The Basics of Chemistry

G. Bény

A. Deschuytere

C. De Jonghe

S. De Jonge

Basic Chemistry 2

Chapter 1. The atom

1. Atoms and matter.

All matter is made up of atoms. There are 92 different atoms in nature. Some of them

are rare while others are extremely common in the universe. All these atoms are also

present on earth (table 1)

Table 1 The abundance of atoms in the earth’s crust.

Atom

Presence in the

earth’s crust

(in %)

Oxygen 45,5

Silicon 27,2

Aluminium 8,3

Iron 6,2

Calcium 4,7

Magnesium 2,8

All other atoms 5,3

Atoms are formed in the interior of stars.

Besides the 92 so-called natural atoms there are also the synthetic atoms made by man.

They are formed during processes in nuclear reactors or in particle accelerators.

1. The structure of an atom.

A. Elementary particles.

The atoms themselves are made up of even smaller particles. They are called elementary

particles. These are: the electron, the proton and the neutron. Table 2 gives more

information about the mass and charge of these particles.

Table 2 The properties of elementary particles in atoms.

Mass (in g) Charge (in C)

Proton 1,67262 x 10-24 +1,6022 x 10-19

Neutron 1,67493 x 10-24 0

Electron 9,10939 x 10-28 - 1,6022 x 10-19

B. The charge of the elementary particles.

As shown in table 2 protons and electrons carry a charge, protons being positive and

electrons being negative. Neutrons are neutral particles. Atoms always contain the same

number of protons and electrons. They are therefore always neutral. An oxygen atom

contains 8 protons and 8 electrons.

In some cases atoms can either loose or receive electrons thereby forming charged

particles called ions (see below).

Basic Chemistry 3

The charge of the electron is the smallest charge value found in nature. All other charges

are multiples of this smallest value. This value is therefore called the elementary charge

and is about 1,6 x 10-19 C. An electron has a charge of -1 elementary charge (or just -1)

while a proton has a charge of +1 elementary charge (or +1).

C. The mass of elementary particles.

Protons and neutrons have almost the same mass, while electrons have a much smaller

mass. The electrons represent only a very small part of the total mass of an atom. Their

mass is therefore usually neglected when calculating atomic mass.

Protons and neutrons (the heavy particles) make up the nucleus of an atom. They are

therefore called nucleons. This means that almost the entire mass of an atom is in its

nucleus. The electrons are in an almost massless space around the nucleus.

Not all atoms have the same size but the radius of a medium atom corresponds to about

100 pm (where 1 pm - picometer - corresponds to 10-12 m). The radius of the nucleus is

about 5 x 10-3 pm.

D. The composition of atoms.

Natural matter is made up of 92 different atoms. These atoms differ in the number of

protons found in the nucleus, this number being called atomic number (symbol Z). The

atomic number varies between 1 for the lightest atom (hydrogen) to 92 for the heaviest

one (uranium).

Atoms with an atomic number higher than 92 are all man-made and are called synthetic

atoms.

Although atoms can be identified by their atomic number Z it is more usefull to

distinguish them by using a symbol and a name. The complete list of the atoms with their

atomic number, symbol, name and possible other information can be found in the

periodic table.

2. Isotopes.

Almost all atoms, defined by their atomic number, have isotopes. Isotopes are atoms that

have the same atomic number (same number of protons) but a different number of

neutrons (in the nucleus). The total number of particles in the nucleus (sum of protons

and neutrons) is called the mass number (symbol A). This means that isotopes have the

same atomic number but a different mass number. Some examples can be found in table

3.

Isotopes can either be very stable or can disappear by radioactive decay. This decay can

be fast or slow depending on the isotope concerned.

Basic Chemistry 4

Table 3 Isotopes of selected atoms (not all existing isotopes have been included).

Atomic number (Z) Name (Symbol) Mass number of the

isotopes

Abundance (in %)

1 Hydrogen (H) 1 99,985

2 0,015

6 Carbon (C) 12 98,89

13 1,11

20 Calcium (Ca) 40 96,97

42 0,64

43 0,14

44 2,1

46 0,003

48 0,18

92 Uranium (U) 235 0,72

238 99,27

It has been shown that the relative abundance of the isotopes of a particular atom is

constant whatever the origin of this atom.

The chemical properties of the isotopes of a particular atom are the same. This is so

because the chemical properties are determined by the number of electrons in an atom

and not by the composition of the nucleus. Because isotopes have the same atomic

number they also have the same number of electrons and as a consequence also the same

chemical properties.

3. Representation of an atom.

In most cases atoms are represented by their symbol. In some cases however more

information can be included. When a specific isotope is mentioned its mass number can

be included. The mass number is added on the upper left side of the symbol. This is

shown in the following examples.

Example 1 The symbolic representation of some isotopes

The isotope of uranium with mass number 238: 238U (pronounced “uranium 238”)

The isotope of hydrogen with mass number 2: 2H.

The isotope of carbon with mass number 14: 14C

If the atomic number is included it is mentioned on the lower left side of the symbol.

The isotopes of hydrogen have their own name.

Table 4 The isotopes of hydrogen

Isotope Name 1H Hydrogen 2H Deuterium 3H Tritium

Basic Chemistry 5

4. Atomic mass.

A. Absolute atomic mass.

The mass of an atom is equal to the sum of the mass of all the elementary particles

present in that atom.

Example 2 What is the mass of the 2H hydrogen isotope? What part of that mass

corresponds to the contribution of the electrons?

This isotope contains 2 nucleons (1 proton, 1 neutron) and 1 electron.

Mass of the hydrogen atom 2H = mass proton + mass neutron + mass electron.

Mass 2H = 1,67262 x 10-24 g + 1,67493 x 10-24 g + 9,10939 x 10-28 g.

Mass 2H = 3,34846 x 10-24 g

Contribution of the electron = (9,10939 x 10-28 g/ 3,34846 x 10-24 g) x 100% = 0,0272 %

As can be seen from this calculation, the contribution of the electron to the total mass of

an atom is very small. Usualy this contribution is neglected.

B. The atomic mass unit.

The mass of an atom, expressed in gram, is a very small number. For that reason a new

unit was introduced in order to express such small masses in a more practical way. This

unit is the atomic mass unit (amu). It is defined as 1/12 of the mass of a 12C-isotope.

Since this isotope is made up of 6 protons and 6 neutrons it means that the atomic mass

unit is the mean mass of a proton and a neutron

The value of the amu (rounded) = 1,6 x 10-24 g.

The mass of any atom (or isotope) can be expressed as a multiple of the atomic mass unit.

Example 3 What is the mass of the 2H-isotope expressed in amu?

The mass of this isotope (see higher) = 3,34846 x 10-24 g

Mass 2H expressed in amu = 3,34846 x 10-24 g/1,6 x 10-24 g/amu = 2 amu (rounded)

C. Mean atomic mass.

The mass of a particular atom is in fact the mean mass of all its isotopes taking into

account the abundance and mass of those isotopes.

Example 4 The calculation of the mean atomic mass of chlorine.

Chlorine has the following isotopes:

35Cl: an atomic mass of 34,9688 amu and an abundance 75,53 %

37Cl: an atomic mass of 36,965 amu and an abundance 24,47 %

The mean atomic mass of chlorine =

34,9688 amu x 75,53/100 + 36,965 amu x 24,47/100 = 35,45 amu

This can be done for any atom in the periodic table.

Basic Chemistry 6

D. Relative mean atomic mass.

The relative atomic mass (symbol Ar) is defined as a number that says how many times

an atom is heavier than the amu. This number (it has no units) is given in the periodic

table.

Example 5 What is the mass of an aluminium atom?

