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Chemistry: A Primer
by
Mark J Donohue
Atoms
Chemistry is the study of matter and its properties, the changes that matter undergoes and the energy
associated with these changes. Matter is the substance that objects are made of and matter exists in
three familiar forms – solids, liquids and gases. The basic building blocks of all matter are atoms.
The atom is an electrically neutral, spherical entity consisting of
subatomic particles (smaller than the atom) called protons,
neutrons and electrons. These subatomic particles are
distinguished from each other by their different electrical
charges and masses.
Protons are positively charged particles (+) while electrons are
negatively charged particles (-). Protons (+) and electrons (-) are
electrically attracted to each other like the poles of a magnet.
Neutrons are neutral (o), meaning they do not carry a charge and
are not electrically attracted to either protons or electrons.
An atom is constructed in a way that neutrons and protons are contained in the nucleus (center) of an
atom. Because protons reside in the nucleus, the nucleus always has a positive charge. The nucleus of an
atom, despite containing only positively charged protons (which ordinarily should repel each other) is
held together by what is called the nuclear force or the strong force. This force opposes and overcomes
the force of repulsion between the protons and holds the nucleus together, while the energy associated
with this force is called atomic binding energy.
Electrons are found spinning around the outside of the nucleus.
Electrons do not follow a fixed path or orbit but instead form a
negatively charged cloud (-) that surrounds the nucleus. This
cloud is formed due to electrons spinning around the nucleus at
intense speeds - billions of times per second.
The number of electrons surrounding an atom equals the
number of protons in the nucleus. Because each atom contains
the same number of positively charged protons (+) and the
same number of negatively charged electrons (-), they will
balance each other. As a result each atom is electrically neutral,
meaning its total charge is zero. For example:
Name of Atom
# of Protons
# of Electrons
Hydrogen
1
1
Magnesium
12
12
Sulfur
16
16
Iodine
53
53
Uranium
92
92
Even though the exact location of electrons cannot be predicted, electrons do move about in certain
regions called electron shells. The electron shells surround the nucleus like layers of an onion. Each
electron shell is only capable of holding a certain number of electrons. The shell nearest the nucleus is
the first shell. The first shell is capable of holding a maximum of two electrons. The first shell is the first
to “fill up”, after which electrons will begin to fill the second shell. When the second shell becomes full,
electrons will then begin to fill the third shell and so on. For example:
Shell No.
1st
2nd
3rd
4th
Max. # of Electrons
2
8
18
32
There are as many as seven electron shells.
The negatively charged electrons (-) are held in their shells by the attractive electrical charge from the
positively charged protons (+). The first shell is very stable with the least amount of energy. But as the
distance from the nucleus increases, energy levels also increase as well as an electron’s instability.
The most stable electron configuration for an atom is one in which its outer shell is filled with its
maximum number of electrons.
Protons and neutrons have nearly equal masses - amount
of matter. In fact, the nucleus is incredibly dense and
contributes 99.97% of the atoms mass, but only occupies
about 1 ten-trillionth of its volume. Electrons on the other
hand are much smaller - it takes the mass of 1836
electrons to equal the mass of just one neutron. But the
electron cloud of rapidly moving negatively charged
electrons, surrounding an atom’s nucleus, accounts for
almost all the volume of any one atom.
Elements
Atoms are the smallest division of all substances - they cannot be separated or broken down into a
simpler substance by ordinary chemical means. The different types of atoms are referred to as chemical
elements or just elements. An element consists of only one kind of atom and is considered a pure
substance. Each element is designated by a chemical symbol – one or two letters of the element’s
English name. For example:
English Name
Chemical Symbol
Hydrogen
Magnesium
Sulfur
Chromium
H
Mg
S
Cr
Some of the symbols come from the original Latin names for the elements and thus do not match the
modern English names. For example:
English Name
Latin Name
Sodium
Potassium
Iron
Tin
Natrium
Kalium
Ferrium
Stannum
Chemical Symbol
Na
K
Fe
Sn
An element’s name and chemical
symbol is listed in a chart called the
periodic table. The periodic table,
simply put, is a list of all the elements
(types of atoms) that make up all
matter in the known universe.
