Atoms and Molecules

Earlier Atomic Models and Subatomic Particles

The Atomic Model 

Atom Is Divisible

Do you recall Dalton’s atomic theory? Dalton postulated in his theory that an atom is indivisible. However, the later discoveries of protons and electrons proved this to be erroneous. 

In 1886, while carrying out an experiment in a gas discharge tube, E. Goldstein discovered positively charged radiations which led to the discovery of the subatomic particles called protons. Later, in 1897, J. J. Thomson discovered another type of subatomic particle—the negatively charged electron. Consequent to these discoveries, an atom was no longer indivisible; rather, it became a sum total of differently charged subatomic particles.

We know that an atom is neutral. It is made up of an equal number of oppositely charged particles—protons and electrons. Now, the question that arises is this:

How are the subatomic particles arranged inside an atom?

Many scientists performed varied experiments to develop different models for the structure of an atom. The first such model was proposed by J. J. Thomson. His atomic model is compared to a plum pudding and a watermelon; hence, it is known by the names ‘the plum-pudding model’.

The Plum-Pudding Model of an Atom 

Let us understand Thomson’s atomic model with the help of a slice of a watermelon. The slice consists of a red edible portion with embedded black seeds. Now, if we liken this watermelon to an atom, then (as per Thomson’s model) the positive charge in the atom is spread all over the red edible part; and the negatively charged particles, like the seeds, are embedded in this positively charged space.

In the same way, we can liken an atom to a plum pudding. In this case, the positive charge is spread all over the pudding, while the negatively charged particles are embedded like plums in this positively charged space. 

According to Thomson’s atomic model:

1. An atom consists of a positively charged sphere with electrons embedded in it.

2. The negative and positive charges present inside an atom are equal in magnitude. Therefore, an atom as a whole is electrically neutral.

 

 

Cathode Rays

J.J Thomson discovered that there are small particle present in the atom and that atom is divisible. J.J Thomson and his colleagues conducted experiments using discharge tube apparatus.
A discharge tube apparatus consists of a glass tube of about 15 cm length and 3 cm in diameter, filled with gas at low pressure. The tube is connected with the vacuum pump and two metal electrodes are fitted to the ends of the tube.

Low pressure was created inside the tube and high voltage was applied to the electrodes of the tube. This produced greenish glow at the anode end of the tube. The greenish glow at anode was produced due to the emission of the streams of rays from the cathode. These rays are known as cathode rays. Cathode rays will emit with blue glow.
When J.J Thomson placed a light paddle wheel inside the tube in the path of the cathode rays, the wheel started rotating. This led him to conclude that cathode rays are particulate in nature.

Properties of Cathode Rays

When J.J Thomson applied an electric field in the direction parallel to the path of cathode rays , he observed that the rays were deflected towards the anode.

This observation led to the conclusion that cathode rays are negatively charged.

When the above experiment was conducted with different gases, same observation were made and he named these negatively charged particles as electrons.

An electron is lighter than hydrogen atom and has very small mass in comparison to the mass of an atom.

Thus, J.J Thomson’s experiment and discovery of electron proved that atom is divisible and is made up of sub – atomic particles.

 

Sir Joseph John Thomson (1856−1940) was a British physicist. He is known for the discovery of electrons and for his model of an atom, popularly known as ‘the plum-pudding model’. He received the Nobel Prize in Physics in 1906 for discovering electrons and for his research on conduction in gases. In 1912, while working on the composition of canal rays, he and his colleague (F. W. Aston) found the first evidence for isotopes of neon.

 

Eugen Goldstein (1850−1930) was a German physicist. He is known for the discovery of canal rays which led to the discovery of protons. He also investigated comets using gas discharge tubes. His experiments established that a small object (like a ball) placed in the path of cathode rays produces emissions, flaring outward just like in case of a comet’s tail.

 

Canal Rays

After J.J Thomson’s discovery of atom another question arose that: if electrons are present inside the atom, then how is atom electrically neutral? Does this mean that there are positively charged particle also present inside the atom? 

To find out the answers to such questions, Goldstein conducted an experiment similar to that of J.J Thomson’s but with some modifications, for example he used perforated cathode in the discharge tube.

It was observed during the experiment that some rays were travelling in the direction opposite to that of cathode rays. Goldstein named these rays as anode rays.
When he applied an electric field in the direction parallel to that of the rays he observed that rays deflected towards cathode, thereby he concluded that anode rays are positively charged.
However, the deflection of anode rays in the discharge tube was found to be very less than that of cathode rays, because the emission of cathode rays was not dependent on the nature of the gas taken in the discharge tube. The deflection was seen highest for the hydrogen gas, when taken in the discharge tube.
The positive particles of hydrogen were found to be lightest and were named protons. Their mass is approximately equal to 1840 times that of electron. This mass is assumed as 1 atomic mass unit. The charge on a proton (+1) is equal to charge on an electron in magnitude (-1).

 

Why the Plum-Pudding Model Failed

The plum-pudding model of an atom was unable to explain the findings of Rutherford’s experiment while studying radioactivity.

In an experiment with gold foil, Rutherford bombarded the gold foil with alpha particles. With Thomson’s model as the basis, Rutherford expected small deviations; however, his findings were different from what was expected. 

As we go further into this lesson, we will learn more about Rutherford’s gold-foil experiment, his observations and his conclusions. We will also learn about the atomic model that he came up with on the basis of his conclusions.

Set up for Rutherford’s experiment:

1. A thin gold foil, approximately 1000 atoms thick, was taken. Gold was chosen for its high malleability.
2. A detector screen with a small slit (for emission of radiation from the atom) was placed around the foil.
3. A source of alpha particles was kept in front of the foil.
4. The foil was bombarded with fast-moving alpha particles.

The set-up for Rutherford’s gold-foil experiment is shown in the figure.

 

Rutherford’s Expectations and Observations

What Rutherford expected?

Rutherford expected that the alpha particles would pass straight through the foil and only a small fraction of alpha particles would be deflected. This expectation was in compliance with Thomson’s atomic model.

 

 

What Rutherford observed?

