissolve the precipitate
of stannous hydroxide, which was formed. To prepared solution add 1-2 drops of
silver nitrate. Observe the formation of a black precipitate of metallic silver. Write
equations of the reactions.
5. Conclusions and Interpretations. Lesson Summary
Topic 15
d-Elements of IIB group. Chemical properties of Zinc,
Cadmium, Mercury
1. Objectives
The group IIB elements have multiple effects on biological organisms as
cadmium and mercury are toxic while zinc is required by most plants and animals
in trace amounts.
Zinc is an essential trace element, necessary for plants, animals, and
microorganisms. It is typically the second most abundant transition metal in
organisms after iron and it is the only metal which appears in all enzyme classes.
There are 2-4 grams of zinc distributed throughout the human body, and it plays
"ubiquitous biological roles". About 10% of human proteins potentially bind zinc,
in addition to hundreds which transport and traffic zinc.
Mercury and cadmium are toxic and may cause environmental damage if they
enter rivers or rain water. This may result in contaminated crops as well as the
bioaccumulation of mercury in a food chain leading to an increase in illnesses
caused by mercury and cadmium poisoning.
2. Learning Targets
To study the structural features of the IIB group elements. To know equations
of the chemical reactions that characterize the properties of the IIB group elements
and their compounds.
To study the biological role of elements and toxic effects of their compounds;
their use in medical practice.
3. Self Study Section
3.1. Syllabus Content
General characteristics of the elements of group IIB. Physical and chemical
properties of simple substances.
Zinc. General characteristics. Chemical activity of simple substance. Acid-base
and redox characteristics of zinc compounds. Zinc salts, their solubility and
hydrolysis. Complex compounds of zinc with ammonia, water and hydroxide ions.
Zinc-containing enzymes. Use of zinc compounds in medicine and pharmacy.
84
Cadmium and its compounds compared to similar compounds of zinc.
Mercury. General characteristics, properties that differ from zinc and cadmium.
Reaction of mercury with sulfur, nitric acid and iron (III) chloride. Mercury
nitrates. Hydrolysis. Basic salts. Mercury(I) and mercury(II) compounds. Acid-base
and redox characteristics, the ability to form complexes. Calomel and mercury
chloride, their reaction with ammonia. The toxic effects of cadmium and mercury
compounds. Use of mercury in medicine and pharmacy.
3.2. Theoretical Backgrounds
Group IIB of the Periodic Table includes Zinc, Cadmium and Mercury. The
group IIB elements Zn, Cd, Hg are also called the transition elements. These
elements are found in different proportions in the Earth’s crust: it has been
estimated that zinc is present to the extent of 80 parts per million (compared with
70 for copper and 16 for lead). The estimate for cadmium is only 0.15;
commercially, it is always found associated with zinc or zinc–lead ores and is
produced only as a by-product of zinc and lead smelting. The proportion of
mercury in the Earth’s crust is estimated at 0.08 parts per million. All important
mercury deposits consist of mercuric sulfide, known as the mineral cinnabar. These
metals tend to have properties characteristic of elements with full subshells, which
includes having low melting and boiling points (due to the weak metallic bonding
of the ns2 electrons).
Zinc, cadmium and mercury are in the same group in the periodic table because
they all have similar arrangements of electrons in the outermost shells. However,
the inner electron structure of mercury differs from that of zinc and cadmium, and
therefore its chemical properties differ also; the properties of zinc and cadmium,
though, are very similar. Zinc, cadmium, and mercury can lose the two electrons in
the outermost shell to form dipositive ions, Me2+ (in which Me represents a
generalized metal element), thereby exposing the next innermost shell with a stable
configuration in each case of 18 electrons.
Zinc exhibits only the +2 oxidation state. It can give up two electrons to form an
electrovalent compound; e.g., zinc carbonate ZnCO3. It may also share those
electrons, as in zinc chloride, ZnCl2, a compound in which the bonds are partly
ionic and partly covalent.
Cadmium compounds are mainly ionic, but cadmium also forms complex ions
with ligands (atoms, ions, or molecules that donate electrons to a central metal ion);
e.g., the complex ion with ammonia NH3, having the formula [Cd(NH3)4]2+, or with
the cyanide ion, the formula [Cd(CN)4]2−. Differing from zinc and mercury,
cadmium can form the complex ions represented by the formulas [CdCl3]− and
[CdCl4]2− in solution.
Mercury in its +2 and +1 oxidation states forms the ions Hg2+ and [Hg2]2+,
respectively. In the latter, two electrons are shared in a covalent bond between the
two metal atoms. The [Hg2]2+ ion shows little tendency to form complexes, whereas
the Hg2+ ion does form them. In contrast to compounds of mercury in the +2 state,
which are usually covalent, all the common salts of mercury in the +1 state are
85
ionic, and the soluble compounds - e.g., mercurous nitrate, Hg2(NO3)2 - show
normal properties of ionic compounds.
Electron configurations of elements:
Zn 4s23d10
Cd 5s24d10
Hg 6s24 f145d10
Mercury is the metal which under normal conditions is in the liquid state
(melting temperature of -39 0C).
Zinc and cadmium both react readily with oxygen on heating to form oxides
ZnO and CdO. The corresponding reaction of mercury, though thermodynamically
stable, is very slow. All three element of group IIB react with the halogens, and
elements such as sulphur, chlorine and phosphorus.
Zn and Cd with nonmetals form: ZnS, ZnCl2, Zn3P2, Zn3N2, ZnC2, HgCl2, HgS.
Cadmium sulphide (CdS) exists naturally as the mineral Greenockite.
Zn2+ and Cd2+ both form water soluble oxo-salts with nitrate, sulphate and
perchlorate ions. The metal ions are similar to Mg2+ and so these salts have similar
properties to the analogous magnesium salts. When dissolved in water, aqua ions
are formed, which are acidic, liberating H+ with solutions containing M(OH)+ ions.
Reaction with sodium hydroxide solution:
Zn + 2H2O + 2NaOH = H2 + Na2[Zn(OH)4] – amphotheric properties of Zn
Cd is practically insoluble in alkalis.
Reactions with acids:
Zn + 2HCl = ZnCl2 + H2↑; Cd + 2HCl = CdCl2 + H2↑;
Zn + 2H2SO4(conc.) = ZnSO4 + SO2↑ + 2H2O;
4Zn + 5H2SO4(conc.) = 4ZnSO4 + H2S + 4H2O;
4Zn + 10HNO3(dil.) = 4Zn(NO3)2 + NH4NO3 + 3H2O;
4Cd + 9HNO3 = NH3 + 4Cd(NO3)2 + 3H2O;
3Cd + 8HNO3(dil.) = 3Cd(NO3)2 + 2NO↑ + 2H2O.
Mercury oxide, HgO, can be produced by the combustion of mercury(I) and
mercury(II) nitrates and can be precipitated by the addition of a hydroxide salt to a
solution of a Hg2+ salt in water. The colour of HgO varies from red to yellow
depending on the particle size of the solid. It is soluble in water and forms what is
thought to be the hydroxide, Hg(OH)2, although this has never been isolated. The
hydroxide is amphoteric - it can act as either an acid or a base, depending on the
conditions - but it is more basic than it is acidic.
Mercury dissolves only in acids-oxidizing agent - H2SO4 (conc.), HNO3:
Hg + 4HNO3 (conc.) = Hg(NO3)2 + 2NO2↑+ 2H2O;
6Hg + 8HNO3 (dil.) = 3Hg2(NO3)2 + 2NO↑ + 4H2O;
Hg + 2H2SO4 = HgSO4 + SO2↑ + 2H2O.
Main compounds. Zinc oxide and zinc hydroxide are amphotheric compounds.
They react with acids, forming corresponding salts. They form zincates with alkalis
at the melting and hydroxocomplexes when react with alkalis solution:
Zn + 2NaOH
t

→
Na2ZnO2 + H2O
86
sodium zincate;
Zn(OH)2 + 2KOH = K2[Zn(OH)4].
potassium tetrahydroxozincate (II)
Salts of mercury show oxidation properties:
HgCl2 + Cu = CuCl2 + Hg;
2HgCl2 + CuCl2 = Hg2Cl2 + CuCl4.
Coordination compounds:
[Zn(NH3)4](OH)2 – tetraamminezinc(II) hydroxide;
K2[Zn(CN)4] – potassium tetratiozincate(II);
K2[HgI4] – potassium tetraiodomercuriate (II) (Nestler’s reactant).
Mercury salts react with NH3 to form white mercury amido chloride:
HgCl2 + 2NH3 = Hg(NH2)Cl↓ + NH4Cl.
Use in medicine. ZnO is used as powders and ointments for external application
on burns, skin infections and for skin protection; ZnSO4 – a topical astringent for
mucous membranes, especially those of the eye.
Cadmium is toxic and has a cumulative property. In medicine, practically is not
applied.
