ICSE Class 9 Chemistry • Chapter 4 • Comprehensive Chapter Notes

INTRODUCTION

• Maharshi Kanada (6th Century B.C., India): Introduced the idea of the smallest unit of matter called paramanu (now called atoms). Paramanus combine to form anu (molecule).
• Democritus (460-370 B.C., Greek): Called the paramanu an "atom", meaning indivisible.
• John Dalton (1808): Proposed the first scientific theory, considering atoms as indivisible, fundamental building blocks of matter.

The main postulates of Dalton's atomic theory are:

1. Matter consists of very small and indivisible particles called atoms.
2. Atoms can neither be created nor be destroyed.
3. The atoms of an element are alike in all respects but they differ from the atoms of other elements.
4. Atoms of an element combine in small numbers to form molecules.
5. Atoms of one element combine with atoms of another element in a simple ratio to form molecules of compounds.
6. Atoms are the smallest units of matter that can take part in a chemical reaction.

Subatomic Particles:

• Gases are normally poor conductors of electricity, but studies of static electricity showed flow of charged particles.
• J.J. Thomson: Discovered negatively charged particles called electrons.
• E. Goldstein: Discovered positively charged particles called protons.
• James Chadwick: Discovered neutral particles called neutrons.
• Electrons, protons, and neutrons are the fundamental particles of an atom.

4.1 DEFINITION OF AN ELEMENT

An element is a substance which is made up of only one type of atoms.

For example: Carbon is an element because it cannot be split up into two or more simpler substances by ordinary methods like heating, breaking or passing electricity.

4.2 DEFINITION OF AN ATOM

Example: Take a small piece of zinc and grind it into smaller pieces. All these pieces show properties of zinc. On grinding them further they break up into very fine particles which still show the properties of zinc. But, there comes a stage when the particles cannot be further subdivided into particles exhibiting properties of zinc. These indivisible particles are the atoms of zinc.

"In other words an atom is the smallest possible unit of an element".

4.3 CONSTITUENTS OF AN ATOM

Discovery of electrons

William Crooke, a British scientist, noted that gases are ordinarily poor conductors of electricity. However, when a high voltage (10,000 volts) charge from an induction coil is applied to tubes filled with gases at very low pressure ($0.01\text{ mm Hg}$), the gases become good conductors of electricity and begin to flow from cathode to anode in the form of rays. Since these rays originate from the negative plate, i.e. the cathode and travel from the cathode towards the anode, they are called cathode rays.

Later J.J. Thomson studied the characteristics and the constituents of cathode rays. The apparatus used by him is called a discharge tube or a cathode ray tube shown in Fig. 4.1.

A discharge tube is a hard glass tube fitted with two metal plates known as electrodes, one of which is connected to the positive terminal of battery and is called anode (positive electrode) while the other electrode which is connected to the negative terminal of the battery is called cathode (negative electrode). It has a side tube through which gas can be pumped out by using a vacuum pump to create vacuum.

Properties of Cathode Rays

1. They travel from the cathode to the anode in straight lines.
2. They cause a greenish yellow fluorescence on a soda-glass screen placed in the tube.
3. They are affected by electric field, i.e. they are inflected towards the positive field and deflected from the negative field. This shows that they carry negative charge. (Fig. 4.2).

4. When a beam of cathode rays is made to fall upon hard metallic targets like tungsten, X-rays are produced.
5. They penetrate through matter.
6. They cause ionization of the gas through which they pass.
7. The ratio of the charge ($e$) to mass ($m$) of the particles constituting cathode rays remains the same ($e/m = 1.76 \times 10^{11}\text{ coulomb/kg}$) irrespective of the nature of the gas taken and of the metal forming the cathode.
8. They produce the shadow of an opaque object placed in their path and make a light paddle wheel rotate.

Thomson concluded that:

(i) cathode rays consist of negatively charged particles, now called electrons.
[The name 'electron,' meaning 'atom of negative electricity,' was given by Johnson Stoney].

(ii) these negatively-charged particles are an integral part of all atoms.

(iii) electrons have both definite mass and definite electric charge, both of which are independent of the nature of the gas in the discharge tube.

