ATOMS AND MOLECULES

SCIENCE AND TECHNOLOGY
Atoms and Molecules
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Matter in our Surroundings
3
ATOMS AND MOLECULES
In the previous chapter you learnt about matter.The idea of divisibility of matter was
considered long back in India around 500 B.C. Maharishi Kanad, an Indian
Philosopher discussed it in his Darshan (Vaisesik Darshan). He said if we go on
dividingmatter,weshallgetsmallerandsmallerparticles.Astagewouldcomebeyond
whichfurtherdivisionwillnotbepossible.Henamedtheseparticlesas‘PARMANU’.
This concept was further elaborated by another Indian philosopher, Pakudha
Katyayan.Katyayansaidthattheseparticlesnormallyexistinacombinedformwhich
gives us various forms of matter.
Around the same era, an ancient Greek philosopher Democritus (460 – 370 BC)
and Leucippus suggested that if we go on dividing matter, a stage will come when
further division of particles will not be possible. Democritus called these individual
particles‘atoms’(whichmeansindivisible).Theseideaswerebasedonphilosophical
considerations. Experimental work to validate these ideas could not be done till the
eighteenth century. However, today we know what an atom is and how it is
responsiblefordifferentpropertiesofsubstances.Inthischapter,weshallstudyabout
atoms and molecules and related aspects like atomic and molecular masses, mole
concept and molar masses. We shall also learn how to write chemical formula of
a compound.
OBJECTIVES
After completing this lesson you will be able to :
state the law of conservation of mass and law of constant proportions;
list important features of Dalton’s atomic theory;
distinguish between atoms and molecules;
define isotopic mass, atomic mass, and molecular mass;

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define the mole concept and molar mass;
represent some molecules with the help of a formula;
apply the mole concept to chemical reaction and show a quantitative
relationship between masses of reactants and products and
solve simple problems based on various concepts learnt.
3.1 LAWS OF CHEMICAL COMBINATIONS
There was tremendous progress in Chemical Sciences after 18th century. It arose
out of an interest in the nature of heat and the way things burn. Major progress was
made through the careful use of chemical balance to determine the change in mass
that occurs in chemical reactions.The great French ChemistAntoine Lavoisier used
the balance to study chemical reactions. He heated mercury in a sealed flask that
contained air.After several days, a red substance mercury (II) oxide was produced.
The gas remaining in the flask was reduced in mass. The remaining gas was neither
able to support combustion nor life. The remaining gas in the flask was identified
asnitrogen.Thegaswhichcombinedwithmercurywasoxygen.Furtherhecarefully
performed the experiment by taking a weighed quantity of mercury (II) oxide.After
strong heating, he found that mercury (II) oxide, red in colour, was decomposed into
mercury and oxygen. He weighed both mercury and oxygen and found that their
combined mass was equal to that of the mercury (II) oxide taken. Lavoisier finally
came to the conclusion that in every chemical reaction, total masses of all the
reactants is equal to the masses of all the products. This law is known as the
law of conservation of mass.
There was rapid progress in science after chemists began accurate determination of
masses of reactants and products. French chemist Claude Berthollet and Joseph
Proust worked on the ratio (by mass) of two elements which combine to form a
compound. Through a careful work, Proust demonstrated the fundamental law of
definite or constant proportions in 1808. In a given chemical compound, the
proportions by mass of the elements that compose it are fixed, independent
of the origin of the compound or its mode of preparation.
In pure water, for instance, the ratio of mass of hydrogen to the mass of oxygen
is always 1:8 irrespective of the source of water. In other words, pure water contains
11.11% of hydrogen and 88.89% of oxygen by mass whether water is obtained from
well, river or from a pond. Thus, if 9.0 g of water are decomposed, 1.0 g of hydrogen
and8.0gofoxygenarealwaysobtained.Furthermore,if3.0gofhydrogenaremixed
with 8.0 g of oxygen and the mixture is ignited, 9.0 g of water are formed and 2.0
g of hydrogen remains unreacted. Similarly sodium chloride contains 60.66% of
chlorine and 39.34% of sodium by mass whether we obtained it from salt mines or

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by crytallising it from water of ocean or inland salt seas or synthesizing it from its
elements sodium and chlorine. Of course, the key word in this sentence is ‘pure’.
Reproducibleexperimentalresultsarehighlightsofscientificthoughts.Infactmodern
scienceisbasedonexperimentalfindings. Reproducibleresultsindirectlyhintfor
a truth which is hidden. Scientists always worked for findings this truth and in this
manner many theories and laws were discovered. This search for truth plays an
important role in the development of science.
3.2 DALTON’S ATOMIC THEORY
The English scientist John Dalton was by no means the first person to propose the
existence of atoms, as we have seen in the previous section, such ideas date back
to classical times. Dalton’s major contribution was to arrange those ideas in proper
order and give evidence for the existence of atoms. He showed that the mass
relationship expressed by Lavoisier and Proust (in the form of law of conservation
of mass and law of constant proportions) could be interpreted most suitably by
postulating the existence of atoms of the various elements.
In1803,Daltonpublishedanewsystemofchemicalphilosophyinwhichthefollowing
statements comprise the atomic theory of matter:
1. Matter consists of indivisible atoms.
2. All the atoms of a given chemical element are identical in mass and in all other
properties.
3. Different chemical elements have different kinds of atoms and in particular such
atoms have different masses.
4. Atoms are indestructible and retain their identity in chemical reactions.
5. Theformationofacompoundfromitselementsoccursthroughthecombination
of atoms of unlike elements in small whole number ratio.
Dalton’s fourth postulate is clearly related to the law
of conservation of mass. Every atom of an element
hasadefinitemass.Alsoinachemicalreactionthere
is rearrangement of atoms. Therefore after the
reaction, mass of the product should remain the
same. The fifth postulate is an attempt to explain the
law of definite proportions.Acompound is a type of
mattercontainingtheatomsoftwoormoreelements
insmallwholenumberratio.Becausetheatomshave
definitemass,thecompoundmusthavetheelements
in definite proportions by mass.
John Dalton (1766-1844)
Fig. 3.1

