Download Glossary: Chemical bonds

E.Yu. Nevskaya
Chapters in General and Inorganic
Chemistry
Moscow
Peoples’ Friendship University of Russia
2013
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LECTURES
Chapters In Inorganic Chemistry
BASIC CONCEPTS OF CHEMISTRY
Objects of studying of chemistry are substances and their smallest particles –
molecules and atoms.
Molecule is the least particle of substance possessing its chemical properties.
Atom is the least particle of a chemical element possessing its chemical
properties. Atoms are components of molecules.
Chemical element is the kind of atoms characterized by certain set of
properties.
Mole is a unit of measurement of quantity of substances, containing such
amount of molecules, atoms or other structural units, as 12 g of an isotope 12С.
The number of structural units containing in 1 mole of substances, is known
as Avogadro’s number (NA): NA = 6,021023 mol-1.
The mass of 1 mole of a substance expressed in grams, is called a molar
mass of a substance (M, g x mol-1).
 = m / M = N / NA (mole).
Quantity of a substance:
General laws of Chemistry
1. Incorporated gas law:
2
PV P0 V0

T
T0 .
2. Avogadro’s Law: equal volumes of various gases under identical
conditions (temperature and pressure) contain the identical number of
molecules.
Consequence 1. Masses of two identical volumes of various gases under
identical conditions concern as their molar masses:
m1 M 1

D
m2 M2
(relative density of a gas)..
Consequence 2. Under normal conditions (Р0 = 101325 Pa and Т0 = 273 K)
one mole of any gas occupies the volume of 22,4 l (VМ – molar volume).
3. Mendeleyev-Klayperon’s equation:
PV 
m
RT
M
, where R – a universal gas constant (R = 8,314 J x mole-1 x K).
4. Partial pressure of gas is the part of a total pressure of a gas mixture
which is necessary on a share of the given gas. According to Dalton’s
law, partial pressure of gas in a mixture equals the pressure of gas as if it
occupied the total volume under the same temperature.
Рtot. = Р1 + Р2 + … + Рn.
5. The law of mass conservation. Sum of masses of the substances
entering a chemical reaction, equal total mass of the products.
6. The law of equivalents. Substances react with each other in the
quantities proportional to their equivalents:
n1 = n2 (n – number of equivalents).
3
The equivalent is such a quantity of substance which reacts with 1 mole of
hydrogen atoms or replaces them in chemical reactions (the quantity of substance
corresponding to a unit valency).
Equivalent mass (МE) is the mass of one equivalent of the substance,
expressed in grams:
МE = f x M (where f is the factor of equivalence).
Calculation of the factor of equivalence of different classes of inorganic
compounds:
For simple substances and elements in chemical compounds
f = 1 / V (where V is the valency of an element).
For acids and the bases
f = 1 / m (where m is the basidity of an acid or acidity of a base).
For oxides and salts
f = 1 / n x V (where n is the number of metallic atoms in the compound,
and V is the valency of the metal).
Number of equivalents: n = m / МE (for any substance); n = V / VE (for
gaseous substances), VE is the equivalent volume of the gas (the volume occupied
by one equivalent of a gas under normal conditions). For example, under normal
conditions the equivalent volume of hydrogen (МE = 1 g x mole-1) equals 11,2
liters, and equivalent volume of oxygen (МE = 8 g x mole-1) equals 5,6 litres.
QUESTIONS AND PROBLEMS
1. A metal hydride contains 2.02 g of hydrogen and 13.88 g of metal. Calculate
equivalent mass of the metal.
2. While 53.96 g of a metal is oxidized 101.96 g of an oxide is formed. Calculate
equivalent mass of the metal.
3. 4.80 g of Ca and 7.85 g of Zn replace the same amount of hydrogen from an
acid. Calculate equivalent mass of zinc if the equivalent mass of calcium
equals 20.0 gmole-1.
4. Calculate equivalent masses of a metal and sulfur if 4.86 g of the metal form
5.22 g of oxide and 5.58 g of sulfide.
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5. While 0.595 g of an unknown substance reacts with 0.275 g of hydrogen
chloride, 0.440 g of a salt is formed. Calculate equivalent masses of the
substance and the salt.
6. The volume of 2800 ml of hydrogen measured under normal conditions can
reduce 11.75 g of a metal oxide. Calculate equivalent masses of the metal and
its oxide.
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2. GENERAL CLASSES OF INORGANIC SUBSTANCES
OXIDES
Oxides are binary compounds containing oxygen atoms in the –2 oxidation
state.
OXIDES
Salt-forming
Not salt-forming
(indifferent)
Basic
Amphoteric
Metallic
Some
oxides in
metallic
low
oxides in the
oxidation
oxidation
states
states
+2, +3 and +4
(+1, +2)
BeO, ZnO,
Na2O,
PbO, SnO,
Cu2O,
Al2O3,
FeO,
Cr2O3,
MgO
MnO2
Acidic
Non-metallic
Nonoxides in
metallic
oxidation
oxides in
states +3 and
oxidation
above, as far
states +1
as metallic
and+2
oxides in
СО, SiO,
oxidation
N2O, NO
states +4 and
above
В2О3, СО2,
P2O5, PbO2,
CrO3
Increase in oxidation state of a non-metal changes properties of its oxide
from indifferent to acidic:
N2O
NO
N2O3
Indifferent oxides
NO2
N2O5
Acidic oxides
Increase in oxidation state of a metal changes properties of its oxide from
basic via amphoteric to acidic:
MnO
Mn2O3
Basic oxides
MnO2
Amphoteric
MnO3
Mn2O7
Acidic oxides
oxide
Chemical properties of oxides:
1. Basic and acidic oxides are dissolved in water (react with it) in case if
soluble bases and acids are formed::
6
Na2O + H2O = 2NaOH
Only basic oxides of alkaline and
CaO + H2O = Ca(OH)2 alkaline-earth metals
FeO + H2O 
SO3 + H2O = H2SO4 Only acidic oxides which form
P2O5 + H2O (_isi.) = 2НРО3 soluble acids
P2O5 + 3H2O (гор.) = 2Н3РО4
2NO2 + H2O = HNO2 + HNO3
4NO2 + O2 + 2H2O = 4HNO3
СrO3 + H2O = H2CrO4
SiO2 + H2O 
2. Acidic oxides react with basic oxides:
Na2O + СО2 = Na2СО3 If the acidic oxide is
CaO + SO2 = CaSO3 in
the
gaseous
phase, the reaction
takes place under
room temperature
melting
Na2O + SiO2  Na2SiO3 If the acidic oxide is
solid, the reaction
melting
FeO + PbO2  FePbO3 takes
place
undervhigh
temperatures
(melting)
3. Amphoteric oxides can react with both acidic and basic oxides:
melting
ZnO + Na2O  Na2ZnO2
ZnO + CO2 = ZnCO3
melting
Al2O3 + K2O  2KalO2
Al2O3+3SiO2  Al2(SiO3)3
melting
BASES
Bases are substances, which form hydroxyl-anions (OH-) at dissociation.
Number of OH-groups of the base defines its acidity.
Chemical properties of bases:
Alkalis:
1) react with acidic oxides and acids:
2КОН + SO3 = K2SO4 + H2O
Ca(OH)2 + 2HCl = CaCl2 + 2H2O
Reaction of
neutralization
7
2) react with salts (in case if a precipitate is formed):
2NaOH + FeCl2 = Fe(OH)2 + 2NaCl
Ba(OH)2 + Na2SO4 = BaSO4 + 2NaOH
NaOH + BaCl2 
3) react with amphoteric oxides and amphoteric bases:
melting
2NaOH + BeO  Na2BeO2 + H2O
3NaOH + Cr(OH)3  Na3[Cr(OH)6]
solution
4) dissociate in aqueous solutions and change the colour of acid-base
indicators:
NaOH = Na+ + OH–
(methylorange turns yellow, litmus turns blue, and phenolphthalein turns pink)
Insoluble bases:
1) react with acidic oxides and acids (reaction of neutralization):
Fe(OH)3 + 3HNO3 = Fe(NO3)3 + 3H2O
t

2) decompose at heating: Cu(OH)2 
CuO + H2O
Amphoteric bases (bases which correspond to amphoteric oxides):
 react with acidic oxides and acids (reaction of neutralization):
Zn(OH)2 + CO2 = ZnCO3 + H2O
Al(OH)3 + 3HCl = AlCl3 + 3H2O
 react with alkalis:
Sn(OH)2 + 2NaOH = Na2[Sn(OH)4]
3) decompose at heating:
t

2Al(OH)3 
Al2O3 + 3H2O
ACIDS
Acids are substances, which form hydrogen kations (H+) at dissociation.
Basicity of an acid (number of H-atoms) defines the possibility of full or
incomplete _isinfection_e of an acid in reactions with bases:
HCl + NaOH = NaCl + H2O
(for monobasic acids only one reaction of _isinfection_e is possible).
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For the multibasic acids full and incomplete _isinfection_e is possible:
H2СО3 + 2КОН = К2СО3 + 2Н2О (full _isinfection_e)
H2СО3 + КОН = КНСО3 + Н2О (incomplete neutralization)
Chemical properties of acids:
1.
Change in colour of indicators (methylorange and lithmus turn red).
2.
Interaction with active metals: Са + 2HCl = CaCl2 + H2
3.
Interaction with basic and amphoteric oxides:
СuO + 2HNO3 = Cu(NO3)2 + H2O
Al2O3 + 3H2SO4 = Al2(SO4)3 + 3H2O
4.
Interaction with bases (reaction of neutralization):
Cu(OH)2 + 2HCl = CuCl2 + 2H2O
5.
Interaction with salts (reactions of ionic exchange):
H2SO4 + BaCl2 = 2HCl + BaSO4 (precipitate is formed)
2HNO3 + СаСО3 = Ca(NO3)2 + H2O + CO2 (gas is evolved)
HCl + KNO2 = KCl + HNO2 (weak acid is formed)
SALTS
Salts are substances, which form metallic kations or ammonium (NH4+) and acidic
anions at dissociation.
Salts can be considered as products of replacement of hydrogen atoms in an acid
molecule by metals or hydroxyl-groups in a base by an acidic anion.
Chemical properties of salts:
1.
Interaction with metals (a more active metal replaces a less active one):
FeCl2 + Zn = Fe + ZnCl2.
2.
Interaction with non-metals (a more active non-metal replaces a less active
one):
2NaBr + Сl2 = Br2 + 2NaCl.
3.
Interaction with alkalis (reactions of ionic exchange):
MgCl2 + 2NaOH (exsess) = Mg(OH)2 + 2NaCl.
MgCl2 + NaOH (lack) = MgOHCl + NaCl.
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 Interaction with acids (reactions of ionic exchange):
СаCl2 + H2SO4 (lack) = CaSO4 + 2HCl.
CaSO4 + H2SO4 (exsess) = Ca(HSO4)2
5.
Interaction between two salts (reactions of ionic exchange):
CuCl2 + 2AgNO3 = 2AgCl + Cu(NO3)2.
 CHEMICAL KINETICS
One of the basic concepts of chemical kinetics is the concept of a rate of a
chemical reaction.
Rate of a chemical reaction is denoted as number of elementary acts of a
reaction which results transformation of reactants into reaction products, in a unit
time in a unit volume.
In practice, rates of reactions can be measured as a change in concentrations
of substances participating in it for a certain time interval:
v
c

Out of two chemical reactions, that one is of the greatest rate, in which
under identical time more quantity of a substance is formed.
Law of mass action. Collision of molecules should be a necessary condition
for _isinfectio of chemical interaction between molecules. Collision occurs the
more often, than more molecules contains in the given volume, i.e. rate of a
chemical reaction depends on concentrations of reacting substances.
aA  bB  mM  nN
v  k A B 


where k is a rate constant of the chemical reaction, numerically it equals rate of a
reaction at unit concentrations of reacting substances.
,  - simple numbers, usually not more than 3. For simple reactions they
correspond to stoitiometric coefficients of the reaction.
The rate of the reaction does not depend on concentrations of firm
substances, but only on their surface area.
10
CaO + CO2
CaCO3
v  k CO2 
v  k 2
The equation of a reaction frequently does not reflect its mechanism. For
2HI + H2O2  2H2O + I2
example, the reaction
really proceeds in two stages:
1. HI + H2O2  HOI + H2O (slowly)
2. HOI + HI  I2 + H2O (quickly)
Kinetics of the overall reaction it is described by the first (slow) stage.
Expression of speed of this reaction registers as
v  k HI  H 2 O2 
,
instead of
v  k HI  H 2 O2 
2
Temperature dependence of rates of chemical reactions.
Rate of a
chemical reaction depends on number of effective collisions. Effective collision
occurs only between active molecules. Increase in temperature increases number of
active molecules, providing them with necessary activation energy, and the rate of
the reaction increases.
Activation energy is that additional energy which it is necessary to transfer
to system to start chemical reaction.
Vant Hoff’s rule. At increase in temperature on 10  speed of reaction
increases in 2 – 4 times.
T2 T1
v 2  v1   10 ,
v2
is the rate of a reaction at temperature T2,
v1
is the rate of a reaction at temperature T1,
where
 is the temperature coefficient of the reaction which defines change
of the rate of the reaction at temperature change on 10.
As
v
c

so
T2 T1
 1   2   10 , where
time.
11

is the reactional
Chemical equilibrium
If a chemical reaction can proceed only in one direction it is called as
irreversible. The reactions proceeding simultaneously in two directions, are
reversible. Eventually rate of a direct reaction (v  ) decreases, and rate of a back
reaction (v  ) increases until they become equal. So, a chemical equilibrium is
established in the system. The condition of a chemical equilibrium:
v

= v .
In the equilibrium state, reversible reactions are described by an equilibrium
constant K:
aA + bB
K
CCc C Dd
C Aa C Bb


cC + dD
where
CA, CB, CC, CD are concentrations of gaseous or dissolved substances.
Chemical equilibrium is a dynamic one, so it can be shifted according to le
Chateleu’s principle (principle of counteraction): if an equilibrium system is
affected by any factor (change in concentrations, pressure or temperature), the
equilibrium will be shifted in the direction which weakens the external influence.
The increase in temperature shifts the equilibrium towards an endothermic
reaction (the system absorbs heat and increases its internal energy, H>0), and
decrease in temperature shifts the equilibrium towards an exothermic reaction (the
system evolves heat and decreases its internal energy, H <0).
The increase in pressure causes shifting of the equilibrium towards less
quantity of gaseous substances (as pressure is affected only by gaseous
substances), and decrease in pressure shifts the equilibrium towards more quantity
of gaseous substances. In case if quantities of gaseous substances among reactants
and products are same, change in pressure does not cause shifting of chemical
equilibrium.
The increase in concentration of one of reactants causes
shifting of
equilibrium towards formation of products of reaction, and increase in
12
concentration of one of reactional products shifts the equilibrium towards
reactants.
1. Find the value of the rate constant for the reaction A + B  AB if at
concentration of substances A and B equal to 0.05 M and 0.01 M respectively,
the rate of the chemical reaction is 5 x 10  5 M/min.
2. How many times will the rate of the reaction 2A + B  A2B change if the
concentration of A is doubled, and that of B is halved ?
3. What is the temperature coefficient of the reaction rate if the rate grows 15.6
times when the temperature is increased by 30 Kelvins?
4. Find the equilibrium constant of the reaction N2O4


2NO2
if the initial
concentration of the N2O4 was 0.08 M, and by the moment when equilibrium
was established 50% of N2O4 was dissociated.
5. In which direction will the following equilibria shift :
a) 2CO + O2
b) N2 + O2