In the periodic Table one can find the relative atomic mass of aluminium: 27

A (mean) aluminium atom has a mass of:

Mass Al-atoom = Ar(Al) x amu = 27 x 1,6 x 10-24 g = 4,32 x 10-24 g

5. The mole.

Atoms have a very small mass. This means that in practice very large numbers of atoms

will be involved in chemical or physical processes. A drop of water with a volume of

around 0.05 g contains about 5 x 1021 atoms (hydrogen and oxygen). In order to be able

to handle those large numbers the concept of the mole has been introduced. The mole is

defined as an amount that contains 6,02 x 1023 particles. This number is called

Avogadro’s number (symbol NA). It corresponds to the number of atoms in 12 g of the 12C-isotope. Because this number is so large it only makes sense to use it for very small

particles such as electrons, protons or atoms.

Example 6 How many moles of atoms are present in 0.05 g water?

In 0.05 g water there are 5 x 1021 atoms.

The number of moles in that quantity = number of atoms/NA

The number of atoms in 0,05 g water = 5 x 1021 atoms/ 6,02 x 1023 atoms per mole =

0,00831 mol atoms.

6. Molar mass.

The molar mass of a particle (atom, electron …) is the mass of 1 mole (6.02 x 1023 ) of

those particles. The molar mass (MM) is obtained by multiplying the mass of one

particle by the number of Avogadro. The unit of molar mass is g/mol.

Example 7 What is the molar mass of aluminium?

The relative atomic mass of aluminium (from periodic Table) = 27

The molar mass of aluminium is:

MM(Al) = number of atoms in 1 mole x mass of 1 atom

MM(Al) = NA atoms/mol x Ar(Al) x amu

MM(Al) = 6.02x1023 atoms/mol x 27 amu/atom x 1,6 x 10-24 g/amu

MM (Al) = 27 g/mol

Basic Chemistry 7

As seen from this calculation the absolute value of the molar mass is equal to the value of

the relative atomic mass of that atom. In order to calculate the molar mass of an atom it

is therefore sufficient to look up the relative atomic mass in the table and to add the unit

g/mol. The following table gives some examples.

Table 5 A few examples of the molar mass of atoms.

Atom Ar (rounded, from periodic

table)

1 mole of this atom has a

mass of

O 16 16 g

Al 27 27 g

Si 28 28 g

V 89 89 g

U 238 238 g

Note: remember that the different masses of the last column in the table correspond to the

same number of atoms (6,02 x 1023).

7. The periodic table.

In the periodic table atoms are arranged primarily according to their atomic number.

They are also arranged in such a way that atoms that have similar chemical properties are

close together either horizontally or vertically.

The columns in the table are called groups, the rows are called periods.

Atoms in the same group have similar chemical properties. Groups are therefore given a

number and sometimes also a name. The group that starts with fluorine (F) has number

7A and is called the halogen group.

The periodic table is made of main groups (numbered from IA to VIIA and VIII). The

other groups ( transition groups) carry the suffix B.

Table 6 Information about the maingroups of the periodic table.

Number First atom of the group Name

IA Hydrogen Alkali metals

IIA Beryllium Alkaline earth metals

IIIA Boron Borongroup

IVA Carbon Carbongroup

VA Nitrogen Nitrogengroup

VIA Oxygen Oxygengroup

VIIA Fluorine Halogens

VIII Helium Noble gasses

In the periodic table you can find a lot of information about the atoms.

8. The electronic structure of atoms.

Atoms contain a positively charged nucleus (protons) surrounded by negatively charged

electrons. A neutral atom contains the same number of electrons as protons.

The number of electrons in an atom varies from 1 in hydrogen (Z=1) to 92 in uranium

(Z=92). Those electrons have different energies. Some electrons have a low energy and

Basic Chemistry 8

are closer to the nucleus, others have higher energies and are further away from the

nucleus.

These differences in position can be represented by a model where the electrons are put

in concentric layers. Each layer corresponds to an energy level. Electrons on lower

levels are closer to the nucleus and have a lower energy, electrons on higher levels are

further away from the nucleus.

The electrons on the layer that is the farthest away from the nucleus are called valence

electrons. Those electrons are involved in the interactions (such as the formation of

chemical bonds) between atoms.

The number of valence electrons of an atom corresponds to the number of the group that

contains this particular atom

Table 7 The number of valence electrons (VE) of atoms.

Group First atom of the group Number of VE

IA Hydrogen 1

IIA Beryllium 2

IIIA Boron 3

IVA Carbon 4

VA Nitrogen 5

VIA Oxygen 6

VIIA Fluorine 7

VIII Helium 8

9. Ions.

Normal atoms are neutral because they contain the same number of electrons and protons.

In many cases however atoms, during interactions with other atoms, will loose or absorb

electrons. This occurs during the formation of chemical bonds for example.

When atoms loose or absorb electrons they become ions.

Positive ions (cations) are formed when atoms loose one or more electrons. Such ions

have less electrons than protons and carry a net positive charge. The positive charge is

equal to the number of lost electrons.

Negative ions (anions) are formed when atoms absorb one or more electrons. They have

more electrons than protons and have therefore a net negative charge. The charge is

equal to the number of absorbed electrons.

The number of electrons that a particular atoms can loose or absorb depends on the

number of valence electrons (group number)

A. Positive ions.

Positive ions can be formed by the atoms of the groups IA, IIA en IIIA. They forms ions

with a charge of resp. +1, +2, +3. This is shown in the following table. Some ions of

other main groups are also shown.

Basic Chemistry 9

Table 8 Examples of positive ions of atoms from the main groups.

Group Atom Ion

IA H H+

Li Li+

Na Na+

IIA Be Be+2

Mg Mg+2

Ca Ca+2

IIIA Al Al+3

IVA Pb Pb+2 and

Pb+4

Sn Sn+2 and

Sn+4

Atoms of the transition groups also form positive ions. Some of these ions are shown in

the following table. As can be seen some atoms can from several differently charged

ions.

Table 9 Veel voorkomende ionen van de nevengroepen

Group Atom Ion(s)

IB Cu Cu+ and Cu+2

Ag Ag+

Au Au+ and Au+3

IIB Zn Zn+2

Cd Cd+2

Hg Hg2+2 and Hg+2

VIB Cr Cr+3

VIIB Mn Mn+2

VIIIB Fe Fe+2 and Fe+3

Co Co+2

Ni Ni+2

B. Negative ions.

Negative ions are mostly formed by atoms on the right side of the periodic table. The

most important ones are those formed by the halogens (group 7A). The negative ions of

these atoms are in fact the acid rests of the corresponding binary acids.

Basic Chemistry 10

10. Exercises.

1. Imagine we could enlarge an atom in such a way that the nucleus is as large as a

basketball. How large would that atom be?

2. Imagine that this basketball has the same density as the nucleus of a hydrogen atom.

What would be the mass of this ball?

3. Compete the following table

Table 10 .

Symbol Z A Number of

protons

Number of

neutrons

Number of

electrons

1 3

H+ 2

Cs 55 133

Bi 209

56 138 56

Sn 70

Zn+2 34

17 37 18 238U

4. Faraday’s constant (F) is equal to the charge of 1 mole of electrons. Calculate its

value.

5. Calculate the contribution of the mass of the electrons to the total mass of a 203Hg-

atom.

6. How many valence electrons are present in the 12C-isotope?

Chapter 2. The molecule.

1. Introduction.

A molecule is a particle that contains several atoms. These atoms stick together by

chemical bonds. In this chapter we will discuss how these chemical bonds are formed

and therefore how they give rise to molecules. A molecule is described by its formula

that indicates which and how many atoms are part of the molecule.