As of 2008, there have been observed
117 elements (atoms). Ninety-two of
these elements occur naturally. The
remaining twenty-five are manmade
and have been synthesized via nuclear
processes. Twenty-six of the elements in the periodic table are called essential elements. Essential
elements are those present in the human body and are necessary for life. Four of the essential elements
constitute 96% of the body’s mass and are called major elements. The four major elements in the
human body with their chemical symbol and percentage of total body mass are:
Element
Chemical Symbol
% of Total Body Mass
Oxygen
Carbon
Hydrogen
Nitrogen
O
C
H
N
65%
18.5%
9.5%
3.2%
Eight other elements, referred to as lesser elements contribute 3.8% of the body’s mass. These lesser
elements with their chemical symbol and percentage of total body mass are:
Element
Calcium
Phosphorus
Potassium
Sulfur
Sodium
Chlorine
Magnesium
Iron
Chemical Symbol
Ca
P
K
S
Na
Cl
Mg
Fe
% of Total Body Mass
1.5%
1.0%
0.35%
0.25%
0.2%
0.2%
0.1%
0.005%
An additional fourteen elements, though they are present in tiny amounts, are essential in some way to
human cell function and are called trace elements. Together they account for the remaining 0.2% of the
body’s mass. The trace elements with their chemical symbol are: Aluminum (Al), Boron (B), Chromium
(Cr), Cobalt (Co), Copper (Cu), Fluorine (F), Iodine (I), Manganese (Mn), Molybdenum (Mo), Selenium
(Se), Silicon (Si), Tin (Sn), Vanadium (V), and Zinc (Zn).
The periodic table, while listing an element’s name and chemical symbol, is actually organized or
arranged according to an element’s atomic number. The atomic number identifies how many protons
are found in the nucleus of a particular element. All protons are the same, regardless of the element
they are a part of. What differentiates one element from another is the number of protons within the
nucleus. For example:
Element
Hydrogen
Oxygen
Sulfur
Chromium
Selenium
Number of Protons
1
8
16
24
34
Atomic Number
1
8
16
24
34
The atomic number never changes, meaning that the number of protons in one atom of any given
element is constant and distinguishes that element from all others.
All atoms of an element are identical in atomic number but not in atomic mass number. The atomic
mass number, which is listed in most (but not all) periodic tables, is the total amount of mass in the
nucleus of an atom. Because the nucleus consists of only protons and neutrons, and because the mass
of each proton and neutron is one atomic mass unit (amu), calculating the atomic mass is done by
simply adding the number of protons to the number of neutrons.
The atomic mass number is the result of the fact that every
element has more than one form of itself. What
differentiates one form of a particular element from another
form of the same element is the number of neutrons in the
nucleus. This is also why the atomic mass number of an
element is an average and why not all the atomic mass
numbers are whole numbers. For example:
 All carbon (C) atoms have 6 protons in the nucleus,
but only 98.9% of naturally occurring carbon atoms
have 6 neutrons in the nucleus – atomic mass
number: 6 protons + 6 neutrons = 12. A small
percentage of carbon atoms (1.1%) have 7 neutrons
– atomic mass number: 6 protons + 7 neutrons = 13.
And even fewer (less than 0.01%) have 8 neutrons –
atomic mass number 6 protons + 8 neutrons =14.
Isotopes
As mentioned, the number of protons in an element remains the same, but the number of neutrons in
the atom of an element is variable. An atom that has a different number of neutrons compared to the
number of protons in its nucleus, will not only change the atomic mass of the element, but will now
become an isotope of the element. Presently, there are over 2,000 known isotopes of the elements.
The symbol for an isotope can be written either with the atomic mass number written as a superscript in
the upper left of the chemical letter symbol. Or it can be written with the name of the element
appearing first followed by the atomic mass number (the number “1” is usually assumed and omitted).
For example:

Hydrogen-1 or (1H) is known simply as H or hydrogen - a typical hydrogen atom has one proton
in the nucleus and one orbiting electron, but it has no neutrons.

Hydrogen-2 or (2H), adding one neutron to hydrogen’s one proton = an atomic mass of 2.

Hydrogen-3 or (3H), adding two neutrons to hydrogen’s one proton = an atomic mass of 3.
It is important to note that all isotopes of a single element have the same chemical properties because
these properties are not affected by the number of neutrons, but rather by the electron configuration
(number of electrons and shells). Therefore, all isotopes of an element have nearly identical behavior,
even though they have different masses.