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The Dalton’s atomic theory not only explained the laws of conservations of mass
and law of constant proportions but also predicted the new ones. He deduced the
law of multiple proportions on the basis of his theory. The law states that when
two elements form more than one compound, the masses of one element in
these compound for a fixed mass of the other element are in the ratio of small
whole numbers. For example, carbon and oxygen form two compounds: Carbon
monoxide and carbon dioxide. Carbon monoxide contains 1.3321 g of oxygen for
each 1.000g of carbon, whereas carbon dioxide contains 2.6642 g of oxygen for
1.0000 g of carbon. In other words, carbon dioxide contains twice the mass of
oxygen as is contained in carbon monoxide (2.6642 g = 2 × 1.3321 g) for a given
mass of carbon.Atomic theory explains this by saying that carbon dioxide contains
twice as many oxygen atoms for a given number of carbon atoms as does carbon
monoxide. The deduction of law of multiple proportions from atomic theory was
important in convincing chemists of the validity of the theory.
3.2.1 What is an Atom?
As you have just seen in the previous section that an atom is the smallest particle
ofanelementthatretainsits(elements)chemicalproperties.Anatomofoneelement
is different in size and mass from the atoms of the other elements. These atoms were
considered‘indivisible’byIndianandGreek‘Philosophers’inthebeginningandthe
name ‘atom’as mentioned earlier, emerged out of this basic philosophy. Today, we
know that atoms are not indivisible. They can be broken down into still smaller
particles although they lose their chemical identity in this process. But inspite of all
these developments atom still remains a building block of matter.
3.2.2 What is the size of the atom?
Atomsareverysmall,theyaresmallerthananythingthatwecanimagineorcompare
with. In order to have a feeling of size of an atom you can consider this example:
One teaspoon of water (about 1 mL) contains about three times as many
atoms as Atlantic ocean contains teaspoons of water. Also more than millions
of atoms when stacked would make a layer barely as thick as this sheet of paper.
Atoms of different elements not only differ in mass as proposed by Dalton but also
they differ in size. Now question is why should we bother for the size, mass and
other properties of an atom? The reason is simple, every matter we see around us
is made of atoms. Is it rectangular, circular or spherical? It is difficult to imagine the
real shape of an atom but for all practical purposes it is taken as spherical in size
and that is why we talk of its radius. Since size is extremely small and invisible to
our eyes, we adopt a scale of nanometer (1nm = 10–9 m) to express its size.
You can have a feeling of its size from the following table (Table 3.1).

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Table 3.1 : Relative sizes
Radius (in m) Example
10–10 Atoms of hydrogen
10–4 Grain of sand
10–1 Water melon
0.2×10–1 Cricket ball
You can not see atoms with your naked eyes but by using modern techniques, we
can now produce magnified image of surface of elements showing atoms. The
techniqueisknownasScanningTunneling
Microscopy (STM) (Fig. 3.2)
3.2.3 Atomic Mass
Dalton gave the concept of atomic mass.
According to him, atoms of the same
element have same atomic masses but
atoms of different elements have different
atomic masses. Since Dalton could not
weighindividualatoms,hemeasuredrelative
massesoftheelementsrequiredtoformacompound.Fromthis,hededucedrelative
atomic masses. For example, we can determine by experiment that 1.0000 g of
hydrogen gas reacts with 7.9367 g of oxygen gas to form water. If we know formula
of water, we can easily determine the mass of an oxygen atom relative to that of
hydrogen atom.
Dalton did not have a way of determining the proportions of atoms of each element
forming water during those days. He assumed the simplest possibility that atoms of
oxygen and hydrogen were equal in number. From this assumption, it would follow
that oxygen atom would have a mass that was 7.9367 times that of hydrogen atom.
This in fact was not correct.We now know that in water number of hydrogen atoms
is twice the number of oxygen atoms (formula of water being H2O). Therefore,
relative mass of oxygen atom must be 2×7.9367 = 15.873 times that of hydrogen
atom.After Dalton, relative atomic masses of several elements were determined by
scientists based on hydrogen scale. Later on, hydrogen based scale was replaced
by a scale based on oxygen as it (oxygen) was more reactive and formed a large
number of compounds.
In 1961, C-12 (or 12
6C) atomic mass scale was adopted. This scale depends on
measurement of atomic mass by an instrument called mass spectrometer. Mass
spectrometer invented early in 20th century, allows us to determine atomic masses
Fig. 3.2: Image of Copper surface
by STM technique. Atom can be seen in
magnified image of surfae