 

 

 

 

 

 

Rutherford’s findings were contrary to his expectation. He observed that:

1. Most of the fast-moving alpha particles passed straight through the gold foil.

2. Some particles were deflected through the foil by small angles.

3. One out of every 12000 particles rebounded, i.e., they got deflected by an angle of 180°.

What Rutherford Concluded from His Observations

Rutherford then carefully studied his observations and made the following conclusions.

1. Most alpha particles passed through the gold foil without any deflection. This indicates that most of the space inside an atom is empty.
2. Very few particles suffered a deflection from their path. This means that positive charge occupies very little space inside an atom.
3. Only a small fraction of particles underwent a 180° deflection. This shows that the entire positive charge and mass of an atom are present within a very small volume inside the atom.

A Quick Recap

 

Explore

 

Rutherford’s Atomic Model

Based on his conclusions in the gold-foil experiment, Rutherford devised his own atomic model. The major features of Rutherford atomic model or the nuclear model of an atom are as follows:

1. An atom consists of a nucleus at its centre and all the protons are present inside this nucleus.
2. Electrons reside outside the nucleus and revolve around the nucleus in well-defined orbits.
3. The size of the nucleus is very small as compared to the size of the atom. As per Rutherford’s calculations, the nucleus is 10⁵ times smaller than the atom.
4. Since the mass of the electrons is negligible as compared to the mass of the protons, almost all the mass of the atom is concentrated in its nucleus.

Drawbacks of Rutherford’s Model

 

Ernest Rutherford (1871−1937) was a British chemist and physicist. He is known as ‘the father of nuclear physics’. He discovered radioactive half-life. He proved that alpha radiations are nothing but helium ions. He was awarded the Nobel Prize in Chemistry in 1908 for his work on ‘the disintegration of elements’ and ‘the chemistry of radioactive substances’. He was the first scientist to split an atom in a nuclear reaction. The element ‘rutherfordium’ (atomic number 104) is named after him.

Example 1: 

What would have been observed if neutrons had been used to bombard the gold foil?

1. The observations of the experiment would have remained the same in spite of the change in the nature of the bombarding particles.
2. The neutrons would have suffered no deflection from the subatomic particles.
3. All the neutrons would have been absorbed by the gold atoms.
4. All the neutrons would have rebounded.

Solution: 

The correct answer is B.

Neutrons do not carry any charge; so, they do not suffer any repulsion. Hence, if neutrons had been used to bombard the gold foil, no deflection would have occurred. It is also possible that some neutrons would have been absorbed by the nucleus.

Example 2: 

State whether the following statements are true (T) or false (F).

1. Increasing the energy of the alpha particles will lead to more deflection.____
2. Speed of the alpha particles can be increased by increasing their energy.____
3. Use of aluminium sheet will lead to the same result as in case of gold foil.____

Solution:

1. T: Increasing the energy of the alpha particles will cause them to strike closer to the nucleus. Consequently, they will suffer greater deflection.
2. T: The kinetic energy of the alpha particles is directly related to their velocity. So, increasing their energy will result in an increase in the speed of the particles.
3. F: The positive charge on the nucleus in case of an aluminium foil is much smaller as compared to that on the gold nucleus. So, the result will vary.

Example 3: 

One of the postulates of Rutherford’s atomic model is that

1. an atom consists of a positively charged sphere.
2. an atom has its mass concentrated in its nucleus.
3. the nucleus of an atom is composed of electrons and protons.
4. the mass of an atom is the sum of the masses of all electrons and protons.

Solution:

The correct answer is B.

The postulates of Rutherford’s atomic model are as follows:

1. An atom consists of a nucleus at its centre and all the protons are present inside this nucleus.
2. Electrons reside outside the nucleus and revolve around the nucleus in well-defined orbits.
3. The size of the nucleus is very small as compared to the size of the atom. As per Rutherford’s calculations, the nucleus is 10⁵ times smaller than the atom.
4. Since the mass of the electrons is negligible as compared to the mass of the protons, almost all the mass of the atom is concentrated in its nucleus.

Rutherford also noticed that the actual mass of the nucleus was much more higher than the sum of the masses of protons and electrons. This lead him to predict that nucleus contains some kind of neutral particle whose mass must be equal to that of proton. 

This was experimentally proved by James Chadwick in the year 1932. He proved that nucleus of atom contains an additional neutral particle and called them neutrons. The mass of these neutrons is equal to that of protons.

Atomic Theory and Chemical Combination

Dividing Matter

Matter cannot be divided infinite number of times.

 

For example, if we keep chopping a log of wood into smaller and smaller pieces, then we will reach a point when the wood will not be divisible any further. Minute particles of wood will remain and these will not be visible to the naked eye. This is true for all forms of matter.The same was believed by the early Indian and Greek philosophers. In India, around 500 BC, an Indian philosopher named Maharshi Kanad called these smallest particles ‘parmanu’. The word ‘atom’ is derived from the Greek word ‘atomos’ which means ‘indivisible’. It was the Greek philosopher Democritus who coined the term. However, for these ancient thinkers, the idea of the minute indivisible particle was a purely philosophical consideration. 

By the end of the eighteenth century, scientists had begun to distinguish between elements and compounds. Two French chemists named Antoine Lavoisier and Joseph Proust observed that elements combine in definite proportions to form compounds. On the basis of this observation, each of them proposed an important law of chemical combination. The laws proposed by them helped Dalton formulate his atomic theory.

Dalton’s Atomic Theory 

In the early nineteenth century, an English chemist named John Dalton proposed a theory about atoms. Known as ‘Dalton’s atomic theory’, it proved to be one of the most important theories of science. The various laws of chemical combination also supported Dalton’s theory. Dalton asserted that ‘atoms are the smallest particles of matter, which cannot be divided further’. He published his atomic theory in 1808 in his book A New System of Chemical Philosophy. The postulates of Dalton’s atomic theory are as follows:

• All matter is made up of very tiny particles. These particles are called atoms.

• An atom cannot be divided further, i.e., atoms are indivisible.

• Atoms can be neither created nor destroyed in a chemical reaction.

• All atoms of an element are identical in all respects, e.g. in terms of mass, chemical   properties, etc.

• Atoms of different elements have different masses and chemical properties.

• Atoms of different elements combine in small whole-number ratios to form compounds.