Despite the fact that the atomic mass has very little bearing on chemical reactions, the physical
properties of atoms do greatly depend on its mass. The stability of each atom’s nucleus depends on the
ratio of protons to neutrons. Many isotopes have a ratio of protons to neutrons that render them
unstable. Unstable isotopes are called radioisotopes and the energy they emit is called radiation. While
stable isotopes do not decay (lose their energy) and are not radioactive. For example:
 Taking our example of naturally occurring carbon from the
previous page. Carbon exists naturally with 6, 7 or 8 neutrons.
These carbon isotopes have atomic masses of 12, 13 and 14. The
isotopes are called (12C) or carbon-12, (13C) or carbon-13 and (14C)
or carbon-14.
Carbon-12 and carbon-13 are stable isotopes and are not
radioactive. But carbon-14 is unstable and radioactive, decaying
with a half-life of about 5,700 years.
 Uranium (U) has three naturally occurring isotopes. These are (234U) or uranium-234, (235U) or
uranium-235, and (238U) or uranium-238. Since each atom of uranium has 92 protons, simply
subtracting the number of protons from the atomic mass number will yield the number of
neutrons in each element.
In this case the isotopes must have 142, 143 and 146 neutrons respectively. Or put another way
uranium has no stable isotopes and is radioactive. Uranium is present in the earth's crust and
decay’s extremely slow. For example the half life for uranium – 235 is 704 million years.
Of particular interest is uranium-238 which naturally occurs in most types of
granite and soil in varying amounts and has a half life of 4.5 billion years. As
uranium-238 undergoes radioactive decay it produces the by-product radon.
Radon (Rn), element – 86, is a colorless radioactive gas that accumulates in
homes and businesses which may cause health problems in higher
concentrations.
The terms radioisotopes and radiation tend to strike fear in the minds of most people. However, both
artificial (manufactured) and naturally occurring radioisotopes are important and are used in medicine.
The use of radioisotopes in medicine is a specialty known as
nuclear medicine. For example:
 Thallium-201 (201Tl) is used in imaging the inside of the
human body for diagnostic procedures while other
radioisotopes are used therapeutically to treat
various forms of cancer.
Ions
When the electron configuration of a single element does change, either by gaining an electron or losing
an electron, it becomes an ion of that element. An ion causes an atom, which normally has a neutral
charge, to develop either a positive (+) or negative (-) charge. This is the result of an unequal number of
protons and electrons. The process of gaining or losing electrons from a neutral atom or molecule is
called ionization.
In an atom, if protons outnumber electrons, it’s called a cation with a positive charge (+). If the electrons
are more numerous, it is an anion with a negative charge (-). An ion consisting of a single atom is called
a monatomic ion
An ion of an element is symbolized by writing its chemical symbol followed by the number of its positive
(+) or negative (-) charges as a superscript to the upper right (the number “1” is assumed and omitted).
For example:
 Potassium (K+) is an ion that has one positive charge (+) because it has given up one electron. Or
in other words – the potassium atom has lost 1 of its electrons and therefore there is now one
more positively charged proton (+) in the nucleus than there are negatively charged electrons (-)
spinning around it. This means the potassium atom will have a 1+ or just (+) positive charge
making it a cation (+).
 Calcium (Ca2+) is an ion that has two positive charges because it has given up two electrons. Or
in other words - the calcium atom has lost 2 of its electrons and therefore there are now two
more positively charged protons (+) in the nucleus than there are negatively charged electrons (-)
spinning around the nucleus. This means the calcium atom will have a 2+ positive charge making
it a cation (+).
 Other monatomic cations (+): hydrogen (H+), lithium (L+), magnesium (Mg2+), aluminum (Al3+)
 Sulfide (S2-) is an ion that has two negative charges because it has gained two electrons. In other
words - there are two more negatively charged electrons (-) spinning around the nucleus than
there are positively charged protons (+) in the nucleus. This means the sulfur atom will have a 2negative charge making it an anion (-).
 Other monatomic anions (-): hydride (H-), fluoride (F-), oxide (O2-), nitride (N3-)
In review – atoms consist of:
 Electrons – an atom that gains or loses electrons becomes an ion of the same element.
 Neutrons – an atom that gains or loses neutrons becomes an isotope of the same elements
 Protons – an atom that gains or loses protons becomes a different element.