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precisely.The masses of atoms are obtained by comparison with C-12 atomic mass
scale. In fact C-12 isotope is chosen as standard and arbitrarily assigned a mass
of exactly 12 atomic mass units. One atomic mass unit (amu), therefore, equals
exactly one twelfth of mass of a carbon–12 atom,Atomic mass unit (amu) is now-
a-days is written as unified mass unit and is denoted by the letter ‘u’.
The relative atomic mass of an element expressed in atomic mass unit is called its
atomic weight. Now-a-days we are using atomic mass in place of atomic weight.
Further, you have seen that Dalton proposed that masses of all atoms in an element
are equal. But later on it was found that all atoms of naturally occurring elements
are not of the same mass. We shall study about such atoms in the following section.
Atomic masses that we generally use in our reaction or in chemical calculations are
average atomic masses which depend upon relative abundance of isotopes of
elements.
3.2.4 Isotopes and Atomic Mass
Dalton considered an atom as an indivisible particle. Later researches proved that
an atom consists of several fundamental particles such as : electrons, protons and
neutrons.Anelectronisnegativelychargedandaprotonispositivelychargedparticle.
Number of electrons and protons in an atom is equal. Since charge on an electron
is equal and opposite to charge of a proton, therefore an atom is electrically
neutral. Protons remain in the nucleus in the centre of the atom, and nucleus is
surrounded by negatively charged electrons.
1p 6p 8p
(a) (b) (c)
Fig. 3.3: Arrangement of electrons around nucleus in (a) hydrogen,
(b) carbon and (c) oxygen atoms
The number of protons in the nucleus is called atomic number denoted by Z. For
example in Fig. 3.3, there are 8 protons in the oxygen nucleus, 6 protons in carbon
nucleus and only one proton in hydrogen nucleus. Therefore atomic numbers of
oxygen, carbon and hydrogen are 8,6 and 1 respectively. There are also neutral

Atoms and Molecules
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particles in the nucleus and they are called ‘neutrons’. Mass of a proton and of a
neutron is nearly the same.
Total mass of the nucleus = mass of protons + mass of neutrons
Total number of protons and neutrons is called mass number (A). By convention
atomic number is written at the bottom of left corner of the symbol of the atom of
a particular element and mass number is written at the top left corner. For example,
symbol 12
6C indicates that there is a total of 12 particles (nucleons) in the nucleus
of a carbon atom, 6 of which are protons. Thus, there must be 12-6 = 6 neutrons.
Similarly 16
8O indicates 8 protons and 16 nucleons (8 protons + 8 neutrons). Since
atomiselectricallyneutral,oxygenhas8protonsand8electronsinit.Further,atomic
number (Z) differentiates the atom of one element from the atoms of the other
elements.
An element may be defined as a substance where all the atoms have the
same atomic number.
Butthenucleiofalltheatomsofagivenelementdonotnecessarilycontainthesame
number of neutrons. For example, atoms of oxygen, found in nature, have the same
number of protons which makes it different from other elements, but their neutrons
(in nucleus) are different. This is the reason that the masses of atoms of the same
element are different. For example, one type of oxygen atom contains 8 protons and
8 neutrons in one atom, second type 8 protons and 9 neutrons and the third type
contains 8 protons and 10 neutrons. We represent these oxygen atoms as 16
8O, 17
8O
and 18
8O respectively. Atoms of an element that have the same atomic number
(Z) but different mass number (A) are called isotopes. In view of difference in
atomic masses of the same element, we take average atomic masses of the elements.
This is calculated on the basis of the abundance of the isotopes. Atomic masses
of some elements are provided in Table 3.2.
Example 3.1 : Chlorine is obtained as a mixture of two isotopes 35
17Cl and 37
17Cl .
These isotopes are present in the ratio of 3:1. What will be the average atomic mass
ofchlorine?
Solution : 35
17Cl and 37
17Cl are present in the ratio of 3:1 i.e. out of four atoms,
3 atoms are of mass 35 and one atom of mass 37. Therefore,
Average atomic mass =
35 3 37 1 142
4 4
× + ×
= = 35.5 u
Thus, average atomic mass of chlorine will be 35.5u.

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Table 3.2 : Atomic mass* of some common elements
Elements Symbol Mass (u) Elements Symbol Mass (u)
Aluminium Al 26.93 Magnesium Mg 24.31
Argon Ar 39.95 Manganese Mn 54.94
Arsenic As 74.92 Mercury Hg 200.59
Barium Ba 137.34 Neon Ne 20.18
Boron B 10.81 Nickel Ni 58.71
Bromine Br 79.91 Nitrogen N 14.01
Caesium Cs 132.91 Oxygen O 16.00
Calcium Ca 40.08 Phosphorus P 30.97
Carbon C 12.01 Platinum Pt 195.09
Chlorine Cl 35.45 Potassium K 39.1
Chromium Cr 52.00 Radon Rn (222)**
Cobalt Co 58.93 Silicon Si 23.09
Copper Cu 63.56 Silver Ag 107.87
Fluorine F 19.00 Sodium Na 23.00
Gold Au 196.97 Sulphur S 32.06
Helium He 4.00 Tin Sn 118.69
Hydrogen H 1.008 Titanium Ti 47.88
Iodine I 126.90 Tungsten W 183.85
Iron Fe 55.85 Uranium U 238.03
Lead Pb 207.19 Vanadium V 50.94
Lithium Li 6.94 Xenon Xe 131.30
Zinc Zn 65.37
* Atomic masses are average atomic masses covered upto second decimal places.
In practice, we use rounded figures.
** Radioactive
INTEXT QUESTIONS 3.1
1. Name the scientists who proposed the law of conservation of mass and law of
constant proportions.
2. 12gofmagnesiumpowderwasignitedinacontainerhaving20gofpureoxygen.
After the reaction was over, it was found that 12 g of oxygen was left unreacted.
Show that it is according to law of constant proportions.
2Mg + O2 ⎯→ 2MgO