• In a given compound, the relative numbers and types of atoms are constant.

John Dalton (1766−1844) was born into the poor family of a weaver in Eaglesfield, England. He was colour-blind from childhood. He became a teacher when he was barely twelve years old. By the time he was nineteen, he had become the principal of a school. In 1793, Dalton left for Manchester to teach physics, chemistry and mathematics at a college. Elected a member of the Manchester Literary and Philosophy Society in 1794, he became its president in 1817 and remained in that position until his death. During his early career, he identified the hereditary nature of red−green colour blindness. In 1803, he postulated the law of partial pressures (known as Dalton’s law of partial pressures). He was the first scientist to explain the behaviour of atoms in terms of relative atomic weight. He also proposed symbolic notations for various elements.

Law of Chemical Combination 

Lavoisier and Proust proposed two laws that sought to explain the chemical combinations of elements. These laws are the law of conservation of mass and the law of constant proportions, respectively.

Let us first study about the law of conservation of mass.

According to this law, ‘mass can be neither created nor destroyed.’ In other words, for a chemical reaction taking place in a closed system, the total mass of the reactants is the same as that of the products.

For example, an unmixed solution of barium chloride and aluminium sulphate weighs the same as that of the mixed solution of the two.

Closed System

A system is said to be closed if there is no exchange of mass across the boundaries of the system and the surroundings. However, it can exchange energy with its surroundings.

 

Law of Conservation of Mass 

 

Example

Law of Constant Proportions 

This law is also known as the law of definite proportions. According to this law, ‘in a compound, elements are always present in definite proportions by masses.’ We know that compounds are composed of two or more elements. So, according to this law, the proportions in which elements are present in a compound remain the same, irrespective of its method of preparation. Let us understand this law with the help of a couple of examples.

Pure water obtained from any source (well, river, lake or sea) and from any country (India, Russia or America) will always contain two hydrogen atoms and one oxygen atom. The atoms of hydrogen combine with the atom of oxygen in the ratio of 1 : 8 by mass to form water. The ratio by the number of atoms for water will always be 2 : 1. The mass of a water molecule is 18 g. So, a molecule of water contains 2 g of hydrogen and 16 g of oxygen. 

Ammonia contains one nitrogen atom and three hydrogen atoms. Irrespective of the source from which ammonia is obtained, it will always contain nitrogen and hydrogen in the ratio of 14 : 3 by mass. The mass of an ammonia molecule is 17 g. So, a molecule of ammonia contains 14 g of nitrogen and 3 g of hydrogen. Similarly, 34 g of ammonia contains 28 g of nitrogen and  6 g of hydrogen.

Laws of Chemical Combination and Dalton’s Atomic Theory

The law of conservation of mass can be explained using the first and third postulates of Dalton’s atomic theory.

The law of conservation of mass states that mass can be neither created nor destroyed. According to this law, in case of a chemical reaction taking place in a closed system, the total mass of the reactants equals that of the products.

Mass is the amount of matter in something. As per the first postulate of Dalton’s atomic theory, all matter is made up of atoms. The third postulate of the same theory asserts that atoms can be neither created nor destroyed in a chemical reaction, i.e., the total number of atoms and their mass should remain the same before and after the reaction. This is the same as the law of conservation of mass.

Laws of Chemical Combination and Dalton’s Atomic Theory

The law of constant proportions follows directly from the sixth and seventh postulates of Dalton’s atomic theory.

The sixth postulate of Dalton’s atomic theory states that atoms of different elements combine in small whole-number ratios to form compounds. The seventh postulate states that the relative numbers and types of atoms in a compound are constant. This has the same meaning as the law of constant proportions, which states that the elements in a compound are always present in definite proportions by masses.

Now, we know that a sample of carbon dioxide (no matter how it is prepared) is made up of carbon and oxygen. A molecule of carbon dioxide comprises 1 carbon atom and 2 oxygen atoms. The mass of a carbon dioxide molecule is 44 g. The relative atomic masses of carbon and oxygen are 12 u and 16 u respectively. As per the law of constant proportions, in carbon dioxide, carbon and oxygen always combine in the ratio of 3 : 8 by mass.

Atoms, Molecules and Ions

Atoms: An Overview 

When we talk about atoms, two questions usually strike our mind... 

Let us go through this lesson to find the answers to these questions. We will also learn how to represent different atoms in symbolic forms. So, in short, we are going to study:

• Size of an atom
• Representation of atoms
• Atomic mass

Size of an Atom 

• The size of an isolated atom cannot be measured; however, we can estimate its size by assuming that its radius is half the distance between adjacent atoms in a solid.
• Atoms are very small in size. They are so small that it is not possible to see them even under a powerful optical microscope.
• The size of an atom is indicated by its radius, called the atomic radius.
• Since an atom is very small, we need a very small unit for reporting the atomic radius; thus, the radius of an atom is often expressed in nanometre.

    

 

                                          

Surfaces of Silicon Atoms

Atoms cannot be seen with the naked eye, but the use of modern techniques has enabled us to see the surfaces of atoms. The magnified image of the surfaces of silicon atoms is shown in the following figure.

 

Size of an Atom 

Hydrogen atom is the smallest of all atoms. The given figure shows the atomic radii of some elements.

Classical Representation of Atoms

• A large number of elements are known to us today. It would be cumbersome to refer to them by their names all the time in our studies. For the sake of convenience, we need symbols that represent these elements. Toward the end of the nineteenth century, scientists felt this need to assign standard characteristic symbols to the elements. 
• John Dalton was the first scientist to use symbols to represent different elements. Dalton’s proposed symbols for some elements are shown in given figure.

• Each symbol proposed by Dalton represents an atom of the respective element. For example, if someone wanted to represent two hydrogen atoms, then he would have to draw the symbol of hydrogen atom twice as shown.

