Atoms and Molecules
Introduction:-
Ancient Indian and Greek philosophers have always wondered about the unknown and unseen form of matter.
The idea of divisibility of matter was considered long back in India, around 500 BC.
An Indian philosopher Maharishi Kanad, postulated that if we go on dividing matter (padarth), we shall get smaller and smaller particles.
Ultimately, a stage will come when we shall come across the smallest particles beyond which further division will not be possible.
He named these particles Parmanu.
Another Indian philosopher, Pakudha Katyayama, elaborated this doctrine and said that these particles normally exist in a combined form which gives us various forms of matter.
Around the same era, ancient Greek philosophers - Democritus and Leucippus suggested that if we go on dividing matter, a stage will come when particles obtained cannot be divided further.
Democritus called these indivisible particles atoms (meaning indivisible).
All this was based on philosophical considerations and not much experimental work to validate these ideas could be done till the eighteenth century.
By the end of the eighteenth century. scientists recognised the difference between elements and compounds and naturally became interested in finding out how and why elements combine and what happens when they combine.
Antoine L. Lavoisier laid the foundation of chemical sciences by establishing two important laws of chemical combination.
3.1 Laws of Chemical Combination
The following two laws of chemical combination were established after much experimentations by Lavoisier and Joseph L. Proust.
3.1.1 LAW OF CONSERVATION OF MASS:-
Law of conservation of mass states that mass can neither be created nor destroyed in a chemical reaction.
3.1.2 LAW OF CONSTANT PROPORTIONS:-
Lavoisier, along with other scientists, noted that many compounds were composed of two or more elements and each such compound had the same elements in the same proportions, irrespective of where the compound came from or who prepared it.
In a compound such as water, the ratio of the mass of hydrogen to the mass of oxygen is always 1:8, whatever the source of water.
Thus, if 9 g of water is decomposed. 1 g of hydrogen and 8 g of oxygen are always obtained.
Similarly in ammonia, nitrogen and hydrogen are always present in the ratio 14:3 by mass. whatever the method or the source from which it is obtained.
This led to the law of constant proportions which is also known as the law of definite proportions.
This law was stated by Proust as "In a chemical substance the elements are always present in definite proportions by mass".
The next problem faced by scientists was to give appropriate explanations of these laws. British chemist John Dalton provided the basic theory about the nature of matter.
Dalton picked up the idea of divisibility of matter, which was till then just a philosophy. He took the name 'atoms' as given by the Greeks and said that the smallest particles of matter are atoms.
His theory was based on the laws of chemical combination.
Dalton's atomic theory provided an explanation for the law of conservation of mass and the law of definite proportions.
According to Dalton's atomic theory, all matter, whether an element, a compound or a mixture is composed of small particles called atoms.
The postulates of this theory may be stated as follows:
All matter is made of very tiny particles called atoms, which participate in chemical reactions.
Atoms are indivisible particles, which cannot be created or destroyed in a chemical reaction.
Atoms of a given element are identical in mass and chemical properties.
Atoms of different elements have different masses and chemical properties.
Atoms combine in the ratio of small whole numbers to form compounds
The relative number and kinds of atoms are constant in a given compound.
You will study in the next chapter that all atoms are made up of still smaller particles.
3.2 What is an Atom?
The building blocks of all matter are atoms.
How big are atoms?
Atoms are very small, they are smaller than anything that we can imagine or compare with.
More than millions of atoms when stacked would make a layer barely as thick as this sheet of paper.
If atoms are so insignificant in size, why should we care about them?
This is because our entire world is made up of atoms.
We may not be able to see them, but they are there, and constantly affecting whatever we do.
Through modern techniques, we can now produce magnified images of surfaces of elements showing atoms.
3.2.1 WHAT ARE THE MODERN DAY SYMBOLS OF ATOMS OF DIFFERENT ELEMENTS?
Dalton was the first scientist to use the symbols for elements in a very specific sense.
When he used a symbol for an element he also meant a definite quantity of that element. that is, one atom of that element.
Berzilius suggested that the symbols of elements be made from one or two letters of the name of the element.
In the beginning, the names of elements were derived from the name of the place where they were found for the first time.
For example, the name copper was taken from Cyprus.
Some names were taken from specific colours.
For example, gold was taken from the English word meaning yellow.