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3.3 WHAT IS A MOLECULE?
Dalton proposed in his hypothesis that atoms react to form a molecule which he said
as ‘compound atoms’. Today we know what a molecule is. A molecule is an
aggregate of two or more than two atoms of the same or different elements
in a definite arrangement. These atoms are held together by chemical forces or
chemical bonds. (You will study details of molecules in unit of chemical bonding)
An atom is the smallest particle of a substance but can not exist freely.
Contrary to this, a molecules can be considered as the smallest particle of
an element or of a compound which can exist alone or freely under ordinary
conditions. A molecule of a substance shows all chemical properties of that
substance. To describe the chemical composition of a molecule we take the help of
symbols of elements and formulas (described in sec 3.5). Oxygen molecule, with
which we are familiar, is made of two atoms of oxygen and therefore it is a diatomic
molecule (represented by O2), hydrogen, nitrogen, fluorine, chlorine, bromine and
iodine are other examples of diatomic molecules and are represented as H2, N2, F2,
Cl2 , Br2 and I2 respectively (Fig. 3.4).
H H N N F F
Cl Cl Br Br I I
Fig. 3.4 : Representation of diatomic molecules
Someotherelementsexistasmorecomplexmolecules.Phosphorusmoleculeconsists
of four atoms (denoted by P4) whereas sulphur exists as eight atom molecule (S8)
at ordinary temperature and pressure Fig. 3.5. A molecule made of four atoms is
tetratomicmolecule.Normally,moleculesconsistingofmorethanthreeorfouratoms
are considered under the category of polyatomic molecules. Only a few years back,
a form of carbon called buckminsterfullerene having molecular formula C60 was
discovered which you will study later on in you higher classes.
Structure of phosphorus molecule P4 Structure of a sulphur molecule S8
Fig. 3.5 : Molecules of phosphorus and sulphur
Molecules of compounds are composed of more than one kind of atoms.Afamiliar
example is a water molecule which is composed of more than one kind of atoms.
P
P
P
P S S S
S
S
S
S
S

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In one water molecule, there are two atoms of hydrogen and one atom of oxygen.
It is represented as H2O.Amolecule of ammonia consists of one nitrogen atom and
three hydrogen atoms.Amolecule of ethyl alcohol (C2H5OH) is composed of nine
atoms (2 atoms of carbon, 6 atoms of hydrogen and one atom of oxygen) Fig. 3.6.
Molecule of water Molecule of ammonia Molecule of ethyl alcohol
Fig. 3.6 : Molecules of water, ammonia and ethyl alcohol
3.3.1 Molecular Mass
You have just read that a molecule can be represented in the form of a formula
popularly known as molecular formula. Molecular formula may be of an element
or of a compound. Molecular formula of a compound is normally used for
determining the molecular mass of that substance. If a substance is composed
of molecule (for example : CO2, H2O or NH3), it is easy to calculate the molecular
mass. Molecular mass is the sum of atomic masses of all the atoms present in that
molecule. Thus the molecular mass is the sum of atomic masses of all the atoms
present in that molecule. The molecular mass of CO2 is obtained as
C 1×12.0 u = 12.0 u
2 O 2 × 16.0 u = 32.0 u
Mass of CO2 = 44.0 u
Hence, we write molecular mass of CO2 = 44.0 u
Similarly, we obtain molecular mass of NH3 as follows:
N 1×14.0 u = 14.0 u
3 H 3 × 1.08 u = 3.24 u
Mass of NH3 = 17.24 u
Molecular mass of ammonia, NH3 = 17.24 u
For substances which are not molecular in nature, we talk of formula mass. For
example,sodiumchloride(denotedbyformula,NaCl)isanionicsubstance.Forthis,
wewillcalculateformulamass,similartomolecularmass.Incaseofsodiumchloride,
NaCl;
Formula mass = mass of 1 Na atom + mass of 1 Cl atom
= 23 u + 35.5 u = 58.5 u
You will learn more about the molecular and ionic compounds in detail later.