Mercury and its salts are highly toxic, but some preparations are used in
medicine as an antiseptic, diuretics and laxatives: HgCl2 – mercury chloride or
mercuric chloride (poison) in a concentration of 1:1000 as an antiseptic; HgNH4Cl
– in dermatology; HgO – in ophthalmology as an ointment.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Write the ground-state electron configuration and orbital notation for each of
the following atoms and ions: Zn, Hg, Zn2+ and Hg2+.
2. Write equations of the reactions of Zn, Cd, Hg with dilute and concentrated
solutions of hydrochloric, sulfuric and nitric acids.
3. How to obtain mercury (II) oxide from mercuric chloride? Write equation of the
reaction.
4. How does zinc react with alkalis, ammonia, and ammonium chloride? Write
equations of the reactions.
87
5. Write equations of the reactions of thermal decomposition of Zn(NO3)2,
Cd(NO3)2 and Hg(NO3)2 salts.
6. Write molecular and ionic equations of hydrolysis of zinc nitrate and mercury
(II) nitrate.
7. Write equations of the reactions:
а)HgCl2 + NH3(dil) → ...;
b) Hg2Cl2 + NH3(dil) →…;
c) Hg(NO3)2 + КІ(excess) →… .
8. Write the formulas of zinc subgroup compounds which are used in medicine.
b) Problems to Solve
1. Mercury chloride solution is prepared by dissolution of 1g of salts in 2 kg of
water. What is the molarity and mass percentage of prepared solution? (density ρ
= 1.0 g/ml).
Answer: 0.0018 М; 0.05 %.
2. Calculate solubility of mercury (II) chloride in mol/l, g/l and g/100g of Н2О.
Ksp(HgCl2) = 1.32·10–18.
Answer: 6.9⋅10–7 mol/l; 3.2⋅10–4 g/l; 3.2⋅10–5 g/100 g.
3. Calculate concentration of Cadmium ions in 0.1 М К2[Cd(CN)4] solution which
contains, in addition 6.5 g/l of КCN.
Answer: 7.8⋅10-15 mol/l.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: “Studying the chemical properties of d-elements
of IIB group”.
• When carrying out experiments with mercury salts one must be careful
because they are poisonous!
4.2.1. Red-ox properties of Zinc and Mercury.
а) Into four test-tubes put pieces of zinc and add: into the first one - water, into
the second - a solution of hydrochloric acid, into the third - a solution of nitric acid,
into the fourth - a concentrated alkali solution. What happens in the test-tubes?
Why at dissolving in nitric acid very diluted gas is not released? Write equations of
the reactions.
b) Dip a copper coin into a solution of dilute nitric acid and wash it with water.
Drip a salt of mercury (II) on it and after a few minutes dry with filter paper. Write
equation of the reaction.
4.2.2. Acid-base properties of Zn, Cd and Hg compounds.
To solutions of zinc, cadmium, mercury (I) and mercury (II) salts add few drops
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of sodium hydroxide. Mark the color of formed precipitates. Add an excess of
sodium hydroxide. Which precipitates are dissolved? Why? Write equations of the
reactions.
4.2.3. Preparation and properties of mercury (I) chloride.
а) To a solution of mercury (I) nitrate add a small amount of diluted
hydrochloric acid. What is observed?
b) Divide obtained precipitate into 3 parts, and add to the first test-tube
concentrated solution of hydrochloric acid, to the second - concentrated nitric acid
solution and ammonia solution to the third one. In which case the dissolution of the
precipitate is observed? Write equations of the reactions.
4.2.4. Preparation and properties of IIB group sulfides.
Add to the solutions of zinc, cadmium and mercury salts ammonium sulfide
solution. Note the color of the precipitates and write equations of the reactions.
Determine the solubility of the precipitates in solutions of hydrochloric and nitric
acids (at heating). Write equations of the reactions.
4.2.5. Hydrolysis of zinc, cadmium and mercury salts.
а) On a strip of universal indicator paper put 1 drop of solutions of zinc and
cadmium salts. Compare the pH of solutions. Write equations of the reactions.
b) Dissolve crystals of mercury (II) nitrate in a few drops of water. Explain the
turbidity. Determine the pH of solution by using a litmus paper. Write equations of
the hydrolysis reaction in molecular and ionic forms.
4.2.6. Properties of zinc, cadmium and mercury coordination compounds.
а) To 1–2 ml of zinc sulfate (ZnSO4) solution in a test-tube add a solution of
ammonium hydroxide (NH4OH). To obtained precipitate add an excess of
ammonium hydroxide solution. What is observed? Write equations of the reactions.
b) To 1–2 ml of cadmium sulfate (CdSO4) solution in a test-tube add
ammonium hydroxide solution (NH4OH). To received precipitate add an excess of
ammonium hydroxide solution. What is observed? Write equations of the reactions.
c) To 1 ml of mercury (II) chloride (HgCl2) solution in a test-tube add
ammonium tiocyanide (NH4CNS) solution. Add an excess of ammonium tiocyanide
solution to obtained precipitate. What is observed? Write equations of the
reactions.
d) Repeat the experiment 4.2.6.b but instead of potassium thiocyanate solution
take potassium iodide solution. Write equations of the reactions. What is the color
of obtained precipitate and solution? Leave solution for the next experiment.
4.2.7. Preparation of Nessler’s reagent.
To obtained in the experiment 4.2.6b potassium tetraiodomercurate (II) add
KOH solution. The resulting solution is called Nessler’s reagent and is used for
determination of ammonium ions in solution. Add Nessler’s reagent into the testtube with ammonium salt solution. What is observed? What is the color of
precipitate?
89
5. Conclusions and Interpretations. Lesson Summary
Topic 16
d-Elements of VIB group. Chromium elements family
1. Objectives
Chromium is a naturally occurring element found in rocks, animals, plants, soil,
and in volcanic dust and gases. Chromium (III) is a component of most soils. In
areas of serpentine and peridotite rocks, chromite is the predominant chromium
mineral. Chromium (III) is an essential trace element and a daily intake of 50 to
200 micrograms per day is recommended for normal glucose, protein, and fat
metabolism. The body can reduce chromium (VI) to chromium (III) and this
detoxification leads to increased levels of chromium (III). Chromium (VI)
compounds are highly toxic for both acute and chronic exposures and can cause
cancer.
Molybdenum is an essential trace element for several enzymes important to
animal and plant metabolism. Molybdenum is essential to humans. It is needed for
at least three enzymes. Sulfite oxidase catalyses the oxidation of sulfite to sulfate,
necessary for metabolism of sulfur amino acids. Sulfite oxidase deficiency or
absence leads to neurological symptoms and early death. Xanthine oxidase
catalyses oxidative hydroxylation of purines and pyridines including conversion of
hypoxanthine to xanthine and xanthine to uric acid. Aldehyde oxidase oxidises
purines, pyrimidines, pteridines and is involved in nicotinic acid metabolism. Low
dietary molybdenum leads to low urinary and serum uric acid concentrations and
excessive xanthine excretion.
2. Learning Targets
To study chemical reactions those characterize the properties of chromium,
molybdenum and tungsten and their most important compounds.
To know the biological role of these elements and their application in the
economy.
3. Self Study Section
3.1. Syllabus Content
General characteristics of d-elements of VI group. Chromium compounds in
nature. Simple substance and its chemical activity. Chromium carbonyl.
Chromium (II) compounds and their acid-base and redox characteristics.
Chromium (III) compounds and their acid-base and redox characteristics. The
ability to form complexes. Identification reaction of Cr3+ ion. Chromium (VI)
compounds – oxide and dichromic acid. Chromates and dichromates, their acidbase and redox properties. Chromium peroxide.
Molybdenum and Tungsten, general characteristics. The ability to form isopoly- and hetero-polyacids, redox properties of the compounds. Biological role of
90
chromium and molybdenum. Use of chromium, molybdenum and tungsten
compounds in pharmaceutical analysis and medicine.
3.2. Theoretical Backgrounds
Chromium is found in the center of the periodic table, a chart that shows how
chemical elements are related to each other. Elements in Groups 3 through 12 are
known as the transition metals. These elements all have similar physical and
chemical properties. They have a bright, shiny surface and high melting points.
Chromium is a fairly active metal. It does not react with water, but reacts with most
acids. It combines with oxygen at room temperature to form chromium oxide
(Cr2O3). Chromium oxide forms a thin layer on the surface of the metal, protecting
it from further corrosion (rusting).
Molybdenum does not occur free in nature; it is usually found in molybdenite
ore, MoS2, and wulfenite ore, PbMoO4. Molybdenum is also recovered as a byproduct of copper and tungsten mining. It is a silvery-white metal of the chromium
group. It is very hard and tough, but it is softer and more ductile than tungsten. It
has a high elastic modulus. Of the readily-available metals, only tungsten and
tantalum have higher melting points.
Molybdenum is a transition metal in Group 6 of the Periodic Table between
chromium and tungsten. Although molybdenum is sometimes described as a ‘heavy
metal’ its properties are very different from those of the typical heavy metals,
mercury, thallium and lead. It is much less toxic than these and other heavy metals.
Its low toxicity makes molybdenum an attractive substitute for more toxic
materials.