Properties of Electrons

1. Electrons from all sources are alike, having identical mass.
2. They are a constituent part of all atoms.
3. The mass of an electron is $1/1837$ the mass of a hydrogen atom (or $9.108 \times 10^{-31}\text{ kg}$).
4. An electron carries unit negative charge of magnitude $-1.602 \times 10^{-19}\text{ coulombs}$.
5. The electron is extremely small; its radius is less than $1 \times 10^{-15}\text{ m}$.

Atoms are found to be electrically neutral, so they must contain, in addition, particles that are positively charged, such that the total negative charge of the electrons is equal to the total positive charge. This realization led to the discovery of positively charged subatomic particles protons.

4.4 DISCOVERY OF PROTONS

Goldstein noticed another set of rays travelling in a direction opposite to that of the cathode rays, i.e. from anode towards cathode, when a perforated cathode was used in the above discharge tube (shown in Fig. 4.3). He called these rays as canal rays since these rays passed through holes or 'canals' in the cathode. These rays were named as positive rays or anode rays.

Properties of Anode Rays

1. Anode rays travel in a straight line.
2. They consist of minute material particles and hence produce mechanical effects.
3. They are made up of positively charged particles.
4. Positive rays are deflected by electric and magnetic fields but in a direction opposite to that of the cathode rays. This means that these rays consist of positively charged particles called the protons.
5. These rays produce fluorescence on a zinc sulphide screen.
6. Their $e/m$ i.e. charge to mass ratio, differs from gas to gas. Its value is much less than that of an electron and is maximum when hydrogen is taken in the discharge tube.

Properties of Protons

1. A proton possesses a unit positive charge +1 (plus one) of the value $1.602 \times 10^{-19}\text{ coulombs}$.
2. Its mass is the same as that of a hydrogen atom (1 a.m.u., i.e., 1837 times the mass of an electron, which is $1.672 \times 10^{-24}\text{ g}$).
3. The proton resides in the central part of an atom, i.e. in the nucleus.

After the discovery of electrons and protons J.J. Thomson proposed 'plum pudding' model of the atom.

According to this model:

1. An atom is considered to be a sphere of uniform positive charge and electrons are embedded into it.
2. The total positive charge is equal to the total negative charge so that an atom as a whole is electrically neutral.
3. The mass of an atom is considered to be uniformly distributed.

4.5 DISCOVERY OF NUCLEUS

In 1911, Lord Rutherford, a scientist from New Zealand, directed a stream of alpha-particles (An $\alpha$-particle is a doubly charged helium ion $\text{He}^{2+}$, containing two protons and two neutrons. It is formed by removing 2 electrons from helium atom) towards a very thin (one millionth of a centimetre) gold foil. He selected a gold foil because he wanted as thin a layer as possible and gold is the most malleable metal.

He observed that:

1. Most of the alpha particles passed straight through the foil.
2. Some alpha particles were slightly deflected from their straight path.
3. Very few (nearly one in ten thousand) alpha particles were either deflected by very large angles or completely bounced back.

Later, Rutherford generalised these results of alpha particles scattering experiment and suggested a model of the atom that is known as Rutherford's Atomic Model.

4.5.1 Rutherford's atomic model

According to this model:

1. The atom contains a large empty space (extra-nuclear space). This is why most of the $\alpha$-particles pass through the metal foil without deviating from their path.
2. There is a positively charged mass at the centre of the atom, known as the nucleus, in which the entire mass of the atom is concentrated. Therefore, the nucleus is the densest part of the atom. The electrically positive nature of the nucleus is supported by the fact that the positively charged $\alpha$-particles approaching the nucleus get deflected. This proved that the nucleus contained protons.
3. The size of the nucleus is very small compared to the size of an atom. This is proved by the fact that most of the $\alpha$-particles pass straight through the gold foil.
4. The electrons revolve around the nucleus at large distances from each other and from nucleus (in circular orbits found in the empty space of the atom).
5. An atom as a whole is electrically neutral, because the total positive charge of the nucleus is balanced by the total negative charge of the electrons, i.e., the number of protons and electrons in an atom are equal.
6. Electrons revolve around the nucleus in close circular paths (ORBITS). The force of attraction between the negatively charged electrons and the positively charged nucleus is counter balanced by the centrifugal force acquired by the revolving electrons.