2CO2
2NO
H 0  566kJ
H 0  180 kJ
( 1 ) when the temperature is lowered; ( 2 ) when the pressure is increased; ( 3 )
when the concentration of oxygen is increased ?
13
4. STRUCTURE OF ATOMS
Glossary: Atoms, elements, and ion
see also
http://antoine.frostburg.edu/chem/senese/101/atoms/glossary.shtml
alpha particle. (42He)
A particle that is commonly ejected from radioactive nuclei, consisting of
two protons and two neutrons. Alpha particles are helium nuclei. Alpha
particles have a mass of 6.644 655 98×10-27kg or 4.001 506 1747 atomic
mass units. [1998 CODATA values]
alpha ray. ( -ray) alpha radiation.
A stream of alpha particles. Alpha rays rapidly dissipate their energy as they
pass through materials, and are far less penetrating than beta
particles and gamma rays.
Anion. Compare with cation.
An anion is a negatively charged ion. Nonmetals typically form anions.
Anode. Compare with cathode.
The electrode at which oxidation occurs in a cell. Anions migrate to the
anode.
Atomic mass unit. (amu,u) amu; _isinf.
A unit of mass equal to 1/12 the mass of a carbon-12 nucleus, which is 1.660
538 73 × 10-27 kg ± 0.000 000 13 × 10-27 kg [1998 CODATA values].
Abbreviated as amu or u. Sometimes called the _isinf, after John Dalton,
architect of the first modern atomic theory.
Atomic nucleus. Nucleus; nuclei; atomic nuclei.
A tiny, incredibly dense positively charged mass at the heart of the atom.
The nucleus is composed of protons and neutrons (and other particles). It
contains almost all of the mass of the atom but occupies only a tiny fraction
of the atom’s volume.
14
Atomic number. (Z)
The number of protons in an atomic nucleus.
Atomic theory.
An explanation of chemical properties and processes that assumes that tiny
particles called atoms are the ultimate building blocks of matter.
Atomic weight. Atomic mass.
The average mass of an atom of an element, usually expressed in atomic
mass units. The terms mass and weight are used interchangeably in this case.
The atomic weight given on the periodic table is a weighted average
of isotopic masses found in a typical terrestrial sample of the element.
Atom. Compare with molecule and ion.
An atom is the smallest particle of an element that retains the chemical
properties of the element. Atoms are electrically neutral, with a positively
charged nucleus that binds one or more electrons in motion around it.
Beta particle. (ß-)
An electron emitted by an unstable nucleus, when a neutron decays into
a proton and an electron. In some cases, beta radiation consists of positrons
(“antielectrons” which are identical to electrons but carry a +1 charge.”)
Note that beta particles are created in nuclear decay; they do not exist as
independent particles within the nucleus.
Brownian motion. Brownian movement.
Small particles suspended in liquid move spontaneously in a random
fashion. The motion is caused by unbalanced impacts of molecules on the
particle. Brownian motion provided strong circumstantial evidence for the
existence of molecules.
Cathode ray.
A negatively charged beam that emanates from the cathode of a discharge
tube. Cathode rays are streams of electrons.
Cathode. Compare with anode.
The electrode at which reduction occurs.
15
Cation. Compare with anion.
A cation is a positively charged ion. Metals typically form cations.
Chemical change. Reaction; chemical reaction. Compare with physical change.
A chemical change is a dissociation, recombination, or rearrangement of
atoms.
compound Compare with element and mixture.
A compound is a material formed from elements chemically combined in
definite proportions by mass. For example, water is formed from chemically
bound hydrogen and oxygen. Any pure water sample contains 2 g of
hydrogen for every 16 g of oxygen.
Deuterium. (D, 2H)
An isotope of hydrogen that contains one neutron and one proton in its
nucleus.
Electric charge. Charge.
A property used to explain attractions and repulsions between certain
objects. Two types of charge are possible: negative and positive. Objects
with different charge attract; objects with the same charge repel each other.
Electron. (e-) Compare with proton and neutron.
A fundamental _isinfectio of matter, having a negative charge of 1.602 176
462 × 10-19 coulombs ± 0.000 000 063 × 10-19 coulombs and a mass of 9.109
381 88 × 10-31 kg ± 0.000 000 72 × 10-31 kg [1998 CODATA values].
element Compare with compound and mixture.
An element is a substance composed of atoms with identical atomic
number . The older definition of element (an element is a pure substance
that can’t be decomposed chemically) was made obsolete by the discovery
of isotopes.
Element symbol.
An international abbreviation for element names, usually consisting of the
first one or two distinctive letters in element name. Some symbols are
abbreviations for ancient names.
16
Group.
 A substructure that imparts characteristic chemical behaviors to a
molecule, for example, acarboxylic acid group. (also: functional
group). 2. A vertical column on the periodic table, for example,
the halogens. Elements that belong to the same group usually show
chemical similarities, although the element at the top of the group
is usually atypical.
Heavy water. (D2O)
Water that contains 2H, rather than 1H. Heavy water is about 11% denser
than ordinary water.
Ion.
An atom or molecule that has acquired a charge by either gaining or losing
electrons. An atom or molecule with missing electrons has a net positive
charge and is called a cation; one with extra electrons has a net negative
charge and is called an anion.
Isotope. Isotopic; isotopy. Compare with isomer, allotrope, isobar, and isotone.
Atoms or ions of an element with different numbers of neutrons in
their atomic nucleus
. Isotopes have the same atomic number but different mass number
. Isotopes have very similar chemical properties but sometimes differ greatly
in nuclear stability.
Isotopic abundance. Compare with natural abundance. The fraction of atoms of a
given isotope in a sample of an element.
Isotopic mass. Isotopic masses. The mass of a single atom of a given isotope,
usually given in Daltons.
IUPAC
International Union of Pure and Applied Chemistry, an organization which
sets international standards for chemical nomenclature , atomic weights ,
and the names of newly discovered elements.
Law of conservation of mass.
17
There is no change in total mass during a chemical change. The
demonstration of conservation of mass by Antoine Lavoisier in the late 18 th
century was a milestone in the development of modern chemistry.
Law of definite proportions.
When two pure substances react to form a compound, they do so in a
definite proportion by mass. For example, when water is formed from the
reaction between hydrogen and oxygen, the ‘definite proportion’ is 1 g of H
for every 8 g of O.
law of multiple proportions.
When one element can combine with another to form more than one
compound, the mass ratios of the elements in the compounds are simple
whole-number ratios of each other. For example, in CO and in CO2, the
oxygen-to-carbon ratios are 16:12 and 32:12, respectively. Note that the
second ratio is exactly twice the first, because there are exactly twice as
many oxygens in CO2 per carbon as there are in CO.
mass number. (M,A) Compare with atomic number and atomic weight.
The total number of protons and neutrons in an atom or ion. In nuclide
symbols the mass number is given as a leading superscript. In isotope names
(e. g. carbon-14, sodium-23) the mass number is the number following the
element name.
metal. Metallic. Compare with nonmetal and metalloid.
A metal is a substance that conducts heat and electricity, is shiny and reflects
many colors of light, and can be hammered into sheets or drawn into wire.
Metals lose electrons easily to formcations. About 80% of the known
chemical elements are metals.
Natural abundance. Compare with isotopic abundance.
The average fraction of atoms of a given isotope of an element on Earth.
Neutral.
18
 having no net electrical charge. Atoms are electrically neutral; ions
are not. 2. A solution containing equal concentrations of H+ and
OH-.
Neutron. (n, 10n) Compare with proton and electron.
An elementary particle found the atomic nucleus of all stable atoms except
the hydrogen-1 atom. Neutrons have no charge and have a mass of
1.008665 daltons.
Nonmetal. (metal,metalloid) non-metal.
A nonmetal is a substance that conducts heat and electricity poorly, is brittle
or waxy or gaseous, and cannot be hammered into sheets or drawn into wire.
Nonmetals gain electrons easily to form anions. About 20% of the known
chemical elements are nonmetals.
Nuclear binding energy.
Energy needed to break an atomic nucleus into separate protons and
neutrons.
Nucleon. Compare with proton, neutron and atomic nucleus.
A proton or a neutron in the atomic nucleus.
Nuclide symbol. Compare with atomic nucleus, nuclide and element symbol.
A symbol for an nuclide that contains the mass number as a leading
superscript and theatomic number as a leading subscript. For ions, the ionic
charge is given as a trailing superscript. For example, the nuclide symbol for
the most common form of the chloride ion is3517Cl-, where 35 is the mass
number, 17 is the atomic number, and the charge on the ion is -1. The atomic
number is sometimes omitted from nuclide symbols.
Nuclide. Compare with atomic nucleus and nuclide symbol.
An atom or ion with a specified mass number and atomic number. For
example, uranium-235 and carbon-14 are nuclides.
Periodic table.
An arrangement of the elements according to increasing atomic number that
shows relationships between element properties.
19
Period.
Rows in the periodic table are called periods. For example, all of the
elements in the second row are referred to as ‘second period elements’. All
elements currently known fall in the first seven periods.
Polymorph. Polymorphism; polymorphic. Compare with isotope and allotrope.
Solid substances that occur in several distinct forms. Polymorphs have
different chemical and physical properties. allotropes are polymorphs of
elements.
Proton. (p+) Compare with electron and neutron.
An elementary particle found the atomic nucleus with a positive charge
equal and opposite that of the electron. Protons have a mass of
1.007276 daltons.
Radioactivity. Radiation; radioactive.
Spontaneous emission of particles or high-energy electromagnetic
radiation from the nuclei of unstable atoms. “Radiation” refers to the
emissions, and “radioactive source” refers to the source of the radiation.
Stoichiometry.
1. Ratios of atoms in a compound. 2. Ratios of moles of compounds in a
reaction. 3. A branch of chemistry that quantitatively relates amounts of
elements and compounds involved in chemical reactions, based on the law
of conservation of mass and the law of definite proportions.
x-ray spectrum. X-ray spectra.
Atom is a complicated particle consisting of positively charged atomic
nucleus and electronic shells where negatively charged electrons are located.
The positive charge of a nucleus equals the sum of negative charges of
electrons.
The nucleus itself consists of positively charged protons and uncharged
neutrons. Protons, neutrons and электроны carry the name «elementary particles».
20
Using relative units, one can state that electrons are –1 charged, and protons
are +1 charged.
The mass of an atom expressed in nuclear mass units, is called a relative
nuclear mass or mass number of an atom, Мr. It equals sum of masses of all
elementary particles of the atom. As mass numbers of protons and neutrons are
equal (1 a.m.u.), and masses of electrons are neglectably small (approximately
2,000 times less than corresponding masses of protons and neutrons) mass number
of an atom equals sum of number of protons and neutrons.
Symbols of chemical elements are usually represented in the way:
a X c
b
,
where X
is the symbol of the element;
a is the mass number of the element;
b
is an element serial number in the Periodic Table (equals number of
protons in the atom);
c is the charge of the ion.
Natural chemical elements exist in the form of a mixture of isotopes.
Isotopes are atoms of the same chemical element with identical number of
protons, but different mass numbers (number of neutrons). For example, natural
35
chlorine exists in the form of two isotopes: 17 Cl (17 protons and 18 neutrons), and
37
17
Cl
(17 protons and 20 neutrons). Relative masses of elements in periodic system
of elements, are average masses of natural isotopes.
Structure of electronic shells
Electrons possess so called wawe dualism (simultaneously properties of a
particle and a wave).
In this connection, for the description of properties электрона enter special
function which name state function of an electron or wave function, . It is
entered in such a manner that the square of its module is proportional to probability
21
to find out a particle (electron) in the given place at the appointed time (probability
density). See also http://en.wikipedia.org/wiki/Atomic_nucleus.
Wave function of an electron is called «orbital». It characterizes energy and
the
form
of
spatial
distribution
of
an
electronic
cloud.
See
also
http://en.wikipedia.org/wiki/Atomic_orbital.
Quantitative parities in the theory of a structure of atom are defined by the
Shrodinger’s wave equation:
h 2   2  2  2 
 2  2  2  2   U  E
8 m  x
y
z 
where U is potential energy of the electron;
Е is full energy of the electron;
m is the mass of the electron;
x, y, z are _isinfe co-ordinates of the electron;
 is the wave function;
Consequence of the decision of Shrodinger’s equation is the set of four
quantum numbers which _isinfection behaviour of the electron in the atom.
N, principle quantum number, it defines the general stock of energy of the
electron, i.e.energetic shell.
N = 1,2,3 …
l,
azimutal quantum number, defines the form of electronic orbital
(subshell).
L = 0,1,2 … (n-1).
If l = 0 the orbital is called s-orbital (spheric movement of the electron). At l
= 1 we have p-orbital (double-lobe movement form). Movement forms of d – and f
– orbitals (l = 2 and 3 accordingly) have even more complicated character.
The number of orbitals at an energetic level coincides with its number. So,
for the first shell (n = 1) there is only one subshell (l = 0), that is 1s-orbital.
Similarly for n = 2 (the second shelll) two subshells exist (l = 0, 1) or 2s, 2porbitals; for the third shell (n = 3, l = 0, 1, 2) 3s, 3p, and 3d-orbitali exist etc.
22
ml is the magnetic quantum number, it _isinfection_ projection of the
magnetic moment of the electron on an external magnetic field, that is defines the
_isinfe orientation of the electronic orbital. Its values are defined by azimutal
quantum number:
ml = l; (l-1); (l-2) … 0
For the orbital quantum number l = 0, magnetic quantum number has one
possible value (ml = 0), that is only one way of orientation of s-orbital in space is
possible. Similarly we receive, that for p-orbitals (l = 1, ml =-1, 0, +1) there are
three possible ways of orientation (along co-ordinate axes), for d-orbitals there are
five possible ways of orientation (l = 2, ml =-2,-1, 0, +1, +2).
To specify the concept electronic orbital, we can state that it represents a set
of positions of electrons in the atom. Conditionally nuclear orbitals are designated
in the form of cages (energetical cells):
1
s
2
s
3
s
2
p
3
3
p
d
ms , the spin quantum number, defines the moment of spinning of the
electron. As there are only two ways of spinning (clockwise and anticlockwise),
the magnetic quantum number can accept two values:
ms =  l.
Conditionally, electrons having different values of spin quantum number, are
designated by opposite directed arrows: .
Placing of electrons in atoms
If the atom is in the ground state (does not possess superfluous energy) its
electrons occupy the lowest energetic orbitals. Energy of an electron in
multielectronic atoms depends not only on its attraction to a nucleus, but also from
23
repultion from other electrons. Mutual influence leads to that energy of the electron
depends not only on principle, but also on azimutal quantum number.
Klechkovsky’s rules
1. The increase in energy of electronic subshells goes as increase in the sum
of the principle and azimutal quantum numbers (n+l).
2. In case of equality of the sum (n+l) the increase in energy of subshell goes
as increase in the principle quantum number.
Filling of orbitals by electrons occurs according to the arrows in a following
order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p etc.
Pauli’s exclusive principle
In one atom there can not be two electrons with an identical set of quantum
numbers. Because everyone electronic orbital is _isinfection_ by a set of three
quantum numbers (principle, azimutal and magnetic), electrons of the same orbital
can differ only by the value of spin quantum number (ms = ). A consequence
of a Pauli’s principle of is that one orbital can contain not more than two electrons.
In connection with the aforesaid at the first energetic shell not more than two
electrons can exist:

1s
Or 1s2;
And the second energetic shell can maximally contain 8 electrons:

  
2s
2p
Or 2s22p6 etc.
The maximum number of electrons at any shell N = 2n2, where n is the principle
quantum number.
Hund’s rule
In a subshell electrons fill orbitals so that the total spin quantum number
becomes maximum (orbitals of a subshell are first filled by one electron each and
only after all orbitals are filled, pairing of electrons takes place).
For example, four electrons on a p-subshell can be arranged in two different
ways:
24




Or

(ms) = + 1
(ms) = 0
As in the first case the total spin number is more,the first electronic structure
is realized.
Electronic formulae of atoms and ions
The number of electrons in atom is defined by an element serial number in
periodic table. Using the above rules and principles, for a sodium atom of (11
electrons) the following electronic formula is received:
11Na:

1s


1s
2s
1s22s22p63s1





2s
2p
3s
The electronic formula of a Ti atom:
2 2
6 2
6 2
2
22Ti: 1s 2s 2p 3s 3p 4s 3d

2p




3s
3p
 

 
4s
3d
If one electron do not suffice to full or half-full d-subshell (d10 or d5configurations), it is transmitted from the next s-subshell. As a result, the electronic
formula of Cr atom looks like
24Cr:
1s22s22p63s23p64s13d5, instead of
24Cr:
1s22s22p63s23p64s23d4, and atom of copper – 29Cu: 1s22s22p63s23p64s13d10, instead
of 29Cu: 1s22s22p63s23p64s23d9.
The number of electrons in a negatively charged ion – anion – exceeds
number of electrons in a neutral atom: 16S2 – 1s22s22p63s23p6 (18 electrons).
At formation of a positively charged ion – cation – electrons are detouched
first from the outer shell: 24Cr3 +: 1s22s22p63s23p64s03d3 (21 electrons).
Electrons in atom can be divided into two types: internal and external
(valence). Internal electrons occupy completely filled subshells, have low values of
energy and do not participate in chemical transformations of elements.
Valence electrons are all electrons of the outer electronic shell as far as
electrons of not-filled subshells.
Valence electrons take part in formation of chemical bonds. Unpaired
electrons have special activity. The number of unpaired electrons defines the
valence state of a chemical element.
25
In case if not filled orbitals are available on the highest energetic level,
unpairing of valence electrons may occur, and the valence state of the atom
increases (formation of an excited state of the atom takes place).
For example, valence electrons of sulfur are 3s23p4:

16S

3s


3p
3d
In the ground state the S atom has 2 unpaired electrons, so its valence state is
II.
At an expense of some energy one of paired electrons of sulfur can be
translated on an empty d-orbital, that corresponds to the first excited state of the
atom:
16S*


3s



3p
3d
In this case, the S atom has four unpaired electrons, and its valence state
equals IV.
One of 3s- paired electrons can also be moved to a free 3d-orbital:
16S
** 

3s




3p
3d
In such a condition, the S atom of sulphur possesses the valence state VI.
If the outer electronic shell has no free orbitals or subshells, unpairing of
electrons is not possible and the atom can have onlt one valence state (example –
O-atom):
8О

2s



2p
See also http://en.wikipedia.org/wiki/Electron_configuration.
26

THE PERIODIC LAW AND PERIODIC TABLE OF
CHEMICAL ELEMENTS
Glossary: The periodic table
actinide.
Elements 89-102 are called actinides. Electrons added during the Aufbau
construction of actinide atoms go into the 5f subshell. Actinides are unstable
and undergo radioactive decay. The most common actinides on Earth are
uranium and thorium.
Alkali metal. (alkaline earth metal) alkali metal element.
The Group 1 elements, lithium (Li), sodium (Na), potassium (K), rubidium
(Rb), cesium (Cs), and francium (Fr) react with cold water for form strongly
alkaline hydroxide solutions, and are referred to as “alkali metals”.
Hydrogen is not considered an alkali metal, despite its position on some
periodic tables.
Alkaline earth.
An oxide of an alkaline earth metal, which produces an alkaline solution in
reaction with water.
Alkaline earth metal. (alkali metal)
The Group 2 elements, beryllium (Be), magnesium (Mg), calcium (Ca),
strontium (Sr), barium (Ba), and radium (Ra) form alkaline oxides and
hydroxides and are called “alkaline earth metals”.
Amphoteric. Ampholyte.
A substance that can act as either an acid or a base in a reaction. For
example, aluminum hydroxide can neutralize mineral acids ( Al(OH)3 + 3
HCl = AlCl3 + 3 H2O ) or strong bases ( Al(OH)3 + 3 NaOH = Na3AlO3 + 3
H2O).
atomic radius. Metallic radius; covalent radius; atomic radii. Compare with ionic
radius.
27
One half the distance between nuclei of atoms of the same element, when
the atoms are bound by a single covalent bond or are in a metallic crystal.
The radius of atoms obtained from covalent bond lengths is called the
covalent radius; the radius from interatomic distances in metallic crystals is
called the metallic radius.
Block.
A region of the periodic table that corresponds to the type of subshell (s, p,
d, or f) being filled during the Aufbau construction of electron
configurations.
Congener.

Elements belonging to the same group on the periodic table. For
example, sodium and potassium are congeners. 2. Compounds
produced by identical synthesis reactions and procedures.
First ionization energy. (IE,IP) first ionization potential. Compare with second
ionization energy,adiabatic ionization energy, vertical ionization
energy, electronegativity, and electron affinity.
The energy needed to remove an electron from an isolated, neutral atom.
Group.