Molecules and Compounds
When two or more atoms are chemically bonded together they create units known as molecules. A
molecule may contain two or more atoms of the same element such as an oxygen molecule (O2) or a
hydrogen molecule (H2). But most molecules contain two or more atoms of different elements and are
known as compounds (i.e. water H2O).
In fact only a few elements occur free in nature. The overwhelming
majority of elements occur in chemical combination with other
elements as molecules and compounds. Most of the atoms in the
human body are joined into compounds.
Chemical formulas were developed to identify, with chemical symbols, the composition of a substance.
There are three types of chemical formulas – empirical, molecular and structural. It is beyond the scope
of this report to detail the differences and will therefore refer to them collectively as chemical formulas.
Each element in a molecule is represented by its chemical symbol followed by a lower right numeric
subscript indicating the number of atoms present of that element. A subscript is only used when more
than one atom is being represented (the number “1” is assumed and omitted). For example:
Molecules – two or more atoms of the same element:
Name
Chemical Formula
Oxygen
Hydrogen
Nitrogen
O2
H2
N2
Description
a gas made up of 2 oxygen atoms
a gas made up of 2 hydrogen atoms
a gas made up of 2 nitrogen atoms
O2
Molecules that are also compounds – two or more atoms of a different element:
Name
Chemical Formula
Water
H2O
Carbon dioxide
CO2
Calcium bromide
CaBr2
Sodium chloride
NaCl
Description
made up of 2 atoms of hydrogen
and 1 atom of oxygen.
made up of 1 atom of carbon and 2 atoms of
oxygen.
made up of 1 atom of calcium and 2 atoms of
bromine.
made up of 1 atom of sodium and 1 atom of
chlorine.
H2O
NaCl
Chemical Bonds
The forces that bind the atoms into molecules and compounds are called chemical bonds. The
determining factor in the linking or binding of atoms is the number of electrons in the outermost shell of
an element. The outer most shell of an atom is referred to as the valence shell. An atom whose valence
shell is filled with the maximum amount of electrons possible (for that shell) is said to be chemically
stable. Meaning it is very unlikely this atom will bond with other atoms. For example:
 Neon has an atomic number of 10, with 2 electrons in the first shell and 8 electrons in the
second shell, which is this elements valence shell. The potential maximum number of electrons
in the second shell of any atom is 8. Therefore, neon’s valence shell is full making it chemically
stable.
The atoms of most biologically important elements do not have their outer most valence shell full.
Making these elements unstable, but because atoms seek to reach a state of maximum stability, an
atom will try to fill its outer shell. Elements do this by chemically bonding and sharing their electrons
with other unstable elements. There are three different ways in which chemical bonding takes place:
1) covalent bonds, 2) ionic bonds, and 3) hydrogen bonds.
Covalent Bonds
Covalent bonds are formed when atoms share one or more pairs of electrons from their outer shell.
Electrons associate in pairs – two electrons that occupy the same “orbit” but have opposite “spins”. In
covalent bonds each atom contributes one electron to each pair. The greater the number of electrons
shared between two atoms the stronger the covalent bond. Neither of the combining atoms loses or
gains electrons during this bonding, it is strictly an act of sharing.
Again using carbon, below is an example of the formation of a covalent bond.
 An atom of carbon (C), with an atomic number of six, has six protons in the nucleus and
therefore six electrons surrounding the nucleus. Two electrons fill the first electron shell, while
the outer most second shell has an additional four electrons. Because the potential maximum
number of electrons in the second shell is eight, that means carbon has four unoccupied places
in the second shell.
An atom is most stable when its outer electron shell is filled, therefore carbon will bind with
other atoms to find the four electrons it needs to stabilize its outer shell. A good match would
be with four individual hydrogen (H) atoms which each are missing one electron in their outer
shell. The result - four covalent bonds are formed between one carbon and four hydrogen atoms
to form the molecule methane (CH4) - see image below.
Sometimes adjacent atoms share two or three pairs of electrons rather than just one. In these cases the
bond is called a double bond or triple bond. Covalent bonds do not break apart in the presence of water,
but rather the input of energy is needed to break them apart. Covalent bonds are the most common
chemical bonds in the human body and form most of the body’s structures.
There are several combinations of atoms called functional groups that occur repeatedly in biological
molecules. The atoms in the functional groups tend to move from molecule to molecule as a single unit
rather than as single atoms. Functional groups usually attach to molecules by single covalent bonds.