Compounds of molybdenum which are commonly encountered have
molybdenum in its highest oxidation state, VI, for example molybdenum trioxide,
MoO3, sodium molybdate, Na2MoO4.2H2O, and ammonium di- and
heptamolybdate, (NH4)2Mo2O7 and (NH4)6Mo7O24.4H2O. In aqueous solution
molybdenum(VI) is present as the simple molybdate, [MoO4]2- ion which is like
sulfate or, depending on the concentration and pH as a polymeric polymolybdate
ion. The lower oxidation state, IV, is found in the commonest ore of molybdenum
the disulfide, MoS2. Molybdenum(IV) also forms an oxide, MoO2. The redox
chemistry of molybdenum-oxygen compounds, as in selective oxidation catalysts
and molybdenum oxidase enzymes, has molybdenum cycling between oxidation
states (VI) and (IV).
Chromium in activity series of metals is placed near the zinc. It shows the +2,
+3 and +6 oxidation states. Molybdenum and Tungsten are located just before
hydrogen and react with strong oxidizing acid and do not react with diluted acids
according to the equations:
Cr + 2HCl = CrCl2 + H2↑;
Mo + 3H2SO4(s) = H2MoO4 + 3SO2 + 2H2O;
W + 2HNO3 + 8HF = H2[WF8] + 2NO↑ + H2O.
Cr(OH)2 is typical base that shows reducing properties, as well as all
compounds of chromium (II):
91
4CrCl2 + O2 + 4HCl = 4CrCl3 + 2H2O.
Chromium is one of the few metals that has the property of passivation.
Chromium will spontaneously react to form a thin layer of oxide that protects the
metal against further corrosion. This surface is hard and nonreactive. This makes
chromium ideal for electroplating other metals to protect them from oxidizing, and
because of it’s hardness, it is used to harden the surface of many objects, such as
metal tools. Chromium (III) compounds are stable, Cr2O3 is found in nature, it
shows amphoteric properties:
Cr2O3 + 2KOH = 2KCrO2 + H2O;
Cr2O3 + 3SO3 → Cr2(SO4)3.
Chromium (III) compounds often react as reducing agents in redox reactions:
Cr(SO4)3 + 3Br2 + 16KOH = 2K2CrO4 + 6KBr + 3K2SO4 + 8H2O;
Cr2O3 + KClO3 + 4KOH = 2K2CrO4 + KCl + 2H2O.
Oxidation state +6 is typical for all 3 elements. These compounds include
anhydrides, which are soluble in water or in alkalis to form acid and salt:
CrO3 + H2O → H2CrO4;
MoO3 + 2NaOH → Na2MoO4;
WO3 + 2KOH → Na2WO4.
Acids of these elements form isopoly acids when react with the strong acids:
H2CrO4 H2SO

4 → H2Cr2O7 H2SO

4 → H2Cr4O13;
Na2MoO4 H2SO

4 → Na2Mo2O7 H2SO

4 → Na2Mo3O10;
Na2WO4 H2SO

4 → Na2W2O7 H2SO

4 → H2W2O7.
Chromate, CrO42-, is a salt of chromic acid. This salt is associated with a yellow
color in basic conditions, for example potassium chromate. Dichromate, Cr2O72-, is
a salt of dichromic acid. This salt is associated with a strong orange color in acidic
conditions, for example potassium dichromate. However, compounds of chromate
or dichromate with heavy metals usually display a red color. Dichromate is a strong
oxidizing agent but it is a bad precipitating agent. Chromate on the other hand is
used as a precipitating agent but it is a poor oxidizing agent.
K2Cr2O7 + 3H2S + 4H2SO4 = Cr(SO4)3 + 3S + K2SO4 + 7H2O;
K2Cr2O7 + 14HCl = 2CrCl3 + 2KCl + 3Cl2 + 7H2O.
Potassium dichromate reacts with reducing agents to form superacids:
K2Cr2O7 + 4H2O2 + H2SO4 = K2SO4 + 2H2CrO6 + 3H2O.
blue colour
O
O
+ H2O → H2CrO6 or H2Cr2O12.
Cr
O
O
O
This reaction is used to determine cromium (VI) ions.
Ammonium molibdate (NH4)2MoO4 is used to determine phosphate-ions in
solution:
12(NH4)2MoO4+ H3PO4 + 21HNO3 → (NH4)3[P(Mo3O10)4] + 21NH4NO3 +12H2O.
92
Health effects of chromium: One radioactive isotope of chromium is used in
medical research, chromium-51. This isotope is used as a tracer in studies on blood.
A tracer is a radioactive isotope whose presence in a system can easily be detected.
The isotope is injected into the system at some point. Inside the system, the isotope
gives off radiation. That radiation can be followed by means of detectors placed
around the system.
A common use of chromium-51 is in studies of red blood cells. The isotope can
be used to find out how many blood cells are present in a person’s body. It can be
used to measure how long the blood cells survive in the body. The isotope can also
be used to study the flow of blood into and out of a fetus (an unborn child).
People can be exposed to chromium through breathing, eating or drinking and
through skin contact with chromium or chromium compounds. The level of
chromium in air and water is generally low. In drinking water the level of
chromium is usually low as well, but contaminated well water may contain the
dangerous chromium(IV); hexavalent chromium. For most people eating food that
contains chromium(III) is the main route of chromium uptake, as chromium(III)
occurs naturally in many vegetables, fruits, meats, yeasts and grains. Various ways
of food preparation and storage may alter the chromium contents of food. When
food in stores in steel tanks or cans chromium concentrations may rise.
Chromium(III) is an essential nutrient for humans and shortages may cause
heart conditions, disruptions of metabolisms and diabetes. But the uptake of too
much chromium(III) can cause health effects as well, for instance skin rashes.
Chromium(VI) is a danger to human health, mainly for people who work in the
steel and textile industry.
The health hazards associated with exposure to chromium are dependent on its
oxidation state. The metal form (chromium as it exists in this product) is of low
toxicity. The hexavalent form is toxic. Adverse effects of the hexavalent form on
the skin may include ulcerations, dermatitis, and allergic skin reactions. Inhalation
of hexavalent chromium compounds can result in ulceration and perforation of the
mucous membranes of the nasal septum, irritation of the pharynx and larynx,
asthmatic bronchitis, bronchospasms and edema. Respiratory symptoms may
include coughing and wheezing, shortness of breath, and nasal itch.
Molybdenum is an important alloying agent which contributes to the
hardenability and toughness of quenched and tempered steels. It also improves the
strength of steel at high temperatures. It is used in certain heat-resistant and
corrosion-resistant nickel-based alloys. Ferro-molybdenum is used to add hardness
and toughness to gun barrels, boilers plates, tools, and armor plate. The metal is an
essential trace element in plant nutrition. Molybdenum sulfide is used as a
lubricant, particularly at high temperatures where oils would decompose.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
93
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Write the electron configurations of Cr, Mo and W atoms and specify the
possible valences of these elements.
2. Write the reactions between dilute and concentrated hydrochloric, sulfuric and
nitric acids with metallic chromium, molybdenum and tungsten.
3. Write the formulas of oxides and hydroxides of Cr, Mo, W and show their acidbase properties.
4. How the redox properties of chromium compounds change, depending on its
oxidation state?
5. Write the equations of the following reactions:
a) CrCl3 + H2O2 + KOH → …;
b) K2Cr2O7 + Na2S + H2SO4 → …;
c) Fe(CrO2)2 + O2 + Na2CO3 → … .
6. Illustrate isopoly- and peroxy- acids (or peracids) using compounds of
Chromium.
b) Problems to Solve
1. What volume of 1 M sodium chromate solution can be obtained after fusing
sodium carbonate with 112 g of iron chromite?
Answer: 1.0 L.
2. Determine the molecular weight and empirical formula of tungsten fluoride, if
the density of tungsten fluoride steam is heavier than density of an air in 10.3
times.
Answer: WF6.
3. Calculate solubility of BaCrO4 in water (Ksp(BaCrO4) = 1,6⋅10–10) and in 0.001
М HCl solution (Кacid (H2CrO4) = 5.76⋅10–8).
Answer: 1.26⋅10–5 mol/l; 1.66⋅10 mol/l.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
94
4.2. Experimental Guidelines: «Chemical properties of chromium, molybdenum
and tungsten compounds».
4.2.1. Preparation of chromium (III) oxide.
Put sample of ammonium dichromate (1.2 g) in the porcelain cup and gently
heat to the beginning of the reaction. Weigh the porcelain cup with powder of
chromium (III) oxide after the cooling and calculate the percentage yield. Write
equation of the reaction.
4.2.2. Preparation and properties of chromium (III) hydroxide.
Prepare precipitate of chromium (III) hydroxide by careful adding (dropwise
with stirring) of sodium hydroxide solution to 3.4 ml of chromium (III) salt
solution. Divide obtained precipitate into two test-tubes. Add the excess of sodium
hydroxide solution to the first test-tube and hydrochloric acid solution to the second
one. Note the difference in the color of solutions. Write equations of the reactions.