Rutherford's model of atomic structure is similar to the structure of the solar system. Just as in the solar system, the Sun is at the centre (having the maximum mass) and the planets revolve around it, similarly in an atom the nucleus contains the main mass and the electrons revolve around it in orbits or shells.

4.5.2 Drawback of Rutherford's atomic model

The comparison of electrons with planets in the solar system is the main drawback of Rutherford's atomic model. According to the classical laws of mechanics and electrodynamics, if an electrically charged particle is in motion, it inevitably radiates energy. Thus, an electron, when moving round the nucleus continually, should radiate energy, i.e. loses energy. As a result, it should be gradually pulled towards the nucleus and end up colliding with it. This should result in the total collapse of the atom (Fig. 4.7).

If it was so, the atom should be highly unstable and hence matter would not exist in the form that we know. However, we know that an atom is structurally stable.

Thus Rutherford's model failed to explain the stability of an atom.

4.6 BOHR'S ATOMIC MODEL

In 1913, Niels Bohr, a Danish physicist, explained the causes of the stability of the atom in a different manner.

Main postulates of Bohr's atomic model

1. The electrons revolving around the nucleus are confined to certain fixed orbits called shells or energy levels, each of which is associated with a fixed amount of energy.
2. While it is revolving around the nucleus in an orbit, an electron neither loses nor gains energy.
3. An electron revolving in a particular orbit, on gaining a certain amount of energy, jumps to the next orbit and vice versa. Since each orbit is associated with a fixed amount of energy, Bohr called it an energy level.

For convenience, these energy levels are labelled K, L, M, N or I, II, III, IV, etc. The orbit closest to the nucleus is the K shell. It has the least amount of energy and the electrons present in it are called K electrons, and so on with the successive shells and their electrons.

4.7 DISCOVERY OF NEUTRONS

By now we know that an atom contains electrons and protons, and that the atomic mass of an electron is negligible. Therefore, an atom of helium, which contains 2 protons should have a

$\text{mass } = 2 \times 1\text{ a.m.u.} = 2\text{ a.m.u.}$

But the atomic mass of a helium atom was found to be approximately $4.0\text{ a.m.u.}$ It was, therefore, proposed that, in the nucleus of an atom, there must be another particle. This particle should not possess any electrical charge and must be equal in mass to the proton.

In 1932 Chadwick discovered these particles by bombarding light nuclei like beryllium with alpha ($\alpha$) particles i.e. helium nuclei.

${}_{4}\text{Be}^{9} + {}_{2}\text{He}^{4} (\alpha) \rightarrow {}_{6}\text{C}^{12} + {}_{0}n^{1} \text{ (neutron)}$

These particles are found to be neutral, so named neutrons.

Properties of Neutrons

1. This particle was not found to be deflected by any magnetic or electric field, proving that it is electrically neutral.
2. Its mass is equal to $1.676 \times 10^{-24}\text{ g}$ (1 amu).

Table 4.1 : Comparison of Dalton's atomic theory with the modern atomic theory

Dalton's atomic theory	Modern atomic theory
(i) Atoms are indivisible particles.	(i) Atoms are divisible into sub-atomic particles like protons, electrons and neutrons.
(ii) Atoms can neither be created nor destroyed.	(ii) Atoms can be created and destroyed by nuclear fusion and fission.
(iii) The atoms of an element are alike in all respects, but they differ from the atoms of other elements.	(iii) The atoms of an element may not be alike in all respects, as is seen in the case of isotopes.
(iv) Atoms follow the laws of chemical combination to form compounds.	(iv) In the formation of organic compounds, the laws of chemical combination are not always followed.

* An alpha particle consists of two protons and two neutrons. It is same as Helium atom that has lost its both electrons.

4.8 ATOM — ITS STRUCTURE

The atom is built up of a number of sub-atomic particles*. The three sub-atomic particles of great importance in understanding the structure of an atom are electrons, protons and neutrons, the properties of which are given in Table 4.2.