Modern Representation of Atoms

• Many of the symbols proposed by Dalton were difficult to draw and remember. Therefore, an alternative method of representing elements was required.
• Another scientist named JÓ§ns Jacob Berzelius suggested that letters of the alphabet can be used as symbols to represent the elements. The modern symbols of elements are based on this idea.
• The International Union of Pure and Applied Chemistry (IUPAC) approves the names and symbols for the elements.
• The modern symbol of an element is made up of one or two letters of the English or Latin name of that element.
• As a rule, the first letter of a symbol is always written as a capital letter and the second as a small letter.
• The modern representation of atoms is more convenient and meaningful than the classical representation. 
• To conclude, the symbols of the various elements are significant as:
  • They represent distinct elements.
  • They represent single atoms of the elements.

Modern Representation of Atoms

Table mentioned below shows the modern representation of atoms:

Elements	Symbols	Elements	Symbols
Aluminium	Al	Iron (from Latin: ferrum)	Fe
Argon	Ar	Lead (from Latin: plumbum)	Pb
Calcium	Ca	Magnesium	Mg
Carbon	C	Nitrogen	N
Chlorine	Cl	Oxygen	O
Copper (from Latin: cuprum)	Cu	Potassium (from Latin: kalium)	K
Fluorine	F	Silicon	Si
Gold (from Latin: aurum)	Au	Silver (from Latin: argentum)	Ag
Hydrogen	H	Sodium (from Latin: natrium)	Na
Iodine	I	Zinc	Zn

 

IUPAC The International Union of Pure and Applied Chemistry (IUPAC) is an international non-governmental organisation established in 1919.  IUPAC nomenclature, in general, is a system of naming chemical compounds and describing the science of chemistry. Some man-made elements are given temporary three-letter symbols.

Atomic Mass 

• Every atom has some characteristic mass of its own and this is known as atomic mass.
• All the atoms of an element have the same atomic mass. Atoms of different elements have different atomic masses.
• Determination of atomic mass:

       It is difficult to determine the mass of an individual atom; so, the mass of an atom is        ascertained in relation to the mass of C-12 isotope. Thus,        atomic mass is in fact relative atomic mass.

• C-12 is an isotope of carbon and its mass is used as a standard reference to calculate the relative atomic masses of all elements.
• The IUPAC adopted one-twelfth of the mass of a C-12 isotope as the standard unit to measure relative atomic masses. It named this unit as atomic mass unit (amu) or unified atomic mass unit (u).

So,

In simple words, Atomic mass is a term which gives the total mass of protons and neutrons in an atom. Also, the atomic mass is measured with respect to mass of 1/12 the mass of one carbon atom.

Atomic Mass 

The atomic masses of some common elements are given in the following table.

Elements	Atomic masses (u)	Elements	Atomic masses (u)
Hydrogen	1	Chlorine	35.5
Helium	4	Potassium	39
Carbon	12	Calcium	40
Nitrogen	14	Argon	40
Oxygen	16	Iron	56
Fluorine	19	Copper	63.5
Neon	20	Zinc	65
Sodium	23	Bromine	80
Magnesium	24	Silver	108
Sulphur	32	Gold	197

Atomic Mass Unit

Atomic mass unit (1 u) is defined as exactly one twelfth the mass of an atom of carbon-12. Atomic mass unit is only number and it has no units. On the basis of above unit, the atomic mass of carbon atom is 12 amu.

Relative atomic mass or atomic weight

Definition with respect to hydrogen:

It is the ratio of mass of one atom of an element to the mass of an atom of hydrogen taken as unity.

Definition with respect to carbon:

It is the ratio of mass of one atom of an element to 1/12th mass of an atom of carbon.

Relative molecular mass or molecular weight

Definition with respect to hydrogen:

It is the ratio of mass of one molecule of a substance to the mass of an atom of hydrogen taken as unity.

Definition with respect to carbon:

It is the ratio of mass of one atom of an element to 1/12th mass of an atom of carbon.

Gram Molecular Volume

The volume occupied by 1 gram molecule of a dry gas at S.T.P is called gram molecular volume. The experimental value of 1 gram molecular volume of a gas is 22.4 litres at S.T.P.

The relationship between the mole, Avogadro’s number, and mass is summarised as follows:

Symbols are very important in chemistry as they are used to represent different  elements. The usage of symbols to represent elements has been in trend from the ancient Greek time. Greeks used symbols to represent the four elements, earth, air, fire and water.
In the era of alchemists, different materials were represented using pictorial symbols.
For example,

Element	Symbol
Nickel
Arsenic
Antimony
Water

 

Do You Know? The process of converting a less valuable metal into a more valuable metal like gold is called alchemy, and the men who started this process are called alchemists.

System for determining symbols for different elements:
Certain rules have been framed to determine symbols of different elements. They are as follows:
1. The symbols of common elements, mainly non-metals, are the first letter of their respective names.
For example, the symbols of oxygen and fluorine are O and F respectively.

2. If the name of an element shares the initial letter with another element, then the first and the second letter of its  name  are used as the symbol.
For example, symbols used for Barium and Beryllium are Ba and Be respectively.

3. If the first two letters of the names are the same, then the first and third lettes are used as the symbol.
For example, symbols used for Magnesium and Manganese are Mg and Mn respectively.

4. Symbols of some elements are based on their old or Latin names. There are 11 elements whose symbols are derived from their Latin names.
For example, symbol of sodium is Na, which is derived from its Latin name Natrium.

5. If the symbol of any element is a single letter, it should be written in capital.

6. If the symbol of any  element has two letters then the first one should be in capital followed by small letter.

Significance of symbols:
1. Symbol of an element signifies the name of the element.
2. It also signifies that one atom of that element is present.

A Brief Introduction to Molecules and Ions

Most atoms are not stable in free state. So, they combine with other atoms to form molecules.

For example:

A water molecule is formed when two hydrogen atoms combine with one oxygen atom. An oxygen molecule is formed when two oxygen atoms combine with each other.

                             

Some atoms are charged. Such charged atoms and molecules are called ions.

A positively charged ion is called cation.

A negatively charged ion is called anion.

 

In this lesson, we are going to study about:

• Molecules and molecular compounds
• Ions and ionic compounds

Molecules 

The term ‘molecule’ originates from the French word ‘molécule’, which means ‘extremely minute particle’. It was coined by the French philosopher and mathematician Rene Descartes in the early seventeenth century. 