Keep sodium chromate for the next experiment!
4.2.3. Chromium (III)-ion as a reducing agent.
To a solution of sodium chromite add bromine water and boil the mixture (the
experiment is carried out in a fume cupboard). What is the color of the solution?
Write equation of the oxidation reaction of sodium chromite into sodium chromate.
4.2.4. Transformation of chromate-ion in dichromate-ion.
To potassium chromate solution add dilute sulfuric acid. Note a change in color
of the solution. To the obtained solution of potassium dichromate add solution of
potassium hydroxide until its color changes. Write equations of the reactions of
transition of CrO42– ions into Cr2O72– ions and vice versa.
4.2.5. Oxidizing properties of chromium (VI) compounds.
To 2.3 cm3 of dilute solution of hydrogen peroxide add dilute sulfuric acid, 1-2
3
cm of ethyl ether and a few drops of potassium dichromate. Carefully mix the
mixture. What is in the ether and the water layer? Write equation of the reaction
and structural formula of chromium peroxide.
4.2.6. Reaction of identification of phosphate ions.
To 3-5 cm3 of phosphoric acid or its salt add a solution of (NH4)2MoO4 and
solution of nitric acid, and heat. What is observed? Write equation of the reaction.
5. Conclusions and Interpretations. Lesson Summary
Topic 17
d-Elements of VIIB group. Manganese elements family
1. Objectives
Manganese is one of the chemical elements that has both positive and negative
effects on living organisms. A very small amount of the element is needed to
95
maintain good health in plants and animals. The manganese is used by enzymes in
an organism. Enzymes are necessary to keep any cell operating properly. If
manganese is missing from the diet, enzymes do not operate efficiently. Cells begin
to die, and the organism becomes ill.
Fortunately, the amount of manganese needed by organisms is very small. It is
not necessary to take extra manganese to meet the needs of cells.
The element is a required trace mineral for all known living organisms. In larger
amounts, and apparently with far greater activity by inhalation, it can cause a
poisoning syndrome in mammals, with neurological damage which is sometimes
irreversible.
Typically for a transition metal, rhenium also acts a catalyst for many reactions.
Technetium is used as rust prevention and is used as a medical tool. As a medical
tool it is injected into the body and it goes to certain organs then it gives off
radiation that is easily detected. The amount of radiation given off indicates the
problems in the organs.
2. Learning Targets
To be able to write the equations of red-ox and acid-base reactions which
characterize the chemical properties of manganese and its compounds.
To know its biological role, economic importance and use of manganese
compounds in chemistry and medicine.
3. Self Study Section
3.1. Syllabus Content
General characteristics of manganese. Chemical activity of simple substance.
The ability to form coordination compounds (formation of carbonyles).
Manganese (ІІ) and manganese (ІІІ): acid-base and red-ox properties,
coordinaton compounds formation. Determinaton of Mn2+ ion. Manganese (ІV)
oxide, acid-base and red-ox properties, effect of pH on the redox properties.
Manganese (VІ) compounds. Manganese (VІI) compounds: acidic oxide,
permanganic acid, its salts, red-ox properties, oxidation of organic compounds,
thermal decomposition. The biological role of manganese. Application of
potassium permanganate in pharmaceutical analysis and as antiseptics solutions.
3.2. Theoretical Backgrounds
The most common oxidation states of manganese are +2, +3, +4, +6 and +7. Of
the wide variety of compounds formed by manganese, the most stable occur in
oxidation states +2, +6, and +7. These are exemplified, respectively, by the
manganous salts (with manganese as the Mn2+ ion), the manganates (MnO42−), and
the permanganates (MnO4−). As in the case of titanium, vanadium, and chromium,
the highest oxidation state (+7) of manganese corresponds to the total number of 3d
and 4s electrons. That state occurs only in the oxo species permanganate (MnO4−),
dimanganese heptoxide (Mn2O7), and manganese trioxide fluoride (MnO3F), which
show some similarity to corresponding compounds of the halogens — for example,
in the instability of the oxide. Manganese in oxidation state +7 is powerfully
96
oxidizing, usually being reduced to manganese in the +2 state. The intermediate
oxidation states are known, but, except for some compounds in the +3 and +4
states, they are not particularly important.
The chemistry of rhenium is rather diverse. Among other things, it shows the
largest range of oxidation states of any known element, namely −1, 0, +1, +2 and so
on all the way up to +7 — the last of which is its most common oxidation state. It is
also the metal that led to the discovery of the first metal–metal quadruple bond.
Manganese is placed in Activity series of metals before hydrogen, but rhenium
and technetium after hydrogen. That is why the listed elements do not react with
diluted acids, but react with strong oxidizing acids:
Mn + 2HCl = MnCl2 + H2↑;
Tc (Re) + 7HNO3 (dil.) = HReO4 (TcO4) + 7NO2↑ + 3H2O;
2Re + 7H2SO4 (conc.) = 4HReO4 + 7 SO2↑ + 6 H2O.
Manganese forms few oxides with oxygen:
MnO
Mn2O3
MnO2
MnO3
Mn2O7
basic
amphotheric
acidic
Oxidation state +2. Mn(OH)2 is base, MnSO4, Mn(NO3)2 are manganese
compounds which show reducing properties, for example:
2MnSO4 + 5PbO2 + 6HNO3 = 2HMnO4 + 2PbSO4 + 3Pb(NO3)2 + 2H2O.
Oxidation state +4. MnO2 shows amphoteric properties, in red-ox reactions can
be both oxidizing and reducing agent:
MnO2 + CaO 
→ CaMnO2;
MnO2 + 4HCl = MnCl2 + Cl2 + 2H2O;
2MnO2 + 3PbO2 + 6HNO3 = 2HMnO4 + 3Pb(NO3)2 + 2H2O.
Oxidation state +7 typical for the whole group of elements Mn, Re, Tc.
Corresponding compounds - HMnO4, HReO4. Potassium permanganate is a strong
oxidising agent, its oxidation-reduction potential depends on pH:
– in acidic medium MnO4– it is reduced to Mn+2:
2KMnO4 + 5H2O2 + 3H2SO4 = K2SO4 + 2MnSO4 + 5O2 + 8H2O;
– in neutral medium it is reduced to MnO2:
2KMnO4 + 3Na2SO3 + HOH = 2MnO2 + 3Na2SO4 + 2KOH;
– in basic medium it is reduced to Mn+6:
2KMnO4 + Na2SO3 + 2KOH = 2K2MnO4 + Na2SO4 + H2O.
Manganese is an important trace element in nutrition, although exposure to the
element is toxic in higher quantities. The deep-purple compound potassium
permanganate (KMnO4) is used for disinfecting, deodorizing, and bleaching and as
an analytical reagent. Technetium is used as rust prevention and is used as a
medical tool. As a medical tool it is injected into the body and it goes to certain
organs then it gives off radiation that is easily detected. The amount of radiation
given off indicates the problems in the organs.
t
97
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. What is the electron configuration of manganese?
2. Show the difference and similarity of the electron shells structures of atoms of
VIIA and VIIB groups elements using as an example chlorine and manganese
element. Indicate how it affects on the properties of oxides and hydroxides.
3. Write down the chemical reaction between manganese and the halogen fluorine.
4. How reactive is manganese with water?
5. Draw the structural formulas of manganese oxides and corresponding
hydroxides and write the examples of the reactions that characterize the acidbase properties of these compounds.
6. Where are Mn, Tc, Re located in a Table of Standart Electrode Potentials?
Write equations of the reactions of listed elements with acids.
7. What is one of manganese’s main uses in technology?
8. Write the following equations of the reactions:
→ …;
a) MnO2 + NaOH 
b) MnO2 + H2SO4 (conc.) → …;
t
c) MnO2 + KOH + KСlO3 
→ … .
9. How does pH influence on red-ox properties of potassium permanganate?
10. Complete the following equations of the reactions:
a) KMnO4 + H2O2 + … → MnSO4 + K2SO4 + O2 + …
b) KMnO4 + MnSO4 + … → MnO2 + …
c) KMnO4 + K2SO3 + … → K2MnO4 + K2SO4 + … .
b) Problems to Solve
1. What volume of potassium permanganate containing 15.804 g of KMnO4 in 1
dm3 of solution is required for oxidation of 50 ml of 0.1 N hydrogen sulfide
solution in acidic medium?
Answer: 10 ml.
t
98
2. What volume of solution with concentration of potassium permanganate 0.1046
mol-eq/ml is required for oxidation in acid medium of iron (II) sulfate, prepared
by dissolution of 0.1242 g of iron in dilute sulfuric acid?
Answer: 21.2 ml.
3. What mass of potassium permanganate is required to obtain 10 ml of chlorine
(standard condition) at its interaction with concentrated hydrochloric acid?
Answer: 28.2 g.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of manganese compounds».
4.2.1. Preparation and properties of manganese (II) hydroxide.
Into a test-tube pour 2-3 ml of manganese sulphate solution and add an alkali
solution up to precipitate formation. Divide the precipitate into two parts and
examine its dissolving in acid and alkali. Write equations of the reactions.