2. The chemical bond.

A. Definition.

The chemical bond is an interaction between atoms that results in the fact that these

atoms stay together and form a more or less permanent structure (the molecule).

Chemical bonds can be broken thus allowing chemical reactions to occur. During

chemical reactions chemical bonds are broken and new chemical bonds are formed with

the original atoms.

In these processes the valence electrons play a vital role.

Two types of chemical bonds can be distinguished: the ionic bond and the covalent bond.

B. The covalent bond.

The best way to understand the formation of a covalent bond is to study what happens

when two hydrogen atoms bind together to form a hydrogen molecule (H2). Imagine that

two hydrogen atoms (each containing 1 proton and 1 electron) are at a very large distance

from each other. The only interactions that exist in this situation are the attraction forces

between the proton of an atom and its own electron. Those interactions define the

starting energy of this particular system (see picture, extreme right)

When these atoms get closer to each other (from right

to left in the figure), new interactions start to appear.

The proton of one atom will also start to attract the

electron of the other atom. These attraction forces

decrease the energy of the system and pull the atoms

together.

When the atoms are very close to each other the

protons of the atoms start to repulse each other. This

increases the energy of the system (far left). The

figure shows that there is a distance between the atoms

at which the energy of the system has reached its smallest value. At this distance the

hydrogen atoms have formed a chemical bond. In hydrogen this distance has a value of

about 74 pm. It is called the bond length.

The above discussion shows that the chemical bond is formed because both protons

attract both the electrons of the hydrogen atoms. These electrons are therefore called the

common electron pair or binding pair. The chemical bond whereby an electron pair is

attracted by both atoms is called a covalent bond. In drawings the common electron pair

is usualy represented by a horizontal line between the atoms ( H—H).

Basic Chemistry 2

C. The polarity of a covalent bond.

When a covalent bond is formed between two identical

atoms (such as hydrogen) the common electron pair is

attracted by both atoms with the same strength. This

pair will be distributed in a symmetrical way between

both atoms.

When two different atoms bind together (such as

hydrogen and fluorine) each atom will attract the

common electron pair with a different force. The

elctron pair will be displaced towards the atom that

attracts the pair more strongly.

The atoms that pulls the pair towards itself will therefore receive a slightly negative

charge, whereas the other one will receive a slightly positive

charge. Those charges are smaller than 1 and are represented

by the symbols - or +.

The covalent bond that is formed in this way is called a polar

covalent bond. It has two poles ( a negative and a positive

one). It is also called a dipole. The strenght of the dipole is

given by the dipole moment. This is calculated as the absolute

value of the charge of one pole (both poles have the same

absolute value) multiplied by the distance between the poles

(the bond length).

The bond between identical atoms is called non-polar because of its symmetry.

D. Electronegativity.

In order to determine the polarity of a covalent bond, one needs to know which one of the

atoms pulls more strongly on the common electronpair. This is given by the

electronegativity (EN) of an atom. This value, between 0 and 4, is usualy mentionned in

the periodic table. Following table shows some of those values.

Table 1 The elektronegativity of some atoms.

H

2,2

Li

1,0

Be

1,5

B

2,0

C

2,5

N

3,0

O

3,5

F

4,0

Na

0,9

Al

1,5

Si

1,8

P

2,1

S

2,5

Cl

3,0

K

0,8

Br

2,8

As the difference in electronegativity between two

atoms increases so the polarity of the covalent bond

will also increase. When two identical atoms bind

this difference is zero and the bond is non-polar.

E. The ionic bond.

Basic Chemistry 3

The ionic bond is an extreme case of a polar bond. It is formed when the difference in

electronegativity between the binding atoms is so strong that the binding electrons are

completely displaced towards one of the atoms. This atom will therefore become a

negative ion while the other one becomes a positive ion. These two oppositely charged

ions will attract each other and form an ionic bond. An example is the formation of the

ionic bond between sodium and chlorine.

Example 1 How is the bond between sodium and chlorine formed?

Sodium is an atom with 1 valence electron and a low electronegativity.

Chlorine is an atom with seven VE and a very high electronegativity.

The chlorine atom pulls one electron from sodium and gets a charge of -1. The sodium

atom gets a charge of +1.

The Cl- ion and the Na+ ion attract each other because they have opposite charges.

When the difference of electronegativity between two atoms is equal to 1.7 or larger the

bond between those atoms will be ionic. This means that most ionic bonds are formed

between atoms with a very low EN (left in the table) and atoms with a high EN (on the

right in the table).

Example 2 What is the bond between H en O?

The electronegativity of these atoms is (see table):

EN(H) = 2.2

EN(O) = 3.5

The difference EN = 1.3

EN is larger than zero but smaller than 1.7.

The bond between H and O is a polar covalent bond.

3. The molecular formula.

The formula of a molecule describes what atoms and in what numbers are present in a

molecule.

Example 3 What is the composition of a molecule of sulfuric acid (H2SO4)?

One molecule of sulfuric acid contains 2 hydrogen atoms, 1 sulfur atom and 4 oxygen

atoms. These atoms are bound together by chemical bonds..

Note: the molecular formula gives no information on the arrangement of the atoms in this

molecule or about the spatial structure of the molecule.

Basic Chemistry 4

4. Molecules and ions.

Water (H2O) and kitchen salt (NaCl) are very different

compounds. Water is made up of a large quantity of separate

particles (molecules) each made up of two hydrogen atoms

and one oxygen atom. The bonds

between those atoms are covalent

bonds. Kitchen salt on the other

hand is made up entirely of

positively charged sodium atoms

(ions) and negatively charged chlorine atoms (ions). Those ions

are arranged in a lattice and kept together by coulomb attraction

forces (ionic bonds). This means that the formula H2O is a real

representation of a water particle whereas NaCl only gives the

ratio of these ions in the compound. In the next part of this text we will work with the

formula NaCl as if it were a real molecular formula.

5. Molecular mass.

A. Absolute molecular mass.

The mass of a molecule is equal to the sum of the mass of all its atoms.

Example 4 What is the mass of a watermolecule?

A water molecule (H2O) is made up of 2 hydrogen atoms and one oxygen atom.

Mass water molecule = mass oxygen atom + 2x mass hydrogen atom.

Mass water molecule = Ar(O) x amu + 2 x Ar(H) x amu

Mass water molecule = 16 amu + 2 x 1 amu

Mass water molecule = 18 amu

Mass water molecule = 18 x 1,6 x 10-24 g = 2,88 x 10-23 g

B. Relative molecular mass.

In the same way as with atoms the mass of a molecule can be given by a number that says

how many times its mass is heavier than the amu. It is called the relative molecular mass

(Mr). It is equal to the sum of the relative atomic masses of the atoms that make up this

particular molecule. The relative molecular mass of water is:

Mr (H2O) = Ar(O) + 2 x Ar(H) = 18.

6. Molar mass of molecules.

The molar mass of a molecule is the mass of 1 mol of these molecules. It can be obtained

by calculating the relative molecular mass and adding the unit g/mol

Basic Chemistry 5

Example 5 What is the molar mass of water?

Mass of one molecule water = 18 amu.

The mass of 1 mole of water = mass of 1 molecule x NA

Molar mass (H2O) = 18 x amu x NA

MM(H2O) = 18 g/mol

7. The oxidation state of an atom in a molecule.

When atoms bind together to form chemical bonds they use their valence electrons. They

will (partially) lose or (partially) absorb electrons. The oxidation state is a number that

describes this process of donating or absorbing electrons. The oxidation state of any

atom in a molecule can be calculated by using a few simple rules. The oxidation states

are usually written as roman numericals in order to distinguish them from ionic charges.