In view of John Dalton’s laws of definite and multiple proportions, the existence of molecules was accepted by many chemists since the early nineteenth century. However, it is the work of Jean Baptiste Perrin on the Brownian motion (1911) of particles of liquids and gases which is considered to be the final proof of the existence of molecules.

Atomicity of Molecules

The number of atoms constituting a molecule is known as its atomicity. The given table lists the atomicity of some common elements. 

Elements	Atomicity
Helium (He), Neon (Ne), Argon  (Ar)	Monoatomic (1 atom per molecule)
Oxygen (O2), Hydrogen (H2), Nitrogen (N2) Chlorine (Cl2), Fluorine (F2)	Diatomic (2 atoms per molecule)
Phosphorus (P4)	Tetratomic (4 atoms per molecule)
Sulphur (S8)	Polyatomic (8 atoms per molecule)

Buckminsterfullerene is an allotrope of carbon in which sixty carbon atoms are bonded together.

 

Ions

A positively charged ion is called cation, while a negatively charged ion is called anion.

There are many ions which are polyatomic ions.

The given table lists the symbols and atomicity of some common ions.

Cations	Symbols	Atomicity	Anions	Symbols	Atomicity
Aluminium	Al3+	Monoatomic	Bromide	Br−	Monoatomic
Ammonium		Polyatomic	Carbonate		Tetra-atomic
Calcium	Ca2+	Monoatomic	Chloride	Cl−	Monoatomic
Cuprous ion	Cu+	Monoatomic	Fluoride	F−	Monoatomic
Cupric ion	Cu2+	Monoatomic	Hydride	H−	Monoatomic
Hydrogen	H+	Monoatomic	Hydroxide	OH−	Diatomic

Ions

The given table lists the symbols and atomicity of some other common ions.

Cations	Symbols	Atomicity	Anions	Symbols	Atomicity
Ferric ion	Fe3+	Monoatomic	Iodide	I−	Monoatomic
Magnesium	Mg2+	Monoatomic	Nitrate		Tetra-atomic
Nickel	Ni2+	Monoatomic	Nitride	N3−	Monoatomic
Potassium	K+	Monoatomic	Nitrite	NO−2	Tetra-atomic
Silver	Ag+	Monoatomic	Oxide	O2−	Monoatomic
Sodium	Na+	Monoatomic	Phosphate		Polyatomic
Zinc	Zn2+	Monoatomic	Sulphate		Polyatomic
Hydrogen carbonate	HCO−3	Polyatomic	Sulphite		Tetra-atomic

Ionic Compounds 

The compounds which are formed by the combination of cations and anions are known as ionic compounds. 

For example:

• Zinc oxide (ZnO): It is formed when a zinc ion (Zn²⁺) combines with an oxide ion (O²⁻). 
• Magnesium chloride (MgCl₂): It is formed when a magnesium ion (Mg²⁺) combines with two chloride ions (Cl⁻).
• Potassium bromide (KBr): It is formed when a potassium ion (K⁺) combines with a bromide ion (Br⁻).
• Sodium chloride (NaCl): It is formed when a sodium ion (Na⁺) combines with a chloride ion (Cl⁻). The structure of NaCl crystals is shown in the given figure. You can see that there is a group of Na⁺ and Cl⁻ ions combined with each other.

Ionic Compounds 

 

 

Example 1: 

Find the atomicity of each of the following ions.

i) S²⁻

_ii)

_iii)

iv) OH⁻

Solution:

Ions	Atomicity
S2-	1
	5
	5
OH−	2

Example 2: 

Identify the anions and cations present in the following compounds.

Compounds	Anions	Cations
NaCl
KMnO4
NaOH
KBr
NH4OH

Solution:

Compounds	Anions	Cations
NaCl	Cl−	Na+
KMnO4		K+
NaOH	OH−	Na+
KBr	Br−	K+
NH4OH	OH−

Example 3: 

Give the symbols and valence numbers for the following ions.

Ions	Symbols	Valence numbers
Ammonium
Carbonate
Sulphate
Chloride
Phosphate

Solution:

Ions	Symbols	Valence numbers
Ammonium		+1
Carbonate		−2
Sulphate		−2
Chloride	Cl−	−1
Phosphate		−3

Chemistry of Compounds

Molecular Formula: A Brief Overview

Just like each atom has a unique symbol, each compound has a unique molecular formula.

The molecular formula of a compound provides information about the names and numbers of atoms of the different elements present in a molecule of that compound.

Molecular formula is a chemical formula that indicates the kinds of atoms and the numbers of each kind of atom in a molecule of a compound.

Examples

• The molecular formula of glucose is C₆H₁₂O₆. One molecule of glucose contains 6 atoms of carbon, 12 atoms of hydrogen and 6 atoms of oxygen.
• The molecular formula of water is H₂O. One molecule of water contains 2 atoms of hydrogen and 1 atom of oxygen.

Salient features of chemical formula:

• Compounds are formed when two or more elements combine chemically. Hence, compounds can also be represented using symbols.
• The notation used for representing any compound is called chemical formula of that compound.
• Each compound has a unique chemical formula.
• The chemical formula of any compound tells us about : The different elements which combine to form the compound and the number of atoms of each element present in a molecule of the compound
• For example, H₂O is the chemical formula of water. This denotes that there are two atoms of hydrogen and one atom of oxygen present in one molecule of water.

Chemical Formulae

Let us understand the information derived from chemical formulae by taking the example of carbon dioxide. The chemical formula of carbon dioxide is CO₂. Using this formula, we can derive the following information about carbon dioxide.

• Two elements are present in carbon dioxide: carbon(C) and oxygen (O).
• CO₂ represents one molecule of carbon dioxide.
• Since one atom of carbon combines with two atoms of oxygen, the valency of carbon is twice that of oxygen.
• CO₂ is a neutral molecule. It has no charge.
• The relative atomic masses of carbon and oxygen are 12 u and 16 u respectively. So, the ratio by mass between carbon and oxygen is 12 : 32, i.e., 3 : 8.

Writing Chemical Formulae

To write the chemical formula of a compound, one should have prior knowledge of two things.

• The symbols of the constituent elements.
• The combining capacity of the atom of each element constituting the compound.

The number of atoms of other elements with which one atom of an element combines is decided by the valency of that element.