A substructure that imparts characteristic chemical behaviors to
a molecule, for example, acarboxylic acid group. (also: functional
group). 2. A vertical column on the periodic table, for example,
the halogens. Elements that belong to the same group usually show
chemical similarities, although the element at the top of the group
is usually atypical.
Halogen. Group VIIA; group 18.
An element of group VIIA (a. k. a. Group 18). The name means “salt
former”; halogens react with metals to form binary ionic compounds.
Fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At) are
known at this time.
Ionic radius. Compare with atomic radius.
28
The radii of anions and cations in crystalline ionic compounds, as
determined by consistently partitioning the center-to-center distance of ions
in those compounds.
Ionization energy. (IE,IP) ionization potential. Compare with adiabatic ionization
energy, vertical ionization energy, electronegativity , and electron affinity.
The energy needed to remove an electron from a gaseous atom or ion.
Lanthanide contraction.
An effect that causes sixth period elements with filled 4f subshells to be
smaller than otherwise expected. The intervention of the lanthanides
increases the effective nuclear charge, which offsets the size increase
expected from filling the n=6 valence shell. As a consequence, sixth period
transition metals are about the same size as their fifth period counterparts.
Lanthanide. Compare with actinide and inner transition metals.
Elements 57-70 are called lanthanides. Electrons added during the Aufbau
construction of lanthanide atoms go into the 4f subshell.
Main group elements.
Elements of the s and p blocks.
Metal. Metallic. Compare with nonmetal and metalloid.
A metal is a substance that conducts heat and electricity, is shiny and reflects
many colors of light, and can be hammered into sheets or drawn into wire.
Metals lose electrons easily to formcations. About 80% of the known
chemical elements are metals.
Nonmetal. (metal, metalloid) non-metal.
A nonmetal is a substance that conducts heat and electricity poorly, is brittle
or waxy or gaseous, and cannot be hammered into sheets or drawn into wire.
Nonmetals gain electrons easily to form anions. About 20% of the known
chemical elements are nonmetals.
Periodic law.
29
The periodic law states that physical and chemical properties of the elements
recur in a regular way when the elements are arranged in order of increasing
atomic number.
Periodic table.
An arrangement of the elements according to increasing atomic number that
shows relationships between element properties.
Periodic trend.
A regular variation in element properties with increasing atomic number that
is ultimately due to regular variations in atomic structure.
Period.
Rows in the periodic table are called periods. For example, all of the
elements in the second row are referred to as ‘second period elements’. All
elements currently known fall in the first seven periods.
Second ionization energy. (IE,IP) second ionization potential. Compare with first
ionization energy, adiabatic ionization energy, vertical ionization
energy, electronegativity, and electron affinity.
The energy needed to remove an electron from an isolated +1 ion. The third
ionization energy would be the energy required to remove an electron from
an isolated +2 ion, and so on.
Transition metal. Transition element; outer transition element.
An element with an incomplete d subshell. Elements which have
common cations
with incomplete d subshells are also considered transition metals. Elements
with incomplete f subshells are sometimes called “inner transition elements”.
In 1869 the Russian chemist Dmitry Mendeleyev has shown that properties
of simple substances, and also forms and properties of chemical compounds of
elements are in periodic dependence on nuclear scales of elements. As
expression of this periodic law, the table which reflects the law was served. See
also http://en.wikipedia.org/wiki/Atom#Valence_and_bonding_behavior
30
In 1914 the English scientist G.Mozli has shown, that the charge of a
nucleus of an atom is numerically equal to an element serial number in the periodic
table, so properties of elements and their compounds are in periodic dependence
on a nucleus charge of the atom.
The periodic table of elements reflects electronic structures of atoms. Each
period (a horizontal series of periodic table) begins by an element in which
electrons start occupying a new electronic shell with the principle quantum number
which equals the number of the period).
Groups (vertical columns) contain elements with identical number of
valence electrons which equals the group number. Groups A contain s-elements
(valence electrons occupy s-subshells). In case if valence electrons are on s – and
p-subshells, they carry the name of p-elements. Elements with not fulfilled d – or
f-subshells, are known as d – and f-elements. They occupy groups B of the
periodic table.
Change of properties of chemical elements in the periods and groups of
the periodic table
Chemical properties of elements are illustrated by interactions of their
atoms.
Properties of chemical elements can be divided into metallic (reducing, i.e.
properties to lose electrons) and nonmetallic (_isinfect, i.e. properties to gain
electrons).
Properties of chemical elements depend on strengths of attraction of valence
electrons to a positively charged nucleui of atoms and are defined by following
characteristics:
Ionization energy (Ei) is an energy which is necessary for spending for a
separation and removal an electron from atom, an ion or a molecule. Ionization
energy is a measure of metallic (reducing) properties of elements: the lower the Ei,
31
the stronger the metallic properties are. In groups at increase in a serial number of
an element, the ijnization energy decreases, and in period – increases.
Li
Na Mg Al
Si
P
S
Cl
Ar
K Ca
Еi(eV) 5.39 5.14 7.64 5.98 8.15 10.4 10.4 13.01 15.8 4.3 6.1
Energy of electron affinity (Ea) is an energy which is allocated at joining an
electron to an atom or a molecule. It _isinfection_ non-metallic (_isinfect)
properties of elements: the greater the value Ea, the stronger the non metallic
properties are. In the periods from left to right energy of electron affinity and nonmetallic (_isinfect) properties of elements increase, and in groups from up to down
they decrease.
F
Cl
Br
I
O
N
C
B
Be
Li
ЕA (eV) 3.62 3.82 3.54 3.24 1.48 0.20 1.13 0.30 -0.19 0.54
The half-sum of _isinfecti energy and energy of electron affinity is called
electronegativity of atom. It increases with increase in non-metallic properties of
elements.
In the periodic table, non-metallic elements are settled down in groups A and
occupy its right top part. Metallic elements of groups A are in the left bottom part
of periodic table. All elements of groups B possess metallic properties.
32
6. CHEMICAL BONDS
Glossary: Chemical bonds
alkane. Paraffin. Compare with hydrocarbon and alkene.
A series of organic compounds with general formula CnH2n+2. Alkane
names end with –ane. Examples are propane (with n=3) and octane (with
n=8).
Antibonding orbital. Antibonding; antibonding molecular orbital.
A molecular orbital that can be described as the result of destructive
interference of atomic orbitals on bonded atoms. Antibonding orbitals have
energies higher than the energies its constituent atomic orbitals would have
if the atoms were separate.
Average bond enthalpy. Compare with bond enthalpy.
Average enthalpy change per mole when the same type of bond is broken in
the gas phase for many similar substances
axial. An atom, bond, or lone pair that is perpendicular to equatorial atoms, bonds,
and lone pairs in a trigonal bipyramidal molecular geometry.
Bond energy. Compare with bond enthalpy.
Energy change per mole when a bond is broken in the gas phase for a
particular substance.
Bond enthalpy. Compare with average bond enthalpy.
Enthalpy change per mole when a bond is broken in the gas phase for a
particular substance.
Bond length.
The average distance between the nuclei of two bonded atoms in a stable
molecule.
Bond order.
1. In Lewis structures, the number of electron pairs shared by two
atoms. 2. In molecular orbital theory, the net number of electron pairs
in bonding orbitals (calculated as half the difference between the number of
33
electrons in bonding orbitals and the number of electrons in antibonding
orbitals.
Chemical bond. Bond; bonding; chemical bonding.
A chemical bond is a strong attraction between two or more atoms. Bonds
hold atoms inmolecules and crystals together. There are many types of
chemical bonds, but all involve electrons which are either shared or
transferred between the bonded atoms.
Covalent bond. Covalent; covalently bound. Compare with covalent
compound and ionic bond.
A covalent bond is a very strong attraction between two or more atoms that
are sharing their electrons. In structural formulas, covalent bonds are
represented by a line drawn between the symbols of the bonded atoms.
Electric dipole. Dipole.
An object whose centers of positive and negative charge do not coincide.
For example, a hydrogen chloride (HCl) molecule is an electric dipole
because bonding electrons are on average closer to the chlorine atom than
the hydrogen, producing a partial positive charge on the H end and a partial
negative charge on the Cl end.
Electric dipole moment. (µ) dipole moment.
A measure of the degree of polarity of a polar molecule. Dipole moment is a
vector with magnitude equal to charge separation times the distance between
the centers of positive and negative charges. Chemists point the vector from
the positive to the negative pole; physicists point it the opposite way. Dipole
moments are often expressed in units called Debyes.
electronegativity Compare with ionization energy and electron affinity.
Electronegativity is a measure of the attraction an atom has for bonding
electrons. Bonds between atoms with different electronegativities are polar,
with the bonding electrons spending more time on average around the atom
with higher electronegativity.
Enthalpy of atomization. ( Hat) atomization enthalpy; heat of atomization.
34
The change in enthalpy that occurs when one mole of a compound is
converted into gaseous atoms. All bonds in the compound are broken in
atomization and none are formed, so enthalpies of atomization are always
positive.
Free radical.
A free radical is a molecule with an odd number of electrons. Free radicals
do not have a completed octet and often undergo vigorous redox reactions.
Free radicals produced within cells can react with membranes, enzymes, and
genetic material, damaging or even killing the cell. Free radicals have been
implicated in a number of degenerative conditions, from natural aging to
Alzheimer’s disease.
Geometric isomer.
Geometric isomers are molecules that have the same molecular formula and
bond connections, but distinctly different shapes.
Hydrogen bond. Hydrogen bonding.
An especially strong dipole-dipole force between molecules X-H...Y, where
X and Y are small electronegative atoms (usually F, N, or O) and ... denotes
the hydrogen bond. Hydrogen bonds are responsible for the unique
properties of water and they loosely pin biological polymers like proteins
and DNA into their characteristic shapes.
Incomplete octet.

An atom with less than eight electrons in its valence shell. 2. An
atom with less than eight total bonding and nonbonding electrons
in a Lewis structure, for example, B in BH3 has an incomplete
octet.
Inductive effect. Inductance effect.
An inductive effect is the polarization of a chemical bond caused by the
polarization of an adjacent bond. (Field effects are polarization caused by
nonadjacent bonds).
Inert pair. Inert pair effect.
35
Valence electrons in an s orbital penetrate to the nucleus better than
electrons in p orbitals, and as a result they’re more tightly bound to the
nucleus and less able to participate in bond formation. A pair of such
electrons is called an “inert pair”. The inert pair effect explains why
common ions of Pb are Pb4+ and Pb2+, and not just Pb4+ as we might expect
from the octet rule.
Infrared spectroscopy. IR spectroscopy.
A technique for determining the structure (and sometimes concentration) of
molecules by observing how infrared radiation is absorbed by a sample.
Ionic bond. Ionically bound; ionic bonding. Compare with covalent bond.
An attraction between ions of opposite charge. Potassium bromide consists
of potassium ions (K+) ionically bound to bromide ions (Br-). Unlike
covalent bonds, ionic bond formation involves transfer of electrons, and
ionic bonding is not directional.
Ionic compound. Salt. Compare with covalent compound and ionic bond.
A compound made of distinguishable cations and anions, held together by
electrostatic forces.
Lewis structure. Electron dot structure; dot structure.
A model pioneered by Gilbert N. Lewis and Irving Langmuir that represents
the electronic structure of a molecule by writing the valence electrons of
atoms as dots. Pairs of dots (or lines) wedged between atoms represent
bonds; dots drawn elsewhere represent nonbonding electrons.
Lone pair. Nonbonding pair; unshared pair.
Electrons that are not involved in bonding.
Molecular geometry.

The three-dimensional shape of a molecule. For example,
methane (CH4) has a tetrahedral molecular geometry. 2. The study
of molecular shapes.
Molecular orbital. Compare with atomic orbital and orbital.
36
A wavefunction that describes the behavior of an electron in a molecule.
Molecular orbitals are usually spread across many atoms in the molecule,
and they are often described as a combination of atomic orbitals on those
atoms.
Multiple bond.
Sharing of more than one electron pair between bonded atoms. A double
bond consists of two shared pairs of electrons; a triple bond consists of three
shared pairs.
Octet.
A set of eight valence electrons.
Octet rule.
A guideline for building Lewis structures that states that atoms tend to gain,
lose, or sharevalence electrons with other atoms in a molecule until they
hold or share eight valence electrons. The octet rule almost always holds for
carbon, nitrogen, oxygen, and fluorine; it is regularly violated for other
elements.
Pi bond. ( bond) Compare with sigma bond.
In the valence bond theory, a pi bond is a valence bond formed by side-byside overlap of p orbitals on two bonded atoms. In most multiple bonds, the
first bond is a sigma bond and all of the others are pi bonds.
Polar bond. Compare with covalent bond and ionic bond.
A bond involving electrons that are unequally shared. Polar bonds can be
thought of as intermediate between the extremes represented by covalent
bonds and ionic bonds.
Polar molecule. Polar. Compare with covalent compound, ionic
compound and polar bond.
An asymmetric molecule containing polar bonds. H2O, NH3, and HCl are
examples of polar molecules. Non-examples are CO2, CCl4, and BCl3 which
contain polar bonds but are nonpolar because they have symmetric
shapes. Alkanes are usually asymmetric but are nonpolar because they
37
contain no polar bonds. Polar molecules are electric dipoles and they attract
each other via dipole-dipole forces.
Resonance.
Description of the ground state of a molecule with delocalized electrons as
an average of several Lewis structures. The actual ground state doesn’t
switch rapidly between the separate structures: it is an average.
Resonance effect. Mesomeric effect.
If electron density at a particular point in a molecule is higher or lower than
what you’d expect from a single Lewis structure, and various canonical
structures can be drawn to show how electron delocalization will explain the
discrepancy, the difference in electron density is called a “resonance effect”
or “mesomeric effect”.
Sigma bond. ( bond) Compare with pi bond.
In the valence bond theory, a sigma bond is a valence bond that is
symmetrical around the imaginary line between the bonded atoms. Most
single bonds are sigma bonds.
Triple bond. ( )
A covalent bond that involves 3 bonding pairs. In the valence bond theory,
one of the bonds in a triple bond is a sigma bond and the other two are pi
bonds . For example, the central bond in acetylene is a triple bond: H-C CH.
valence.
The number of hydrogen atoms that typically bond to an atom of an element.
For example, in H2O, oxygen has a valence of 2; carbon in CH4 has a
valence of four.
Valence bond.
In the valence bond theory, a valence bond is a chemical bond formed by
overlap of half-filledatomic orbitals on two different atoms.
Valence electron.
38
Electrons that can be actively involved in chemical change; usually electrons
in the shell with the highest value of n. For example, sodium’s ground
state electron configuration is 1s2 2s2 2p63s1; the 3s electron is the only
valence electron in the atom. Germanium (Ge) has the ground state electron
configuration 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p2; the 4s and 4p electrons are the
valence electrons.
Valence shell.
The shell corresponding to the highest value of principal quantum number
in the atom. The valence electrons in this shell are on average farther from
the nucleus than other electrons; they are often directly involved in chemical
reaction.
The structure of chemical compounds basically is defined by the nature of
chemical bonds.
The chemical bond arises at the interaction of atoms causing formation of
chemically steady two-or multinuclear system (molecule, crystal, etc.). The
formation of a chemical bond is connected with the general decrease in energy of a
system of co-operating particles.
The major characteristics of a chemical bond are bond energy, bond length,
bond angles.
Bond energy is a quantity of energy allocated at formation of a chemical
bond. The more the bond energy, the stronger the molecule is.
Bond length is a distance between nuclei of atoms in a molecule. Bond
lengths are caused by sizes of reacting atoms and degree of overlap of their
electronic shells.
On a way of formation three principal types of a chemical bond are
distinguished. These are ionic, co-valent, and metallic.
Co-valent bond
The chemical bond between the atoms carried out by shared electron pairs is
called a co-valent bond. It arises between identical atoms forming gaseous
39
binuclear molecules, and also in the condensed state with participation of nonmetallic atoms.
There are two basic concepts of description of a co-valent bond:
1. Method of valence bond (VВ).
2. Method molecular orbitals (МО).
Both methods mutually supplement and do not exclude each other as they
use
various
ways
of
approaches.
See
also
http://en.wikipedia.org/wiki/Chemical_bond
Method of valence bonds (VB)
According to the VB method, valency can be considered as number of
formed shared electronic pairs. From the point of view of the exchange
mechanism, valency of an element is defined by number of non-paired electrons.
Atoms can form a limited number of chemical bonds according to their
valency. It corresponds to saturability of a co-valent bond.
Depending on number of unpaired electrons, atoms can form one, two or
three co-valent bonds, i.e. a co-valent bond may be simple, double or triple.
The strongest chemical bonds arise in a direction of a maximal overlap of
atomic orbitals. As the orbitals have different spatial dispositions, therefore covalent bonds are _isinfection_ by orientations.
Depending on directions of overlapping one can distinguish ,  and bonds.
-bonds are formed when two atomic orbitals overlap along an axis
connecting nuclei of atoms.
-bonds are formed when two atomic orbitals are site-overlapped.
-bonds arise at overlapping of two d-orbitals, located in parallele planes.
The hybridization process also influences orientation of a co-valent bond.
See also http://en.wikipedia.org/wiki/Orbital_hybridisation#sp2_hybrids
Hybridization is a mixing of of different subshells of an atom, electrons of
which participate in formation of equivalent chemical bonds.
40
Depending on hybridization type hybrid orbitals have different position in
space.
Sp – linear, an interbond angle 180 
sp2 – triangular, an interbond angle 120 

sp3 – tetrahedral, an interbond angle  109 

Shared electron pairs in a molecule are shifted to a more electronegative
atom, thus a co-valent bond possesses a property of polarity. Molecules formed by
identical atoms (Cl2, H2, etc.), have non-polar bonds. The more the difference of
electronegativities of the two atoms forming a chemical bond, the more polar it is.
In case if an exchange mechanism of formation of a co-valent bond takes
place, one of atoms (donor) delivers a pair of electrons, and another (acceptor) – a
not-filled orbital. So, a co-ordinate bond is formed.
For example, formation of an ammonium ion (NH4 +) involves realization of
both mechanisms of formation of a co-valent bond.
41
2
2
3
The electronic formula of a nitrogen atom 7N 1s 2s 2p .


  
1s
2s
2p
It
has five electrons on the valence shell of which three electrons are unpaired and
form co-valent bonds with three H-atoms (exchange mechanism). The lone
electron pair of nitrogen participates in formation of a co-ordinate bond with an ion
Н+ (nitrogen represents itself as the donor, and ion Н + as an acceptor of electrons).
The method of valence bonds allows to distinguish concepts of valency and
oxidation state.
Valency of an atom _isinfection_ its ability to form co-valent chemical
bonds.
Oxidation state is a conditional charge on atom in a molecule if to assume,
that shared electronic pairs are completely shifted to a more electronegative atom.
For example, valency of nitrogen in molecule NH3 is III (three co-valent
bonds). Since electronegativity of nitrogen exceeds that of hydrogen, all three
formed shared electronic pairs are shifted towards nitrogen, giving to it a negative
charge (oxidation state –3).
The VB method ВС theoretically predicts structures and properties of many
molecules and ions. Therefore it cannot explain existence of some molecular ions
(He2+, O22- ets). By means of this method it is impossible to explain magnetic
properties of some molecules, for example: О2 and В2.
See also http://en.wikipedia.org/wiki/Chemical_bond#Covalent_bond
Method of molecular orbitals (МО)
According to МО method, the molecule consists of a set of nuclei and
electrons disposed on molecular orbitals.
Main tenets of the MO method:
1. Each electron in a molecule occupies a certain energetic level (molecular
orbital, MO) which is _isinfection_ by a molecular -function and a
corresponding set of quantum numbers.
2. The total number of formed МО equals the number of initial atomic orbitals.
3. Filling of MO occurs according th all the principles presented for atomic
orbitals.
42
4. МО it is considered as a linear combination of atomic orbitals (MO –
LCAO).
Let’s consider, for example, formation of molecule АВ. Valence electrons
of each atom are on p-orbitals.
If the wave function of the isolated atom A is А, and for the atom B it is
В, so according to the МО method:
АВ=С1АС2В
where С1 and С2 are coefficients considering the income of each atom in formation
of a molecular orbital.
(p*.)
E
_

А
В
+
А
(p)
The new molecular orbital with a lower energy (p) is known as a binding
orbital. As its energy is lower than the energy of an atomic orbital, electrons on it
stabilize the molecule.
The MO with a higher energy (p*) is called an antibinding orbital.
Electrons on it tend to destroy the molecule.
Stability of a molecule is described by a bond order, B.O.
B.O. = ½ (No of binding electrons – No of antibinding electrons)
If B.O. = 0 the number of electrons on binding orbitals equals the number of
electrons on antibinding electrons. Such molecule is unstable and breaks up to
initial atoms (does not exist).
Conditions for formation of MO from AO are the following:
close values of energies of overlapping atomic orbitals;
 considerable overlap of AO (formation of  and -types of MO);
 identical spatial disposition of AO (рх – рх, instead of рх – рz).
43
On level power molecular orbital two-nuclear molecules settles down in a
following order: s(1s)  s(1s)  s(2s) s*(2s) x(2px)  z(2pz) = y(2py)
 z*(2pz) = y*(2py)  x*(2px)
Examples of the description of molecules using the MO method (energetic
diagrammes of molecules and molecular ions):

Molecule Н2
Energetic diagrammes shows transformation of atomic orbitals
into
molecular orbitals.
2. Molecule Не2
As each He-atom has two electrons paired on 1s-atomic orbitals, in the He2
molecule both binding and antibinding orbitals contain same number of electrons.
B.O. = 0, so the molecule does not exist.
44
3. Molecular ion Не2+
One can suppose that a molecular ion Не2+ can be formed if a Не-atom interacts
with the Не+-ion. Totally, 3 electrons are present in the space, two of which
occupy a s-binding orbital and the rest one is deposited on the s*-antibinding
orbital. B.O.= ½, i.e. the molecular ion Не2+ exists and forms a semi-bond from
the point of view of the VB-method. The existence of a Не+-ion was proved
experimentally, aqnd it was found that the bond between two nuclei is twice
weaker than that in the H2-molecule.