4.2.2. Reducing properties of manganese (II) compounds.
To 2-3 cm3 of manganese (II) salt pour 1 cm3 of 2 M potassium (sodium)
hydroxide solution. To the obtained precipitate add 5-6 drops of bromine water.
What is observed? Write equation of the reaction.
4.2.3. Preparation of slightly soluble salts of manganese (II).
In three test-tubes obtain: a) manganese chromate, b) manganese carbonate, c)
manganese sulfide by the reaction of manganese (II) sulfate with potassium
chromate, sodium carbonate and ammonium sulfide respectively. What is
observed? To each precipitate add 2-3 ml of hydrochloric acid and conclude about
the solubility of the obtained manganese (II) salts in acid. Write equations of the
reactions.
4.2.4. Properties of manganese (IV) oxide.
а) Oxidizing properties. Into a test-tube put 1 g of MnO2 and add 1-2 drops of
concentrated hydrochloric acid solution. Using moistened with water iodine-starch
paper observe evolution of chlorine. Write equation of the reaction.
b) Reducing properties. In the crucible/cup melt 3-4 granules of potassium
hydroxide, a few crystals of potassium nitrate and a bit of MnO2 (avoid the excess
of MnO2). Observe the color of the melt. Cool it and dissolve in water. Write
equation of the reaction.
4.2.5. Oxidizing properties of potassium permanganate, depending on the pH of
the medium.
Into three test-tubes pour 1-2 ml of: sulfuric acid into the first, water into the
second and sodium hydroxide into the third. Then, to each tube add 2 ml of sodium
99
sulfite and 3-5 drops of potassium permanganate. Note changes in color and write
equations of the reactions.
4.2.6. Potassium permanganate oxidation with ethanol.
Into two test-tubes pour 0.5 ml of potassium permanganate. Into one of
them, add 1 ml of sulfuric acid solution, and 1 ml of sodium hydroxide solution –
into the second. To both tubes add 0.5 ml of ethanol and heat the solutions.
Observe changes in color of solutions. Write equations of the reactions of
potassium permanganate reduction with ethanol in acidic and basic mediums.
5. Conclusions and Interpretations. Lesson Summary
Topic 18
d-Elements of VIIIB group. Iron and Its Compounds
1. Objectives
Iron is believed to be the tenth most abundant element in the universe. Iron is
also the most abundant (by mass, 34.6%) element making up the Earth; the
concentration of iron in the various layers of the Earth ranges from high at the inner
core to about 5% in the outer crust. Most of this iron is found in various iron
oxides, such as the minerals hematite, magnetite, and taconite. The Earth’s core is
believed to consist largely of a metallic iron-nickel alloy.
Iron is of critical importance to plants, humans, and animals. It occurs in
hemoglobin, a molecule that carries oxygen in the blood. Hemoglobin picks up
oxygen in the lungs, and carries it to the cells. In the cells, oxygen is used to
produce energy the body needs to survive, grow, and stay healthy.
An iron deficiency (lack of iron) can cause serious health problems in humans.
For instance, hemoglobin molecules may not form in sufficient numbers. Or they
may lose the ability to carry oxygen. If this occurs, a person develops a condition
known as anemia. Anemia results in fatigue. Severe anemia can result in a lowered
resistance to disease and an increase in heart and respiratory (breathing) problems.
Some forms of anemia can even cause death.
2. Learning Targets
To be able to write equations of red-ox and acid-base reactions which
characterize the chemical properties of iron and its compounds.
To know biological role of iron, its economic importance and use of iron
compounds in chemistry and medicine.
3. Self Study Section
3.1. Syllabus Content
General characteristic of iron, its ionic state, coordination number. Occurrence
in nature. Chemical activity of iron. Complex formation ability. Corrosion of iron
products.
100
The compounds of iron (II) - acid-base and red-ox properties. Complex
compounds with cyanide and thiocyanate ions, porphyrins. Biological role of
hemoglobin. Iron (III) compounds. Iron (III) oxide and hydroxide. Iron (III)
chloride and its hydrolysis. Complex compounds of iron (III). Determination of
Fe2+ and Fe3+ cations. Iron (VI) compounds. Preparation of ferrates and their
oxidizing properties.
Application of iron and its compounds in medicine.
3.2. Theoretical Backgrounds
Pure iron is chemically reactive and corrodes rapidly, especially in moist air or
at elevated temperatures.There are four allotropes of iron known as 'ferrites'. These
are designated α-, β-, γ-, and δ- with transition points at 770, 928, and 1530 °C. The
α- and β- ferrites have the same crystal structure, but when the α- form becomes the
β- form, the magnetism disappears.
The most common iron ore is hematite (Fe2O3 mostly). Iron is also found in
magnetite (Fe3O4) and taconite (a sedimentary rock containing more than 15% iron
mixed with quartz).
Chemical element iron shows oxidation states +2, +3, +6 in its compounds. Iron
metal reacts in moist air by oxidation to give a hydrated iron oxide. This does not
protect the iron surface to further reaction since it flakes off, exposing more iron
metal to oxidation. This process is called rusting and is familiar to any car owner.
Finely divided iron powder is pyrophoric, making it a fire risk.
On heating with oxygen, O2, the result is formation of the iron oxides Fe2O3 and
Fe3O4:
4Fe + 3O2 → 2Fe2O3;
3Fe + 2O2 → Fe3O4.
Air-free water has little effect upon iron metal. However, iron metal reacts in
moist air by oxidation to give a hydrated iron oxide. This does not protect the iron
surface to further reaction since it flakes off, exposing more iron metal to oxidation.
This process is called rusting and is familiar to any car owner.
Iron reacts with excess of the halogens F2, Cl2, and Br2, to form ferric, that is,
Fe(III), halides:
2Fe + 3F2→ 2FeF3 (white);
2Fe + 3Cl2→ 2FeCl3 (dark brown);
2Fe + 3Br2 → 2FeBr3 (reddish brown).
Iron metal dissolves readily in dilute sulphuric acid in the absence of oxygen to
form solutions containing the aquated Fe(II) ion together with hydrogen gas, H2:
Fe + H2SO4 → FeSO4 + H2.
In practice, the Fe(II) is present as the complex ion [Fe(OH2)6]2+.
If oxygen is present, some of the Fe(II) oxidizes to Fe(III):
Fe2+ – 1e– → Fe3+.
The strongly oxidizing concentrated nitric acid, HNO3, reacts on the surface of
iron and passivates the surface.
Fe (II) compounds. Oxide FeO and hydroxide Fe (OH)2 have basic properties,
101
they are insoluble in water, and highly soluble in mineral acids.
Fe(OH)2 is easily oxidized by atmospheric oxygen:
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3.
white
brown
The salts of iron (II) oxidize into the salts of iron (III), for example:
2KMnO4 + 10FeSO4 + 8H2SO4 → 5Fe2(SO4)3 + 2MnSO4 + K2SO4 + 8H2O;
3FeCl2 + 4HNO3 → 2FeCl3 + Fe(NO3)3 + NO + 2H2O;
4FeSO4 + O2 + 2H2O → 4FeOHSO4.
Fe(III) compounds. Fe2O3 oxide shows amphotheric properties:
Fe2O3 + 3H2SO4 → Fe2(SO4)3 + 3H2O;
Fe2O3 + Ca(OH)2 → Ca(FeO2)2 + H2О;
Fe(OH)3 + 3HCl → FeCl3 + 3H2O;
Fe(OH)3 + 3NaOH → Na3[Fe(OH)6].
All salts of iron (III) in solution hydrolyze:
Fe3+ + HOH ⇄ FeOH2+ + H+, pH < 7.
Fe(VI) compounds are obtained at melting of iron (III) oxide with oxidizing
agents in basic medium:
Fe2O3 + 3KNO3 + 4KOH → 2K2FeO4 + 3KNO2 + 2H2O.
The elements of iron subgroup form a cationic, anionic and neutral complexes.
Typical coordination numbers are 6 and 4. Among cyanide complexes the most
important are K3[Fe(CN)6] and K4[Fe(CN)6] salts. These salts are used in the
qualitative analysis.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Write down electron configuration of an atom of iron.
2. Write the formulas of iron oxides and show their acid-base properties by writing
the corresponding equations of the reactions.
3. Where is iron placed in a Table of Standart Electrodes Potentials? Write down
equations of the reactions of iron with hydrochloric, sulfuric and nitric acids
(concentrated and diluted).
4. Complete equations of the reactions:
102
a) FeSO4 + KCN →…;
b) Fe3O4 + H2SO4→…;
d) K4[Fe(CN)6] + Cl2 → …;
c) Fe2(SO4)3 + H2S → …;
e) FeSO4 + K2Cr2O7 + H2SO4 → …; f) K3[Fe(CN)6] + H2O2 + KOH →… .
5. How do the redox properties of iron compounds change depending on the
oxidation state of elements? Write equations of the reactions.