For example, both hydrogen (H) and chlorine (Cl) have a valency of 1. Therefore, one atom of hydrogen reacts with one atom of chlorine to form one molecule of hydrogen chloride (HCl).

The valency of an ion is equal to the charge on it.

Chemical Formulae

The valencies of some common ions are given in the following table.

Names of ions	Symbols	Valencies	Names of ions	Symbols	Valencies
Aluminium	Al3+	3	Sulphite		2
Ammonium		1	Bromide	Br−	1
Calcium	Ca2+	2	Carbonate		2
Copper(II)	Cu2+	2	Chloride	Cl−	1
Hydrogen	H+	1	Hydride	H−	1

Chemical Formulae

The valencies of some common ions are given in the following table.

Names of ions	Symbols	Valencies	Names of ions	Symbols	Valencies
Iron(II)	Fe2+	2	Hydrogen carbonate		1
Iron(III)	Fe3+	3	Hydroxide	OH−	1
Magnesium	Mg2+	2	Nitrate		1
Nickel	Ni2+	2	Nitrite		1
Potassium	K+	1	Oxide	O2−	2
Silver	Ag+	1	Phosphate		3
Sodium	Na+	1	Sulphate		2
Zinc	Zn2+	2	Sulphide	S2−	2

Chemical Formulae

The following rules need to be kept in mind while writing the chemical formulae of compounds.

•The valencies or charges on the ions must be balanced. The charge on a cation must be equal in magnitude to the charge on an anion so that the opposite charges cancel each other out and the net charge of the molecule becomes zero.

Examples

• In case of CaO, the valency of Ca is +2 and that of O is −2. These are then crossed over and the compound formed is CaO.

Formula of calcium oxide

Symbols Ca O

Charges 2+ 2−

• The charge on Mg²⁺ is +2 and that on Cl⁻ is −1. Thus, one Mg²⁺ ion combines with two Cl⁻ ions to form a molecule with the formula MgCl₂.
• In case of a compound consisting of a metal and a non-metal, the symbol of the metal is written first.

Chemical Formulae

Example

• In calcium chloride (CaCl₂) and zinc sulphide (ZnS), calcium and zinc are metals, so they are written first; chlorine and sulphur are non-metals, so they are written after the metals.
• In case of compounds consisting of polyatomic ions, the polyatomic ions are enclosed in brackets before writing the number to indicate the ratio.

Example

• In case of aluminium sulphate, to balance the charges, two ions combine with one Al³⁺ ion. Thus the formula for aluminium sulphate is Al₂(SO₄)₃. Here, the brackets with the subscript 3 indicate that three sulphate ions are joined to two aluminium ions.

Formula of aluminium sulphate

Symbols Al SO₄

Charges 3+ 2−

Chemical Formulae

 

Example 1:

Give two examples each of molecules having one atom, two atoms and three atoms.

Solution:

Molecules having one atom (/monatomic molecules): Argon (Ar) and Neon (Ne)

Molecules having two atoms (/diatomic molecules): Nitrogen (N₂) and Oxygen (O₂)

Molecules having three atoms (/triatomic molecules): Nitrogen dioxide (NO₂) and carbon dioxide (CO₂)

Example 2:

The valencies of a few ions are provided below.

H⁺ = 1, = 2, Br⁻ = 1, Mg²⁺ = 2 and K⁺ = 1

Write the formulae for magnesium bromide, magnesium sulphate, hydrogen bromide and potassium sulphate.

Solution:

•Magnesium bromide: MgBr₂

•Magnesium sulphate: MgSO₄

•Hydrogen bromide: HBr

•Potassium sulphate: K₂SO₄

Example 3:

Write the names of the following compounds.

i)H₂CO₃

ii)KNO₃

iii)(NH₄)₃PO₄

iv)Na₂CO₃

v)Al(NO₃)₃

vi)NaHCO₃

Solution:

i)H₂CO₃: Hydrogen carbonate

ii)KNO₃: Potassium nitrate

iii)(NH₄)₃PO₄: Ammonium phosphate

iv)Na₂CO₃: Sodium carbonate

v)Al(NO₃)₃: Aluminium nitrate

vi)NaHCO₃: Sodium hydrogen carbonate

Mass of a Molecule-An Overview

The molecular mass of a substance is the sum of the atomic masses of all the atoms present in a molecule of that substance. Molecular mass is expressed in atomic mass unit (u).

To calculate the molecular mass of a substance, the masses of all its constituent atoms are added. For example:

Molecular mass of glucose, C₆H₁₂O₆

= 6 (Atomic mass of C) + 12 (Atomic mass of H)
+ 6 (Atomic mass of O)

= 180 u

Molecular Mass and Its Calculation

In the early twentieth century, scientists performed a number of experiments to conclude that an atom can be further divided into three types of  subatomic particles— electrons, protons and neutrons. The neutrons and protons of an atom provide it its mass. This is called atomic mass.

As mentioned before, the atoms constituting a molecule provide it its characteristic mass. The molecular mass of a substance is the sum of the atomic masses of all the atoms present in a molecule of that substance. This is also called ‘relative molecular mass’ and its unit is atomic mass unit (u). 

So, to calculate the mass of a molecule, we need to know the masses of the individual atoms present in that molecule.

Let us consider a molecule A_aB_bC_cD

Here,

• A, B, C and D are the atoms of four arbitrary elements. 
• a, b and c are the respective numbers of atoms of elements A, B and C present in the molecule.
• There is only one atom of element D in the molecule.

The molecular mass of A_aB_bC_cD is calculated as follows:

(a × Atomic mass of A) + (b × Atomic mass of B) + (c × Atomic mass of C)
+ (1 × Atomic mass of D)

Example 1: 

Calculate the formula unit mass of sodium hydroxide. The atomic masses of sodium, hydrogen and oxygen are 23 u, 1 u and 16 u respectively.

Solution: 

The chemical formula of sodium hydroxide is NaOH.

It is given that:

Atomic mass of sodium (Na) = 23 u

Atomic mass of oxygen (O) = 16 u

Atomic mass of hydrogen (H) = 1 u

A formula unit of sodium hydroxide contains one atom each of sodium, oxygen and hydrogen.