Now-a-days. IUPAC (International Union of Pure and Applied Chemistry) is an international scientific organisation which approves names of elements, symbols and units.
Many of the symbols are the first one or two letters of the element's name in English.
The first letter of a symbol is always written as a capital letter (uppercase) and the second letter as a small letter (lowercase).
For example:-
hydrogen, H
aluminium, Al and not AL
(ii) cobalt, Co and not CO.
Symbols of some elements are formed from the first letter of the name and a letter appearing later in the name.
Examples are: (i) chlorine, Cl (ii) zinc, Zn etc.
Other symbols have been taken from the names of elements in Latin, German or Greek.
For example, the symbol of iron is Fe from its Latin name ferrum,
sodium is Na from natrium.
potassium is K from kalium.
Therefore, each element has a name and a unique chemical symbol.
3.2.2 ATOMIC MASS
The most remarkable concept that Dalton's atomic theory proposed was that of the atomic mass.
According to him, each element had a characteristic atomic mass.
The theory could explain the law of constant proportions so well that scientists were prompted to measure the atomic mass of an atom.
Since determining the mass of an individual atom was a relatively difficult task.
Relative atomic masses were determined using the laws of chemical combinations and the compounds formed.
Let us take the example of a compound. carbon monoxide (CO) formed by carbon and oxygen.
It was observed experimentally that 3 g of carbon combines with 4 g of oxygen to form CO.
In other words, carbon combines with 4/3 times its mass of oxygen.
Suppose we define the atomic mass unit (earlier abbreviated as 'amu', but according to the latest IUPAC recommendations.
It is now written as 'u' - unified mass) as equal to the mass of one carbon atom, then we would assign carbon an atomic mass of 1.0 u and oxygen an atomic mass of 1.33 u.
However, it is more convenient to have these numbers as whole numbers or as near to a whole numbers as possible.
While searching for various atomic mass units, scientists initially took 1/ 16 of the mass of an atom of naturally occurring oxygen as the unit.
This was considered relevant due to two reasons:
oxygen reacted with a large number of elements and formed compounds.
this atomic mass unit gave masses of most of the elements as whole numbers.
However, in 1961 for a universally accepted atomic mass unit, carbon-12 isotope was chosen as the standard reference for measuring atomic masses.
One atomic mass unit is a mass unit equal to exactly one-twelfth (1/12) the mass of one atom of carbon-12.
The relative atomic masses of all elements have been found with respect to an atom of carbon-12.
Imagine a fruit seller selling fruits without any standard weight with him. He takes a watermelon and says, "this has a mass equal to 12 units" (12 watermelon units or 12 fruit mass units).
He makes twelve equal pieces of the watermelon and finds the mass of each fruit he is selling, relative to the mass of one piece of the watermelon.
Now he sells his fruits by relative fruit mass unit (amu), as in Fig. 3.4.
Similarly, the relative atomic mass of the atom of an element is defined as the average mass of the atom, as compared to 1/12 the mass of one carbon-12 atom.
3.2.3 HOW DO ATOMS EXIST
Atoms are not able to exist independently.
Atoms form molecules and ions.
These molecules or ions aggregate in large numbers to form the matter that we can see, feel or touch.
3.3 What is a Molecule?
A molecule is in general a group of two or more atoms that are chemically bonded together, that is, tightly held together by attractive forces.
A molecule can be defined as the smallest particle of an element or a compound that is capable of an independent existence and shows all the properties of that substance.
Atoms of the same element or of different elements can join together to form molecules.
3.3.1 MOLECULES OF ELEMENTS
The molecules of an element are constituted by the same type of atoms.
Molecules of many elements, such as argon (Ar), helium (He) etc. are made up of only one atom of that element.
But this is not the case with most of the non metals.
For example,
A molecule of oxygen consists of two atoms of oxygen and hence it is known as a diatomic molecule, O2
If 3 atoms of oxygen unite into a molecule, instead of the usual 2. we get ozone, O3,
The number of atoms constituting a molecule is known as its atomicity.
Metals and some other elements, such as carbon, do not have a simple structure but consist of a very large and indefinite number of atoms bonded together.
Let us look at the atomicity of some non-metals.
3.3.2 MOLECULES OF COMPOUNDS
Atoms of different elements join together in definite proportions to form molecules of compounds.
Few examples are given in Table 3.4.