6. What products are formed at mixing of the following solutions: а) FeCl3 and
Na2CO3; b) Fe2(SO4)3 and (NH4)2S? Write equations of the reactions.
7. Give examples of coordination compounds of iron in biology and medicine.
b) Problems to Solve
1. Will a precipitate form at mixing of equal volumes of 0.0001 M iron (II)
sulphate and sodium sulfide solutions? Ksp(FeS) = 3.7⋅10–18.
Answer: precipitate will form.
2. Calculate mass of iron (III) oxide obtained at thermal decomposition of 18 g of
the corresponding hydroxide, if the degree of decomposition is 91%.
Answer: 12.25 g.
3. What volume of iron (III) chloride solution with a concentration of 0.1 N should
be taken to release 0.48 g of iodine from solution of potassium iodide?
Answer: 37.5 ml
4. Calculate mass of metal (in tones), which can be extracted from 700 tons of iron
ore containing 35% of impurities, if the mass of pure iron in melted metal is
95%.
Answer: 478.95 t.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: “d-elements of VIIIB group. Fe and its
compounds”
4.2.1. Reactions of identification of iron (II) and iron (III) ions.
а) To 0.5 ml of Mohr’s salt solution pour 5-10 drops of red prussiate of potash
K3[Fe(CN)6]. What is the color of the precipitate formed? Name the obtained
compound.
b) To 0.5 ml of iron (III) salt solution pour 5-10 drops of yellow prussiate of
potash K4[Fe(CN)6]. What is the color of the precipitate formed? Name the
obtained compound.
c) To 0.5 ml of iron (III) salt solution pour 5-10 drops of ammonium or
potassium thiocyanate. Repeat the same experiment with Mohr’s salt solution.
What is the color of the solution in first and the second case? Write equations of
the processes.
4.2.2. Reactions of iron with acids.
103
Into two test-tubes pour 1-2 cm3 of dilute acid solution: hydrochloric into the
first, and nitric - into the second. Into each tube add several pieces of iron wire.
Observe what happens in each tube. Then, to each tube add 5-7 drops of
ammonium thiocyanate. In which tube a characteristic color of iron (III) ions is
formed? Which of acids ions are oxidizing agents in each reaction? Write equations
of the reactions.
4.2.3. Preparation and properties of iron (II) hydroxide.
To 1-2 cm3 of Mohr’s salt solution (NH4)2Fe(SO4)2 add 0.5-1 ml of sodium
hydroxide solution to obtain a precipitate. Observe how color of the precipitate
changes with time. Write equations of the reactions.
4.2.4. Reducing properties of iron(II) compounds.
Into two test-tubes pour 1-2 cm3 of a solution of iron sulphate or Mohr’s salt,
add 0.5 cm3 of dilute sulfuric acid. Then, add bromine water into the first tube, and
potassium permanganate – into the second. What is observed? How to prove
whether iron (II) ions remain in solution? Write equations of the reactions.
4.2.5. Preparation and properties of iron (III) hydroxide.
Into two test-tubes pour 0.5 cm3 of iron (III) salt solution. Into each tube add 510 drops of sodium hydroxide solution. What is the color of precipitate? Into one
of tubes pour 0.5-1 cm3 of 2 M hydrochloric acid solution and 1-2 cm3 of
concentrated sodium hydroxide - into the second. In which tube the precipitate is
dissolved? Write equations of the reactions.
4.2.6. Оxidizing properties of iron (III) compounds.
Into two test-tubes pour 0.5-1 cm3 of iron (III) chloride solution. Into the first
tube pour 1-2 cm3 of hydrogen sulfide, and into the second - 0.5-1.0 cm3 of
potassium iodide. What changes are observed? Write equations of the reactions.
4.2.7. Hydrolysis of iron (II) and iron (III) salts.
On two strips of universal indicator paper put drop by drop 0.5 N iron (II)
sulfate and iron (III) chloride solutions. Determine the pH of solutions. Write
equations of the reactions.
5. Conclusions and Interpretations. Lesson Summary
Topic 19
d-Elements of VIIIB group. Cobalt and Nickel Compounds.
Platinum metals
1. Objectives
There are 10 radioactive isotopes of cobalt that are currently known. Cobalt-60
is one of the most commonly used radioactive isotopes and is used in medicine to
find a treat certain diseases including the Schilling test which determines if a
patients body is making and using vitamin B-12 effectively. Co-57,58 are also used
104
for the same purpose. Cobalt-60 is also used to treat cancer, because the radiation it
gives off kills cancer cells. Cobalt is a trace mineral that the human body needs in
only small amounts. When trace metals are absent in a diet this leads to health
problems. Animals use trace minerals to make essential enzymes which function as
catalysts. Enzymes are needed for living cells to function properly. For example
cobalt is needed for the natural production of B-12 vitamins. B-12 vitamins ensure
that enough red blood cells are produced in the human body.
Cobalt compounds are commonly used to make colored glass, glazes, paints,
rubber, inks, cosmetics, and pottery. These compounds compounds include: cobalt
oxide, cobalt potassium nitrite, cobalt aluminate, and cobalt ammonium phosphate.
Cobalt compounds can also be used as catalyst.
Nickel is an essential trace element for many species. Enzymes known as
hydrogenases in bacteria contain nickel. Nickel is also important in plant ureases.
2. Learning Targets
To study chemical reactions that characterize the properties of d-elements of
VIII group.
To know their biological role and application in medicine and pharmacy, as
well as their national economic importance.
3. Self Study Section
3.1. Syllabus Content
Cobalt and Nickel. Valence states. Chemical activity. The most important
compounds of cobalt (II), cobalt (III) and nickel (II). Characteristics of redox
properties. Hydrolysis of cobalt (II) and nickel (II) salts. Complex compounds with
cyanide, thiocyanate and fluoride ions. Aqua-complexes. Vitamine В12. Reactions
of Со2+ and Ni2+ cations identification. Chugaev elimination.
The biological significance and chemical basis of application of cobalt and
nickel compounds in medicine and pharmacy.
Platinum metals, general characteristics of simple substances and their
interaction with acids. Physical properties and applications of platinum metals.
Complex compounds of platinum (II) and platinum (IV), coordination numbers,
structure, oxidation reactions, reduction reactions and replacement. Oxides of
osmium (VIII) and ruthenium (III). Chemical basis of application of platinum
group metals compounds in medicine.
3.2. Theoretical Backgrounds
Cobalt is a hard ferromagnetic, silver-white, hard, lustrous, brittle element. It is
a member of group VIII of the periodic table. Like iron, it can be magnetized. It is
similar to iron and nickel in its physical properties. The element is active
chemically, forming many compounds. Cobalt is stable in air and unaffected by
water, but is slowly attacked by dilute acids. Cobalt and nickel have very similar
chemical properties. Both are magnetic and both resist corrosion.
Nickel is a relatively unreactive element. At room temperature, it does not
combine with oxygen or water or dissolve in most acids. At higher temperatures, it
105
becomes more active. For example, nickel burns in oxygen to form nickel oxide
NiO. It also reacts with steam to give nickel oxide and hydrogen gas.
Chemical elements Co and Ni show oxidation states +2 and +3. The most stable
cobalt and nickel compounds have a valence - II.
Co and Ni are placed before hydrogen in a Table of Standard Electrode
Potentials. Chemical activity of methals increases at heating. All metals react with
oxygen, halogens, sulfur, nitrogen, carbon and hydrogen. These metals displace
hydrogen from dilute acids, concentrated solution of HNO3 and H2SO4 passive
cobalt and nickel.
Me(II)compounds. MeO - oxides and Me(OH)2 - hydroxides show basic
properties, insoluble in water and highly soluble in mineral acids.
Me(II) compounds can be oxidized with oxidizing agent to compounds of
Me(III):
Me2+ – 1e– → Me3+.
Reducing properties decrease in a row Fe – Co – Ni. Co(OH)2 is oxidized by
the oxygen slowly, Ni(OH)2 is oxidized in a basic medium:
2Ni(OH)2 + 2NaOH + Br2 → 2Ni(OH)3 + 2NaBr.
Me(III)compounds. Oxidation state +3 is not typical for Co and Ni. Under the
action of acids on bases of Me(ОН)3 salts of Me(II) are formed:
2Co(OH)3 + 6HCl → 2CoCl2 + Cl2↑ + 6H2O;
4Ni(OH)3 + 4H2SO4 → 4NiSO4 + O2↑ + 10H2O.
Co(III) compounds form a lot of the complexex with different kind of ligands:
[Co(NH3)5Cl]Cl2, [Co(NH3)5H2O]Cl3, [Co(NH3)3(NO2)3].
Co and Ni form complex compounds with carbon (II) oxide - carbonyles:
[Ni(CO)4], [Co2(CO)8].
Health effect: Like many other metals, small quantities of cobalt are essential to
the survival of many animals, including humans. Cobalt is a main ingredient of
Vitamin B12, whose deficiency can cause the brain and nerves to function
abnormally. A little cobalt in soil helps to improve the health of grazing animals
substantially. However, exposure to large amounts of cobalt can be toxic, while the
metal in powdered form is extremely flammable, and thus a fire hazard.
