Structure of the Atom

The present state of atomic theory is characterized by the fact that we not only believe the existence of atoms to be proved beyond a doubt, but also we even believe that we have an intimate knowledge of the constituents of the individual atoms.

• Niels Bohr

1.0Introduction

In the study of the chemical classification of matter, we have learnt that matter is made up of atoms and molecules. These are in fact, the building blocks of matter. However, the major breakthrough regarding the composition of matter was made by Dalton. Now a question arises, whether atom is really the smallest portion of matter as has been stated by Dalton. Moreover, in what respect the atoms of different elements differ from each other. This indeed posed a major challenge to the scientists of that era. This mystery was solved towards the end of the 19th century by scientists like J. J. Thomson, Goldstein, Rutherford, Chadwick, Bohr and many more.

The surface of the copper penny is made of copper atoms represented as they would be seen through the lens of a very powerful electronic microscope.

As a result of researches carried by them, it was established that atom is not as simple as was predicted by

Dalton.

It is made up of a number of particles out of which the most important are electrons, protons and neutrons.

This was just the beginning. The present picture of the atom is altogether different.

The matter which we suppose to be the main constituent of the universe is built out of small selfcontained buildingblocks, the chemical atoms.

In the present chapter, we shall mainly focus on the composition of atom and how sub-atomic particles (like electron, proton and neutron) were discovered and how they are arranged within the atom i.e. to study the "model of atom".

2.0Discovery of Electrons

The electrons were noticed and identified for the first time by J. J. Thomson in 1897 in an experiment which was performed in discharge tube. The discharge tube was filled with air (or any other gas) under study and the following observations were recorded on the basis of the experiments performed. The emission of light stopped at 10000 volt when the pressure inside 0.0001 mm Hg the tube was further reduced to about 0.01 mm mercury. The walls of the discharge tube opposite to the cathode started glowing with a greenish light called fluorescence. At high voltage and low pressure the bonds between the molecules break down and the bonds between atoms also break down. Charged particles are separated out from it. On further investigations, it was found that the fluorescence is due to the bombardment of the walls of the tube by the rays emitted from the cathode. These were known as cathode rays. Thomson repeated the experiment by replacing air with a number of gases and by also changing the metal which constituted cathode.

Properties of cathode rays

• These rays travels in straight line away from cathode with very high velocity nearly equal to the speed of light.
• A shadow of metallic object placed in the path is casted on the wall opposite to cathode.
  
• Cathode rays cause green fluorescence on glass surface, i.e the glass surface on which the cathode rays strike show a coloured shine.
• Cathode rays consist of matter particles and possess energy by the virtue of their mass and velocity. Cathode rays set a paddle wheel into motion when it is placed in the path of these rays. This is due to the impact of the particles of the cathode rays on the blades of the paddle wheel.
  
  Cathode ray particles strike the blades of the paddle wheel and set it into motion.
• These rays are deflected by the electric or magnetic fields when the rays are passed between two electrically charged plates, these are deflected towards the positively charged plates. It shows that cathode rays carry negative charge. These particles carrying negative charge were called negatrons by Thomson. The name negatron was changed to "Electron" by Stoney.
  
  Deflection of cathode ray towards positive plate of the electric field
  
  Deflection of cathode ray in the magnetic field
• These rays produce heat energy when they collide with the matter. It shows that cathode rays possess kinetic energy which is converted into heat energy when stopped by matter.
• These rays affect the photographic plate.
• Cathode rays can penetrate the thin foil of solid materials.
• Cathode rays can ionize the gases through which they pass.
• The nature of cathode rays is independent of

(a) The nature of cathode. (b) The gas in discharge tube.

Determination of charge and mass on electron

The charge on the electron was determined by R. A. Millikan with the help of oil drop experiment. The value was found to be $1.6\times 10^{-19}$ coulomb (C). The mass of the electron was found to be $9.1\times 10^{-31}\mathrm{kg}$. This was regarded as negligible and was actually nearly $\frac{1}{1840}$ times the mass of hydrogen atom. The charge/mass ratio is also called e/m ratio.