So, formula unit mass of NaOH = (23 + 16 + 1) u = 40 u

Example 2: 

Calculate the formula unit mass of potassium sulphate. The atomic masses of potassium, oxygen and sulphur are 39 u, 16 u and 32 u respectively.

Solution:

The chemical formula of potassium sulphate is K₂SO₄.

It is given that:

Atomic mass of potassium (K) = 39 u

Atomic mass of sulphur (S) = 32 u

Atomic mass of oxygen (O) = 16 u

A formula unit of potassium sulphate contains two atoms of potassium, one atom of sulphur and four atoms of oxygen.

So, molecular mass of K₂SO₄ = [(2 × 39) + (1 × 32) + (4 × 16)] u 

= (78 + 32 + 64) u 

= 174 u

Example 3:

Calculate the formula unit mass of [Co(NH₃)₅Br]SO₄._­­

Solution:

Atomic mass of cobalt (Co) = 59 u

Atomic mass of nitrogen (N) = 14 u

Atomic mass of hydrogen (H) = 1 u

Atomic mass of bromine (Br) = 80 u

Atomic mass of sulphur (S) = 32 u

Atomic mass of oxygen (O) =16 u

The formula unit mass of [Co(NH₃)₅Br]SO₄ is calculated as follows:

(Atomic mass of Co) + {5 × [(Atomic mass of N) + (3 × Atomic mass of H)]} 

+ (Atomic mass of Br) + (Atomic mass of S) + (4 × Atomic mass of O) 

= 59 u + {5 × [14 u + (3 × 1) u]} + 80 u + [32 u + (4 × 16) u] 

= (59 + 85 + 80 + 96) u

= 320 u

Molecular Mass of Elements

 

 

 

Formula Unit and Formula Unit Mass

• Molecular mass and formula unit mass are basically the same. Molecular mass is the mass of a molecule, while formula unit mass is the mass of an ionic compound.
• The term formula unit is used for substances whose constituent particles are ions. For example, the formula unit of calcium oxide is CaO.
• The formula unit mass of a substance is the sum of the atomic masses of all the atoms in the formula unit of a compound. 

Let us try to calculate the formula unit mass of nitric acid.

The chemical formula of nitric acid is HNO₃.

Atomic mass of hydrogen (H) = 1 u

Atomic mass of nitrogen (N) = 14 u

Atomic mass of oxygen (O) = 16 u

In a formula unit of HNO₃, there are one hydrogen (H) atom, one nitrogen (N) atom and three oxygen (O) atoms.

Thus, formula unit mass of HNO₃ = (1 × 1) u + (1 × 14) u + (3 × 16) u

= (1 + 14 + 48) u

= 63 u

Properties of Atoms and Molecules

Atomic Number and Mass Number 

In the 1830s, representation of elements and compounds was a major concern for chemists.

Many symbolic notations for elements were devised during this period. Gradually, the representations became standardized. Currently, the general symbolic notation for an element is:

. Now, take for example the specific symbolic notations for oxygen and nitrogen.

Element	Symbolic notation
Oxygen
Nitrogen

Wondering what these symbolic notations represent? Go through this lesson to find out.

You know that the symbolic notation of oxygen is . In this notation, the letter ‘O’ symbolises the element ‘oxygen’; the number ‘16’ represents the mass number of oxygen; and the number ‘8’ indicates the atomic number of oxygen. 

Thus, in the general symbolic notation of an element, the letter ‘E’ is the symbol of the element, the letter ‘A’ is its mass number, and the letter ‘Z’ is its atomic number.

The atomic number is the number of protons present in the nucleus of an atom. It is denoted by Z. 

The total number of the protons and the neutrons present in the nucleus of an atom is known as mass number. It is denoted by A.

Atomic Number and Mass Number 

 

Symbolic Notations of Some Elements

Elements	Symbolic notations	Symbols	Atomic numbers	Mass numbers
Hydrogen		H	1	1
Helium		He	2	4
Lithium		Li	3	7
Beryllium		Be	4	9
Boron		B	5	11
Carbon		C	6	12
Nitrogen		N	7	14
Oxygen		O	8	16
Fluorine		F	9	19
Neon		Ne	10	20

Symbolic Notations of Some Elements

Elements	Symbolic notations	Symbols	Atomic numbers	Mass numbers
Sodium		Na	11	23
Magnesium		Mg	12	24
Aluminium		Al	13	27
Silicon		Si	14	28
Phosphorus		P	15	31
Sulphur		S	16	32
Chlorine		Cl	17	35
Argon		Ar	18	40
Potassium		K	19	39
Calcium		Ca	20	40

Relation between Atomic Number and Mass Number  

Mass number (A) of an atom = Number of protons + Number of neutrons

Therefore, Mass number (A) = Atomic number (Z) + Number of neutrons

Therefore, Number of neutrons = A - Z

Hence, the number of neutrons can be calculated if the atomic number and mass number of an element are known.

An atom of sodium contains 11 protons and 12 neutrons. Can you calculate the mass number of sodium atom?

Now, mass number (A) = number of protons + number of neutrons

Therefore, mass number of sodium atom = 11 + 12 = 23

Hence, the mass number of sodium is 23.

An atom of carbon is represented as. Can you tell the number of neutrons and protons present in carbon atom?

It is seen from the symbolic notation of carbon that the atomic number and mass number of carbon atom is 6 and 12 respectively.

Now, number of neutrons = mass number − atomic number = 12 − 6 = 6

Since the number of protons is equal to the atomic number of that element. Thus, the number of protons present in a carbon atom is 6.

Example 1: 

What is the symbol of the element sodium?

1. Na
2. N
3. So
4. S

Solution: 

The correct answer is A.

The symbol of sodium is Na. It is derived from the Latin name for the element, i.e., ‘natrium’.

Example 2: 

What is the atomic number of an element having five protons and six neutrons?

1. 11
2. 9
3. 6
4. 5

Solution: 

The correct answer is D.

The atomic number of an element is the number of protons or electrons present in an atom of the element. Since an atom of the given element has five protons, its atomic number is 5.

Example 3: 

What is the number of neutrons in an element having 39 protons and 89 as its mass number?