Table 2 Rules for determining the oxidation state (OT) of an atom.

OT of atoms not bound to an other (different) atoms = 0

OT of hydrogen in a molecule is usually +I

OT of oxygen in a molecule is usually –II

OT of the atoms of groups IA, IIA and IIIA are +I, +II and +III resp.

The sum of the OT of the atoms in a molecule multiplied by the number of

atoms = 0

The sum of the OT of the atoms in an ion multiplied by the number of atoms =

the charge of the ion

Using these rules it is possible to calculate the OT of all the atoms in most of the

molecules and ions in this course.

Example 6 What are the oxidation states of the atoms in H2SO4?

The oxidation states of H and O are resp +I and –II.

The sum of these oxidation states are therefore = 2 x (+I) + 4 x (-II) = -VI

Because the sum of the OT’s must be equal to zero (molecule) the OT of S = +VI.

Answer: OT(H) = +I, OT(O) = -II and OT(S) = +VI

Example 7 What are the oxidation states of the atoms in NH4+?

The oxidation state H = +I which gives a total 4 x (+I) = +IV

Because the ion has a charge +1, the sum of all OT’s should be = +I

Answer: The OT of N = -III.

Basic Chemistry 6

8. Exercises.

1. What kind of bonds are the following: H-Cl, N-H, O-O, K-Cl? Are they polar covalent,

non-polar covalent or ionic?

2. Arrange the following bonds in order of increasing polarity and also give the direction

of the polarity: C-H, H-H, H-Br, H-F en B-H

3. How many atoms are there in a molecule Ca3(PO4)2?

4. Calculate the mass of a propane molecule (C3H8)

5. What is the molar mass of sulfuric acid (H2SO4)?

6. How many moles is 1 kg of water?

7. How many grams of sulfuric acid contain the same number of molecules as 500 g of

propane?

8. How many grams of K is there in 150 g of KNO3?

9. Calculate the oxidation state of every atom in the following particles: K2SO4, HNO3,

CrO4-2, KMnO4, HSO4

-.

Chapter 3. Naming chemical compounds.

1. Classification of chemical compounds.

Usually chemical compounds are classified according to their properties. In this chapter

we will discuss what are the important classes of chemical compounds and how they are

named.

2. Important classes of chemical compounds.

In this course we will mostly be working with following important classes:

- acids,

- bases and hydroxides,

- salts,

- oxides.

The compounds of each class will have very specific properties that can be observed

during chemical reactions.

A. Acids.

A.1. Properties of acids.

Acids are chemical compounds that have been known (and used) by man for a long time.

This is certainly so because of the typical taste they have. Some examples of such

“ancient” acids are

- Acetic acid formed when wine turns acidic,

- Lactic acid that is produced during the acidification of milk,

- The acid that is formed in the stomach and that can sometimes be tasted.

Acids are compounds that are able to produce a positively charged hydrogen ion (H+, a

proton). Their formula therefore contains one or more hydrogen atoms and can be

represented in a general way as:

HnA

In this formula n is equal to 1, 2 or 3. When n=1 the acid is called monoprotic, when n is

larger than 1 the acid is called polyprotic.

In the formula A is called: the acid rest.

The classification of acids is base on the composition of the acid rest. This acid rest

always contains at least one non-metal (or a metal that behaves as a non-metal).

Examples are chlorine, sulfur, phosphorus, manganese a.s.o.

When the acid rest does not contain oxygen it is called a binary acid. When it contains

oxygen it is called an oxoacid or a ternary acid.

Rem.: there are many compounds that contain hydrogen atoms that are not acidic. An

example is methane (CH4). A hydrogen atom that has no acidic properties is called a

non-acidic hydrogen.

Basic Chemistry 2

A.2. Binary acids

The acid rest of binary acids contains a non-metal (and no oxygen). The name of such an

acid is given as follows:

Name of a binary acid = hydrogen + non-metal + -ide.

The name of the acid rest is derived from the name of the acid without hydrogen.

The following binary acids and their acid rests are very common and important.

Table 1 Some important binary acids and their acid rests.

Formula Name Acid rest Name acid rest

HF hydrogen fluoride F- fluoride(ion)

HCl hydrogen chloride Cl- chloride(ion)

HBr hydrogen bromide Br- bromide(ion)

HI hydrogen iodide I- iodide(ion)

H2S hydrogen sulfide HS-

S2-

hydrogen sulfide(ion)

sulfide(ion)

HCN hydrogen cyanide CN- cyanide(ion)

In english binary acids in solution are also given another name: HCl in solution is also

called hydrochloric acid (more examples in the handbook table 2.5)

Sometimes the word –ion is added to the name of the acid rest to make clear that it is

indeed a negatively charged particle. When not all hydrogen atoms have been removed

the number of remaining hydrogen atoms is added to the name (example is the

hydrogensulfide ion).

Rem.: hydrogen cyanide is sometimes called a pseudo-binary acid because its acid rest

contains two non-metals.

A.3. Oxoacids.

The acid rest of oxoacids contains a non-metal with at least one oxygen atom. For these

acids two names are commonly used. They are formed as follows:

Name a: name of non-metal + ic acid. (HClO3: chloric acid).

Name b: hydrogen non-metal + -ate. (HClO3: hydrogen chlorate)

The problem with oxoacids is that very often two or more oxoacids contain the same non-

metal but a different number of oxygen atoms. In order to name these acids we start with

the normal name (above) and change it following some simple rules ( see also handbook

page 49)

Basic Chemistry 3

Example 1 What are the names of the different oxoacid containing chlorine?

There are four different oxoacids with chlorine: HClO, HClO2, HClO3, en HClO4. The

normal acid is HClO3. It is named in the usual way.

HClO3 is called hydrogen chlorate or chloric acid.

HClO4 contains one more oxygen atom and is called hydrogen perchlorate or perchloric

acid.

HClO2 contains one less oxygen atom and is called hydrogen chlorite or chlorous acid.

HClO contains even less oxygen and is called hydrogen hypochlorite or hypochlorous

acid.

The same principles apply to other acids. What acid carries the normal name is not

determined by its formula but by the history of their discoveries.

The names of the acid rests are derived from name b.

The following table shows the names and formulas of some important acids.

Basic Chemistry 4

Table 2 Lijst met belangrijke oxozuren.

Non-metal

in acidrest

Formula Name Acidrest Name of the acidrest

Carbon (C) H2CO3 Carbonic acid

Hydrogen carbonate

HCO3- Hydrogen carbonate(ion)

CO32- Carbonate(ion)

Nitragen

(N)

HNO3 Nitric acid

Hydrogen nitrate

NO3- Nitrate(ion)

HNO2 Nitrous acid

Hydrogen nitrite

NO2- Nitrite(ion)

Fosfor (P) H3PO4 Phosphoric acid

Hydrogen phosphate

H2PO4- Dihydrogen phosphate(ion)

HPO42- Monohydrogen phosphate(ion)

PO43- Phosphate(ion)

H3PO3 Phosphorous acid

Hydrogen phosphite

H2PO3- Dihydrogen phosphite(ion)

HPO32- Monohydrogen phosphite(ion)

PO33- Phosphite(ion)

Arseen

(As)

H3AsO4 Arsenic acid

Hydrogen arsenate

H2AsO4- Dihydrogen arsenate(ion)

HAsO42- Monohydrogen arsenate(ion)

AsO43- Arsenate(ion)

H3AsO3 Arsenous acid

Hydrogen arsenite

H2AsO3- Dihydrogen arsenite(ion)

HAsO32- Monohydrogen arsenite(ion)

AsO33- Arsenite(ion)