Several examples of functional groups are:





Carboxyl (acid)
Hydroxyl
Amino
Phosphate
Sulfhydryl
COOH
OH
NH2
H2PO4
SH
contained in aldehyes like formaldehyde
contained in ethanol or glucose
contained in amino acids
contained in nucleic acids and ATP
stabilize the structure of proteins
Ionic Bonds
Ionic bonds form via the force of attraction that holds together oppositely charged ions. That is to say
ionic bonds are the chemical bonds between positively charged cations (+) and negatively charged (-)
anions. The resulting cations and anions attract each other through electrostatic forces and form an
ionic compound. This typically occurs between metals and non-metals. For example:
 Sodium (Na) – atomic number 11, is an alkali-metal and has one outer shell
electron. If a sodium atom loses this electron the total number of positively
charged protons (11) will now exceed the number of negatively charged
electrons (10). The sodium atom now becomes a positively charged cation
with a charge of 1+ or (Na+).
But if chlorine (Cl) – atomic number 17, a non-metal with its seven outer
shell electrons accepts one electron from a neighboring sodium atom it will
now have more negatively charged electrons (18) than positively charged
protons (17). The chlorine atom now becomes a negatively charged anion
with a charge of 1- or (Cl-) and will also now be referred to as chloride.
The resulting positive and negative charges attract each other and form an
ionic bond or an ionic compound, which in this case is sodium chloride
(NaCl)… a.k.a. table salt.
NaCl
An ion consisting of a single atom is called a monatomic ion, but in molecules or compounds with two or
more atoms it is called a polyatomic ion. A majority of which are negatively charged anions (-).
Polyatomic ions containing oxygen are called oxyanions. For example:
 Sulfate (SO42-) is an example of a polyatomic anion which is also an oxyanion. Sulfate consists of
one atom of sulfur (S) and four atoms of oxygen (O4) with the molecule carrying 2 more
electrons than there are protons giving it a (2- ) charge, making it an anion.
 Other polyatomic anions which are also oxyanions are: peroxide (O22-), nitrate (NO3-),
sulfite (SO32-), bicarbonate (HSO3-), chromate (CrO42-)
 A couple of polyatomic cations are: ammonium (NH4+), mercury(I) (Hg22+)
In the human body, ionic bonds are found mainly in teeth and bones where
they give strength to the tissue. Ions are also found dissolved in body fluids as
electrolytes. An ionic compound that breaks apart into cations and anions
when dissolved in a solution is called an electrolyte. This is because the
solution will now be able to conduct an electric current. Electrolytes are
critical for controlling water movement within the body, maintaining acid-base
levels and producing nerve impulses.
A tip to help further distinguish the difference between an ionic and covalent bond is that – most
covalent compounds consists of molecules, while no molecules exist in a sample of an ionic compound
only individual atoms. For example:
 In a covalent compound a cup of water contains a collection of individual water molecules, each
surrounded by other water molecules.
 In an ionic compound like sodium chloride (table salt), there is a continuous array of oppositely
charged individual sodium and chloride ions (atoms), not a collection of “sodium chloride
molecules”.
Hydrogen Bonds
Hydrogen bonds are the forces of attraction between hydrogen atoms and nearby oxygen, nitrogen or
fluorine atoms. Hydrogen bonds may occur between atoms in neighboring molecules or between atoms
in different parts of the same molecule.
Hydrogen bonds are weak when compared to ionic and covalent bonds. Thus, they cannot bind atoms
into molecules. However, hydrogen bonds do establish important links between molecules, such as
water molecules or between different parts of large molecules such as proteins and DNA. In this role
hydrogen bonds add strength and stability to large molecules. For example:

Hydrogen bonds cause large biological molecules such as proteins to fold back on themselves,
creating a three dimensional shape that is essential for the function of proteins.
Hydrogen bond folding a protein (above)
&
Hydrogen bond in DNA structure (right)
Chemical Reactions
The actual process of the making and breaking of chemical bonds takes place via chemical reactions.
Chemical reactions form new products that have different chemical properties than the original reacting
material - reactant. All chemical reactions involve a change in substance and a change in energy.
However, neither matter nor energy is created or destroyed in a chemical reaction – only changed.