Molecule Li2
Energetic diagrammes shows transformation of atomic orbitals
into
molecular orbitals.
5. Molecule О2
Electronic configuration of valence shell of O-atom is (2s22p4). The interaction of
two s-orbitals of two oxygen atoms is analogue to previous cases.
p-Orbitals can form both  and  bonds.
45
All 12 electrons of the two O-atoms occupy lower molecular orbitals. As two p*
molecular orbitals are of same energy, O2-molecule has two unpaired electrons and
possesses paramagnetic properties.
Bond order B.O. = 2 which correlates with the double bond from the VBmethod.
See also http://en.wikipedia.org/wiki/Molecular_orbital_diagram
Ionic bonds
Ionic bond represents an electrostatic interaction between ions of opposite
charges. Ionic bond can be considered as a limiting case of a polar co-valent bond
where the difference of electronegativities of the two atoms forming a chemical
bond exceeds 2). Usually it is considered that a ionic bond is formed at interaction
of typical metals and typical non-metals.
Energy of ionic bonds depends upon:
1. energy of electrostatic interaction between ions, i.e. it increases with increase in
charges of ions and reduction of their radii;
46
2. energy of electronic affinity of non-metals which increases at increase in nonmetallic properties of elements;
3. _isinfecti energy of atoms.
Example: formation of a molecule of sodium chloride:
Na + Cl  NaCl
Na  Na + + e
Еi = 495 kJ
Cl + e  Cl -
Еa = 345 kJ
Na + + Cl -  NaCl Ecolomb = 585 kJ
Еbond = Еcolomb + Еa – Еi = 435 kJ
Representation
of
ionic
bonding
between
lithium and fluorine to
form lithium fluoride. Lithium has a low ionization energy and readily gives up
its lone valence electron to the fluorine atom, which has a positive electron
affinity and accepts the electron that was donated by the lithium atom. The end
result is that lithium is isoelectronic with helium and fluorine is isoelectronic
with neon. Electrostatic interaction between the two atoms forms an ionic bond.
Ionic bonds are not directed and not saturable. That defines a great stability
of ionic crystals.
See also http://en.wikipedia.org/wiki/Ionic_bond
47
INTERACTION OF MOLECULES
(THE CONDENSED STATE OF SUBSTANCES)
Chemical stability of molecules is shown only in systems, where distance
between molecules is much more than their sizes (r10-9m). That corresponds to
a gaseous state of a substance.
In case if the distance between molecules makes about 10 -9 m (condencad
state which may be liquid or solid) , arise forces of van der Waals which have
electrostatic nature and are subdivided on:
1) orientational (dipole – dipole);
2) inductional (dipole – not polar molecule);
4) _isinfecti (dispersive interaction of instantly induced dipoles of
polarizable molecules).
48
Hydrogen bonds have intermediate character between intermolecular
interaction and a co-valent bond. It is a kind of interaction between positively
_isinfect atom of hydrogen and negatively _isinfect atoms with high
electronegativity (F, O, N, S, etc.).
At the expense of the small size of an H-atom, it has ability to enter
electronic shells of other atoms where there is an interaction which is intermediate
between electrostatic interaction and co-valent bond (interaction with lone electron
pairs of non-metallic atoms).
Hydrogen bond is indicated as Х – Н … Y (X, Y =F, O, N, S).
See also http://en.wikipedia.org/wiki/Hydrogen_bonds
Model of hydrogen bonds (1) between molecules of water
An example of intermolecular hydrogen bonding in a self-assembled _isinf
complex reported by Meijer and coworkers.[ The hydrogen bonds are the dotted
lines.
49
QUESTIONS AND PROBLEMS
1. What types of bonds can be attributed to the chemical?
2. What are the two main approaches to the consideration of the chemical
bond you know? What is the difference?
3. Define the valence and degree of oxidation.
4. What are the differences between simple covalent, donor-acceptor, dative,
metallic, ionic bonds?
5. Classified as intermolecular bonds?
6. What is electronegativity? How is data electronegativity calculated? What
is the electronegativity of the atoms that form the bond and allow us to judge? How
does the electronegativity of the atoms of elements in moving in the periodic table
Mendeleev top to bottom and left to right?
7. What rules should be guided by the consideration of the molecular
structure of MO LCAO?
8. Using the method of valence bonds, explain the structure of the hydrogen
compounds of elements of the 2nd period.
9. The dissociation energy of the molecules in a series of Cl2, Br2, I2
decreases (239 kJ/mol, 192 kJ/mol, 149 kJ/mol, respectively), but the energy of
dissociation of F2 (151 kJ / mol) is much smaller than the energy of dissociation of
Cl2, and from the general pattern. Explain the facts.
10. Why, under normal conditions, CO2 – gas and SiO2 – solid, H2O – liquid
and H2S – gas? Try to explain the physical state of matter.
11. Using the MO LCAO method, explain the appearance and
characteristics of the chemical bond in B2, C2, N2, F2, LiH, CH4.
12. Using the theory of valence electron pair repulsion, determine the form
of molecules of oxygen compounds of elements of the 2nd period.
50
7. SOLUTIONS
Solutions are homogeneous systems of variable composition.
Solutions consist at least of two components – solvent and solute. Solvents
are accepted to be that substances which keep their aggregate states or which are of
greater amounts.
Amount (mass) of the solute in a mass or volume unit of the solution is
called concentration of a solution.
The most widespread concentration units of solutions are rhe following:
Mass fraction represents a mass of substance in 100 g of a solution:

m( solute)
m( solution)
(100%)
Molar concentration (molarity) is a number of moles of a solute in one
liter of a solution:
CM 
 ( solute)
V( solution)

m( solute)
M ( solute)  V( solution)
Equivalent (normal) concentration is a number of equivalents of a solute in
one liter of a solution:
CE 
n( solute)
V( solution)

m( solute)
M E ( solute)  V( solution)
Solubility is an ability of one substance to be dissolved in other under the set
conditions. Quantitatively it is expressed by solubility factor, s. It equals
concentration of the saturated solution under the given conditions.
Solubility of substances depends on temperature and pressure: for liquid and
solid solutes it increases at rise in temperature, for gases – at fall of temperature
and pressure increase.
Physical and chemical processes in solutions
Interaction between molecules and ions of molecules of solute and solvent
can consist of the several processes proceeding consistently or simultaneously.
51
1. Molecular dissociation of a solute:
(АВ)k
k AB
2. Interaction of molecules of solute with molecules of solvent (formation
of solvates):
AB + (n+m) S

AB  (n+m) S
Electrolytic dissociation (splitting of a solute into solvated
ions):
AB  (n+m) S
Ax +  nS + Bx  mS
Substances which can form ions while being dissolved, are known as
electrolytes.
The quantitative characteristic of electrolytic dissociation is known as
degree of dissociation:
 = Сdis / Сtot,
where Сdis is the consentration of dissociated part of the electrolyte, and Сtot is its
total concentration.
According to the value of the degree of dissociation, the electrolytes can be
devided into two groups:
1. Strong electrolytes (> 0.3 or 30 %). Among strong electrolytes there are
some strong acids (HCl, H2SO4, HNO3, HclO4, HBr, HI), alkalis (soluble bases
such as NaOH, KOH, Ca(OH)2, Ba(OH)2, etc.) and practically all salts. In solutions
strong electrolits are practically completely broken up to ions (dissociation is
irreversible and complete):
2
Al2 (SO4)  2 Al3 + + 3 SO4
2. Weak electrolytes ( 0.03 or 3 %). Among weak soluble electrolytes
there are weak acids, ammonium _isinfect NH4OH, and water itself. Dissociation
52
of weak electrolytes is a reversible and stepwise process which is _isinfection_ by
stepwise and overall equilibria constants (dissociation constants).
For example, dissociation of phosphoric acid is a three-step process:


H + + H2PO4 ;
1-st step: H3PO4
2-nd step: H2PO4
[H  ]  [H2 PO4 ]

[H  ]  [H PO4
H + + HPO4
2
;
2
H + + PO4
3
2

3
]
2
;
K3 =
]
=6108
[H2 PO4 ]
K2 =
[H  ]  [ PO4
3-rd step: HPO4
=8103
[H3 PO4 ]
K1 =
[ HPO4 ]
=21012
Overall process:
[ H  ]3  [ PO4
H3PO4
where
[H+],
3H + + PO4

[H2PO4 ],
concentrations of ions;
3
;
[HPO4
K=
2
],
[H3 PO4 ]
[PO4
3
],
3
]
= К1К2К3=11021
[H3PO4]
are
equilibrium
К1, К2, К3 – stepwise dissociation constants; and K is the
overall dissociarion constant.
See also http://en.wikipedia.org/wiki/Acid-base_reaction_theories
QUESTIONS AND PROBLEMS
1. The dissociating constant of butyric acid C3H7COOH is 1.5  10
5
.
Calculate the degree of its dissociation in a 0.005 M solution.
2. What is the hydrogen ion concentration [H+] in an aqueous solution of
formic acid if  = 0.03?
3. Calculate the concentration of acetate ions in 0.1 M solution of acetic acid
in presence of 0.01 M HCl.
4. Calculate the ionic strength and the activities of the ions in a solution
containing 0.01 molel-1 of Ca(NO3)2 and 0.01 mol/l of CaCl2.
53

REACTIONS OF IONIC EXCHANGE
Reactions of ionic exchange in solutions occur between ions of strong
electrolytes and molecules of weak electrolytes and insoluble substances. They
proceed towards formation of precipitates, gases,
and molecules of weak
electrolytes.
 Na2SO4 + 2HNO2
2NaNO2 + H2SO4 
soluble
strong
soluble
weak
(reaction in a molecular form)

 

2Na + + 2NO2 + 2H + + SO42 
2Na + + SO42 + 2HNO2
(full ionic form of the reaction)

 2HNO2
2NO2 + 2H + 
(net ionic form)
Properties of chemical compounds in solutions are defined by the character
of their dissociation:

HCl  H + + Cl (acids form hydrogen ions Н + while dissociation);
NaOH  Na + + OH
NaCl  Na + + Cl


(bases dissociate to produce ions of hydroxide OH );
(salts form metallic cations and anions of acids).
The main reaction which reflects acidic and basic properties is the reaction
of neutralization (acids interact with bases to produce salts):
 Na2SO4 + 2H2O
2NaOH + H2SO4 

 

2Na + + 2OH + 2H + + SO42 
2Na + + SO42 + 2H2O

 2H2O
2OH + 2H + 
There are electrolits which can participate in chemical reactions both as
bases and as acids. Such electrolytes are called amphoteric. Among them there are
Zn(OH)2, Pb(OH)2, Sn(OH)2, Be(OH)2, Al(OH)3, Cr(OH)3, As(OH)3, and some
others. These substances are capable to react both with acids and with the bases,
forming salts as products of reaction of:
Al(OH)3 + 3HCl  AlCl3 + 3H2O
54
Al(OH)3 + 3H+  Al3+ + 3H2O
Al (OH)3 + 3NaOH  Na3 [Al(OH)6]
Al(OH)3 + 3OH–  [Al(OH)6]3–
Sn(OH)2 + 2HCl  SnCl2 + 2H2O
Sn(OH)2 + 2H+  Sn2+ + 2H2O
Sn(OH)2 + 2NaOH  Na2 [Sn(OH)4]
Sn(OH)2 + 2OH–  [Sn(OH)4]2–
9. DISSOCIATION OF STRONG ELECTROLYTES
Strong Acids

Perchloric acid HclO4

Hydriodic acid HI

Hydrobromic acid HBr

Hydrochloric acid HCl

Sulfuric acid H2SO4

Nitric acid HNO3

Chloric acid HclO3

Bromic acid HbrO3

Perbromic acid HbrO4

Periodic acid HIO4
Strong Bases

Potassium hydroxide KOH

Barium hydroxide Ba(OH)2

Caesium hydroxide CsOH

Sodium hydroxide NaOH

Strontium hydroxide Sr(OH)2

Calcium hydroxide Ca(OH)2

Rubidium hydroxide RbOH
55

Magnesium hydroxide Mg(OH)2
Salts

Sodium chloride

Potassium nitrate

Magnesium chloride

Sodium acetate
In solutions of strong electrolytes owing to their full dissociation,
concentration of ions is great, therefore properties of such solutions will depend on
degree of interaction of ions as with each other, and with polar molecules of
solvent. So, concentrations of ions are replaced by their activities.
Activity is a visual concentration of an ion involving its interaction with
other ions of the solution:
a = fC (f – activity factor).
If f = 1 ions are free and do not co-operate among themselves (a=C).
If f <1 ions co-operate (a <C). The less the activity factor, the more
interaction between ions exists in the solution.
The activity factor depends on total concentration of all the ions in a
solution (ionic strength of a solution):
 = 1/2  Ci Zi2,
where  - ionic strength; Ci – concentration of ions in a solution; Zi – charges of
ions.
log f  0.5Z 2

 1
(Debaue-Huckel’s equation)
For diluted solutions of strong electrolytes with <<1,
log f  0.5Z 2 
Example. Calculate ionic strength of a solution containing 0.02 mol/l of
CaCl2 and 0.05 mol/l of Na2SO4.
As soluble salts are strong electrolytes, they are fully dissociated in the
solution:
CaCl2  Ca2+ + 2Cl–
56
Na2SO4  2Na+ + SO42–
[Ca2+] = C(CaCl2) = 0.02 M;
[Cl–] = 2C(CaCl2) = 0.04M;
[Na+] = 2C(Na2SO4) = 0.1M;
[SO42–] = C(Na2SO4) = 0.05M.
 = 1/2{[Ca2+]Z2(Ca2+) + [Cl–]Z2(Cl–) + [Na+]Z2(Na+) + [SO42–]Z2(SO42–)}
=
1
(0,02  22 + 0,04 12 + 0,1 12 + 0,05  22) = 0,21
2
mol/l.
10. DISSOCIATION OF WEAK ELECTROLYTES
While dissociation of weak electrolytes takes place, an equilibrium is
established.
CH3COO + H+
CH3COOH
Thus, if the total concentration of the electrolyte equals C, and degree of its
dissociation is ,
Cdis = ∙ C
[CH3COO-] = [H +] = C
[CH3COO-] = C – C
K
H  CH COO  


3
CH 3COOH 
C 2 2
C 2 2
C 2


C  C C 1    1  
For 1 K=C2

and
K
C
The resulted equation expresses the Ostwald’s dilution law according to
which degree of dissociation of a weak electrolyte increases with dilution of a
solution.
Addition of common ions in a solution of a weak electrolyte causes shifting
of equilibrium of the reaction towards reduction of dissociation (effect of a
common ion).
57
Electrolytic dissociation of water
Water is a weak electrolyte which dissociates according to the equation:
Н2О
Н + + OH.
At the temperature of 22oC, the equilibrium of the dissociation process
establishes such that:
[H +]  [OH ] = KH2O = 1014
(ionic product of water).
In a neutral solution [H +] = [OH] =
1014
= 107 mol/l.
In an acidic solution [Н +]> [OH];
[H +]> 107 mol/l.
In an alkaline solution [H +] <[OH];
[H +] <107 mol/l.
Knowing concentration of one of the ions, for example [Н+] and ionic
product of water, it is possible to calculate concentration of another type of ions
[OH].
The negative logarithm of concentration of of hydrogen ions (or the negative
logarithm of activity of ions of hydrogen) is named рН:
рН = – log [H+]
In neutral solutions at 22оС рН = 7.
In acidic solutions рН < 7.
In alkaline solutions рН > 7.
Acid-base indicators are substances, changing colouring in certain area of
value pH a solution. Weak organic acids or bases, which molecules and ions have
different colouring can be indicators.
Methylorange
58
Phenolphtalein
Litmus
Area of transition of colouring of some indicators
The indicator
Colour
Area of
transition of
colouring, рН
Methylorange
Phenolphtalein
Litmus
acidic
form
red
colourless.
red
alkaline
form
yellow
Red
dark blue
3.2 – 4.5
8.2 – 10.0
6.0– 9.,0
Example 1. Calculate pH of a 0.01 M solution of NaOH.
As NaOH is a strong electrolyte, it is fully dissociated in solutions:
NaOH  Na+ + OH
[OH] = C(NaOH) = 102 M.
[H+] = 10[OH] = 102 M.
pH = –log [H+] = 12.
Example 2. Calculate pH of a 0.01 M solution of CH3COOH.
59
Acetic acid is a weak electrolyte so it dissociates to a very small extent:
CH3COOH
CH3COO + H+
[H+] = C(CH3COOH)
According to the Ostwald’s dilution law,

[H  ] 
K dis (CH 3COOH )
,
C (CH 3COOH )
so
K (CH 3COOH )
 C (CH 3COOH )  K (CH 3COOH )  C (CH 3COOH )
C (CH 3COOH )
As K(CH3COOH) is a table value which equals 1.75105, so
[ H  ]  1.75  10 5  0.01  4.18  10 4
pH = –log [H+] = –log (4.1810) = 4 – log(4.18) = 3.38.
Example 3. Calculate pH of a 0.01 M solution of CH3COOH in presence of
0.1 M solution of CH3COONa.
CH3COO + H+
CH3COOH
(weak electrolyte)
CH3COONa  CH3COO + Na+
(strong electrolyte)
As the amount of acetate anions produced by a salt (sodium acetate)
extremely exceeds the amount of acetate anions obtained within dissociation of the
acid,
K(CH COOH) 
3
[H  ] 
[CH COO ]  [H ] C (salt )  [ H  ]
3