Zwavel (S) H2SO4 Sulfuric acid

Hydrogen sulfate

HSO4- Hydrogen sulfate(ion

SO42- Sulfate(ion)

H2SO3 Sulfurous acid

Hydrogen sulfite

HSO3- Hydrogen sulfite(ion)

SO32- Sulfite(ion)

H2S2O3 Hydrogen thiosulfate HS2O3- Hydrogen thiosulfate(ion)

S2O32- Thiosulfate(ion)

Basic Chemistry 5

Chlorine

(Cl)

HClO4 Perchloric acid

Hydrogen perchlorate

ClO4- Perchlorate(ion)

HClO3 Chloric acid

Hydrogen chlorate

ClO3- Chlorate(ion)

HClO2 Chlorous acid

Hydrogen chlorite

ClO2- Chlorite(ion)

HClO Hypochlorous acid

Hydrogen

hypochlorite

ClO- Hypochlorite(ion)

Bromine

(Br)

HBrO4 Perbroomacid

Hydrogen perbromate

BrO4- Perbromate(ion)

HBrO3 Bromic acid

Hydrogen bromate

BrO3- Bromate(ion)

HBrO2 Bromous acid

Hydrogen bromite

BrO2- Bromite(ion)

HBrO Hypobromous acid

Hydrogen

hypobromite

BrO- Hypobromite(ion)

Iodine (I) HIO4 Periodic acid

Hydrogen periodate

IO4- Periodate(ion)

HIO3 Iodic acid

Hydrogen iodate

IO3- Iodate(ion)

HIO2 Iodous acid

Hydrogen iodite

IO2- Iodite(ion)

HIO Hypoiodous acid

Hydrogen hypoiodite

IO- Hypoiodite(ion)

Some acids contain a metal instead of a non-metal. You can find them in the following

table. The table also contains acetic acid which is an organic acid. It has a different

structure but behaves in the same way as many inorganic acids.

Table 3 Oxoacids with deviating composition.

Non-metal

in acidrest

Formula Name Acidrest Name of the acidrest

Manganese

(Mn)

HMnO4 Permanganic acid

Hydrogen

permanganate

MnO4- Permanganate(ion)

Chromium

(Cr)

H2CrO4 Chromic acid

Hydrogen chromate

HCrO4- Hydrogen chromate(ion)

CrO42- Chromate(ion)

H2Cr2O7 Dichromic acid

Hydrogen dichromate

HCr2O7- Hydrogen

dichromate(ion)

Cr2O72- Dichromaat(ion)

Carbon (C) CH3COOH Acetic acid

Hydrogen acetate

CH3COO- Acetate(ion)

Basic Chemistry 6

B. Hydroxides and bases.

B.1. Hydroxides.

Hydroxides are compounds made up of positively charged metal ion and one or more

OH-groups. The OH-group is called the hydroxide group It has a charge of –1.

Name of the hydroxides: name of the metal + hydroxide.

The number of OH-groups in the compound is determined by the charge of the metal.

If several ions exist of the same metal (such as Fe2+ and Fe3+) then the name of the

hydroxide should be made more specific. This can be done by giving the charge of the

ion or the number of OH-groups. Some examples are given in the following table.

Table 4 Some metal hydroxides.

Formula Name

NaOH Sodium hydroxide

Ba(OH)2 Barium hydroxide

Fe(OH)2 Iron(II) hydroxide*

Iron dihydroxide

Fe(OH)3 Iron(III) hydroxide

Iron trihydroxide

Al(OH)3 Aluminium hydroxide

* pronounced: iron two hydroxide.

B.2. Difference between base and hydroxide.

As mentionned above hydroxides are defined by the presence of one or more OH-groups

in the formula.

Bases on the other hand are defined by their chemical properties. They react with acids

for example.

Some hydroxides are bases (react in the same way as bases) some are not. The

hydroxides that have a basic behaviour are the ones where the metal is from group IA or

group IIA.

There are also compounds that are bases but that have a formula that is completely

different. Ammonia (NH3) is such a compound.

Table 5 Examples of bases and hydroxides.

Compound Is a…

Sodium hydroxide Base

Iron(II) hydroxide Hydroxide (is not a base)

Ammonia Base

Calcium hydroxide Base

The ammonium ion (NH4+) is the positive ion that is associated with the base ammonia.

Basic Chemistry 7

B.3. Salts.

Salts are compounds made up of a positive ion (metal or ammonium) and a negative

group (acid rest). These two groups are combined in such a way that the total charge of

the molecule is zero. The formation of a salt can be seen as the replacement of one or

more hydrogens from an acid by another positive group. When not all the acidic

hydrogens have been replaced the salt is called an acidic salt.

Name of the salt: name of the positive group + name of the acid rest.

The following table shows some examples.

Table 6 Some examples of salts with their name.

Formula Name

KCl Potassium chloride

Na2SO4 Sodium sulfate

NaHSO4 Sodium hydrogensulfate

Ca3(PO4)2 Calcium phosphate

NH4Cl Ammonium chloride

FeSO4 Iron(II) sulfate

Fe2(SO4)3 Iron(III) sulfate

NaH2PO4 Sodium dihydrogenphosphate

B.4. Oxides.

Oxides are compounds that combine an element with oxygen. Most atoms can form one

or more oxides. They can for example be formed during combustion reactions.

Name: name of the atom + oxide.

If an atom can form many oxides the name should be made more specific either by giving

the exact composition of the oxide or by giving the charge of the atom.

Oxides are often separated in metal oxides and non-metal oxides

The formula of metal oxides can be deduced very easily from the charge of the metal

(usualy group number)

In this text we will only be considering non-metal oxides in which the oxidation state of

the non-metal is the same as in the known oxoacids

Table 7 Some examples of metal oxides.

Group Formula Name

group I Na2O Sodium oxide

group II MgO Magnesium

oxide

group III Al2O3 Aluminium

oxide

Transition MnO2 Manganese(IV) Manganese dioxide

Basic Chemistry 8

atoms oxide

FeO Iron(II) oxide

Fe2O3 Iron(III) oxide

HgO Mercury(II)

oxide

Monomercury

monooxide

Table 8 Some examples of non-metal oxides with their corresponding oxoacids.

Group Formula Name Name Oxoacid

group IV CO2 Carbon(IV) oxide Carbon dioxide H2CO3

group V N2O5 Nitrogen(V) oxide Dinitrogen pentaoxide HNO3

group VI SO2 Sulfur(IV) oxide Sulfur dioxide H2SO3

SO3 Sulfur(VI) oxide Sulfur trioxide H2SO4

group VII Cl2O7 Chlorine(VII) oxide Dichlorine heptaoxide HClO4

Table 9 Enkele andere bestaande niet-metaaloxiden.

Group Formula Name group I H2O (hydrogen oxide) water

group IV CO Carbon(II) oxide Carbon monoxide

group V N2O Nitrogen(I) oxide Dinitrogen

monoxide

group VIII XeO3 Xenon(VI) oxide Xenon trioxide

3. Exercises

1. Name the following compounds:

FeO, K2Cr2O7, As2S3, Ba(NO3)2, KClO3, AgCl, LiOH, KNO2, H2S, KMnO4.

2. Write the formula of the following compounds:

Aluminium oxide, copper(I) sulfate, dicopper sulfate, sodium nitrite, iron(III) oxide,

tin(IV) chloride, barium carbonate, ammonium chloride, dinitrogen trioxide, potassium

hydrogensulfate.

Chapter 4. Properties of compounds.

1. Introduction.

When studying the behavior of chemical compounds one can consider two major

possibilities: the compound can be pure or it can be mixed with other compounds such as

when dissolved in water.