The electron determine the chemical properties of an element and arrangement of electrons strongly influence an atom's magnetic properties. $\frac{\mathrm{e}}{\mathrm{m}}=\frac{1.6\times 10^{-19}\mathrm{C}}{9.1\times 10^{-31}\mathrm{kg}}=1.76\times 10^{11}\mathrm{C}/\mathrm{kg}$.

3.0Discovery of protons

Anode Rays or Canal Rays

It has been established that electron is a negatively charged particle present in all the atoms. As an atom is electrically neutral, there must be some positively charged particle present in the atom to neutralize the negative charge of the electrons. It has been confirmed by experiment done by "Goldstein". He discovered the existence of a new type of rays in discharge tube.

He carried out experiment in discharge tube containing perforated cathode and it was placed somewhere in the centre of the tube. These rays moved towards cathode and passed through the perforation in the cathode. Initially these rays were called canal rays because they passed through the canals or holes of the cathode. The rays were named anode rays as they were directed away from the anode. Anode rays consist of positively charged particles, therefore, these were also named as positive rays.

Characteristics of anode rays

• Anode rays travel in straight lines. These rays rotate the light paddle wheel placed in their path. Anode rays are deflected by magnetic or electric field. In electric field they deflect towards negatively charged plate. This indicates that these rays are positively charged.
• The anode rays affect photographic plate.
• The nature of anode rays depends upon the type of gas used. The charge (e) to mass (m) ratio (e/m) of anode rays particle is different for different gases. The value of e/m is maximum for hydrogen gas.
• The positive rays obtained from hydrogen are made up of some type of positive particles. These particles are known as protons.
• Mass of proton is $1.67\times 10^{-24}\mathrm{g}$. or $1.67\times 10^{-27}\mathrm{kg}$ Mass of 1 H atom $=1837\times$ (Mass of $1{\mathrm{e}}^{-}$)

Explanation for the origin of anode rays

In case of hydrogen gas when enclosed in the discharge tube, these positively charged residues were called protons. Actually the name proton arises from the word protium which represents a hydrogen atom. It is the positively charged residue left when an electron is knocked out of the hydrogen atom or protium. Anode rays consists of positively charged residual particles (protons) are formed by the removal of one or more electrons from the gaseous atoms.

4.0Structure of atom

The discovery of electrons and protons was made by Thomson and Goldstein along with their co-workers brought a revolution in the field of science, particularly chemistry. The Dalton's concept of indivisibility of atom was discarded. Now the major challenge before the scientists was to determine the relative positions of these fundamental particles in an atom.

(i) Thomson's model of an atom The first model of the atom was first of all proposed by Sir J.J. Thomson.

Thomson proposed that an atom consist of a uniform sphere of positive charge (protons) in which the negatively charged (electrons) are distributed more or less uniformly.

The positive charges due to protons and negative charges due to electrons balance each other. As a result, an atom as a whole is electrically neutral or it has no net charge.

Limitation

Thomson's model of an atom could only explain the main characteristics of the atom at that time.

Drawbacks of the Thomson model

(i) An important drawback of this model is that the mass of the atoms is considered to be evenly spread over that atom. (ii) It is static model. It does not explain the movement of electron. (iii) It did not have any experimental support. However, his prediction that an atom is electrically neutral and has no net charge is

One of the early model of the atom was the plum pudding model, in which the electrons were pictured as embedded in a positively charged spherical cloud, much as raisins are distributed in an oldfashioned plum pudding. still accepted. This was indeed a big contribution towards the structure of the atom.

Rutherford's model of an atom (Discovery of nucleus)

In 1911, scientist "Ernest Rutherford" gave a new picture for the structure of atom by his $\alpha$-particle scattering experiment and proposed the structure of atom.

Rutherford performed the famous $\alpha$-particle scattering experiment in order to know the relative positions of electrons and protons in an atom.