1. 45
2. 50
3. 55
4. 60

Solution:

The correct answer is B.

We know that:

Mass number = Number of protons + Number of neutrons

In case of the given element:

Mass number = 89

Number of protons = 39

So,

89 = 39 + Number of neutrons

=> Number of neutrons = 89 -39 = 50

Example 4: 

What is the symbol of the element having 22 neutrons and 40 as its mass number?

1. Al
2. Mg
3. Ar
4. Ca

Solution:

The correct answer is C.

The given element has:

Mass number = 40

Number of neutrons = 22

We know that: 

Mass number = Number of protons + Number of neutrons

So,

40 = Number of protons + 22

=> Number of protons = 40 - 22 = 18

Also,

Atomic number = Number of protons = 18

Argon is the element having 18 as its atomic number and 40 as its mass number. The symbol of argon is Ar.

• Water is the major constituent of the human body. It is made up of two elements: hydrogen and oxygen.
• Almost all the mass of our body is made up of the following six elements.

1. Oxygen (65%)
2. Carbon (18%)
3. Hydrogen (10%)
4. Nitrogen (3%)
5. Calcium (1.5%)
6. Phosphorus (1%)

• Some of the other elements found in our body are:
• Sulphur (0.25%)
• Sodium (0.15%)
• Magnesium (0.05%)
• Zinc (0.7%)

The Periodic Table

The periodic table is a table classifying all the known elements.

It is divided into 18 columns (called groups) and 7 rows (called periods).

The elements are arranged in the rows or periods by order of increasing atomic number.

The elements in the columns or groups display similar chemical and physical properties. This feature of the periodic table makes it easy to study the vast number of elements.

The periodic table is shown in the figure.

 

Isotopes

In 1910, Frederick Soddy recorded the existence of elements having different atomic masses to show similar properties. These elements came to be known as isotopes.
Isotopes are defined as atoms of same element having the same atomic number, but different mass numbers. These atoms contain an equal number of protons and electrons, but a different number of neutrons.
For example, in nature, hydrogen is found in three forms with different mass numbers, namely protium ( H11), deuterium( H12), and tritium (  H13 ). These are the three naturally occurring isotopes of hydrogen. The atomic number of each isotope is 1, but the mass number varies i.e. it is 1, 2, and 3 respectively. Some other examples of isotopes include C- 12 and C-14, which are isotopes of carbon, and Cl-35 and Cl-37, which are isotopes of chlorine.
 

However, it was F. W. Aston who, in 1919, discovered various stable isotopes for a number of elements. In that year, he was successful in proving the existence of two isotopes of neon:
neon-20 and neon-22. 

The isotopes of an element are species with different numbers of neutrons but the same numbers of protons and electrons.

Isotopes are pure substances just like elements and compounds.

We know that the atomic mass of an element is the sum total of its protons and neutrons. Now, in case of an element having isotopes, the number of neutrons does not remain fixed. The atomic mass of such an element includes the atomic masses of its isotopes. This gives rise to the term ‘average atomic mass’. The average atomic mass of an element with isotopes takes into account the atomic masses of all its isotopes with respect to their abundance in nature.

Applications of Isotopes

1. An isotope of uranium exhibits nuclear fission properties. It is used in nuclear reactions as a fuel.
2. An isotope of cobalt is used for treating cancer.
3. An isotope of iodine is used for treating goitre.
4. An isotope of carbon is used in radiocarbon dating to determine the age of an organic sample.
5. An isotope of calcium is used in biomedical research on cellular functions and bone formation in mammals.
6. An isotope of iron is used for detecting sulphur in air.
7. An isotope of hydrogen (tritium) is used in estimating the age of water bodies and the rate of their replenishment through precipitation.

Relationship between Atomic Masses and Isotopes of an Element

 

Average Atomic Mass of an Element

 

Simulating Isotopes

 

 

 

Isobars and Isotones

Isobars

Isobars are elements having the same mass number but different atomic numbers. Take, for example,. They have the same mass number (i.e., 13) but different atomic numbers (i.e., 6 and 7).

One can also say that isobars have an equal number of nucleons but different numbers of protons, neutrons and electrons.

Here are two more examples of isobars: 

1. 
2. 

Isotones

Isotones are atoms of different elements having the same number of neutrons.

Here are some examples isotones: 

1. 
2. 
3. 

Example 1: 

Which of the following atomic pairs are isotopes of each other?

A.  
B.  
C.  
D.  

Solution:

The correct answer is B.

Isotopes are those species of the same element which possess the same number of protons but different numbers of neutrons. In other words, isotopes have the same atomic number but different atomic masses. Among the given pairs, only is an atomic pair with the same number of protons

 

Example 2: 

are the three stable isotopes of magnesium. The average atomic mass of magnesium is 24.3. The relative abundance of  and are 78.7% and 11.2% respectively. What percentage of magnesium existing in nature is ?

A.   10%
B.   12%
C.   14%
D.   16%

Solution:

The correct answer is A.

Let us take:

Average atomic mass of magnesium = M

Atomic mass and relative abundance of = m₁ and p 

Atomic mass and relative abundance of = m₂ and q 

Atomic mass and relative abundance of = m₃ and r

Now, we know that:

M = 24.3

m₁ = 24 and p = 78.7%

m₂ = 25

m₃ = 26 and r = 11.2%

The average atomic mass of magnesium can be found as:

Hence, 10% of the magnesium existing in nature is . 

Radioactivity

Radioactivity is the phenomenon wherein the nucleus of an unstable atom loses energy after emitting ionized particles or radiations. In this process, the original atom transforms into a new stable atom. 

In nature, there are many unstable isotopes that emit radiations. The common types of radiations are:

1. Alpha radiation 

   1. It consists of a stream of positively charged particles that are generally helium ions, having atomic mass as 4 and carrying a +2 charge. 
   2. It results in a decrease in atomic number by 2 and a decrease in mass number
      by 4.

2. Beta radiation

   1. It consists of a stream of electrons. So, it is negatively charged.
   2. It results in an increase in atomic number by 1, but mass number remains unchanged.

3. Gamma radiation

   1. It is a photon (packet of light) with very high energy.
   2. It does not result in a change in either atomic number or mass number.