3.3.3 WHAT IS AN ION?
Compounds composed of metals and non metals contain charged species.
The charged species are known as ions.
lons may consist of a single charged atom or a group of atoms that have a net charge on them.
An ion can be negatively or positively charged.
A negatively charged ion is called an 'anion' and
the positively charged ion, a 'cation.
Take, for example, sodium chloride (NaCl). Its constituent particles are positively charged sodium ions (Na+) and negatively charged chloride ions (Cl-).
Ions may consist of a single atom or a group of atoms that have a net charge on them.
A group of atoms carrying a charge is known as a polyatomic ion (Table 3.6).
We shall learn more about the formation of ions in Chapter 4.
3.4 Writing Chemical Formulae
The chemical formula of a compound is a symbolic representation of its composition.
The chemical formulae of different compounds can be written easily.
For this exercise, we need to learn the symbols and combining capacity of the elements.
The combining power (or capacity) of an element is known as its valency.
Valency can be used to find out how the atoms of an element will combine with the atom(s) of another element to form a chemical compound.
The valency of the atom of an element can be thought of as hands or arms of that atom.
Human beings have two arms and an octopus has eight.
If one octopus has to catch hold of a few people in such a manner that all the eight arms of the octopus and both arms of all the humans are locked, how many humans do you think the octopus can hold?
Represent the octopus with O and humans with H.
Can you write a formula for this combination?
Do you get OH4, as the formula? The subscript 4 indicates the number of humans held by the octopus.
The valencies of some common ions are given in Table
The rules that you have to follow while writing a chemical formula are as follows:
the valencies or charges on the ion must balance.
when a compound consists of a metal and a non-metal, the name or symbol of the metal is written first.
For example: calcium oxide (CaO), sodium chloride (NaCl). iron sulphide (FeS), copper oxide (CuO) etc.. where oxygen, chlorine, sulphur are non metals and are written on the right whereaswhereas calcium, sodium, iron and copper are metals, and are written on the left.
in compounds formed with polyatomic ions, the number of ions present in the compound is indicated by enclosing the formula of ion in a bracket and writing the number of ions outside the bracket.
For example, Mg (OH)2.
In case the number of polyatomic ion is one, the bracket is not required.
For example, NaOH.
3.4.1 FORMULAE OF SIMPLE COMPOUNDS
While writing the chemical formulae for compounds, we write the constituent elements and their valencies as shown below.
Then we must crossover the valencies of the combining atoms.
The simplest compounds, which are made up of two different elements are called binary compounds.
Examples
1. Formula of hydrogen chloride
Symbol :
Valency:
Formula of the compound would be HCl
2. Formula of hydrogen sulphide
Symbol :
Valency :
Formula: H2S
3. Formula of carbon tetrachloride
Symbol
Valency
Formula: CCI4
4. Formula of magnesium chloride
Symbol
Charge
Formula: MgCl2
Note that in the formula, the charges on the ions are not indicated.
Some more examples:-
Formula for aluminium oxide:
Symbol
Charge
Formula:
Formula for calcium oxide:
Symbol
Charge
Formula
Here, the valencies of the two elements are the same.
You may arrive at the formula Ca2O2.
But we simplify the formula as CaO.
(c) Formula of sodium nitrate:
Symbol
Charge
Formula
(d) Formula of calcium hydroxide:
Symbol
Charge
Formula
Note that the formula of calcium hydroxide is Ca(OH)2 and not CaOH2.
We use brackets when we have two or more of the same ions in the formula.
Here, the bracket around OH with a subscript 2 indicates that there are two hydroxyl (OH) groups joined to one calcium atom.
In other words, there are two atoms each of oxygen and hydrogen in calcium hydroxide.
(e) Formula of sodium carbonate:
Symbol
Charge
Formula
In the above example, brackets are not needed if there is only one ion present.
(f) Formula of ammonium sulphate:
Symbol
Charge
Formula:
3.5 Molecular Mass and Mole Concept
3.5.1 MOLECULAR MASS
The molecular mass of a substance is the sum of the atomic masses of all the atoms in a molecule of the substance.
It is therefore the relative mass of a molecule expressed in atomic mass units (u).
Example 3.1 (a) Calculate the relative molecular mass of water (H2O).
Solution
(b) Calculate the molecular mass of HNO3
Solution:
3.5.2 FORMULA UNIT MASS
The formula unit mass of a substance is a sum of the atomic masses of all atoms in a formula unit of a compound.