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1. Nitrogen forms three oxides : NO, NO2 and N2O3. Show that it obeys law of
multipleproportions.
2. Atomic number of silicon is 14. If there are three isotopes of silicon having 14,
15 and 16 neutrons in their nuclei, what would be the symbol of the isotope?
3. Calculatemolecularmassofthecompoundswhoseformulasareprovidedbelow:
C2H4, H2O and CH3OH
3.4 MOLE CONCEPT
When we mix two substances, we get one or more new substances. For example,
when we mix hydrogen and oxygen and ignite the mixture, we get a new substance-
water. This can be represented in the form of a chemical equation,
2H2 (g) + O2 (g) ⎯→ 2H2O (l)
In the above equation, 2 molecules (four atoms) of hydrogen react with 1 molecule
(2 atoms) of oxygen and give two molecules of water. We always like to know how
many atoms/molecules of a particular substance would react with atoms/molecules
of another substance in a chemical reaction. No matter how small they are. The
solution to this problem is to have a convenient unit. Would you not like to have
a convenient unit? Definitely a unit for counting of atoms/molecules present in a
substance will be desirable and convenient as well. This chemical counting unit of
atoms and molecules is called mole.
The word mole was, apparently introduced in about 1896 byWilhelm Ostwald who
derived the term from the Latin word ‘moles’ meaning a ‘heap’ or ‘pile’. The mole
whosesymbolis‘mol’istheSI(internationalsystem)baseunitformeasuringamount
of substance. It is defined as follows:
A mole is the amount of substance that contains as many elementary
entities (atoms, molecules, formula unit or other fundamental particles)
as there are atoms in exactly 0.012 kg of carbon-12 isotope.
In simple words, mole is the number of atoms in exactly 0.012 kg (12 grams) of
C-12.Although mole is defined in terms of carbon atoms but the unit is applicable
to any substance just as 1 dozen means 12 or one gross means 144 of anything.
Mole is scientist’s counting unit like dozen or gross. By using mole, scientists
(particularly chemists) count atoms and molecules in a given substance. Now it is
experimentally found that the number of atoms contained in exactly 12 g of C-12

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is 602,200 000 000 000 000 000 000 or 6.022×1023. This number is called
Avogadro’snumber inhonourofAmedeoAvogadro,anItalianlawyerandphysicist.
When this number is divided by ‘mole’ it becomes a constant and is known as
Avogadro’s constant denoted by symbol, NA = 6.02×1023 mol–1. We have seen that
Atomic mass of C = 12 u
Atomic mass of He = 4 u
We can see that one atom of carbon is three times as heavy as one atom of helium.
On the same logic 100 atoms of carbon are three times as heavy as 100 atoms of
helium. Similarly 6.02×1023 atoms of carbon are three times as heavy as 6.02×1023
atoms of helium. But 6.02×1023 atoms of carbon weigh 12 g, therefore 6.02×1023
atoms of helium will weigh 1/3 × 12g = 4 g. We can take a few more examples
of elements and can calculate the mass of one mole atoms of that element.
3.4.1 Molar Mass
Mass of one mole of a substance is called its molar mass. A substance may be
an element or a compound. Mass of one mole atoms of oxygen means mass of
6.02×1023 atoms of oxygen. It is found that one mole atoms of oxygen weighs 16.0
g. When we say one mole molecules of oxygen that means 6.02×1023 molecules
of oxygen (O2). One mole molecules of oxygen will weigh 32.0 g. Thus,
Mass of one mole atoms of oxygen = 16 g mol–1
Mass of one mole molecules of oxygen = 32 g mol–1
When it is not clear whether we are asking for one mole of atoms or one mole of
molecules then we take natural form of that substance. For example, one mole of
oxygen means one mole of oxygen molecules as oxygen occurs in the form of
moleculesinnature.Incaseofcompounds,thesamelogicisapplicable.Forexample,
one mole of water means one mole molecules of water which weighs 18 g.
Numerically one mole of a substance is equal to atomic or molecular mass of that
substance expressed in grams.
Remember, molar mass is always expressed in the unit of g/mol or g mol–1.
For example,
Molar mass of nitrogen (N2) = 28 g mol–1
Molar mass of chlorine (Cl2) = 71 g mol–1
Table 2.3 Provides molecular and molar mass of a few common substances.

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Table 3.3 : Molecular and molar masses
Formula Molecular Mass (u) Molar mass (g/mol)
O2(oxygen) 32.0 32.0
Cl2(chlorine) 71.0 71.0
P4 (phosphorus) 123.9 123.9
CH4 (methane) 16.00 16.0
NH3 (ammonia) 17.0 17.0
HCl (hydrochloric acid gas) 36.5 36.5
CO2 (carbon dioxide) 44.0 44.0
SO2 (sulphur dioxide) 64.0 64.0
C2H5OH (ethyl alcohol) 46.0 46.0
C6H6 (benzene) 78.0 78.00
Example 3.2 : How many grams are there in 3.5 mol of oxygen?
Solution : For converting mole into mass in grams and vice-versa, we always need
a relationship between mass and mole.
Molar mass of oxygen (O2) = 32 g mol–1
Therefore, number of grams of oxygen in 3.5 mol of it
= 3.5 mol of oxygen × 32.0 g mol–1
= 112.0 g of oxygen
Example 3.3 : Find out number of molecules in 27 g of water.
Solution: Mole concept provides a relationship between number of particles and
their mass. Thus it is possible to calculate the number of particles in a given mass.
Number of mole of H2O =
2
2
Mass of water (H O)
Molar mass of H O
= –1
27g 3
218 g mol
= mol = 1.5 mol
Since 1 mol of water contains 6.02×1023 molecules.
Therefore, 1.5 mol of water contains = 6.02 × 1023 molecules mol–1 ×1.5 mol
= 9.03 × 1023 molecules of water