[CH COOH]
C (acid )
3
K (CH 3COOH )  C (acid ) 1.75  10 5  0.01

 1.75  10 6
C (salt )
0.1
pH = –log [H+] = –log (1.7510) = 6 – log(1.75) = 5.76.
60
11. EQUILIDRIA IN SOLUTIONS WITH PRECIPITATES
SOLUBILITY
According to an IUPAC definition, solubility is the analytical composition of
a saturated solution expressed as a proportion of a designated solute in a designated
solvent. Solubility may be stated in units of concentration, molality, mole fraction,
mole ratio, and other units.
The solubility of a given solute in a given solvent typically depends on
temperature.
In the saturated solution of a sparingly soluble electrolyte, a dynamic
equilibrium between a solid phase and a solution is established:
Ca2 + + CO32 
CaCO3 (s.)
As the concentration of a solid phase is constant, the equilibrium constant of
the reaction may be written as following:
Ksp = [Ca2 +]  [CO32 ]
where Ksp is called a solubility product and [Ca2 +] and
[CO32 ] represent
concentrations of corresponding ions in a saturated solution over the precipitate/
Numerical values of solubility products of substances are presented in special
tables.
61
In the presence of electrolytes with common ions the equilibrium is shifted
towards formation of the precipitate, i.e. common ions decrease solubilities of
precipitates (effect of a common ion).
In the presence of strong indifferent electrolytes which are not containing
common ions, mobility of ions in a solution decreases and the equilibrium shifts
towards increase the cfoncentration of ions in the solution, i.e. strong indifferent
electrolytes increase solubilities of precipitates (salt effect).
Precipitates are usually formed during reactions of ionic exchange. At first
moment of time after we start mixing the solutions, concentrations of ions are
small and the saturation of solution by ions is not reached.
Condition for formation of a precipitate: product of concentrations of ions
in a solution should exceed the value of solubility product of the expected
precipitate. For example,
[Ca2 +] [CO32 ]> Ksp (CaCO3).
Condition for dissolution of a precipitate: product of concentrations of ions
in a solution should deceed the value of solubility product of the expected
precipitate ([Ca2 +] [CO32 ] <Ksp (CaCO3)).
Correlation between solubility products and solubilities of precipitates
Considering the existence of a precipitate of KxAy type:
KxAy
x Ky + + y Ax 
Expression for solubility product of this substance looks like:
Ksp = [Ky +] x [Ax ] y
If to express the solubility of the electrolyte through “S”, then the solution
will contain cations and anions in the following amounts:
[Ky +] = xS (mol/l),
[Ax ] = yS (mol/l).
So the expression of the solubility product can be expressed as:
62
Ksp = (xS) x (yS) y,
or
Ksp
xxy y
s x y
Example 1. Correlation between Ksp and S for Ag2S
2 Ag + + S2 
Ag2S
[S2 ] = S (mol/l).
[Ag +] = 2S (mol/l),
Ksp = (2S)2 (S) = 4S3,
K sp
4
s3
Example 2.
or
Calculate the solubility of AgCl in excess of 0.01 M solution of
HCl.
AgCl
Ag+ + Cl
s
s
s+0.01  0.01
KSp (AgCl) = [Ag+][Cl] = s  0.01 = 1,61010, s = 1,610 mol/l
Example 3.
Will a precipitate of PbCl2 be formed if to mix equal voloumes of
0.01M solution of Pb(NO3)2 and 0.02M solution of HCl?
Pb(NO3)2 + 2HCl  PbCl2 + 2HNO3
Pb 2+ + 2Cl PbCl2
While mixing equal volumes of solutions, concentrations of reactants will
decrease twice.
Concentrations of ions in the solution are determined as:
[Pb2+] = C(Pb(NO3)2) = 0.01 / 2 = 0.005 mol/l
[Cl] = C(HCl) = 0.02 / 2 = 0.01 mol/l
63
Product of concentrations of ions
Pi = [Pb2+][Cl]2 = 0.005  (0.01)2 = 510
KSp(PbCl2) = 1.6105.
As Pi < Ksp, so the precipitation of PbCl2 will not take place.
QUESTIONS AND PROBLEMS
1. Calculate the solubility product constant of PbBr2 if the solubility of the salt is
1.32  10
2
molel-1.
2. Calculate the solubility of the salt of Al2S2 in the presence of 0.01N Na2S.
3. Calculate the mass of Ag+ in the saturated solution of AgCl.
4. Will a precipitate of silver sulfate be formed if a 1M solution of H 2SO4 is added
to an equal volume of a 0.02 M solution of AgNO3?
5. Calculate the equilibrium constant of the reaction and explain whether the
precipitate of calcium oxalate can be dissolved in acetic acid.
12. DIRECTION OF REACTIONS OF IONIC EXCHANGE
Reactions of an ionic exchange are irreversible only in that case when weak
electrolytes or precipitates are in one part of the equation. If weak electrolytes and
precipitates are present both among reactants and products, reactions of ionic
exchange are reversible and shifted either towards products (reaction proceeds) or
towards reactants (reaction does not proceed). For determination of possibility of a
reaction, the equilibrium constant of the reaction Kr should be calculated.
To calculate the equilibrium constant of the reaction, its numerator
represents product of all the constants of reactants as far as the denominator
contains corresponding constants for reactional products. Among constants there
are overall dissociation constants of weak acids and bases, ionic product of water,
solubility products of precipitates, inctability constants of precipitates.
In the event that the equilibrium constant of the reaction Кr > 1, the
equilibrium is shifted towards products, so the reaction of ionic exchange is
64
possible. If Кr < 105, the equilibrium of the reaction is practically completely
displaced towards reactants, and the forward reaction is impossible. In that case
when the Кr value is concluded in limits from 1 to 10 5, the given reaction is
possible an excess of one of the reactants.
For example,
СaCO3  + 2 CH3COOH
Ca (CH3COO)2 + H2CO3
In an ionic form the reaction may be registered as
СaCO3  + 2 CH3COOH
Ca2 + + 2 CH3COO–- + H2CO3
Kr 
Ksp(CaCO3 ) K 2 (CH 3COOH )
.
K1 ( H 2 CO3 ) K 2 ( H 2 CO3 )
From the corresponding tables it can be found that Ksp (CaCO3) = 4.8x10–9;
K(CH3COOH) = 1.7x10–5; K1(H2CO3) = 4.5x10–7, K2(H2CO3) = 4.8x10–11. So, Kr
= 6.6x10–2. Thus we can conclude that the precipitate of calcium carbonate can be
dissolved in excess of acetic acid.
13. HYDROLYSIS OF SALTS
If the salt dissolved in water contains ions of weak acids or weak bases, the
process of hydrolysis of salt occurs. Hydrolysis of salts is an exchange reaction of
ions of salt with molecules of water leading to formation of molecules and ions of
new weak electrolytes.
Key rules of reactions of hydrolysis:
1. Only anions of weak acids and cations of weak bases which are
components of a salt undergo hydrolysis.
2. Hydrolysis is a stepwise process. At each step one hydrolyzed ion reacts
with only one molecule of water.
3. Under usual conditions hydrolysis only the first step of hydrolysis
proceeds. Hydrolysis increases at heating and dilution of solutions of salts.
4. Hydrolysis is a reversible process, and its equilibrium may be shifted.
Addition of ions common with those evolved during hydrolysis (Н+ or OH–) shifts
65
the equilibrium towards reduction of hydrolysis. Addition of opposite ions shifts
the equilibrium towards increase of hydrolysis.
Types of reactions of hydrolysis
1. Salt is formed by ions of a strong base and a strong acid (for example,
NaCl, KNO3, etc.).
NaCl + H2O 
hydrolysis does not proceed.
2. Salt is formed by ions of a strong base and a weak acid (for example,
Na2CO3, KSCN, etc.).
Na2CO3 + Н2О  the anion of a weak acid is involved in hydrolysis:
CO32– + HOH
HCO3– + OH–
As far as hydroxyd-ions are evolved within the reaction, the medium
becomes alkaline, рН> 7.

Addition to a solution of alkalis (NaOH), containing common ions (OH ),
depresses the process of hydrolysis as far as addition of acids shifts the equilibrium
of the reaction towards products (hydrolysis increases).

Salt is formed by ions of a weak base and a strong acid (for
example, AlCl3, FeSO4, etc.).
AlCl3 + H2O  the cation of a weak base is involved in hydrolysis:
Al3 + + HOH
AlOH2 + + H +
As hydrogen catons are evolved, the medium becomes acidic with рН <7.

Salt is formed by ions of the weak basis and weak acid:
Soluble salts (for example, (NH4) 2CO3, NH4NO2, etc.).
(NH4) 2CO3 + H2O  both cation and anion are involved in hedrolysis:
2NH4+ + CO32– + HOH
NH4OH + HCO3– + NH4+ (рН  7)
Insoluble salts (for example, FeS, ZnSiO3, etc.).
FeS + H2O 
insoluble salts do not undergo hydrolysis.
Some salts of trivalent metals containing anions of very weak acids (Fe2S3,
Al2(CO3)3, etc.) undergo full and irreversible hydrolysis:
Fe2S3+6H2O 2Fe(OH)3+3H2S
66
Quantitative calculations in hydrolysis
Quantitatively hydrolysis is described by a hydrolysis constant (Kh) and
degree of hydrolysis (h).
h
С hydrolyzed.
Kh
C salt
h
С salt
Example 1:
NaCN + H2O
CN– + H2O
NaOH + HCN
OH– + HCN;
pH> 7
It is possible to express a hydrolysis constant through concentrations of ions
in a solution taking into account that concentration of water practically does not
change, or by a rule of calculation of an equilibrium constant of a reversible
reaction:
K
H 2O
[OH  ][ HCN ]
Kh 

K
[CN  ]
acid
Example 2: NH4Cl + H2O
NH4OH + HCl
NH4+ + H2O
Kh 
H+ + NH4OH; pH <7
K
[ H  ][ NH 4 OH ]
H 2O

K
[ NH 4  ]
base
Example 3: NH4CH3COO + H2O
NH4OH + CH3COOH
NH4+ + CH3COO– + H2O
NH4OH + CH3COOH
[ NH 4 OH ]  [CH 3COOH ]
Kh 

K
[ NH  ]  [CH COO  ]
4
3
K
H 2O
K
acid
base
Example 4. Calculate рН of a 0,1 M solution of potassium phosphate.
HPO42– + OH–
PO43– + HOH
K
Kh 
h
Kr
C
K PO
3 4

H 2O
K III H PO
3 4

10 14
 7.69  10 3
12
1.3  10
7.69  10 3
 2.77  10 1
0.1
67

[OH ] = hC = 2.77 10–1  0.1 = 2.7710–2;
[H+] = 10–14 / [OH–] = 10–14 / 2.77 10–2 = 3.61 10–13;
рН = –lg [H+] = 12.44.
Special types of hydrolysis

Hydrolysis of Bi(NO3)3 and SbCl3
Under standard conditions only 1-st step of hydrolysis is possible:
Bi(NO3)3 + HOH
BiOH(NO3)2 + HNO3
Bi3 + + HOH
BiOH2 + + H+
After dilution, the 2-nd step of hydrolysis becomes possible. The process
becomes irreversible because of formation of a precipitate of bithmuth oxonitrate:
BiOH(NO3) 2 + HOH  Bi(OH)2NO3 + HNO3

BiONO3 + H2O
BiOH2 + + HOH + NO3 – BiONO3 + H +
2. Mutual irreversible hydrolysis
Ions Н+ (or OH–) can combine together to form water molecules. So if two
salts of different types of hydrolysis are mixing together, that will cause mutual
strengthening of hydrolysis of both salts and as a result – formation of endproducts of hydrolysis (mutual irreversible hydrolysis). For example, mixing of
solutions Na2CO3 and AlCl3 leads to evolution of СО2 gas and formation of a
precipitate of Al(OH)3:
2AlCl3 + 3Na2CO3 + 3H2O  2Al(OH)3 + 3CO2 + 6NaCl

2Al3 + + 3CO32 + 3H2O  2Al(OH)3 + 3CO2 
Sometimes in similar cases least soluble of possible products of hydrolysis
are precipitated. For example, the solubility of _isinf hydroxocarbonate
(CuOH)2CO3 is less than the one of cupric hydroxide Cu(OH) 2. Therefore mixing
of solutions of CuSO4 and Na2CO3 leads to precipitation of (CuOH)2CO3:
2CuSO4 + 2Na2CO3 + H2O  (CuOH) 2CO3  + CO2  + 2Na2SO4

2Cu2 + + 2CO32 + H2O  (CuOH) 2CO3  + CO2 
68
QUESTIONS AND PROBLEMS
1. Calculate pH of 0.1 M solution of NaOH (assume the dissociation to be
complete).
2. Calculate pH of a 0.01 M solution of acetic acid if the degree of dissociation
of the electrolyte equals 0.042.
3. Calculate pH of an ammonium buffer solution prepared by mixing of equal
volumes of 0.1 M solution of NH4OH and 0.01 M solution of NH4Cl.
4. Which of the salts listed below undergo hydrolysis? Write the net ionic
equations and indicate whether aqueous solutions of salts are neutral, acidic or
basic. NaCN, KNO3, K2S, ZnCl2, NH4NO2, MgSO4.
5. Calculate hydrolysis constant and degree of hydrolysis in 0.1 M solutions of:
a) NH4Cl; b) Na2CO3 (only the first step of hydrolysis should be taken into
consideration).
6. When aqueous solutions of Cr(NO3)3 and Na2S are
mixed
together, a
precipitate is formed and a gas is evolved. Write the molecular and net ionic
equations of the reaction.
14. OXIDATION-REDUCTION REACTIONS
Oxidation-reduction are those reactions which proceed with the change in
oxidation states of chemical elements because of transition of electrons from one
particle (atom, molecule or ion) to another.
The loosing of electrons by an atom attended by an increase in its oxidation
number is called oxidation; the gaining of electrons by an atom attended by a
decrease in its oxidation number is called reduction.
A substance containing an element that undergoes oxidation is called a
reducing agent. These are almost all metals and some non-metals (C, H2 and
others, negatively charged ions of non-metals (S2- I- , N3- and others), cations in
intermediate oxidation numbers (Sn2+, Fe2+ and others), ions containing elements
69
in intermediate oxidation numbers (SO32-, NO2-, SnO22- and others). In laboratories,
such reducing agents as H2, SO2, KI, H3PO3, H2S, HNO2 are usually used.
A substance containing an element that undergoes reduction is called an
oxidizing agent. These are atoms and molecules of some non-metals of high
activity (O2, O3, Cl2 and others) positively charged metallic ions (Fe3+, Cu2+, Hg2+
and others), particles containing ions in their highest oxidation numbers (MnO 4-,
NO3  , SO42-, Cr2O72-, ClO3- and others). The strongest oxidizing agent is electrical
current (oxidation on anode). In laboratories, such oxidizing agents as KmnO4,
K2Cr2O7, HNO3, H2SO4 (conc.), H2O2, PbO2 are used.
Elements in intermediate degrees of oxidation can show both properties of
oxidizers, and properties of reducers (Na2SO3, KNO2, etc.). See also
http://en.wikipedia.org/wiki/Redox
To balance redox-reactions, the half-reaction method is used. In acidic
media molecules of water and hydrogen-ions enter redox half-reactions. In alkaline

media both water molecules and OH ions are available. In neutral media the left
part of the half-equation contains water molecules and the right part contains either

H+ or OH ions.
Some examples of redox half-reactions:
Concentrated sulfuric acid
SO42  + 4H+ + 2e = H2SO3 + H2O
SO42  + 8H+ + 6e = S + 4H2O
Nitric acid

NO3 + 4H+ + 3e = NO + 2H2O

NO3 + 3H+ + 2e = HNO2 + H2O

NO3 + 2H+ + e = NO2 + H2O

NO3 + 10H+ + 8e = NH4+ + 3H2O
Manganese compounds

MnO4 + 8H+ + 5e = Mn2+ + 4H2O

MnO4 + 2 H2O + 3e = MnO2 + 4OH
MnO4

+ e = MnO42 
70

Chromium compounds
Cr2O72  + 14H+ + 6e = 2Cr3+ + 7H2O
CrO42  + 4H2O + 3e = Cr(OH)63  + 2OH

Hydrogen peroxide
H2O2 + 2e = 2OH

H2O2 + 2H+ + 2e = 2H2O
2H+ + O2 + 2e = H2O2
2H2O + O2 + 2e = H2O2 + 2OH

Direction of oxidation-reduction reaction
Oxidation-reduction properties of substance define on their _isinfect ability,
quantitatively expressed through redox-potential .
The standard redox-potential (is defined as the potential of a given
redox-system at concentrations (activities) of all the substances participating in the
electrode process equal unity.
The dependence of a redox-potential on concentrations of substances
participating in electrode processes and on temperature is expressed by the Nernst
equation:
[ Ox]
 = o + 2.3 RT
log
[Re d ]
nF
where
R – the molar gas constant; T – absolute temperature;
F – the
Faraday’s constant; n – number of electrons participating in the electrode process;
[Ox] – concentration of the oxidized form of a substance; [Red] – concentration of
the reduced form of a substance.
In case if T = 297 K (25oC),
 = o +
0.059
n
[ Ox ]
log [Re
d]
The more is the absolute value of redox potential, the stronger are oxidizing
properties of the oxidized form.
The less is the absolute value of redox potential, the stronger are reducing
properties of the reduced form.
71
The possibility of a redox-reaction can be determined from the electromotive
force of the reaction (E):
E = (ox) - (red)
In case if E > 0, the forward redox-reaction is possible. In case if E < 0, the
forward redox-reaction is impossible, and the reaction proceeds in the backward
direction.
Example. Determine the possibility of the reaction
2FeCl3 + 2KI  2FeCl2 + 2KCl + I2
The ferric ion Fe+3 is an oxidizing agent as it decreases its oxidation state
(Fe+3 + e  Fe+2).
Ite iodide ion I  is a reducing agent as it increases its oxidation state (2I  – 2
e  I2).
According to the table of redox-potentials,
0(ox) = 0(Fe+3/Fe+2) = 0.77 V
0(red) = 0(I2/2I  ) = 0.54 V
The electromotive force of the reaction:
E = 0(ox) -0(red) = 0.77 – 0.54 = 0.23 V (>0)
So, the given redox-reaction can proceed.
Classification of oxidation-reduction reactions
All oxidation-reduction reactions can be divided into three groups:
1). Reactions of intermolecular oxidation-reduction are reactions in which
the exchange of electrons occurs between atoms of different molecules: Fe +
CuSO4 = FeSO4 + Cu.
2). Reactions of intramolecular oxidation-reduction are such reactions
where both oxidizing and reducing agents are components of the same molecule:
2КclO3 = 2KCl + 3O2.
3). Reactions of _isinfection_e_g are those reactions in which the same
atom in a molecule acts simultaneously both as an _isinfect and a reducing agent:
3К2MnO4 + 2H2O = 2KmnO4 + MnO2 + 4KOH.
72
Balancing of oxidation-reduction reactions
(method of half-reactions)
To balance redox-reactions, the method of half-reactions is used. It states the
following:
1. The redox-reaction may be considered as a sum of two separate
processes: oxidation and reduction.
2. Molecules and ions of the medium are involved in the transfer of
electrons. This means that in acidic media molecules of water and
hydrogen-ions enter redox half-reactions. In alkaline media both water
molecules and OH- ions are available. In neutral media the left part of the
half-equation contains water molecules and the right part contains either
H+ or OH- ions.
3. Number of electrons gained by an oxidizing agent equals number of
electrons lost by a reducing agent.
Some half-reactions (transformation of an oxidized form into a reduced one)
with participation of the most typical oxidizing and reducing agents are presented
below. Half-reactions for oxidizing forms should be written as presented, as far as
for reducing agents they should be presented in a back direction, i.e. from right to
left.
Concentrated sulfuric acid