The study of pure compounds mainly concerns their different states (solid, liquid,

gaseous) and how they go from one state to the other (melting, evaporating a.s.o.).

In this part of chemistry basics we will be considering what happens when compounds

are mixed with water.

2. Water.

Water is a very important chemical compound. It is pesent on the surface of the earth in

very large quantities but it is also a major constituent of living organisms. It is very often

used as solvent in chemical reactions and processes.

When compounds interact with water two phenomena can be observed. On the one hand

there is their solubility and on the other hand their electrolyte behaviour. Electrolyte

behavior takes into consideration how well a compounds forms ions in a solution. It is

off course related to its solubility.

3. Solubility.

Solubility is defined as the maximum quantity of a compound can be dissolved in a

certain quantity of a solvent (water) at a certain temperature. Solubility is often

expressed in gram/liter but any other unit can also be used. Although the values for the

solubility of different compounds varies enormously, they will usually be divided in two

categories: poorly soluble and highly soluble.

The solubility of poorly soluble compounds (often also called non-soluble) are below a

certain value (such as 1g/l) whereas the others are above that value.

The following table shows the solubility of several important chemical compounds. In

this table a higher placed rule takes precedence over a lower one.

Table 1 Solubility of compounds in water

1. All sodium, potassium and ammonium salts and all nitrates are soluble.

2. All silver, lead(II) and Hg22+ salts are poorly soluble exept the nitrates (above).

3. Alle (per)chlorates, acetates, chlorides, bromides and iodides are soluble exept

exceptions (higher).

4.All carbonates, sulfides and phosphates are non-soluble.

5. All metal oxides and hydroxides are non-soluble except those of sodium, potassium,

lithium.

6. All sulphates are soluble except those of calcium and barium and ions mentionned

above.

Basic Chemistry 2

4. Electrolyte behavior.

When a compound dissolves in water it can either remain in it molecular state or it can

split into ions.

A compound that does not split into ions is called a non-electrolyte. Sugar is an example

A compound that separates into ions is called an electrolyte. Electrolytes can be weak or

strong depending on how well they separate into ions. Strong electrolytes will split for

almost 100 % whereas non-electrolytes will only form a small amount of ions.

The difference between strong and weak electrolytes can be determined by the

conductivity of their solutions. Solutions of non-electrolytes do not conduct electricity,

solutions of weak electrolytes only slightly and solutions of strong electrolytes conduct

electricity very well.

A. Electrolyte behavior of acids.

Most acids dissolve very well in water but are also weak electrolytes. They are called

weak acids. Only some acids are strong electrolytes. The are called strong acids. The

following table shows which acids are strong. All other acids are weak.

Table 2 The strong acids

HI, HBr, HClO4, HCl, H2SeO4, H2SO4, HMnO4, HNO3, H2CrO4, HClO3

B. Electrolyte behavior of hydroxides and bases.

Most metal hydroxides are non-soluble and are therefore weak electrolytes. Soluble

hydroxides (NaOH, LiOH, KOH) are strong electrolytes.

Ammonia (NH3) is very soluble but a weak electrolyte.

C. Electrolyte behavior of salts.

The electrolyte behavior of salts corresponds to their solubility. Soluble salts are strong

electrolytes, non-soluble salts are weak electrolytes.

D. Electrolyte behavior of oxides.

Oxides that dissolve in water will usually react with water and form the corresponding

hydroxides or acids (see next chapter). How many ions are formed in either case depends

on the electrolyte behavior of the compound that is formed by that reaction.

Chapter 5. Chemical Reactions.

1. Definition.

A chemical reaction is a process in which a substance (or substances) is changed into one

or more new substances. A chemical reaction may be simple, but can also be very

complex.

The starting materials in a chemical reaction are called reactants, the substances formed

as a result of the chemical reaction are referred to as the products.

In a chemical equation the reactants are conventionally written on the left and the

products on the right side of the arrow:

reactants products

2. Law of conservation of mass.

During a chemical reaction no matter is produced or destroyed. To conform with this law

of conservation of mass there must be the same number of each type of atom on both

sides of the arrow; that means we must have as many atoms after the reaction ends as we

did before it started.

3. The Reaction Equation.

A chemical equation shows which compounds react with each other, and which products

are formed. A balanced chemical equation is conform to the law of conservation of

mass meaning that the number and type of atoms present in the products must be equal to

the number and type of atoms present in the reactants. A balanced chemical equation

then provides information about the ratio in which reactants and products react.

Stoichiometric coefficients are used to balance chemical equations . These are whole

numbers written before the formulas of the molecules in the chemical equation. You can

change the coefficients, but not the subscripts to balance a reaction.

Example 1 illustrates this method.

Example 1 The reaction equation for the synthesis of ammonia.

Ammonia is an important chemical compound that is formed in the so called Haber

process with hydrogen and nitrogen gas as reactants.

The process can be written as follows:

N2 + H2 NH3

However, the number of hydrogen and nitrogen atoms on the right and left side of the

equation is not the same.

Placing appropriate coefficients brings the equation in balance:

N2 + 3 H2 2 NH3

Basic chemistry 2

4. Classes of Chemical Reactions.

In general chemical equations can be subdivided in two important groups: reactions

during which the oxidation states of the atoms do not change and reactions during which

the oxidation states do change.

Reactions involving varying oxidation states of the atoms are called oxidation-reduction

reactions or simply redox reactions.

A. Reactions without Change in Oxidation State.

A.1. Introduction.

During these reactions the oxidation states of the atoms don’t change. This concept must

be used to check the correctness of the chemical equation. It also sometimes allows us to

predict the compounds that are formed during the reaction e.g. in reactions of oxides (see

below)

A.2. Reactions with Oxides.

(1) In General

In reactions of oxides charged particles (ions) are not involved. As such these reactions

differ from the other reactions that will be discussed in this section (for instance the

reactions between acids and bases) in which ions react with each other.

(2)Oxides with Water

(a)Metal oxides

In general: oxides of alkaline- and earth alkaline metals (group IA en IIA) react with

water to produce hydroxides. They are called the base forming oxides. The other

oxides don’t react.

Example 2 Reactions of metal oxides with water.

Na2O + H2O 2 NaOH

Fe2O3 + H2O no reaction

(b)Non-metal oxides.

In general: during the reaction of a non-metal oxide with water, the corresponding oxo-

acid is formed. That’s why these oxides are called acid forming oxides. The oxidation

state does not change, this allows us to chose the correct oxo-acid of the non-metal

(sometimes different oxo-acids, with different oxidation states of the non-metal, exist)

Example 3 Reactions of non-metal oxides with water.

SO3 + H2O H2SO4 (not H2SO3).

P2O5 + H2O 2 H3PO4

Basic chemistry 3

As shown above, non-metal oxides can be associated with their corresponding oxo-acids.

The same can be done with metal oxides, resulting in the formation of hydroxides. Some

of these oxides however don’t react with water e.g. Fe2O3 which could be associated with

Fe(OH)3.

It is important to remember this because is enables you to understand the way oxides

react with other compounds.

(3)Oxides with Oxides.

In general: when metal-oxides react with non-metal oxides, salts are formed. The acid

rest is derived from the oxo-acid derived from the non-metal oxide.

Example 4 Reaction of oxides with oxides

Na2O + SO3 Na2SO4

(4)Oxides with Acids.

In general: when metal-oxides react with acids, salts and water are produced.

Example 5 Reaction of an oxide with an acid.

Na2O + 2 HCl 2 NaCl + H2O

Fe2O3 + 6 HCl FeCl3 + 3 H2O

Note: this reaction occurs with all metal oxides, contrary to their reaction with water.