Table 4.2 Properties of sub-atomic particles

Particle	Symbol	Charge (electronic unit)	mass (amu)	mass (grams)
1. Electron	${}_{-1}e^{0}$ or $e^-$	$-1$	$1/1840$	$9.1 \times 10^{-28}\text{ g}$
2. Proton	${}_{1}^{1}\text{H}$ or $p^+$	$+1$	$1$	$1.6 \times 10^{-24}\text{ g}$
3. Neutron	${}_{0}^{1}n$ or $n$	$0$	$1$	$1.6 \times 10^{-24}\text{ g}$

There are two structural parts of an atom, the nucleus and the empty space in which there are imaginary paths called orbits.

Nucleus: The protons and neutrons [collectively called nucleons] are found in the central part or nucleus of the atom.

Orbits: It is the imaginary path where electrons revolve around the nucleus of the atom.

* Modern theories about the atom include the existence of over 200 subatomic particles such as positron, neutrino, etc.

4.9 ATOMIC NUMBER [Z]

• An atom of an element has its own characteristic number of protons in its nucleus, which distinguishes it from the atoms of other elements.
• This characteristic number (number of protons) is called the atomic number of the element.
• Atomic number is denoted by $Z$.
• The atomic number is, therefore, equal to the number of electrons in the atom of an element as well, since in a neutral atom,
  no. of protons = no. of electrons.
• Number of protons gives the total positive charge present in the nucleus of an atom.

Mass Number [A]

• Mass number of an element is the total number of protons and neutrons present in its nucleus. Mass number is denoted by A.
• Mass number, however, is a whole number approximation of the atomic mass calculated in atomic mass units.

Symbolic representation of Element	Atomic Number [Z]	Mass Number [A]
${}_{Z}^{A}\text{X}$ X denotes the symbol of the element Superscript A = Mass number Subscript Z = Atomic number	Atomic number [Z] = no. of protons [p] = no. of electrons [e] $\therefore Z = p = e$	Mass number [A] = no. of neutrons [n] + no. of protons [p] $\therefore A = n + p$ $[\because p = Z] \therefore \text{no. of neutrons } [n] = A - Z$
Example : ${}_{17}^{35}\text{Cl}$      Atomic number [Z] = p = e = 17      Mass number [A] = n + p = 35      No. of neutrons [n] = A - Z = 18

4.10 DISTRIBUTION OF ELECTRONS IN THE ORBITS — BOHR-BURY SCHEME

Distribution of electrons into different shells (orbits) of an atom was suggested by Bohr and Bury.

The following rules are followed for writing the number of electrons in different energy levels or shells.

1. The maximum possible number of electrons in a particular shell is given by the formula $2n^2$, where $n$ denotes the serial number of that shell.
   • First or K shell has ($2 \times 1^2$) = 2 electrons
   • Second or L shell has ($2 \times 2^2$) = 8 electrons
   • Third or M shell has ($2 \times 3^2$) = 18 electrons
   • Fourth or N shell has ($2 \times 4^2$) = 32 electrons and so on.
2. Electrons are not accomodated in a given shell, unless the inner shells are filled. In other words the shells are filled in stepwise manner.
3. The outermost shell of an atom cannot accommodate more than 8 electrons, even if it has the capacity to accommodate more electrons. For example electronic arrangement in Calcium having 20 electrons is:
   K (2)    L (8)    M (8)    N (2)
   The maximum capacity of M shell is 18 electrons but only 8 electrons are filled in it. This is due to the fact that the elements become stable only after acquiring 8 electrons in the outermost shell.
4. The outermost shell of a chemically stable atom can accommodate a maximum of $8$ electrons, except Helium atom, which has only one shell, and can thus accommodate a maximum of 2 electrons only.

Structure of Magnesium atom

In lighter elements, up to argon (atomic number 18), each inner shell is completely filled before any electron can occupy an outer shell. However, in elements heavier than argon, the situation changes. Although the third shell can accommodate up to 18 electrons, yet the fourth shell begins to be filled after it has only 8 electrons.

For example, potassium (atomic number 19) has 19 electrons. Its electronic configuration is:

K (2)    L (8)    M (8)    N (1)