SO42 + 8H + + 6e = S + 4H2O

SO42 + 4H + + 2e = H2SO3 + H2O
Nitric acid

NO3 + 10H + + 8e = NH4 + + 3H2O

NO3 + 4H + + 3e = NO + 2H2O

NO3 + 3H + + 2e = HNO2 + H2O

NO3 + 2H + + e = NO2 + H2O
Manganese compounds

MnO4 + 8H + + 5e = Mn2 + + 4H2O

MnO4 + 2 H2O + 3e = MnO2 + 4OH
73


MnO4 + e = MnO42

Chromium compounds

Cr2O72 + 14H + + 6e = 2Cr3 + + 7H2O


CrO42 + 4H2O + 3e = Cr (OH) 63 + 2OH

Hydrogen peroxide
H2O2 + 2e = 2OH

H2O2 + 2H + + 2e = 2H2O
2H + + O2 + 2e = H2O2
2H2O + O2 + 2e = H2O2 + 2OH

Example 1.
Ca + HNO3 (dil) NH4NO3 + ...
Ca0 – 2e Ca2+
4
NO3 + 10H+ + 8e NH4+ + 3H2O
1
4Ca0 + NO3 + 10H+ 4Ca2+ + NH4+ + 3H2O
4Ca + 10HNO3 (dil) 4Ca(NO3)2 + NH4NO3 + 3H2O
Example 2.
KmnO4 + Na2SO3 + H2SO4 
MnO4- + 8Н+ + 5e  Mn+2 + 4Н2О
x2
SO32- + Н2О – 2e  SO42- + 2Н+
x5
2MnO4- + 6Н+ + 5SO32-  2Mn+2 + 3Н2О + 5SO422KmnO4 + 5Na2SO3 + 3H2SO4 = 2MnSO4 + 5Na2SO4 + K2SO4 + 3H2O
74
QUESTIONS AND PROBLEMS
1. Complete the equations of the following reactions and balance them:
(a) K2S + KmnO4 + H2SO4 = S + ....
(b) KI + K2Cr2O7 + H2 SO4 = I2 + ...
( c) K MnO4 + H2O2 = ...
2. Indicate the direction in which the following reactions can proceed
spontaneously:
(a) H2O2 + HclO = HCl + O2 + H2O
(b) H3PO4 + 2HI = H3PO3 + I2 + H2O
3. Can a salt of iron (III) be reduced to a salt of iron (II) in an aqueous
solution by (a) potassium bromide, (b) potassium iodide?
4. Using the table of standard electrode potentials, calculate the equilibrium
constants for the following reactions:
(a) Zn + CuSO4 = Cu + ZnSO4
(b) Sn + Pb(CH3COO)2 = Sn(CH3COO)2 + Pb
15. COMPLEX COMPOUNDS
It is possible to define complex compounds from point of view of their
different features. See also http://en.wikipedia.org/wiki/Coordination_complex
1.
Complex compounds are definite chemical compounds formed by a
combination of individual components without formation of new pairs of
electrons.
Example: 3NaOH + Al(OH)3  Na3[Al(OH)6]
2. Complex compounds are definite chemical compounds in which valence
states of some chemical elements do not equal their oxidation states.
Example: in Na3[Al(OH)6] oxidation state of Al is +3 as its valence state
is six (Al forms chemical bonds with 6 hydroxo-groups)
3. Complex compounds are definite chemical compounds where a coordinate type of a co-valent bond is realized.
75
Example: lone electronic pairs of OH– anions (donors of electrons)
overlap with empty orbitals of the valence shell of Al 3+ cations (acceptor
of electrons).
Structure of complex compounds
In a molecule of a complex compound, one of the atoms, generally
positively charged, occupies the central site (central ion or complexing agent).
Oppositely charged ions or neutral molecules called ligands are coordinated
around the central ion. The complexing agent and ligands form inner sphere of a
complex compound. It is characterized by coordinate bonds which are formed
while overlapping of empty p- and d-orbitals of a central ion and orbitals
containing lone electron pairs of ligands. The ions in the outer sphere are mainly
bonded to the complex ions by forces of electrostatic interaction (ionic bonds).
The total number of coordinate bonds formed by the complexing agent is
known as coordination number of the central ion. It mainly depends upon the
charge of the complexing agent (for monocharged ions it usually equals 1, for
discharged ions – 4 or 6, for tricharged – 6 and above), and the size of an ion (the
larger the central ion, the greater its coordination number is, for lanthanides and
actinides it can reach to 12)..
Cisplatin, PtCl2(NH3)2 A platinum atom with four ligands.
Ligands possess the property of dentation. In accordance with the number of
coordinate bonds formed by a ligand with the central ion, the ligand may be a
monodental, bidental, or polydental. Dentation is defined by number lone
electronic pairs in a molecule of a ligand and their mutual spatial disposition. For
example, the ammonia molecule NH3 has one lone electronic pair belonging to the
N-atom therefore ammonia is a monodental ligand. The water molecule has two
76
lone electronic pairs, and chloride-ion has them four. However because valence
orbitals of oxygen and chlorine are in the sp3-type of hybridization and are located
under a corner 10928 ’ they cannot form chemical bonds with the same central
ion, therefore such лиганды are monodental except the case of polynuclear
complexes where they act as bidental bridge ligands.
Example: Na2 [Cu2Cl6]
Cl
Na2
Cl
Cl
Cu
Cu
Cl
Cl
Cl
Dicharged acidic anions such as SO32–, C2O42– ets. Are usually bidental
chelating ligands.
Example: Na2[Cd(SO3)2]
O
Na2
O
O S
Cd
O
S O
O
Example. For a complex compound K3[Fe(CN)6]
1. Ions of the outer sphere are 3К+
2. Ion of the inner sphere (complex ion) is [Fe(CN)6]3–
3. Central ion (complex-forming ion) is Fe3+
4. Coordination number of the central ion is 6.
5. Ligands are 6CN– anions, both are monodental.
Nomenclature of complex compounds
Names of complex compounds are similar to the names of simple salts. The
order of naiming particles in a complex ion is the following: anionic ligands –
neutral ligands – central ion. Number of ligands is designated with the help of
greek numerals.
Examples:
[Cu(NH3)4]Cl2 – tetraammine copper(II) chloride;
K2 [Cu(OH)4] – potassium tetrahydroxocupprate(II);
[Cr(NH3)3Cl3] – trichloro triammine chromium(III).
Classification of complex compounds
77
There are several types of classification of complex compounds.
1. Depending upon a charge of the inner sphere:
(i)
Cationic complexes (the inner sphere is positively charged –
complex cations). Examples: [Cr(H2O)6]Cl3, [Co(NH3)6]Cl3.
(ii)
Anionic complexes (the inner sphere is negatively charged –
complex anions). Examples: K2[HgI4], Na[Sb(OH)6].
(iii)
Neutral complexes (the inner sphere is not charged).
Examples: [Pt(NH3)2Cl2], [Fe(CO)5].

Depending upon the type of the ligand:
(i)
Aqua-complexes
(ligands
are
water
molecules
–
[Cu(H2O)5]SO4).
(ii)
Ammino-complexes (ligands are molecules of ammonia or
organic ammines – [Ag(NH3)2]Cl).
(iii)
Hydroxy-complexes
(ligands
are
OH–
anions
–
Na2[Sn(OH)4]).
(iv)
Carbonyl-complexes (ligands are molecules of carbon
monoxide – [Fe(CO)5]).
(v)
Acido-complexes (ligands are anions of inorganic acids).
Examples:
chlorocomplexes
K2[HgCl4],
fluorocomplexes
K3[FeF6], cyanocomplexes Kfe[Fe(CN)6], thiocyanocomplexes
K3[Fe(SCN)6], sulphitocomplexes K[Ag(SO3)], etc.
3. Depending upon the nature of a central ion: complexes of copper,
silver, iron, chrome etc.
4. Isomerism
 The arrangement of the ligands is fixed for a given complex, but in some
cases it is mutable by a reaction that forms another stable isomer.
 There exist many kinds of isomerism in coordination complexes, just as in
many other compounds.
78
Stereoisomerism
 Stereoisomerism occurs with the same bonds in different orientations relative
to one another. Stereoisomerism can be further classified into:
Cis–trans isomerism and facial–meridional isomerism
 Cis–trans isomerism occurs in octahedral and square planar complexes (but
not tetrahedral). When two ligands are mutually adjacent they are said to
be cis, when opposite each other, trans. When three identical ligands occupy
one face of an octahedron, the isomer is said to be facial, or fac. In
a fac isomer, any two identical ligands are adjacent or cis to each other. If
these three ligands and the metal ion are in one plane, the isomer is said to be
meridional, or mer. A mer isomer can be considered as a combination of
a trans and acis, since it contains both trans and cis pairs of identical ligands.
cis-[CoCl2(NH3)4]+
trans-[CoCl2(NH3)4]+
fac-[CoCl3(NH3)3]
79
mer-[CoCl3(NH3)3]
Optical isomerism
 Optical isomerism occurs when a molecule is not superposable with its
mirror image. It is so called because the two isomers are each optically
active, that is, they rotate the plane of polarized light in opposite directions.
The symbol Λ (lambda) is used as a prefix to describe the left-handed
propeller twist formed by three bidentate ligands, as shown. Likewise, the
symbol Δ (delta) is used as a prefix for the right-handed propeller twist.[8]
Λ-[Fe(ox)3]3−
Δ-[Fe(ox)3]3−
Λ-cis-[CoCl2(en)2]+
80
Δ-cis-[CoCl2(en)2]+
Structural isomerism
Structural isomerism occurs when the bonds are themselves different. There are
four types of structural isomerism: _isinfecti isomerism, solvate or hydrate
isomerism, linkage isomerism and coordination isomerism.
1. Ionisation isomerism – the isomers give different ions in solution although
they have the same composition. This type of isomerism occurs when the
counter ion of the complex is also a potential ligand. For example
pentaaminebromidocobalt(III)sulphate [Co(NH3)5Br]SO4 is red violet and in
solution gives a precipitate with barium chloride, confirming the presence of
sulphate ion, while pentaaminesulphatecobalt(III)bromide
{Co(NH3)5SO4]Br is red and tests negative for sulphate ion in solution, but
instead gives a precipitate of AgBr with silver nitrate.
2. Solvate or hydrate isomerism – the isomers have the same composition but
differ with respect to the number of solvent ligand molecules as well as the
counter ion in the crystal lattice. For example [Cr(H2O)6]Cl3 is violet
colored, [Cr(H2O)5Cl]Cl2·H2O is blue-green, and [Cr(H2O)4Cl2]Cl·2H2O is
dark green
3. Linkage isomerism occurs with ambidentate ligands that can bind in more
than one place. For example, NO2 is an ambidentate ligand: It can bind to a
metal at either the N atom or an O atom.
4. Coordination isomerism – this occurs when both positive and negative ions
of a salt are complex ions and the two isomers differ in the distribution of
ligands between the cation and the anion. For example
[Co(NH3)6][Cr(CN)6] and [Cr(NH3)6][Co(CN)6]
81
Dissociation of complex compounds and complex ions
Majority of complex compounds are electrolytes. In solutions they dissociate
and form both simple and complex ions (outer and inner spheres). This type of
dissociation is an irreversible process (complex compounds are strong
electrolytes).
[Cu(NH3)4]SO4 → [Cu(NH3)4]2+ + SO42–
(primary dissociation)
The inner sphere of a complex compound undergoes the process of
secondary dissociation and dissociates reversibly and stepwise to split into central
ion and ligands (complex ions are weak electrolytes):
[Cu(NH3)4]2+ → [Cu(NH3)3]2+ + NH3
[Cu(NH3)3]2+ → [Cu(NH3)2]2+ + NH3
[Cu(NH3)2]2+ → [Cu(NH3)]2+ + NH3
[Cu(NH3)]2+ → Cu2+ + NH3
The overall process: [Cu(NH3)4]2+ → Cu2+ + 4NH3
Each of the above processes can be characterized by an equilibrium constant
(stepwise instability constants of a complex ion). The equilibrium constant of the
overall process is called an overall instability constant of a complex ion:
K inst 
[Cu 2 ]  [ NH 3 ] 4
2
[Cu( NH 3 ) 4 ]
Instability constants of complex ions are table values. The less the value of
the overall instability constant, the more stable the complex ion is.
The equilibrium constant of reaction of formation of a complex ion carries
the name of a constant of stability, Kst.:
Cu2+ + 4NH3 → [Cu(NH3)4]2+
2
K st 
[Cu( NH 3 ) 4 ]
1

2
4
K inst
[Cu ]  [ NH 3 ]
Example 1. Calculate concentration of Ag+ cations in 0.1 M solution of
[Ag(NH3)2]Cl.
Primary dissociation does not produce Ag+ cations:
82
[Ag(NH3)2]Cl → [Ag(NH3)2]+ + Cl–
0.1 M
0.1 M
The complex ion undergoes the secondary dissociation:
[Ag(NH3)2]+ → Ag+ + 2NH3
K inst 
[ Ag  ]  [ NH 3 ] 2

[ Ag ( NH 3 ) 2 ]
 6.1  10 8
Suppose the concentration of Ag+ ions to be x, so the concentration of
ammonia is 2x.
6.1  10 8 
6.110–9 = 4x3
x  (2 x) 2
;
0.1
x3
6.1  10 9
 10 3
4
Example 2. Calculate concentration of Ag+ cations in 0.1 M solution of
[Ag(NH3)2]Cl in excess of 1M ammonia.
[Ag(NH3)2]Cl → [Ag(NH3)2]+ + Cl–
0.1 M
0.1 M
[Ag(NH3)2]+ → Ag+ + 2NH3
0.1 M
K inst 
[ Ag  ]  [ NH 3 ] 2

[ Ag ( NH 3 ) 2 ]
x

2x +1  1
x 1
 6.1  10 8
0.1
x = 6.110–9
Reactions of complex compounds

Formation of complex compounds
CuSO4 + 4NH3  [Cu(NH3)4]SO4
AgCl + 2NH3  [Ag(NH3)2]Cl
Al(OH)3 + 3NaOH Na3[Al(OH)6]
CdCl2 + 2Na2SO3(excess) Na2[Cd(SO3)2] + 2NaCl

Destruction of complex compounds
[Cu(NH3)4]SO4 + 4HNO3  CuSO4 + 4NH4NO3
Na3[Al(OH)6] + 3HCl (lack)  Al(OH)3 + 3NaCl + 3H2O
Na3[Al(OH)6] + 6HCl (excess)  AlCl3 + 3NaCl + 6H2O
Na2[Cd(SO3)2] + 2NaOH  Cd(OH)2 + 2Na2SO3
83
(Complex ions are unstable in presence of strong acids and the bases)
CuS  + (NH4) 2SO4 + (NH4)2S
[Cu(NH3)4]SO4 + H2S
(The direction of this reaction can be defined by means of calculation of an
equilibrium constant. Approximately it equals 10–23, so hydrosulphuric acid cannot
destroy tetraammine coppric cation)

Exchange reactions of the outer sphere
K[Sb(OH)6] + NaCl  Na[Sb(OH)6]  + KCl
(Reactions of this type are possible in case of formation of precipitates)

Exchange reactions in the inner sphere
Exchange of ligands:
K3[Fe(SCN)6] + 6KCN
K3[Fe(CN)6] + 6KSCN
Exchange of the central ion:
K2 [SnCl4] + CuCl2
K2 [CuCl4] + SnCl2
(Equilibria of these reactions are shifted towards formation of a more stable
complex ion)
Electronic structure of complex ions
Interaction of lone electronic pairs of ligands with empty valence orbitals of
the central ion of different types leads to their hybridization. For example, the
electronic structure of a complex ion [Cu(NH3)4]2+ can be reflected as following:
29Cu:
29Cu
1s22s22p63s23p63d104s1
2+
: 1s22s22p63s23p63d94s04p04d0
    
3d
4s
4p
4d
4 :NH3
Interaction of one s- and three p-orbitals leads to the sp3-hybridization of the
central ion (tetrahedral complex).
Entering of lone electronic pairs of ligands into valence orbitals of the
central ion leads to their interaction with the electrons of 3d-orbitals. This
interaction is defined by degree of penetration of electrons of ligands on empty
84
orbitals of metallic cations. In connection of force of interaction, ligands may be
arranged in a spectrochemical series and are devided into ligands of weak and
strong field:
CO
CN– > NO2–
> SCN– > H2O > OH–
> NH3
Ligands of a strong field
> F– > Cl–
Ligands of a weak field
Lone electronic pairs of ligands of a strong field deeply enter the valence
electronic shell of the central ion and cause pairing of electrons of the 3d-subshell.
Are as a result intraorbital lowspin complexes are formed (coordination bonds are
formed with participation of internal 3d-orbitals, the formed complex has no or few
unpaired electrons):
27Co:
27Co


1s22s22p63s23p63d74s2
3+

: 1s22s22p63s23p63d64s04p04d0


3d
4s
4p
4d
[Co(NH3)6]3+
  
3d
4s
4p
4d
6 :NH3
d2sp3-hybridization, octahedral complex.
Lone electronic pairs of ligands of a weak field slightly interact with 3delectrons of the central ion and do not cause their pairing. Are as a result, outerorbital high-spin complexes are formed::
[CoF6]3–
 