Along with the thousands, if not millions, of chemical reactions that take place in ones immediate
environment, there are also thousands of chemical reactions taking place every second within the
human body. Obviously, it is impossible to examine all these chemical reactions. Fortunately, in
examining chemical reactions certain patterns or types of chemical reactions form.
There not only are several different types of chemical reactions, but also more than one way of
classifying them. However, below are descriptions of three types of chemical reactions that are most
relevant to the topics in this report. The three types of chemical reactions are: 1) precipitation reactions
2) acid-base reactions and 3) reduction-oxidation reactions.
Precipitation Reactions
Of the many thousands of reactions that occur in the environment and in organisms, the overwhelming
majority take place in an aqueous solution (water based).
Substances that dissolve in liquid are collectively known as solutes. The liquids into which they dissolve
are known as solvents. The combination of solutes dissolved in solvent is called a solution. In biological
solutions, water is the universal solution with the human body measuring in at 60% water. Water makes
up most of the volume of cells and body fluids and is the solvent that carries nutrients, oxygen and
wastes throughout the body.
The degree to which a molecule is able to dissolve in a solvent is its solubility – the more easily it
dissolves, the higher its solubility. Molecules that dissolve readily in water are soluble or hydrophilic –
water loving. Common examples of hydrophilic solutes are salt and sugar. Molecules that do not
dissolve readily in water are insoluble or hydrophobic – water fearing. Common examples of
hydrophobic compounds are fats and vegetable oils.
Precipitation reactions take place when two soluble ionic compounds
react to form an insoluble product – a precipitate (formation of a
solid in a solution). Precipitates form because of the strong
electrostatic attraction between the two soluble ionic compounds and
their removal from solution. When solutions of such ions are mixed
the ions collide and stay together and the resulting substance “comes
out of solution” as a solid.
Precipitation reactions occur all around us, for example:

Sulfide and sulfate minerals form as precipitates in hydrothermal fluids
and other aqueous solutions in the Earth’s crust.

Kidney stones form as precipitates of calcium ions and oxalates

Pipes in homes get clogged because precipitates of magnesium and
calcium oxides have deposited themselves within the pipes.

Precipitation reactions are used for making pigments.
Clogged pipe
Acid-Base Reactions
Acid-base reactions or neutralization reactions are obviouslyreactions that occur when an acid reacts
with a base.
An acid is a substance that produces hydrogen (H+) ions when dissolved in water or a compound who
donates a hydrogen ion to another compound (base). A base is a substance that produces hydroxide
(OH-) ions when dissolved in water or a compound that accepts hydrogen ions.
Acid and bases are also electrolytes and are often categorized in terms of “strength” which refers to
their degree of dissociation into ions in aqueous solution. Strong acids and strong bases dissociate
completely into ions when they dissolve in water. Therefore, they easily conduct electricity and are also
strong electrolytes. In contrast, weak acids and weak bases dissociate very little leaving most of their
molecules remaining intact. As a result they conduct only a small amount of current and are considered
as weak electrolytes.
Examples of the more common strong and weak acids and bases with some of their uses are:
Strong Acids
Hydrochloric acid (HCI)
Stomach or gastric acids, ceramic
tile cleaner
Nitric acid (HNO3)
Oxidizer in liquid fuel for rockets,
clog remover
Sulfuric acid (H2SO4)
Car batteries, mineral processing
Weak Acids
Hydrofluoric acid (HF)
Production of fluorides, use in
dissolving glass
Phosphoric acid (H3PO4)
Tangy taste in foods and beverages,
solvent, rust remover
Acetic acid (HC2H3O2)
Found in vinegar – pickling,
food flavoring, solvent
Boric acid (H3BO3)
Mild antiseptic, insecticide
The many uses
of phosphoric
acid
Strong Bases
Sodium hydroxide (NaOH)
Lye, drain cleaner, detergents,
paper/pulp production
Sodium carbonate (Na2CO3)
Water softener, grease
remover
Potassium hydroxide (KOH)
Many industrial applications
Calcium hydroxide Ca(OH)2
A chemical reagent, a filler
Weak Bases
Ammonia (NH3)
All purpose industrial and household cleaner
Pyridine (C5H5N)
Solvent, reagent, used in pharmaceuticals
When an acid and a base come together they, like precipitation reactions, are driven by the electrostatic
attraction of ions and their removal from solution in the formation of a new product or compound. In
acid-base reactions the ions are the hydrogen (H+) cation and the hydroxide (OH-) anion. They react by
neutralizing each other’s properties and their positive (+) / negative (-) charges. The results are the
production of water (H2O) and a neutrally charged compound called a salt.