Formula unit mass is calculated in the same manner as we calculate the molecular mass.
The only difference is that we use the word formula unit for those substances whose constituent particles are ions.
For example, sodium chloride as discussed above, has a formula unit NaCl.
Its formula unit mass can be calculated as
1 x 23+1 x 35.5= 58.5 u
Example 3.2 Calculate the formula unit mass of CaCl2
Solution:-
1. Calculate the molecular masses of H2, O2, CO2, CH4, CH3OH.
2. Calculate the formula unit masses of ZnO, Na2O. K2CO3 (given atomic masses of Zn= 65 u , Na = 23 u , K = 39 u, C= 12 u and O =16u)
3.5.3 MOLE CONCEPT
Take an example of the reaction of hydrogen and oxygen to form water:
H2 + O2 --> 2H2O
The above reaction indicates that
two molecules of hydrogen combine with one molecule of oxygen to form two molecules of water, or
4u of hydrogen molecules combine with 32 u of oxygen molecules to form 36 u of water molecules.
We can infer from the above equation that the quantity of a substance can be characterised by its mass or the number of molecules.
But, a chemical reaction equation indicates directly the number of atoms or molecules taking part in the reaction.
Therefore, it is more convenient to refer to the quantity of a substance in terms of the number of its molecules or atoms, rather than their masses.
So, a new unit "mole" was introduced.
The mole, symbol mol, is the SI unit of amount of substance.
One mole contains exactly 6.02214076 x 1023 elementary entities.
This number is the fixed numerical value of the Avogadro constant Or Avogadro number.
The amount of substance, symbol n, of a system is a measure of the number of specified elementary entities.
An elementary entity may be an atom, a molecule. an ion, an electron. any other particle or specified group of particles.
The mole is the amount of substance of a system that contains 6.02214076 x 1023 specified elementary entities. 1 mole (of anything)= 6.022 x 1023 in number.
Besides being related to a number, a mole has one more advantage over a dozen or a gross.
This advantage is that mass of 1 mole of a particular substance is also fixed.
The mass of 1 mole of a substance is equal to its relative atomic or molecular mass in grams.
The atomic mass of an element gives us the mass of one atom of that element in atomic mass units (u).
To get the mass of 1 mole of atom of that element, that is, molar mass, we have to take the same numerical value but change the units from 'u' to 'g.
Molar mass of atoms is also known as gram atomic mass.
For example, atomic mass of hydrogen 1u. So, gram atomic mass of hydrogen = 1 g.
1 u hydrogen has only 1 atom of hydrogen and 1 g hydrogen has 1 mole atoms. that is. 6.022 x 10 atoms of hydrogen. Similarly,
16 u oxygen has only 1 atom of oxygen, 16 g oxygen has 1 mole atoms, that is. 6.022 x 10 atoms of oxygen.
To find the gram molecular mass or molar mass of a molecule, we keep the numerical value the same as the molecular mass, but simply change units as above from u to g. For example, as we have already calculated. molecular mass of water (H2O) is 18 u.
From here we understand that 18 u water has only 1 molecule of water. 18 g water has 1 mole molecules of water, that is, 6.022 x 10 23 molecules of water.
Chemists need the number of atoms and molecules while carrying out reactions, and
for this they need to relate the mass in grams to the number. It is done as follows:
1 mole = 6.022 x 10" number= Relative mass in grams.
Thus, a mole is the chemist's counting unit.
The word "mole" was introduced around 1896 by Wilhelm Ostwald who derived the term from the Latin word moles meaning a 'heap' or 'pile.
A substance may be considered as a heap of atoms or molecules. The unit mole was accepted in 1967 to provide a simple way of reporting a large number- the massive heap of atoms and molecules in a sample.
Example 3.3
1. Calculate the number of moles for the following
(1) 52 g of He
(2) 12.044 x 10 number of He atoms
Example 3.4 Calculate the mass of the
following:
(1)0.5 mole of N, gas
(2) 0.5 mole of N atoms
(3)3.011 x 10 number of N atoms
(4) 6.022 x 1023 number of N molecules
Example 3.5 Calculate the number of particles in each of the following:
(i)46 g of Na atoms
(ii) 8 g O2, molecules
(iii)0.1 mole of carbon atom