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1. Work out a relationship between number of molecules and mole.
2. What is molecular mass? In what way it is different from the molar mass?
3. Consider the reaction
C (s) + O2 (g) ⎯→ CO2 (g)
18 g of carbon was burnt in oxygen. How many moles of CO2 is produced?
4. What is the molar mass of NaCl?
3.5 WRITING CHEMICAL FORMULA OF COMPOUNDS
As you are aware, a compound is made of two or more than two elements combined
in a definite proportion by mass (law of constant proportions). Thus, the number of
combining atoms in a compound is fixed. The elements are represented by their
symbols (e.g. H for hydrogen, Na for sodium). Similarly a compound is also
represented by a shorthand notation known as formula or chemical formula. The
formula of a compound indicates (i) elements constituting the compound and (ii)
the number of each constituent element. In other word, the formula of a
compound also represents its chemical composition. The atoms of elements
constitutingacompoundareindicatedbytheirsymbolsandtheirnumberisindicated
as a subscript on the right hand bottom of the symbol. For example, in the formula
of water, H2O, two atoms of hydrogen are indicated as subscript ‘2’, while oxygen
isshownwithoutwritinganysubscript,whichmeansthatthenumberofoxygenatom
is just one.
3.5.1 Valency and Formulation
Every element has a definite capacity to combine with other elements. This
combining capacity of an element is called its valency. You will learn very soon
that this combining capacity of elements depends on the electronic configuration of
elements. Valencies of a few elements are given in Table 3.4.
Table 3.4 : Valency of elements
Elements Symbol Valency Elements Symbol Valency
Hydrogen H 1 Phosphorus P 5
Oxygen O 2 Sodium Na 1
Carbon C 4 Magnesium Mg 2
Nitrogen N 3 Calcium Ca 2
Chlorine Cl 1 Aluminium Al 3
Bromine Br 1 Iron Fe 2
Iodine I 1 Barium Ba 2

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Most of the simple compounds are made of two elements. Such compounds are
called binary compounds. It is easy to write formula of such compounds. When
a metal combines with a non-metal, the symbol of the metal-element is written on
the left hand side and that of the non-metal element on right hand side. (If both are
non-metal,wewritemoreelectronegative*elementontherighthandside).Innaming
a compound, the first element is written as such and the name of the second element
i.e. more electronegative element, changes its ending to ‘ide’. For writing chemical
formula,wehavetowritevalenciesasshownbelowandthencrossoverthevalencies
of the combining atoms. Formula of the compounds resulting from carbon and
chlorine,hydrogenandoxygen,andhydrogenandchlorinecanbewrittenasfollows:
Some other examples for writing formula of compounds CaO, NaCl and NH3 can
also be taken for more clarity.
Thus,wecanwriteformulasofvariouscompoundsifweknowelementsconstituting
them and their valencies.
Valency, as mentioned, depends on the electronic configuration and nature of the
elements. Sometimes an element shows more than one type of valency. We say
element has variable valency. For example nitrogen forms several oxides : N2O,
N2O2, N2O3,N2O4 and N2O5. If we take valency of oxygen equal to 2, then valency
of nitrogen in the oxides will be 1,2,3,4 and 5 respectively.Valencies are not always
fixed. Similar to nitrogen, phosphorus also shows valencies 3 and 5 as reflected in
compounds PBr3 and P2O5. In these compounds, there are more than one atom.
In such cases, number of atoms is indicated by attaching a numerical prefix (mono,
di, tri, etc) as mentioned in Table 2.5.
Table 3.5 : Numerical Prefixes
Number of atoms Prefix Example
1 Mono carbon monoxide, CO
2 Di carbon dioxide, CO2
3 Tri phosphorus trichloride, PCl3
4 Tetra carbon tetrachloride, CCl4
5 Penta Dinitrogen pentoxide, N2O5

63
Atoms and Molecules
Notes
MODULE - 2
Here you would notice that ‘–o’ or ‘–a’ at the end of the prefix is often dropped
before another vowel, e.g. monoxide, pentoxide. There is no gap between numerical
prefix and the name of the element. The prefix mono is usually dropped for the first
element.Whenhydrogenisthefirstelementintheformula,noprefixisaddedbefore
hydrogen irrespective of the number. For example, the compound H2S is named as
hydrogen sulphide and not as dihydrogen sulphide.
Thus, we have seen that writing formula of a binary compound is relatively easy.
However, when we have to write formula of a compound which involves more than
two elements (i.e. of a polyatomic molecule), it is somewhat a cumbersome task.
In the following section we shall consider formulation of more difficult compounds.
You will learn later on that there are basically two types of compounds: covalent
compounds and ionic or electrovalent compounds. H2O and NH3 are covalent
compounds. NaCl and MgO are ionic compounds.An ionic compound is made of
two charged constituents. One positively charged and other negatively charged. In
case of NaCl, there are two ions : Na+ and Cl– ion. Charge of these ions in case
of electrovalent compound is used for writing formula. It is easy to write formula
of an ionic compound only if there is one metal and one non-metal as in the case
of NaCl and MgO. If there are more than two elements in an ionic compound,
formulation will be a little difficult and in that situation we should know charge of
cations and anions.
3.5.2 Formulation of Ionic compounds
Formulationofanioniccompoundiseasywhenweknowchargeofcationandanion.
Remember, in an ionic compound, sum of the charge of cation and anion should be
equal to zero.Afew examples of cations and anions with their charges are provided
in Table 3.6.
Table 3.6 : Charges of some common cations and anions
which form ionic compounds
Anions Charge Cations Charge
Chloride ion, Cl– –1 Potassium ion, K+ +1
Nitrate ion, NO3
– –1 Sodium ion, Na+ +1
Hydroxide ion, OH– –1 Ammonium ion, NH4
+ +1
Bicarbonate ion, HCO3
– –1 Magnesium ion, Mg2+ +2
Nitrite ion, NO2
– –1 Calcium ion, Ca2+ +2
Acetate ion, CH3COO– –1 Lead ion, Pb2+ +2
Bromide ion, Br– –1 Iron ion (ous), Fe2+ +2
Iodide ion, I– –1 Zinc ion, Zn2+ +2