When the term salt is used most people naturally think of table salt (sodium chloride – NaCl) which is
only one example of a salt. The term “salt” is a general term which applies to all the products or
compounds formed from acid-base reactions.
Reduction-Oxidation (Redox) Reactions
The process of moving electrons from one element to another element during chemical reactions is
called reduction and oxidation or redox. Redox chemical reactions occur in the formation of both ionic
and covalent bonds.
Oxidation is commonly defined as the addition of an oxygen molecule to any substance it comes in
contact with to form an oxide. Substances that have the ability to oxidize (add oxygen to) other
substances are said to be oxidative and are known as oxidizing agents, oxidants or oxidizers.
However, it is not necessary for oxygen to be present for oxidation to occur and therefore oxidation is
also defined as the loss of at least one electron when two or more substances interact. Therefore,
oxidation can be defined as either:
1) A molecule that gains oxygen.
2) A molecule that losses an electron.
While an oxidizing agent can be defined as either:
1) A chemical compound that readily transfers oxygen atoms.
2) A substance that gains electrons in a chemical reaction – electron acceptor.
Reduction is the opposite of oxidation, whereby oxygen is removed or in the absence of oxygen at least
one electron is added when substances come into contact with each other. Substances that have the
ability to reduce (take oxygen away from) other substances are said to be reductive and are known as
reducing agents, reductants or reducers. Therefore, reduction can be defined as either:
1) A molecule that losses its oxygen.
2) A molecule that gains an electron.
While a reducing agent can be defined as either:
1) A chemical compound that readily accepts oxygen atoms.
2) A substance that loses electrons in a chemical reaction – electron donor.
It is important to note that reduction and oxidation always occur together - simultaneously. Meaning,
when two substances come in contact with each other the oxidizing agent will transfer its oxygen to the
reducing agent. Once this occurs the oxidizing agent will have lost its oxygen and will have said to be
reduced. While the reducing agent which has now accepted the oxygen will have said to be oxidized. Or
to put another way the number of electrons lost by the reducing agent equals the number of electrons
gained by the oxidizing agent.
It should be obvious that oxygen is a major oxidizing agent. It will directly oxidize all but a few of the
metals and most of the nonmetals as well. Often these direct oxidations form normal oxides such as:
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4Li + O2 -------- 2Li2O (lithium oxide)
2Zn + O2 -------- 2ZnO (zinc oxide)
S + O2 -------- SO2 (sulfur dioxide)
Other common examples of reduction-oxidation reactions are:
Rust
Electrochemistry
Combustion
Respiration
Corrosion, decay
Organic – Inorganic
Chemicals can be divided into two main classes of compounds: inorganic and organic. Inorganic
compounds usually lack carbon, are structurally simple, and are held together by ionic or covalent bonds.
Though there are exceptions to this rule such as carbon dioxide (CO2) and bicarbonate ion (HCO3-), which
both contain carbon yet are considered to be inorganic chemical compounds. Studying the field of
inorganic compounds is referred to as - inorganic chemistry. Examples of inorganic compounds are:
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Water
Salts
Acids
Bases
Organic compounds usually contain hydrogen, but always contain carbon and always have covalent
bonds. Studying the field of organic compounds is referred to as – organic chemistry. Examples of
organic compounds are:
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Carbohydrates
Lipids
Proteins
Nucleic Acids
Adensosine Triphosphate (ATP)
There is another discipline or field of study within chemistry called biochemistry. Biochemistry is the
study of the chemical processes in living organisms.
The purpose of this report was to give a basic understanding of some key concepts in the field of
chemistry.
References
rd
Baird C., Environmental Chemistry 3 Edition, W.H. Freeman Company, 2004
nd
Silberberg M.S., Chemistry: The Molecular Nature of Matter and Change, 2 Edition, The McGraw-Hill Companies
Inc. 2000
Silverthorn D.U., Human Physiology: An Integrated Approach, Pearson Education Inc., 2004
Tortora G., Introduction to the Human Body: The essentials of Anatomy and Physiology, Biological Sciences
Textbooks, 2010