(5)Oxides with Hydroxides.

In general: when non-metal oxides react with hydroxides, salts and water are produced.

The acid rest is derived from the oxo-acid coming from the non-metal oxide.

Example 6 Reactions of Oxides with Hydroxides.

SO3 + Ca(OH)2 CaSO4 + H2O

A.3. Thermolysis.

In general: During thermolysis reactions compounds are dissociated under the influence

of heat. These reactions may not be confused with combustion reactions in which

oxygen is a reactant and heat is released. The thermolysis of salts, oxo-acids and

hydroxides results in the formation of the corresponding metal and/or non-metal oxides

and eventually water. The reactions can be regarded to as the opposite reactions of the

reactions discussed above (oxides with oxides and oxides with water).

Basic chemistry 4

Example 7 Thermolysis

CaCO3 + heat CaO + CO2

Cu(OH)2 +heat CuO + H2O

H2CO3 + heat CO2 + H2O

Note.: the oxidation states of the atoms involved do not change.

A.4. Metathesis Reactions.

(1)Introduction.

Metathesis reactions occur because the ions brought in the reaction medium (aqueous

solution) bind and form a new compound. This compound may be a precipitate, a weak

electrolyte or it might disappear from the solution as a gas. The reactions only occur

when a new compound is formed and as such ions disappear from the solution. If not, the

ions remain in the solution and no reaction takes place.

In general, these reactions can be written as follows:

AB + CD AD + CB

As can be seen from this equation, the two negative groups, represented by B and D

(these may be acid-rests or OH--groups) are exchanged.

The reaction occurs if at least one of the products (compounds AD and/or CB) is a weak

electrolyte, a precipitate or a gaseous compound.

If this as not the case, the reaction may be written as follows:

AB + CD no reaction

(2)Reactions with the Formation of a Precipitate.

During these reactions insoluble compounds (often salts) are formed. These compounds

are produced when ions present in the solution before reaction, combine to produce an

insoluble compound.

Basic chemistry 5

Example 8 Reactions with formation of a precipitate

1. Reaction of silver nitrate with sodium chloride

AgNO3 + NaCl AgCl + NaNO3

This reaction occurs because silver chloride is insoluble

1. Reaction of potassium nitrate with sodium chloride

KNO3 + NaCl no reaction because KCl en NaNO3 are both soluble

3. Reaction of potassium hydroxide with iron(III)chloride.

3 KOH + FeCl3 Fe(OH)3 + 3 KCl

This reaction occurs because iron(III)hydroxide is insoluble

(3)Reactions with the Formation of a Weak Electrolyte.

These reactions occur because a weak acid or water are formed. Remind that the

insoluble salts of the previous paragraph are also weak electrolytes.

Example 9 Formation of weak electrolytes

1. Reaction of iron(II)sulphide with hydrogen chloride

FeS + 2 HCl FeCl2 + H2S

This reaction occurs because hydrogen sulfide is a weak electrolyte.

2. Reaction of sodium hydroxide with nitric acid

NaOH + HNO3 NaNO3 + H2O

This reaction occurs because water is a weak electrolyte

(4)Reactions with Formation of Gaseous Compounds.

In these reactions one of the reaction products is H2CO3, H2SO3 or H2S. H2CO3 and

H2SO3 are unstable and will dissociate at low temperatures (thermolysis). CO2 + H2O

or SO2 + H2O will then be formed.

Example 10 Formation of gaseous compounds.

1. Sodium carbonate with hydrogen chloride

Na2CO3 + 2 HCl 2 NaCl + H2O + CO2 (and not 2 NaCl + H2CO3)

2. Potassium sulfite with sulfuric acid.

K2SO3 + H2SO4 K2SO4 + H2O + SO2

A special case of these reactions are the reactions between ammonium salts and bases. In

these reactions gaseous ammonia is formed (and not ammonium hydroxide, because this

compound does not exist!).

Basic chemistry 6

Example 11 The formation of ammonia

Reaction of ammonium chloride with potassium hydroxide.

NH4Cl + KOH KCl + NH3 + H2O (and not NH4OH).

(5)The Essential Reaction Equation.

In metathesis reactions, some ions will take part in the reaction to produce a precipitate, a

weak electrolyte or a gaseous compound. Other ions however do not take part in the

reaction, they are called the spectator ions.

The essential reaction equation is the equation in which the spectator ions are not written.

Only the ions (or compounds) that are really involved in the reaction are mentioned. To

derive the essential equation, first write down the correct molecular equation. Then write

the ionic equation by dissociating all strong electrolytes in their corresponding ions. The

ions that occur on both sides of the resulting ionic equation are spectator ions and may be

omitted when the essential equation is written. Weak electrolytes or slightly soluble

compounds are written as such in the essential equation (on both sides).

Example 12 Writing an essential equation

Reaction of silver nitrate with sodium chloride

Molecular equation: AgNO3 + NaCl AgCl + NaNO3

Ionic equation: Ag+ + NO3- + Na+ + Cl- AgCl + Na+ + NO3

-

Essential equation: Ag+ + Cl- AgCl

Note that the reactions that do not occur don’t have an essential equation.

B. Reactions with Change in Oxidation State.

B.1. Introduction

When the oxidation state of the atoms changes during reaction, the reaction belongs to

the oxidation-reduction reactions or redox reactions. Increase of the oxidation state of an

atom during reaction is referred to as oxidation, decrease of the oxidation state of an atom

during reaction is called reduction. Oxidation and reduction always occur together in a

redox reaction. The reason for this is that oxidation of an atom means that the atom loses

electrons (becomes more positive), while reduction of an atom means that the atom gains

electrons (becomes more negative). Oxidation-reduction reactions are thus considered

electron transfer reactions.

In this category of reactions, the combustion reaction and the reaction of non-noble

metals with acids, can easily be written.

However for other, more complex oxidation-reduction reactions, it is much more

difficult to write down the reaction equation. In this case a procedure, consisting of

different steps, is used.

B.2. Combustion Reactions.

Combustion reactions are reactions of atoms or compounds with oxygen (O2). Although

these combustion reactions may be very complex, we will assume that during these

Basic chemistry 7

reactions each atom present in the reactant will be turned into an oxide. When an

element may form several possible oxides, only the oxide with highest oxidation state is

written.

Example 13 Some combustion reactions.

1 Combustion of iron

4 Fe + 3 O2 2Fe2O3 (and not FeO)

2. Combustion of methane (CH4)

CH4 + 2 O2 CO2 + 2 H2O

3 Combustion of C6H5NO2Cl

4 C6H5NO2Cl + 37 O2 24 CO2 + 10 H2O + 2 N2O5 + 2 Cl2O7

Note: no ions are involved in a combustion reaction, so no essential equation can be

written.

B.3. Reactions of non-noble metals with acids.

When metals (except noble or semi-noble metals) react with acids a salt and hydrogen

gas are formed (H2). Hydrogen evolves from the reaction mixture.

Example 14 Zinc reacts with hydrogen chloride

Zn + 2 HCl ZnCl2 + H2

The essential reaction is: Zn + 2 H+ Zn2+ + H2

Note: noble metals (Au, Pt, Ag) or semi-noble metals (Cu, Hg) react differently with

acids. These reactions belong to the complex oxidation-reduction reactions.

As can be seen in the previous reactions, the oxidation state of several atoms changes,

some increase (Fe, C, Zn) others decrease (O, H).

B.4. Complex Oxidation-Reduction Reactions.

These reactions must be balanced using the procedure outlined in this section.