3d


4s
4p
4d
6 :F-
sp3d2-hybridization, octahedral complex.
85
QUESTIONS AND PROBLEMS
1. For the following complex compounds indicate: (a) inner and outer spheres; (b)
central ion, its charge and coordinating number; (c) ligands ; (d) write the
reaction of dissociation of complex compounds and complex ions; (e) express
overall instability constants; (f) name the following complex compounds.
[Cd(NH3)4]Cl2 ; K2[Cd(CN)4]
2. Which of the above mentioned complexes is more stable?
3. Calculate the concentration of the Ag+ ions in a 0.1 M solution containing an
excess of 1 mole  l-1 of NH3.
4. In which case will a reaction occur between solutions of the electrolytes
indicated below (exchange of ligands)? Write the equations of these reactions
in molecular and net ionic forms and calculate their equilibrium constants:
(a) K2[HgI4] + KCN =
(b) K[Ag(CN)2] + NH3 =
16. SURWAY OF PROPERTIES OF SOME CHEMICAL ELEMENTS
ALKALINE and ALKALINE EARTH METALS
Alkaline metals Li, Na, K, Rb, Cs, Fr are situated in the IA group of the
periodic table. The outer spheres of these atoms consist of only one electron. So
atoms of alkaline metals have a tendency to lose their valence electrons to be
transformed into positively charged ions. Their oxidation number is +1. The
strength of the attraction of the outer electron to the atom can be valued with the
help of the ionization potential which determines the energy of removal of one
electron from a neutral atom. In the IA group from Li to Fr the value of ionization
potential decreases, the chemical activity of metals increases. Alkaline metals are
strong reducing agents.
In air alkaline metals are easily oxidized, that is why they are stored under
oil.
4Li + O2  2Li2O
2Na + O2  2Na2O2
K + O2  KO2
Peroxides and superoxides are strong oxidizing agents. They are decomposed by
water:
Na2O2 + 2H2O  2NaOH + H2O2
86
2KO2 + 2H2O  2KOH  H2O2 + O2
Alkaline metals are more active than hydrogen (they have negative values of
redox-potentials) so they can replace hydrogen both from acids and water:
2M + 2HCl = 2MCl + H2
2M + 2H2O = 2MOH + H2
In aqueous solutions metal hydroxides behave as strong electrolytes and are
MOH → M+ + OH
fully dissociated:
Almost all salts of alkaline metals are water soluble. Solutions of salts
containing anions of weak acids undergo hydrolysis; they are basic.
Pure alkaline metals are produced by electrolysis of melted salts.
Biological significance of alkaline metals. Lithium is found in the liver
and lungs of animals. Large concentrations of lithium are dangerous for humans.
Particles of dust and smoke containing lithium provoke malignant tumors. Sodium
as NaCl is necessary for the balance of salt exchange in organisms, sodium
bicarbonate is used to lower the acidity of gastric juice. In living organisms
potassium is situated in liver, spleen, it regulates the function of muscle cells and
nervous systems. Rubidium is found in the leaves of plants (beetroot, sugar-cane,
tabacoo, tea, coffee, cocoa). In animal organisms it is localized in muscles
performing large load – heart muscle and pectoral muscles of birds. Cesium is
found in mineral water, plants and living organisms. Compounds of rubidium and
cesium are necessary for the growth of plants.
Alkaline earth metals Ca, Sr, Ba, Ra are situated in the IIA group of the
periodic table. Atoms of these elements have two valence electrons. While losing
them, atoms of alkaline earth elements transfer into positively charged ions and
gain the oxidation number +2.
The presence of non-filled sublevels (d- and f-) makes Ca, Sr, Ba and Ra
more chemically active and having physical properties rather different than those
of Be and Mg.
87
Metals of IIA group are less active than alkaline metals. The growth of the
atomic radius, lowering of the ionization potential and of the electronegativity
define the growth of their chemical activity with increase of charges of the nuclei.
In the air alkaline earth metals easily form oxides:
2Ca + O2  2CaO
2Mg + O2
t
2MgO
They replace hydrogen both from water and from acids:
M +2HCl = MCl2 + H2
M + 2H2O = M(OH)2 + H2
The basic character of oxides and hydroxides increases with the increase of
atomic radii from calcium to radium.
BeO + H2O 
BeO + 2HCl  BeCl2 + H2O
t
BeO + 2NaOH
MgO + H2O 
Na2BeO2
MgO + 2HCl  MgCl2 + H2O
CaO + H2O  Ca(OH)2
MgO + NaOH 
CaO + CO2  CaCO3
CaO + 2HCl  CaCl2 + H2O
Alkaline earth metals form different sparingly soluble salts: carbonates,
phosphates, chromates, sulfates. The solubility of sulfates lowers from calcium to
radium.
CaCl2 + Na2CO3  CaCO3+ 2NaCl
CaCO3 + 2HCl  CaCl2 + H2O + CO2
Ca(HCO3)2
t
CaCO3 + H2O + CO2
Natural water containing soluble salts of calcium and magnum is called hard
water. The presence of hydrocarbonates of calcium and magnum stipulates the
temporary hardness of water. Chlorides and sulfates of these elements cause the
constant hardness. The sum of temporary and constant hardness gives the overall
hardness of water.
Biological and agricultural properties of elements of IIA group.
Beryllium and its compounds are toxic. The beryllium-poisoning may cause the
death.
88
Magnum compounds can be found in algae, fungi, ferns, in the tissues of
animals. Magnum is a complexing ion in chlorophyll.
Calcium is a constituent of the bones of vertebrates, predominantly as
_isinfection_e Ca3(PO4)2. Egg-shells, tests of sea animals, shells mainly consist of
calcium carbonate. Organic salts of calcium play a significant role in metabolism
of plants. The deficiency in calcium leads to stopping of their growth, development
of rhizomes, the leaves cover by brown spots and die off. The animals suffer from
rachitis, the heart activity decreases, the blood coagulability becomes worse.
Calcium ions enter human organisms with milk and meat meals while magnum
ions – with vegetable meal.
Strontium compounds in human organisms are mainly concentrated in
bones, the excess (more than 10-3 %) leads to their fragile.
There are approximately 100 times less barium compounds in human bodies
than those of strontium. In very small quantities barium compounds stimulate
activity of marrow. In large quantities they are very toxic and provoke weakness,
gastric-intestinal diseases, brain disorders. Barium chloride and carbonate are used
in the agriculture as chemical weed-killers and pest-killers.
QUESTIONS AND PROBLEMS
1. How does the value of the ionization potential changes with the change of the
position of an element in the periodic table?
2. Which are the rules of storage and handling of alkaline and alkaline earth
metals?
3. How are compounds of alkaline and alkaline earth metals are used in the
medical practice?
4. Calculate pH of 0.01 M solution of sodium acetate.
5. Calculate the equilibrium constants of the reactions and explain why the
precipitate of barium chromate can be dissolved in hydrochloric acid and can’t
be dissolved in acetic acid.
89
ELEMENTS OF IIIA AND IVA GROUPS
(P – ELEMENTS)
Boron and aluminum are elements of the IIIA group of the periodic table.
Atomic radius of boron is 0.91Å and the one of aluminum is 1.43Å. This great
difference affects chemical properties of these elements. Ionization potential of
boron is greater than that of aluminum. Polarity of B-O chemical bond is small, so
in solutions boron exists as BO2- and BO33- ions (acidic properties). Al-O chemical
bonds have a more polar character, so in solutions aluminum exists both as Al 3+
and AlO2- ions (amphoteric properties).
4B + 3º2  2B2O3
B2O3 + 3H2O  2H3BO3
(К1 = 61010)
4Al + 3º2  2Al2O3
2Al + 6H2O  2Al(OH)3 + 3H2
2Al + 6HCl  2AlCl3 + 3H2
2Al + 6NaOH + 6H2O  Na3[Al(OH)6] + 3H2
Al2O3 + H2O 
Al2O3 + 6HCl  2AlCl3 + 3H2O
Al2O3 + 2NaOH
t
2NaAlO2 + H2O
AlCl3 + 3NH4OH  Al(OH)3 + 3NH4Cl
AlCl3 + 3NaOH  Al(OH)3 + 3NaCl
Al(OH)3 + 3NaOH  Na3[Al(OH)6]
Al(OH)3 + 3HCl  AlCl3 + 3H2O
Salts of boric acid H3BO3 are metaborates (Ba(BO2)2) and tetraborates
(Na2B4O7 – borax).
Alluminium salts undergo hedrolysis. Some of them (Al 2S3, Al2(CO3)3) are
fully decomposed by water: Al2S3+6H2O2Al(OH)3+3H2S
90
Biological activities of boron and aluminum. Physiological activity of
boron is rather high. Together with Mn, Cu, Zn and Mo it is among five most
important microelements. It concentrates in bones, teeth, muscles, marrow, liver
and thyroid gland, can be found in adipose tissues of some animals, in milk and
yolk of eggs.
Boron inhibits the action of amilaze and proteinaze, vitamins B 2 and B12,
reinforces the action of insuline. Boric acid and borax are used in medicine as
anticeptics.
Some compounds of aluminum are also used in medicine: Kal(SO4)2 as
astringent; AlOH(CH3COO)2 for _isinfection; Al2(SO4)3 as coagulant.
Carbon and silicon are elements of IVA group of the periodic table of
elements. Their highest oxidation number is +4.
At a room temperature carbon and silicon are inert elements, their activities
increase with heating. At high temperatures they react with the majority of nonmetals and metals.
С + О2  СО2
2С + О2  2СО
Concentrated nitric and sulfuric acids oxidize carbon into CO2, silicon can
be oxidized by mixture of HNO3 and HF. Silicon can also be dissolved in alkalis:
Si + 2NaOH Na2SiO3 + 2H2
Carbon (II) oxide CO is a non-salt forming oxide and a strong reducing
agent:
3СO + Fe2O3
CO + CuO
t
t
2Fe + 3CO2
Cu + CO2
Carbon dioxide CO2 possesses acidic properties and reversibly dissolves in
water to form a weak carbonic acid:
H2CO3 H+ + HCO3-
CO2 + H2O
Salts of carbonic acid are carbonates and hydrocarbonates:
91
CO2 + Ca(OH)2  CaCO3 + H2O
CaCO3 + H2O  Ca(HCO3)2
In aqueous solutions carbonates and hydrocarbonates undergo hydrolysis:
Na2CO3 + H2O
NaHCO3 + NaOH
Silicic acid is even weaker than carbonic acid, the reactions of hydrolysis of
its salts lead to the formation of polyanions:
2Na2SiO3 + H2O
Na2Si2O5 + 2NaOH
Tin and lead possess metallic properties. They may be +2 and +4 charged.
Their hydroxides have amphoteric character:
SnCl2 + 2NaOH  Sn(OH)2 + 2NaCl
Sn(OH)2 + 2HCl  SnCl2 + 2H2O
Sn(OH)2 + 2NaOH  Na2[Sn(OH)4]
SnCl4 + 4NH4OH  H2SnO3 + 4NH4Cl + H2O
H2SnO3 + 2NaOH + H2O  Na2[Sn(OH)6]
H2SnO3 + 4HCl  SnCl4 + 3H2O
Acidic properties increase with the increase in the oxidation states of metals.
Sn2+ compounds are strong reducing agents as far as compounds containing
Pb4+ are strong oxidizing agents:
SnCl2 + 2FeCl3  2FeCl2 + SnCl4
PbO2 + 4HCl  PbCl2 + Cl2 + 2H2O
QUESTIONS AND PROBLEMS
1. Which are oxidation states of elements of IIIA and IVA groups of the periodic
table? Write their electronic structures and mark valence electrons.
2. Which are chemical properties of oxides of boron, aluminum, carbon and
silicon. Using table data, compare strengths of corresponding acids.
3. Prove amphoteric properties of aluminum hydroxide.
4. Calculate pH of 0.1 M solution of NaHCO3.
5. Calculate the equilibrium constant of a reaction and explain if the precipitate of
calcium carbonate can be dissolved in acetic acid.
92
ELEMENTS OF VA AND VIA GROUPS
(P-ELEMENTS)
Nitrogen and phosphorus are elements of VA group of the periodic table.
They have 5 electrons in the outer shell (valence electrons). They may lose
electrons and form positively charged ions (oxidation numbers from +1 to +5) or
gain electrons and form negatively charged ions (oxidation number -3).
Hydrogen compounds of nitrogen and phosphor are ammonia NH 3 and
phosphine (PH3):
N2 + 3H2  2NH3
The presence of a lone electron pair of nitrogen and phosphor leads to the
possibility of formation of a coordinate bond with a proton. In aqueous solutions
ammonia interacts with water molecules to form ammonium hydroxide which
possesses weak basic properties:
NH3 + H2O
NH4OH
NH4+ + OH-
PH3 + HCl  PH4Cl
Phosphonium salts are not stable in aqueous solutions:
PH4+ + H2O  PH3 + H3O+
Ammonium ion has almost same properties as metallic ions, for example it
can form salts. Ammonium salts decompose while heating:
NH4Cl
NH3 + HCl
In oxides nitrogen has various oxidation states from +1 to +5.
4NH3 + 5º2
Pt , t
4NO + 6H2O
2NO + O2  2NO2
4NO2 + O2 + 2H2O  4HNO3
Nitric acid HNO3 is one of the strongest acids with a high oxidizing
strength. Depending on the nature of a reducing agent and the concentration of the
acid, the NO3- group can gain from 1 to 8 electrons and transfer into NO2, NO,
N2O, N2 or NH4+:
Cu + 4HNO3 (conc.)  Сu(NO3)2 + 2NO2 + 2H2O
4Ba + 10HNO3 (conc.)  4Ba(NO3)2 + N2O + 5H2O
93
3Cu + 8HNO3 (dil.)  3Сu(NO3)2 + 2NO + 4H2O
4Ba + 10HNO3 (dil.)  4Ba(NO3)2 + NH4NO3 + 3H2O
6HNO3 (conc.) + S  6NO2 + H2SO4 + 2H2O
5HNO3 (dil.) + 3P + 2H2O  5NO + 3H3PO4
Salts of nitric acid (nitrates) are water soluble and also possess oxidizing
2KNO3 + C  2KNO2 + CO2
properties:
Nitrous acid HNO2 is a weak acid with redox duality:
2NaNO2 + 2KI + 2H2SO4  I2 + 2NO + K2SO4 + Na2SO4 + 2H2O
(oxidizing agent)
5NaNO2 + 2KmnO4 + 3H2SO4  2MnSO4 + NaNO3 + K2SO4 + 3H2O
(reducing agent)
Nitrous acid exists only in diluted solutions and decomposes at high
concentrations:
3HNO2 →HNO3 + 2NO + H2O
In contrast with nitric acid, phosphoric acid has no oxidizing properties.
Phosphates form soluble complexes with a lot of metal ions.
4Р + 5О2  2Р2О5
P2O5 + H2O  2HPO3
t
P2O5 + 3H2O
Orthophosphoric
acid
forms
2H3PO4
phosphates,
hydrophosphates
and
dihydrophosphates: Na3PO4, Na2HPO4 and NaH2PO4.
Sulfur is situated in the VIA group of the periodic table and has 6 valence
electrons. Its oxidation numbers are +4 , +6 and –2.
Hg + S  HgS
S + O2  SO2
2SO2 + О2  2SO3
Sulphide-anion possesses redicing properties:
2KmnO4 + 3H2SO4 + 5H2S  2MnSO4 + 5S + K2SO4 + 8H2O
Majority of sulphides are not soluble in water and form precipitates of
various colors: ZnS (white), MnS (pink), CdS (yellow), Sb 2S3 (orange), SnS
(brown), CuS (black).
Compounds of tetravalent sulphur possess redox-duality:
H2SO3 + 2H2S  3S + 3H2O
94
(oxidizing agent)
H2SO3 + Cl2 + H2O  H2SO4 + 2HCl
(reducing agent)
Sulphorous acid is a weak dibasic acid not stable in acidic media:
SO2 + H2O  H2SO3
H2SO3
H+ + HSO3–
K1 = 210–2
HSO3–
H+ + SO32–
K2 = 610–8
H2SO3  SO2 + H2O
Sulphites undergo hydrolysis: Na2SO3+HOH
NaHSO3+NaOH
Sulfuric acid H2SO4 (conc.) is a strong acid and oxidizing agent:
Сu + 2H2SO4 (conc.)  СuSO4 + SO2 + 2H2O
3Zn + 4H2SO4 (conc.)  3ZnSO4 + S + 4H2O
4Сa + 5H2SO4 (conc.)  4СaSO4 + H2S + 4H2O
2H2SO4 (conc.) + S  3SO2 + 2H2O
2H2SO4 (conc.) + C  2SO2 + CO2 + 2H2O
Diluted sulphuric acid is a strong acid with no oxidizing properties:
Zn + H2SO4 (dil.)  ZnSO4 + H2
Cu + H2SO4 (dil.) 
Biological importance of sulfur. Sulfur is a composite of most important
aminoacids. Metal sulfates are used in medicine: CaSO4 as a stuff for plasters,
BaSO4
in rontgenoscopy of stomach, MgSO410H2O as purgative. Some
antibiotics have sulfur compounds as composites.
QUESTIONS AND PROBLEMS
1. Describe the electronic structure of ammonia and NH4+ ion.
2. Using the molecular orbital method draw molecular diagrams of N 2 and O2
molecules. Compare stabilities of these molecules.
3. Calculate pH of 0.01 M solution of NH4OH.
95
4. Calculate the solubility of Ag2SO4: (a) in pure water; (b) in presence of 0.1 M
solution of sulfuric acid.
5. Compare solubilities of MnS and CuS in diluted hydrochloric acid.
ELEMENTS OF THE VIIA GROUP
(HALOGENS)
Atoms of halogens have 7 electrons in the outer shell (ns 2np5) which
determine their chemical activity. Halogens are strong oxidizing agents. Their
activities increase with decrease in ionic radii: fluorine is the strongest oxidizing
agent:
F2 + H2  2HF
2F2 + 2H2O  4HF + O2
F2 + 2NaCl  2NaF + Cl2
The properties of fluorine differ from those of other halogens. As its atoms
have no empty d-orbitals in the outer sphere, it can’t exist in excited states, and its
only oxidation number is –1.
Atoms of chlorine, bromine and iodine have a vacant d-orbital in their outer
spheres, so 3 electrons may be unpaired to form 3 excited
states. Possible
oxidation numbers of these elements are -1, +1, +3, +5, +7.
Bond energies in the molecules of halogens increase with decrease in atomic
numbers: ECl-Cl > EBr-Br > EI-I.
ECl-Cl = 57.8 kcal/mol;
EBr-Br = 46.1 kcal/mol;
EI-I = 36.2 kcal/mol
The molecule of fluorine has the minimal bond energy which can be
explained by the features of electronic configuration of fluorine comparing with
other halogens.
Hydrogen compounds of halogens are colorless gases. The bond energies
decrease from HF to HI. Their aqueous solutions possess acidic properties, HI is
the strongest one among them. Reducing properties of halogen hydrides increase
with increase of charge of nuclei:
2KmnO4 + 16HCl  2KCl + 2MnCl2 + 5Cl2 + 8H2O
10KI + 8H2SO4 + 2KmnO4  5I2 + 2MnSO4 + 6K2SO4 + 8H2O
96
Oxygen containing compounds of halogens are strong oxidizing agents.
HclO  HCl + [O]
Interactions of halogens with water can be expressed as:
F2 + H2O → 2HF + [O]
[O] + F2 → F2O
X2 + H2O
HX + HOX (X = Cl, Br, I)
The equilibria of these interactions are shifted to the left. In alkaline
solutions the reactions become irreversible:
Cl2 + H2O  HCl + HclO
2Cl2 + 2Ca(OH)2  CaCl2 + Ca(ClO)2 + 2H2O
HCl, HBr and HI are strong acids whose strengths increase from HCl to HI
because of increasing polarizability of anions. Salts of majority of metals are
water soluble exept Ag-halogenides:
AgNO3 + NaCl  AgCl + NaNO3
AgNO3 + NaBr  AgBr + NaNO3
AgNO3 + NaI  AgI + NaNO3
QUESTIONS AND PROBLEMS
2. Can chlorine, bromine and iodine have an oxidation number +2 ? Explain your
answer.
3. Using the molecular orbitals method compare the stabilities of a molecule F 2

and a molecular ion F2 .
4. Write the formulae of all possible oxides of chlorine and corresponding acids.
Compare the strengths of the acids.
5. Which oxygen containing compounds of chlorine are used in medicine?
6. Which of the following reactions can proceed spontaneously in neutral aqueous
solutions:


a) MnO4 + Cl  MnO2 + Cl2


b) MnO4 + Br  MnO2 + Br2

c) MnO4 + I
97

 MnO + I
2
2
TRANSITIONAL ELEMENTS
(d-ELEMENTS)
d-Elements are situated in B-subgroups of the periodic table. Their valence
electrons are those of s-sublevel of the outer shell and of the unfilled d-sublevel.
The presence of 1 or 2 electrons on the outer shell of all d-elements stipulates their
metallic properties. D-Electrons take part in the formation of chemical bonds, so
different oxidation states are known for d-elements. High values of oxidation
numbers are typical only for d-elements with non-paired d-electrons (first 5
elements of each transitional series). The values of first ionization potentials of delements of one transitional series increases with increase in charge of the nuclei.
d-Elements in the same oxidation states usually have similar properties. For
example, all hydroxides of M(OH)2 and M(OH)3 types are weak bases which are
sparingly soluble in water. Sulfide and carbonate ions form precipitates with M 3+
and M2+ ions. All d-elements are good complexing agents.
Ions of d-elements in their higher oxidation states possess acidic properties
and exist in solutions as anions: VO3-, CrO42-, Cr2O72-, MnO42-, MnO4-.
Chromium(III) compounds possess amphoteric properties:
СrCl3 + 3NaOH  Cr(OH)3 + 3NaCl
Cr(OH)3 + 3HCl  CrCl3 + 3H2O
Cr(OH)3 +3NaOH  Na3[Cr(OH)6]
In alkaline solutions they can be transformed into chromates:
2Na3[Cr(OH)6] + 4NaOH + 3Br2  2Na2CrO4 + 6NaBr + 8H2O
In acidic media yellow chromates are transformated into orange
dichromates:
2CrO42- + 2H+  Cr2O72- + H2O
Both chromates and dichromates are strong oxidizing agents:
K2Cr2O7 + 4H2SO4 + 3K2SO3  Cr2(SO4)3 + 4K2SO4 + 4H2O
Manganese forms several oxides in different oxidation states: MnO and
Mn2O3 (basic), MnO2 (amphoteric), MnO3 and Mn2O7 (acidic).
98
Compounds of Mn(II) are reducing agents:
2Mn(NO3)2 + 16HNO3 + 5NaBiO3  2HmnO4 + 5Bi(NO3)3 + 5NaNO3 + 7H2O
Manganese in its high oxidation states (+4 and +7) are strong oxidizing
agents:
MnO2 + 4HCl  MnCl2 + Cl2 + 2H2O
2KmnO4 + 3H2SO4 + 5Na2SO3  2MnSO4 + 5Na2SO4 + K2SO4 + 3H2O
pink
colorless
2KmnO4 + H2O + 3Na2SO3  2MnO2 + 3Na2SO4 + 2KOH
pink
brown precipitate
2KmnO4 + 2KOH + Na2SO3  2K2MnO4 + Na2SO4 + H2O
pink
green
Iron is a metal of an intermediate activity and can be dissolved in mineral acis:
Fe + 2HCl  FeCl2 + H2
Fe + H2SO4 (dil.)  FeSO4 + H2
Fe + HNO3 (dil.)  Fe(NO3)3 + NO + H2O
Ferrous oxide and hydroxide are of basic character:
FeO + H2O 
FeO + H2SO4  FeSO4 + H2O
FeSO4 + 2NaOH  Fe(OH)2 + Na2SO4
Fe(OH)2 + Н2SO4  FeSO4 + 2H2O
Fe(OH)2 + 2NaOH 
Divalent iron can be easily oxidated into a trivalent state:
4Fe(OH)2 + 2H2O + O2  4Fe(OH)3
In acidic media Fe(III) possess light oxidizing properties:
FeCl3 + 2HI 2FeCl2 + I2 + 2HCl
Both ferrous and ferric ions are goog complexing agents:
FeSO4 + 6KCN  K4[Fe(CN)6] + K2SO4
FeCl3 + 6KCN  K3[Fe(CN)6] + 3KCl
FeCl3 + 3KCNS  Fe(SCN)3 + 3KCl
99
K4[Fe(CN)6] + FeCl3  Kfe[Fe(CN)6] + 3KCl
K3[Fe(CN)6] + FeSO4  Kfe[Fe(CN)6] + K2SO4
Copper is a metal of a low activity:
Cu + H2SO4 (dil.) 
Cu + 2H2SO4 (conc.)  CuSO4 + SO2 + 2H2O
Cu + 4HNO3 (conc.)  Cu(NO3)2 + 2NO2 + 2H2O
3Cu + 8HNO3 (dil.)  3Cu(NO3)2 + 2NO + 4H2O
Divalent copper possess slightly amphoteric properties:
CuSO4 + 2NaOH  Cu(OН)2 + Na2SO4
Cu(OН)2 + Н2SO4  CuSO4 + 2H2O
Cu(OН)2+NaOH(dil.)
Cu(OН)2+2NaOH(conc.)Na2[Cu(OH)4]
Cuppric compounds are good complexing agents and possess oxidizing properties:
CuSO4 + 4NH4OH  [Cu(NH3)4]SO4 + 4H2O
2CuSO4 + 4KI  2CuI + I2 + 2K2SO4
Zinc is an active metal which can be dissolved both in acids and bases
because of the amphoteric properties:
Zn + 2HCl  ZnCl2 + H2
Zn + H2SO4 (dil.) ZnSO4 + H2
4Zn + 5H2SO4 (conc.)  4ZnSO4 + H2S + 4H2O
4Zn + 10HNO3 (conc.)  4Zn(NO3)2 + N2O + 5H2O
4Zn + 10HNO3 (dil.)  4Zn(NO3)2 + NH4NO3 + 3H2O
Zn + 2NaOH + 2H2O  Na2[Zn(OH)4]
ZnSO4 + 2NaOH  Zn(OH)2 + Na2SO4
Zn(OH)2 + H2SO4  ZnSO4 + 2H2O
Zn(OH)2 + 2NaOH  Na2[Zn(OH)4]
In ammonia solutions zinc forms complex compounds:
Zn(OH)2 + 6NH4OH  [Zn(NH3)6](OH)2 + 6H2O
100
Biological significance of d-elements.
All derivatives of Cr (VI) are
strongly toxic and provoke ulcers and lung cancer while breathing. Manganese can
be found in both animals and plants tissues. Addition of small quantities of
manganese to fertilizers increase the crop capacity of some plants (maize, sugarbeet, potatoes and others). Iron is a catalyst of respiratory processes. Human bodies
contain 4g of iron, about 57% as hemoglobin. Soils contain from 1 to 5% of iron
compounds. Deficiency of iron provokes growing leaves pale. Biological role of
cobalt in living organisms is connected with circulatory system. An antiemetic
vitamin B12 contains 4.35% of cobalt. Cobalt compounds also oppress malignant
tumors. Plants usually contain 10-4 % of zinc. Small amounts of zinc are necessary
for normal growth of animals and fruiting of plants. Mercury and its compounds
are very toxic and provoke disturbance of cardiac and stomach activities and
weakening of memory.
QUESTIONS AND PROBLEMS
1. Where are d-elements situated in the periodic table?
2. Microelements and their biological activity.
3. Which is the type of hybridization of valence orbitals of a central ion in
K4[FeF6]?
4. Calculate concentrations of Ni2+ ions in 0.1 M solution of [Ni(NH3)6]Cl2 (a) in
water; (b) in presence of 1M NH3.
5. Calculate equilibria constants of reactions and explain why the precipitate of
CuS can’t be dissolved in diluted HCl but dissolves in the concentrated one (in

the excess of Cl ions a complex ion [CuCl4]2  is formed).
101
CHAPTERS IN INORGANIC CHEMISTRY
TEST 1
(1)
1.
Questions 1-10 - 2 points
1. The Schrodinger Equation (formula)
2. Principal quantum numbers. (determination, possible values)
3. What orbitals are written wrong? 1s3, 3p5, 4s1, 2d4, 3f14, 5f7, 2p6, 3d10
4. Give the electron configurations of following atoms or ions: C, Mg, S 2-, Mn,
Mn2+
5. The formation of σ–bonds, π- bonds (picture)
6. Vant Hoff’s rule (formula)
7. Energetic diagrams of molecule O2 according MO method
8. Molar concentration of solution (formula)
9. Debaue-Huckel’s equation (formula)
10.A hydrolysis constant of NaCN
2.
Problems 11 – 14 – 4 points
11.How many carbon atoms are there in 6 g of pure carbon? In 6 g of carbon
dioxide? In 6 g of sugar, C12H22O11?
12.Complete and balance the following reactions. Write them in full and net
ionic form.
i. CaCl2 + H2SO4 →
ii. Al(OH)3 + KOH →
13.Calculate pH of 0.001M solution of potassium carbonate.
14. Calculate the ionic strength and activity of the ions in a solution containing
0.01 molel-1 of Ca(NO3)2 and 0.01 mol/l of CaCl2.
3.
Problems 2 – 7 points
15.KmnO4 + NaNO2 + H2O →
16.KI + H2O2 + H2SO4 →
102
CHAPTERS IN INORGANIC CHEMISTRY
TEST 1
(2)
1.
Questions 1-10 - 2 points
1. The Schrodinger Equation (formula)
2. Orbital quantum number. (determination, possible values)
3. What orbitals are written wrong? 2s1, 5p7, 4d3, 6f4, 3f7, 5p7, 3p6, 2d10
4. Give the electron configurations of following atoms or ions: N, Na, N3-, Fe,
Fe2+
5. The formation of σs-p–bonds, πp-p- bonds (picture)
6. Law of mass action aA + bB → cC + dD (formula)
7. Energetic diagrams of molecule N2 according MO method
8. Mass fraction (ω) (formula)
9. Ionic strength (formula)
10.A hydrolysis constant of NH4Cl
2.
Problems 11 – 14 – 4 points
11.How many silver atoms are there in 10 g of pure silver? In 10 g of silver
nitrate?
12.Complete and balance the following reactions. Write them in full and net
ionic form.
i. HNO3 + Na2CO3 →
ii. Zn(OH)2 + NaOH →
13.Calculate pH of 0.01M solution of ammonia chloride?
14.4. Calculate the ionic strength and activity of the ions in a solution containing
0.01 molel-1 of NaNO3 and 0.01 mol/l of Ba(NO2)2.
3.
Problems 15-16 – 7 points
15.KmnO4 + NaNO2 + KOH →
16.NaNO2 + H2O2 + H2SO4 →
103
EXAM PAPER No 1
(max. 25 points each)
1. Comparative characteristics of simple substances of elements of IA group.
Major compounds. Application of elements and their compounds.
2. Aromatic compounds. Quantum-chemical maintenance of concept of
aromaticity. Reactions of electrophylic substitution, their mechanism and a
rule of orientations.
3. Conditions for chemical equilibrium. The law of mass action. The equation
of an isotherm of chemical reaction. Temperature dependence of an
equilibrium constant. The theory of a crystalline field. Theory of substantive
provisions. Magnetic properties of coordination compounds
EXAM PAPER No 2
(max. 25 points each)
1. Comparative characteristic of properties of elements of a family of iron (Fe,
Co, Ni) and their major compounds. Application of metals and compounds.
2. Ammines of aliphatic and aromatic series. Comparison of their basic
properties. Reactions of ammines with nitrous acid. Reactions of aromatic
diazocompounds with and without nitrogen allocation. Mechanisms of these
reactions.
3. Wave dualism. De Broil’s equation, Heisenberg’s principle of uncertainty.
Wave function, its physical sense and basic properties. Principle of
superposition.
4. Laws of splitting of d-d transitions depending on number of d-electrons of
the central atom and symmetry of a complex compound.
104
TOPICS FOR CLASSROOM PRESENTATIONS AND DISCUSSION
1.
Features of hydrogen. Hydrogen isotopes. Oxidation-reduction properties.
Application. Water. Molecular structure. Abnomal properties of water. Acidbase and oxidation-reduction properties. Hydrogen peroxide. Molecular
structure. Methods of synthesis, properties and applications.
2.
Comparative characteristics of simple substances of elements of IA group.
Major compounds. Application of elements and their compounds.
3.
Comparative characteristic of simple substances of elements IIA of group.
Major compounds. Application of elements and their compounds. Hardness of
water and its elimination.
4.
Comparative characteristic of elements IIIA of group. Difference of chemistry
of boron from chemistry of А1 – Т1. Features of chemistry of thallium. Major
compounds. Application of In – Т1 and their compounds.
5.
Comparative characteristic of properties of elements of IVA group. Features
of chemistry of carbon and silicon. Major compounds and their application.
6.
Comparative characteristic of properties of elements of VA group. Features of
chemistry of nitrogen and phosphorus. Major compounds of nitrogen –
bismuth. Application of elements and their compounds.
7.
Chemical bond in a molecule of nitrogen. Properties and application.
Ammonium, hydrazine, hydroxylamine.
8.
Features of chemistry of phosphorus. Properties and application. Oxides and
acids of phosphorus and their salts. Synthesis, properties, application.
9.
Comparative characteristic of properties of elements of VIA group. Features
of chemistry of oxygen and sulfur. Major compounds and their application.
10. Oxides, acids and salts of sulfur(IV and VI): properties and application. Bases
of the method of gravimetric an analysis.
11. Comparative characteristic of properties of elements of VIIA group. Features
of chemistry of fluorine and chlorine – iodine. Major compounds and their
application.
105
12. Features of chemistry of chlorine. Application of chlorine. Oxides, hydroxides
(acids) and salts of chlorine: properties and application.
13. Comparative characteristic of properties of elements of VIB group.
Application of metals. Major compounds of elements in their valence states
(II, III, VI). Izopoly – and hereropolycompounds, clusters. Application.
14. Chromium, general characteristic. Oxides, hydroxides and salts of Cr(II, III,
VI). Application of chromium compounds.
15. Comparative characteristic of properties of elements of VIIA group. Major
compounds of Mn – Re. Application of metals and their compounds.
16. General properties of manganese and its application. Oxidation-reduction
properties of compounds of manganese (II – VII).
17. Comparative characteristic of properties of elements of a family of iron (Fe,
Co, Ni) and their major compounds. Application of metals and compounds.
18. Comparative characteristic of properties of elements of IB group. Features of
chemistry of copper, silver and gold. Major compounds. Application of metals
and their compounds.
19. Comparative characteristic of properties of elements of IIB group. Features of
chemistry of zinc (II), cadmium (II), mercury (I, II). Amphoteric character of
zincum and zincum hydroxide. Application of Zn – Hg and their compounds.
20. The first law of thermodynamics. Gess’s law. Dependence of thermal effect of
reaction on temperature.
21. The second law of thermodynamics. Thermodynamic potentials. Criteria of
spontaneity of processes.
22. Gibbs’s phase rules. Diagrams of fusibility of two-component systems.
23. Conditions for chemical equilibrium. The law of mass action. The equation of
an isotherm of chemical reaction. Temperature dependence of an equilibrium
constant.
24. Nernst’s equation for galvanic cells. Types of electrodes.
25. Rate of a chemical reaction. The kinetic equation. Molecularity and reaction
order. Influence of temperature on rates of chemical reactions.
106
26. The theory of active collisions. Explanation of kinetics of monomolecular gas
reactions by means of the theory of active collisions.
27. Superficial tension. Adsorption. Gibbs’s isotherm of adsorption. Surfaceactive and inactive substances.
28. Isotherm of adsorption of gas on a homogeneous surface of a firm body
(isotherm of Langmuer).
29. Wave dualism. De Broil’s equation, Heisenberg’s principle of uncertainty.
Wave function, its physical sense and basic properties. Principle of
superposition.
30. Quantum-mechanical operators and their properties. The operator of
Hamilton.
31. Schrödinger’s equation. Quantum-mechanical description of a free particle
32. Real gases. A statistical conclusion of equation Van-der-Waals.
33. Simple liquids. A method of molecular functions of distribution,
коррелятивные functions, function of radial distribution.
34. Fluctuations of a crystal lattice. Classical model of a thermal capacity.
Einstein's quantum models.
RECOMMENDATIONS FOR STUDENTS
The student is recommended to look through the contents at the course start to
identify the overall learning prospective and goals.
Each module include the lecture synopsis, references list, questions for
revision
topics for classroom presentations and discussion, sources to prepare for
classroom activities.
Students are recommended to look through the lecture synopsis in advance
and identify those matters that seem not to be clear enough, to address the
questions at the lecture itself.
Additional activities comprise case studies and project work.
The students should focus their attention on the reference list that covers the
basic reading, and on the sources for further classroom activities, as well.
107
The students can be allowed to choose additional topics for presentations and
project work that go beyond the drafted limits of the module content scope.
In this case the teacher will consult him or her on possible basic sources for
further individual learning.
The students should understand that all kind of activities within the course
studies require students’ prior individual learning, including reading, analysis and
synthesis through the information under study processing.
The students are
recommended to plan their participation in classroom
discussions by arranging a list of possible questions or suggestions on each topic
specified for classroom presentations and discussion.
Presentation can be prepared by two or three students if the scope of the
theme is should be covered from different angles.
The students are recommended to follow their progress evaluation and should
check how the teacher marks and grades students’ activities after each session.
The students are recommended to pay their attention to midterm and final
assessment forms and contents
in advance, thus preparing step by step to
controlling new knowledge appropriation and enhancement.
Developer:
Associated professor of the Department of General Chemistry
Ph.D.
E. Yu. Nevskaya
108
Contents
COURSE GOALS
5
COURSE IN ACADEMIC PROGRAM STRUCTURE
5
CONTENT OF THE DISCIPLINE
7
DISCIPLINE VALUE AND TYPES OF STUDY
7
PARTS OF DISCIPLINE AND INTERDISCIPLINARY BONDS WITH
11
PROVIDED (SUBSEQUENT) DISCIPLINES
PART OF THE DISCIPLINE AND KINDS OF ACTIVITIES
12
SYSTEM OF THE ESTIMATION OF STUDENTS’ KNOWLEDGES
14
STUDY GUIDE
15
LECTURES CHAPTERS IN INORGANIC CHEMISTRY
17
BASIC CONCEPTS OF CHEMISTRY
17
GENERAL CLASSES OF INORGANIC SUBSTANCES
20
OXIDES
20
BASES
21
ACIDS
22
SALTS
23
CHEMICAL KINETICS
24
STRUCTURE OF ATOMS
28
GLOSSARY: ATOMS, ELEMENTS, AND ION
28
THE PERIODIC LAW AND PERIODIC TABLE OF CHEMICAL
41
ELEMENTS
GLOSSARY: THE PERIODIC TABLE
41
CHEMICAL BONDS
47
GLOSSARY: CHEMICAL BONDS
47
INTERACTION OF MOLECULES (THE CONDENSED STATE OF
62
SUBSTANCES)
SOLUTIONS
65
REACTIONS OF IONIC EXCHANGE
68
109
DISSOCIATION OF STRONG ELECTROLYTES
69
DISSOCIATION OF WEAK ELECTROLYTES
71
EQUILIDRIA IN SOLUTIONS WITH PRECIPITATES
75
DIRECTION OF REACTIONS OF IONIC EXCHANGE
78
HYDROLYSIS OF SALTS
79
OXIDATION-REDUCTION REACTIONS
83
COMPLEX COMPOUNDS
89
SURWAY OF PROPERTIES OF SOME CHEMICAL ELEMENTS
100
ALKALINE AND ALKALINE EARTH METALS
100
ELEMENTS OF IIIA AND IVA GROUPS (P – ELEMENTS)
104
ELEMENTS OF VA AND VIA GROUPS (P-ELEMENTS)
107
ELEMENTS OF THE VIIA GROUP (HALOGENS)
110
TRANSITIONAL ELEMENTS (D-ELEMENTS)
112
QUESTIONS FOR REVISION
116
TOPICS FOR CLASSROOM PRESENTATIONS AND DISCUSSION
119
RECOMMENDATIONS FOR STUDENTS
121
110