Atoms and Molecules
64
Notes
MODULE - 2
Sulphite ion, SO3
2– –2 Copper ion (cupric), Cu2+ +2
Carbonate ion, CO3
2– –2 Mercury ion (Mercuric), Hg2+ +2
Sulphate ion, SO4
2– -2 Iron (ic) ion, Fe3+ +3
Sulphide ion, S2– –2 Aluminium ion, Al3+ +3
Phosphate ion, PO4
3– –3 Potassium ion, K+ +1
Sodium ion, Na+ +1
Suppose you have to write formula of sodium sulphate which is made of Na+ and
SO4
2– ions. For this the positive and negative charge can be crossed over to give
subscripts. The purpose of this crossing over of charges is to find the number of ions
required to equate the number of positive and negative charges.
This gives the formula of sodium sulphate as Na2SO4 . We can check the charge
balance as follows
2Na+ = 2×(+1) = +2
= 0
1SO4
2– = 1×(–2) = –2
Thus the compound, Na2SO4 is electrically neutral.
Now it is clear that digit showing charge of cation goes to anion and digit showing
charge of anion goes to cation. For writing formula of calcium phosphate we take
charge of each ion into consideration and write the formula as discussed above.
( ) ( ) ( )2+ 3–
4 3 4 23 2
Ca PO Ca PO=
Writing formula of a compound comes only by practice, therefore write formulas of
as many ionic compounds as possible based on the guidelines given above.
1. Write the name of the expected compound formed between
(i) hydrogen and sulphur
(ii) nitrogen and hydrogen
(iii) magnesiumandoxygen

65
Atoms and Molecules
Notes
MODULE - 2
2. Propose the formulas and names of the compounds formed between
(i) potassium and iodide ions
(ii) sodium and sulphate ions
(iii) aluminiumandchlorideions
3. Write the formula of the compounds formed between
(i) Hg2+ and Cl–
(ii) Pb2+ and 3–
4PO
(iii) Ba2+ and 2–
4SO
WHAT YOU HAVE LEARNT
According to law of constant proportions, a sample of a pure substance
always consists of the same elements combined in the same proportion by mass.
When an element combines with another element and forms more than one
compound, then different masses of the one element that combine with the fixed
mass of another element are in the ratio of simple whole number or integer.This
is the law of multiple proportions.
John Dalton introduced the idea of an atom as an indivisible particle.An atom
is the smallest particle of an element which shows all the properties of that
element.An atom can not exist freely and therefore remains in a combined state.
A molecule is the smallest particle of an element or of a compound which shows
all properties of that substance and can exist freely under ordinary conditions.
A molecule can be represented in the form of a chemical formula using symbols
of elements that constitute it.
Composition of any compound can be represented by its chemical formula.
Atom of the isotope C-12 is assigned atomic mass unit of 12 and the relative
atomic masses of all other atoms of elements are obtained by comparing them
withit.
The mole is the amount of a substance which contains the same number
of particles (atoms, ions or molecule) as there are atoms in exactly 0.012
kg of carbon-12.
Avogadro’s number is defined as the number of atoms in exactly 0.012 kg
(or 12 g) of C-12 and is equal to 6.02 ×1023. Avogadro’s constant is written
as 6.02 × 1023 mol–1.

Atoms and Molecules
66
Notes
MODULE - 2
Mass of one mole atoms or one mole molecules or one mole of formula unit
of a substance is its molar mass.
Thecompositionofanycompoundcanberepresentedbyitsformula.Forwriting
the formula of a compound, valence or valency of an element is used. This is
normally done in case of covalent compounds.
Valency is the combining capacity of an element and is related to its electronic
configuration.
In ionic compounds, the charge on each ion is used to determine the chemical
formula of the compound.
TERMINAL EXERCISE
1. Describethefollowing:
(a) Law of conservation of mass
(b) Law of constant proportions
(c) Law of multiple proportions
2. What is the atomic theory proposed by John Dalton? What changes have taken
place in the theory during the last two centuries?
3. Write the number of protons, neutrons and electron in each of the following
isotopes :
2
1H , 18
8O , 19
9F , 40
20Ca
4. Boron has two isotopes with masses 10.13 u and 11.01 u and abundance of
19.77% and 80.23% respectively. What is the average atomic mass of boron?
(Ans. 10.81 u)
5. Give symbol for each of the following isotopes:
(a) Atomic number 19, mass number 40
(b) Atomic number 7, mass number 15
(c) Atomic number 18, mass number 40
(d) Atomic number 17, mass number 37
6. How does an element differ from a compound? Explain with suitable examples.
7. Charge of one electron is 1.6022×10–19 coulomb. What is the total charge on
1 mol of electrons?