Cobalt is beneficial for humans because it is a part of vitamin B12, which is
essential for human health. Cobalt is used to treat anaemia with pregnant women,
because it stimulates the production of red blood cells. The total daily intake of
cobalt is variable and may be as much as 1 mg, but almost all will pass through the
body unadsorbed, except that in vitamine B12.However, too high concentrations of
cobalt may damage human health. When we breathe in too high concentrations of
cobalt through air we experience lung effects, such as asthma and pneumonia. This
mainly occurs with people that work with cobalt.
Health effects may also be caused by radiation of radioactive cobalt isotopes.
This can cause sterility, hair loss, vomiting, bleeding, diarrhoea, coma and even
death. This radiation is sometimes used with cancer-patients to destroy tumors.
These patients also suffer from hair loss, diarrhea and vomiting.
106
Nickel can pose a health hazard to certain individuals. The most common health
problem is called nickel allergy. Some people are more likely to develop nickel
allergy than are others. People who are sensitive to nickel may develop a skin rash
somewhat like poison ivy. The rash becomes itchy and may form watery blisters.
Once a person gets nickel allergy, it remains with him or her forever.
Nickel can cause more serious health problems too. For example, people who are
exposed to nickel fumes (dust and gas) breathe in nickel on a regular basis. Long
term nickel exposure may cause serious health problems, including cancer.
3.3. References
1. Ye.Ya.Levitin, I.A.Vedernikova. General and Inorganic chemistry. – Kharkiv:
Publishing House of NUPH “Golden Pages”, 2009.
2. Darrel d. Ebbing. General Chemistry. – Boston: Houghton Mifflin company,
1984.
3. Raymond Chang. Chemistry. – New York: McGraw Hill, 2010. – p. 103–131.
4. John Olmsted III, Gregory M. Williams. Chemistry. The Molecular Science.
Mosby. – 1994.
5. Steven S. Zumdahl. Chemistry (4th Edition). – Houghton Mifflin Company. –
1997.
6. Gary L. Miessler, Donald A. Tarr. Inorganic Chemistry. – Prentice Hall. – 1991
3.4. Test Yourself
а) Review Questions.
1. Write the formulas of cobalt and nickel oxides and show their acid-base
properties by writing the corresponding equations of the reactions.
2. Place of cobalt and nickel in a row of standard redox potentials. Write
equations of the reactions of cobalt and nickel with hydrochloric, sulfuric and
nitric (dilute and concentrated) acids.
3. Write equations of the following reactions:
a) Co2O3 + HCl → …;
b) Ni(OH)3 + 4H2SO4 → … .
4. How do the redox properties of the compounds Co(II), Co(III), Ni(II) and
Ni(III) change depending on the oxidation state of the elements? Write
equations of the reactions.
5. Complex compounds of cobalt and nickel in biology and medicine.
b) Problems to Solve
1. Will a precipitate form at mixing of equal volumes of 0.0001 M solutions of
cobalt (II) chloride and sodium sulfide? Ksp(СоS) = 4.0⋅10–21.
Answer: precipitate will form.
2. What volume of gas will release at dissolution of 11 g of cobalt (III) hydroxide
in concentrated hydrochloric acid? What volume of the acid (d = 1.18 g/ml) will
spend?
107
Answer: 1.12 L; 25.8 ml.
4. Laboratory Activities and Experiments Section
4.1. Practical Skills and Suggested Learning Activities
а) discussion and explanation of main questions of the topic;
b) solving of typical numerical problems;
c) explanation of laboratory work technique.
4.2. Experimental Guidelines: «Chemical properties of VIII group d-elements
and their compounds».
4.2.1. Preparation and properties of cobalt (II) and nickel (II) hydroxides.
Into one of the two tubes pour 1-2 cm3 of cobalt (II) chloride, and into the
second - 1-2 ml of nickel (II) sulfate solution. To each tube pour 0.5-1 ml of
sodium hydroxide solution. How has the color of the solutions changed? Divide
each of the prepared solutions into two parts:
а) to the one part add sodium hydroxide and conclude about their amphotery.
b) to the second part add 5-10 drops of bromine water. How does the color of
the precipitate change?
Write equations of the reactions.
4.2.2. Preparation of ammonium complexes of cobalt (II) and nickel (II).
Into one test pour 0.5-1 cm3 of cobalt (II) chloride, and 0.5-1 ml of nickel (II)
sulfate - into the second. To both tubes add 25% solution of ammonium hydroxide
dropwise to obtain the corresponding hydroxide precipitates. What is the color of
precipitate? To each tube pour 25% solution of ammonium hydroxide to dissolve
the precipitates till formation of complex compounds – ammoniates. How does the
color in the tubes change? Write equations of the reactions. Which of the obtained
complexes is more stable?
5. Conclusions and Interpretations. Lesson Summary
108
Appendixes
AppendixA
Physical Constants
Constant
Symbol
Value
Абсолютний нуль температури
Т
–273,15 °С
Діелектрична проникність
вакууму
Electron charge
Molar volume of gas
ε0
8,85⋅10–12 Ф/м
е
1,602⋅10–19 Кл
Vm
22,41⋅10–3 м3/mol
Avogadro’s number
NA
6,02⋅1023 mol–1
Faradey constant
F
9,648⋅104 Кл/mol
Planck’s constant
h
6,626⋅10–34 J⋅s
Rydberg constant
R
1,097⋅107 m–1
Speed of light (in vacuum)
с
2,997⋅108 m/s
Appendix B
SI Units Співвідношення між позасистемними одиницями та одиницями
СІ
Unit
Value
Symbol
Length
мк
Å
Amount of energy
Energy
cal
еV
Entropy
Mass
Volume
Temperature
Pressure
Time
е.о.
t
g
а.m.u.
L
°С
atm
mm Hg
min
name
micron
angsrtom
calorie
електрон-вольт
ентропійна одиниця
tonne
gram
atomic mass unit
liter
celsius
atmosphere
milimeter of mercury
minute
109
еквівалент СІ
10–6 m
10–10 m
4,18 J
1,602⋅10–19 J
4,18 J/К
103 кg
10–3 кg
1,66⋅10–27 кg
10–3 м3 (1 L)
Т= t + 273,15
1,0133⋅105 Pa
133,32 Pa
60 s
Appendix C
Standart Thermodynamical Functions of Some Substances at 298 К
Substances
Al (s)
Al2O3 (s)
С (гр.)