Several procedures exist to balance these reactions. Here only one them, called the half-

reaction method, will be discussed. When using this method, it is not necessary to know

the exact oxidation states of the atoms, It is sufficient to know which oxidation state

increases and which one decreases. Another advantage of this procedure is that it leads to

half-reactions that are important in electrochemical cells.

In this approach, the overall reaction is divided into two half-reactions, one for oxidation

and one for reduction. The half-reactions are written as essential equations, so the

spectator ions are not mentioned. The equations for the two half-reactions are balanced

separately and then added together to give the overall balanced equation.

(1) Procedure

Basic chemistry 8

- Write the unbalanced equation for the reaction in ionic form. Use the rules for

solubility and electrolyte behaviour.

- Identify the atoms with changing oxidation state

- Separate the equation into two half-reactions.

- Balance each half-reaction for number and type of atoms and charges.

o First, balance the atoms that change in oxidation state

o Then, balance oxygen and hydrogen atoms

o If necessary, balance other atoms

o Add electrons to balance the charge.

- Multiply the coefficients of each half-reaction with a whole number until the

amount of electrons in both half-reactions is equal.

- Add the two half-reactions, the electrons on both sides of the final equation will

cancel. Probably other compounds may be cancelled also. This results in the

essential overall equation.

The procedure used to balance oxygen and hydrogen atoms depends upon the fact

whether the reaction occurs in acidic or basic medium. In acidic medium water

molecules and H+-ions are used, in basic medium H2O and OH--ions are used.

(2) An Oxidation-Reduction Reaction in Acidic Medium

For reactions in acidic medium H2O is added to balance the O atoms and H+ to balance

the H atoms.

Example 15 Reaction in acidic medium.

When iron(II)chloride reacts in acidic medium with potassium permanganate, Fe3+ and

Mn2+ are formed

The unbalanced equation:

FeCl2 + KMnO4 in acidic medium => Fe3+ + Mn2+

Changes:

OS(Fe) increases (oxidation), OS(Mn) decreases (reduction)

Division in two half-reactions:

Oxidation: Fe2+ => Fe3+ (note: Fe2+ and not FeCl2)

Reduction: MnO4- => Mn2+ (note: MnO4

- and not KMnO4)

Balance oxygen atoms:

Oxidation: Fe2+ => Fe3+

Reduction: MnO4- => Mn2+ + 4 H2O

Balance hydrogen atoms:

Oxidation: Fe2+ => Fe3+

Reduction: MnO4- + 8 H+ => Mn2+ + 4 H2O

Basic chemistry 9

Balance the charge (add electrons)

Oxidation: Fe2+ => Fe3+ + 1 e

Reduction: MnO4- + 8 H+ + 5 e => Mn2+ + 4 H2O

Multiplication to equalize the amount of electrons:

Oxidation: 5 x (Fe2+ => Fe3+ + 1 e)

Reduction: 1 x (MnO4- + 8 H+ + 5 e => Mn2+ + 4 H2O)

Add the half-reactions:

Essential equation: 5 Fe2+ + MnO4- + 8 H+ => 5 Fe3+ + Mn2+ + 4 H2O

(3) An Oxidation-Reduction Reaction in Basic Medium.

The procedure to balance a reaction in basic medium is the same as in acidic medium

until just before the electrons must be added. To bring the reaction in basic medium, for

every H+ -ion an equal number of OH--ions is added to both sides of the equation. Where

H+ and OH- -ions appear on the same side, these ions combine to give H2O. If possible

water molecules may be cancelled.

Example 16 A reaction in basic medium.

When B2Cl4 reacts with hydroxide ions, BO2- and hydrogen gas are formed

The unbalanced equation:

B2Cl4 + OH- => BO2- + H2

Changes:

OS(B) increases (oxidation), OS(H) decreases (reduction)

Division into two half-reactions:

Oxidation: B2Cl4 => BO2-

Reduction: OH- => H2

Balance the atoms with changing OS:

Oxidation: B2Cl4 => 2 BO2-

Reduction: 2 OH- => H2

Balance oxygen atoms:

Oxidation: B2Cl4 + 4 H2O => 2 BO2-

Reduction: 2 OH- => H2 + 2 H2O

Balance hydrogen atoms:

Oxidation: B2Cl4 + 4 H2O => 2 BO2- + 8 H+

Basic chemistry 10

Reduction: 2 OH- + 4 H+ => H2 + 2 H2O

Balance chlorine atoms:

Oxidation: B2Cl4 + 4 H2O => 2 BO2- + 8 H+ + 4 Cl-

Reduction: 2 OH- + 4 H+ => H2 + 2 H2O

In basic medium:

Oxidation: B2Cl4 + 4 H2O + 8 OH- => 2 BO2- + 8 H+ + 4 Cl- +8 OH-

Reduction: 2 OH- + 4 H+ + 4 OH- => H2 + 2 H2O + 4 OH-

Combine H+ and OH--ions to give H2O:

Oxidation: B2Cl4 + 8 OH- => 2 BO2- + 4 H2O + 4 Cl-

Reduction: 2 H2O => H2 + 2 OH-

Balance the charge (add electrons)

Oxidation: B2Cl4 + 8 OH- => 2 BO2- + 4 H2O + 4 Cl- + 2 e

Reduction: 2 H2O + 2 e => H2 + 2 OH-

Electrons in both half-reactions are equal

Add the half-reations:

B2Cl4 + 8 OH- + 2 H2O => 2 BO2- + 4 H2O + 4 Cl- + H2 + 2 OH-

Essential equation:

B2Cl4 + 6 OH- => 2 BO2- + 2 H2O + 4 Cl- + H2

(4) Notes.

When in an oxidation- reduction reaction the same atoms oxidizes and reduces this

reaction is called an auto-redox reaction.

The oxidation of an atom can be recognized by an increase of the amount of oxygen

atoms to which it is bound. So the formation of oxides during combustion reactions is

clearly the oxidation.

The reduction of an atom can be recognized by a decrease of the amount of oxygen atoms

to which it is bound.

The essential equation of a redox-reaction can be transformed to a molecular equation by

adding the spectator ions to both sides of the balanced overall equation.

Example 17 Some examples.

Example of an auto-redox reaction: Dibromine reacts in basic medium, bromide ions and

bromate ions are formed.

The transition of iodine molecules to iodate ions is an oxidation (increase of the amount

of oxygen atoms).

The transition of permanganate ions into Mn2+ ions is a reduction (decrease of amount of

oxygen atoms)

The molecular equation for the oxidation of iron with potassium permanganate is

Basic chemistry 11

5 FeCl2 + KMnO4 + 4 H2SO4 => Fe2(SO4)3 + 3 FeCl3+ MnSO4 + KCl + 4 H2O

5. Exercises.

Write the following reaction equations. If possible, also write the essential equation.

1. Sulphuric acid with sodium sulfide.

2. Calcium oxide with sulphur trioxide.

3. Lithium oxide with sulphuric acid.

4. Thermolysis of aluminium carbonate.

5. Ammonium chloride with potassium hydroxide.

6. Silver sulfite with hydrogen chloride.

7. Thermolysis of cupper(I)sulfate.

8. Aluminium oxide with hydrogen chloride.

9. Reaction of chloric acid with zinc.

10. Aluminium oxide with hydrogen fosfate.

11. Reaction of perchloric racid with magnesium.

12. Aluminium carbonate with sulphuric acid.

13. Lead(II)nitrate with potassium chloride.

14. Sulphur dioxide with water.

15. Thermolysis of cupper(II)fosfate.

16. Carbon dioxide with a solution of potassium hydroxide.

17. Reaction of sodium hydroxide with hydrogen arsenate.

18. Lithium oxide with water.

19. Combustion of C4H4S

20. Reaction of sodium sulfate with barium nitrate.