67
Atoms and Molecules
Notes
MODULE - 2
8. How many molecules of O2 are in 8.0 g of oxygen? If the O2 molecules were
completely split into O (Oxygen atoms), how many mole of atoms of oxygen
would be obtained?
9. Assume that human body is 80% water. Calculate the number of molecules of
water that are present in the body of a person whose weight is 65 kg.
10. Refer to atomic masses given in the Table (3.2) of this chapter. Calculate the
molar masses of each of the following compounds :
HCl, NH3, CH4, CO and NaCl
11. Average atomic mass of carbon is 12.01 u. Find the number of moles of carbon
in (a) 2.0 g of carbon. (b) 8.0 g of carbon.
12. Classifythefollowingmoleculesasdi,tri,tetra,pentaandhexaatomicmolecules:
H2, P4, SF4, SO2, PCl3, CH3OH, PCl5, HCl
13. What is the mass of
(a) 6.02×1023 atoms of oxygen
(b) 6.02×1023 molecules of P4
(c) 3.01×1023 molecules of O2
14. How many atoms are present in:
(a) 0.1 mol of sulphur
(b) 18.g of water (H2O)
(c) 0.44 g of carbon dioxide (CO2)
15. Write various postulates of Dalton’s atomic theory.
16. Convert into mole:
(a) 16 g of oxygen gas (O2)
(b) 36 g of water (H2O)
(c) 22 g of carbon dioxide (CO2)
17. What does a chemical formula of a compound represents?
18. Write chemical formulas of the following compounds:
(a) Copper (II) sulphate
(b) Calciumfluoride
(c) Aluminiumbromide
(d) Zincsulphate
(e) Ammoniumsulphate

Atoms and Molecules
68
Notes
MODULE - 2
ANSWERS TO INTEXT QUESTIONS
3.1
(i) Lavoisier proposed the law of conservation of mass and Proust proposed the
law of constant proportions
(ii) Incontainer,12gofOxygenwasleftunreacted.Therefore,amountofunreacted
Oxygen = (20 – 12)g = 08g. Thus 12g of magnesium reacted with 8g of oxygen
in the ratio 12:8. This is what we expected for MgO i.e 24g of Mg reacted
with 16g of Oxygen or 12g of Mg will react with 8g of Oxygen.
3.2
(i) Atomic mass of nitrogen is 14u and that of oxygen is 16u.
In NO, 14g of nitrogen reacted with 16g of oxygen
In NO2, 14g of nitrogen reacted with 32g of oxygen
In N2O3, 28g of nitrogen reacted with 48g of oxygen
or
14g of nitrogen reacted with 24g of Oxygen.
Threfore, amount of oxygen which reacts with 12g of nitrogen in case of NO,
NO2 and N2O3 will be in the ratio of 16:32:24 or 2:4:3. This proves the law
of multiple proportions.
(ii) Atomic number of Si is 14
Mass number of silicon atoms having 14,15 and 16 neutrons will be 28,29 and
30 respectively and therefore symbols of istopes of silicon will be
28
14Si 29
14Si and 30
14Si
(iii) Molecular mass of C2H4 = mass of two atom of carbon + mass of 4 atom of
hydrogen
= 2 × 12u + 4 × 1u = 28u
Molecular mass of H2O = mass of two atoms of hydrogen + mass of one atom
ofoxygen
= 2 × 1u + 1 × 16u = 18u
Molecular mass of CH3OH = mass of one atom of carbon + mass of 4 atoms
of hydrogen + mass of one atom of oxygen
= 1 × 12u + 4 × 1u +1 × 16u = 32u

69
Atoms and Molecules
Notes
MODULE - 2
3.3
1. 1 mole of a substance contains 6.023 × 1023 molecules of that substance i.e
1 mole of a substance = 6.023 × 1023 molecules of that substance.
2. Molecular mass is the sum of atomic masses of all the atoms present in that
molecule.
Molecular mass is the mass of one molecule whereas molar mass is the mass
of 1 mol or 6.023 × 1023 elementary entities (atoms, molecules, ions)
3. C(s) + O2 ⎯⎯→ CO2
1 mol 1 mol 1 mol of CO2 (44g)
of C (12g) of O2(32g)
12 g of carbon gives 1 mole of CO2
18g of carbon will give 1.5 mole of CO2
4. Molar mass of NaCl = ( 23.0 + 35.5 )g mol-1
= 58.5 g mol-1
3.4
1. (i) H2S
(ii) NH3
(iii) MgO
2. (i) KI, Potassium iodide
(ii) Na2SO4, Sodium Sulphate
(iii) AlCl3,AluminiumChloride
3. (i) HgCl2
(ii) Pb3(PO4)2
(iii) BaSO4

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