∆H 0298 ,
kJ/mol
∆G 0298 ,
kJ/mol
S0298 ,
J/(mol·К)
0
28,3
0
–1676,0
50,9
–1582,0
0
5,7
0
ССl4 (l)
–135,4
214,4
–64,6
СО (g)
–110,5
197,5
–137,1
CO2 (g)
–393,5
213,7
–394,4
СаСО3 (s)
–1207,0
88,7
–1127,7
СаF2 (s)
–1214,6
68,9
–1161,9
Ca3N2 (s)
–431,8
105,0
–368,6
СаО (s)
–635,5
39,7
–604,2
Са(ОН)2 (s)
–986,6
76,1
–896,8
0
222,9
0
Сl2O (g)
76,6
266,2
94,2
СlO2 (g)
105,0
257,0
122,3
Cl2 (g)
Сl2O7 (l)
251,0
–
–
Cr2O3 (s)
–1440,6
81,2
–1050,0
CuO (s)
–162,0
42,6
–129,9
Fe (s)
0
27,2
0
FeO (s)
–264,8
60,8
–244,3
Fe2O3 (s)
–822,2
87,4
–740,3
Fe3O4 (s)
–1117,1
146,2
–1014,2
0
130,5
0
HBr (g)
–36,3
198,6
–53,3
HCN (g)
135,0
113,1
125,5
HCl (g)
–92,3
186,8
–95,2
HF (g)
–270,7
178,7
–272,8
HI (g)
26,6
206,5
1,8
HN3 (l)
294,0
328,0
238,8
H2O (g)
–241,8
188,7
–228,6
H2O (l)
–285,8
70,1
–237,3
H2S (g)
–21,0
205,7
–33,8
KCl (s)
–435,9
82,6
–408,0
KClO3 (s)
–391,2
143,0
–289,9
H2 (g)
110
Substances
∆H 0298 ,
kJ/mol
∆G 0298 ,
kJ/mol
S0298 ,
J/(mol·К)
MgCl2 (s)
–641,1
89,9
–591,6
Mg3N2 (s)
–461,1
87,9
–400,9
MgO (s)
–601,8
26,9
–569,6
0
191,5
0
HN3 (l)
294,0
328,0
238,8
NH3 (g)
–46,2
192,6
–16,7
N2 (g)
NH4NO2 (s)
–256,0
–
–
NH4NO3 (s)
–365,4
151,0
–183,8
82,0
219,9
104,1
NO (g)
90,3
210,6
86,6
N2O3 (g)
83,3
307,0
140,5
NO2 (g)
33,5
240,2
51,5
N2O4 (g)
9,6
303,8
98,4
N2O5 (g)
–42,7
178,0
114,1
NiO (s)
–239,7
38,0
–211,6
О3 (g)
142,3
237,7
163,4
O2 (g)
0
205,0
0
P2O5 (s)
–1492,0
114,5
–1348,8
PbO (s)
–219,3
66,1
–189,1
PbO2 (s)
–276,6
74,9
–218,3
0
31,9
0
SO2 (g)
–296,9
248,1
–300,2
SO3 (g)
–395,8
256,7
–371,2
SiH4 (g)
34,7
204,6
57,2
SiO2 (quartz)
–910,9
41,8
–856,7
SnO (s)
–286,0
56,5
–256,9
SnO2 (s)
–580,8
52,3
–519,3
N2O (g)
S (s)
Ti (s)
0
30,6
0
TiCl4 (l)
–804,2
252,4
–737,4
TiO2 (s)
–943,9
50,3
–888,6
WO3 (s)
–842,7
75,9
–763,9
ZnO (s)
–350,6
43,6
–320,7
111
Appendix D
Standart Thermodynamical Functions of Some Organic Substances at 298 K
Fotmule and State
∆H 0298 ,
kJ/mol
S0298 ,
J/(mol·К)
∆G 0298 ,
kJ/mol
СН4 (g)
–74,9
186,2
–50,8
С2Н2 (g)
226,8
200,8
209,2
С2Н4 (g)
52,3
219,4
68,1
С2Н6 (g)
–89,7
229,5
–32,9
С6Н6 (l)
49,0
124,5
172,8
–277,6
160,7
–174,8
СН3СООН (l)
–484,4
159,9
–389,6
СО(NH2)2 (s)
СО(NH2)2 (l)
–333,0
–317,7
104,7
175,7
–196,9
–202,7
С6Н12О6 (s)
С6Н12О6 (l)
–1273,0
–1263,1
212,1
264,0
–910,5
–914,5
С12Н22О11 (s)
С12Н22О11 (l)
–2220,9
–2215,8
360,2
403,8
–1544,3
–1551,4
С2Н5ОН (l)
Appendix E
The main Hafl-Reactions and Standart Red-Ox Potentials Values
Hafl-Reactions
Oxidized Form
S2O82–
2SO42–
2,01
2e
2H2O
1,78
2e–
Pb2+ + 2H2O
1,69
5e
Mn2+ + 4H2O
1,51
2e–
Cl–+ H2O
1,49
6e
Cl– +3H2O
1,45
6e–
Br– + 3H2O
1,44
–
H2O2 + 2H
PbO2 + 4H+
MnO4–
–
Reduced Form
2e–
+
+
+ 8H
ClO +2H+
ClO3– + 6H+
BrO3– +6H+
ClO4– + 8H+
Cl20
Cr2O72– + 14H+
2NO3– + 12 H+
O20 + 4H+
2ІO3– + 12H+
Br20
φ0, В
nе–
–
–
–
–
8e
Cl + 4H2O
1,39
2e–
2Cl–
1,36
6e–
2Cr3+ + 7H2O
1,35
–
10e
N20
4e–
2H2O
10e–
І20 + 6H2O
1,20
2Br–
1,07
–
2e
112
+ 6H2O
1,24
1,23
Hafl-Reactions
φ0, В
nе–
Oxidized Form
Reduced Form
NO2– + 2H+
e–
NO + H2O
1,00
NO3–
–
+
+ 4H
3e–
NO + 2H2O
0,96
ClO + H2O
2e–
Cl– + 2 OH–
0,89
NO3–
NO3–
3+
+
+2H
–
2e
NO2– + H2O
0,84
+2H+
e–
NO2 + H2O
0,78
–
2+
Fe
e
Fe
O20 + 2H+
2e–
H2O2
0,68
MnO4– + 2H2O
3e–
MnO2 + 4OH–
0,57
MnO4–
e–
MnO42–
0,54
2І–
0,54
І20
–
2e
+
–
0,77
–
O2 + 2H
4e
4OH
SO42– + 8H+
6e–
S0 + 4H2O
0,40
0,36
S4O62–
2e–
2S2O32–
0,22
SO42– + 2H+
2e–
SO32– + H2O
0,20
Appendix F
Solubility Product Constant Ksp Values for Feebly Soluble Electrolytes at 25 °С
Electrolyte
AgBr
AgCl
Ag2CrO4
AgI
Ag2S
Ag2SO4
BaCO3
BaCrO4
BaSO4
Electrolyte
Ksp
6·10
–13
1,8·10
–10
4·10–12
1,1·10
–16
6·10–50
Ksp
Fe(OH)3
3,7·10–40
FePO4
1,3·10–22
FeS
5·10–18
HgS
1,6·10–52
MgCO3
2,1·10–5
2·10
–5
Mg(OH)2
1,3·10–11
5·10
–9
MnS
2,5·10–10
1,6·10–10
PbBr2
9,1·10–6
–10
PbCl2
2·10–5
1,1·10
–39
Ba3(PO4)2
6·10
CaCO3
5·10–9
PbCrO4
1,8·10–14
PbCO3
7,5·10–14
CaC2O4
2·10
–9
PbI2
8,0·10–9
CaF2
4·10–11
PbS
2,5·10–27
PbSO4
1,6·10–8
SrCO3
1,1·10–10
CaSO4
Ca3(PO4)2
6,3·10
1·10
–5
–29
113
Appendix G
Dissociation Constants of Some Weak Electrolytes
Formule
К1
К2
Acids
–4
HNO2
4,0⋅10
HAlO2
4,0⋅10–13
H3BO3
5,8⋅10–10
1,8⋅10–13
–9
HOBr
2,1⋅10
H2CO3
4,45⋅10–7
4,5⋅10–11
H2SiO3
2,2⋅10
–10
1,6⋅10–12
H3AsO4
5,6⋅10–3
1,7⋅10–7
K3=2,9⋅10–12
H3AsO3
5,7⋅10–10
3,0⋅10–14
HAsO2
5,8⋅10–10
H2O2
2,6⋅10–12
H2SeO4
1⋅10–3
1,2⋅10–2
–3
H2SeO3
3,5⋅10
H2Se
1,7⋅10–4
–2
H2SO3
1,6·10
H2S
8,9·10–8
5,0⋅10–8
1,0⋅10–11
6,3⋅10–8
1,3⋅10–13
–8
HOCl
5,0⋅10
H3PO4
7,5⋅10–3
6,3⋅10–8
K3=1,3⋅10–12
H3PO3
1,0⋅10–2
3,0⋅10–7
H3PO2
9,0⋅10–2
HF
6,6⋅10–4
HCN
7,2⋅10–10
C6H5COOH
6,3⋅10–5
HCOOH
1,8⋅10–4
C2H5COOH
1,34⋅10–5
CH3COOH
1,75⋅10–5
C3H7COOH
1,54⋅10–5
CH2ClCOOH
1,4⋅10–3
H2C2O4
Al(OH)3
5,4⋅10–2
5,4⋅10–5
K3=1,4⋅10–9
1,3⋅10–4
Fe(OH)2
114
Formule
К2
К1
Acids
1,8⋅10–11 K3=1,4⋅10–
Fe(OH)3
12
5,0⋅10–3
2,5⋅10–3
3,4⋅10–7
Cd(OH)2
Mg(OH)2
Cu(OH)2
NH4OH
Pb(OH)2
Be(OH)2
Cr(OH)3
Zn(OH)2
1,8⋅10–5
9,6⋅10–4
3,0⋅10–8
5,0⋅10–11
K3=1,0⋅10–10
4⋅10–5
Appendix H
Instability Constants of Complex Ions
Complex ion
Instability
Constants
Instability
Constants
Complex ion
[Ag(NO2)2]–
1,8·10–3
[Fe(CN)6]3–
1,00·10–42
[Ag(CN)2]–
1,0·10–21
[HgCl4]2–
6,03·10–16
[Ag(NH3)2]+
5,9·10–8
[Hg(CN)4]2–
Ag(S2O3)2]
3–
[Cd(CN)4]2–
[Cd(NH3)4]
2+
1,00·10
–13
7,66·10–18
7,5·10
–8
[Hg(SCN)4]
2–
[HgI4]2–
[Ni(CN)4]
5,50·10–3
[Ni(NH3)6]2+
[Co(NH3)6]2+
4,07·10–5
[PbI4]2–
[Cu(CN)2]
1,00·10
–24
1,29·10–22
1,38·10–30
2–
[Co(CNS)4]2–
–
3,02·10–42
[Zn(CN)4]
1,00·10–22
9,77·10–9
9·10–5
2–
1,00·10–16
[Cu(CN)4]3–
5,13·10–31
[Zn(CNS)4]2–
5,00·10–2
[Cu(NH3)4]2+
9,3·10–13
[Zn(NH3)4]2+
2,00·10–9
[Fe(CN)6]
4–
1,00·10
–37
115
[Zn(OH)4]
2–
7,08·10–16