Rutherford Experiment

One of the early model of the atom was the plum pudding model, in which the electrons were pictured as embedded in a positively charged spherical cloud, much as raisins are distributed in an old-fashioned plum pudding. $\alpha$ particle are charged particles having 2 units of positive charge and 4 units of mass, that is $\alpha$-particles $\left(2{\mathrm{He}}^4\right)$ are doubly charged helium atoms $\left({\mathrm{He}}^{+2}\right)$. He selected a thin foil of a heavy gold metal because it was only 1000 atoms thick. He bombarded the same with high speed alpha ( $\alpha$ ) particles.

Observation from experiment

When fast moving alpha particles strike very thin gold foil in vaccum, it is found that:

Scattering of alpha particles by the atoms of a gold foil Note : Rutherford selected a gold foil because he wanted as thin a layer as possible.

• Most of the fast moving $\alpha$-particles passed straight through the gold foil, without any deflection from their original path.
• Some of the $\alpha$-particles were deflected by the foil by small angles.
• A very few alpha particles got completely rebound on hitting the gold foil and turned back on their path.

Conclusion of Rutherford experiment

Rutherford concluded from the $\alpha$-particle scattering experiment that

• Most of the space inside the atom is empty because most of the $\alpha$ particles passed through the gold foil without getting deflected.
• As a few alpha particles suffered minor deflections and a very few major deflections, this means that these must have met with some obstruction in their path.
• This obstruction must be (a) Very small: Only a few particles were obstructed by it. (b) Each alpha particle has 4 u mass and is quite heavy. It could easily pass through a light obstruction by pushing it aside.
  
  Ernest Rutherford (c) Positively charged: Alpha particles have positive charge. Since they were repelled or deflected back, the obstruction must also carry some charge i.e. positive charge, (similar charged particles always repel each other). Rutherford regarded this very small, dense and positively charged obstruction in an atom as nucleus (means centre). Therefore, Rutherford predicted that all the protons present in an atom are present in this small space i.e. nucleus. The electrons with negligible mass and negative charge were supposed to be present in this portion around the nucleus known as extra nuclear portion. Rutherford model is also called as planetary model of atom because it is similar to that of the solar system. As in solar system, different planets are revolving around the sun, similarly, in an atom the electrons are revolving around the nucleus. Thus, electrons were called planetary electrons.

On the basis of this experiment, Rutherford put forward the nuclear model of an atom, which has the following features. (i) An atom consists of two parts. These are nucleus and extra nuclear portion. (ii) Nucleus is present in the centre of the atom and is surrounded by extra nuclear portion. (iii) The size of the nucleus is very small as compared to that of the atom.

Note: Electrons in an atom are not stationary. Only the energy of the electrons is stationary in an energy level.

Size of atom $=10^{-8}\mathrm{cm}$ Size of nucleus $=10^{-13}\mathrm{cm}$ $\frac{\text{ atom size }}{\text{ nucleus size }}=\frac{10^{-8}}{10^{-13}}=10^5$ atom size $=10^5\times$ (nucleus size) (iv) The mass of the atom is mainly of the nucleus. All the protons are present in the nucleus. (v) The positive charge on the nucleus is because protons are present. (vi) All the electrons are present in the extra nuclear space around the nucleus. (vii) The total positive charge of the nucleus due to the presence of protons is the same as that of the electrons present in the extra nuclear space. Therefore, the atom as a whole is electrically neutral. (viii) Electrons present in the extra nuclear portion are not stationary. Those which revolve around the nucleus at high speed follow a circular path.

A nuclear atom viewed in cross section. (The symbol $\sim$ means approximately.) This drawing does not show the actual scale. The nucleus is actually much smaller than the atom itself.

Drawback of Rutherford model of atom

According to Maxwell's Radiation theory, any charged particle moving in circular orbits with fast speed emit energy in the form of radiations. Since electrons are also fast-moving charged particles, they must release energy continuously in the form of radiations. With reduced energy, they must be drawn closer to the nucleus. Since the loss of energy is a continuous process, the electrons must come closer and closer to the nucleus of the atom.

Ultimately, they must fall into the nucleus or become its part. In other words, the atom should collapse. However, this actually does not happen and the atom is quite stable.

(iii) Bohr model of an atom The limitation of Rutherford model of atom was explained by Neils Bohr. He stated that although the fast-moving electron is expected to lose energy as radiations, but it does not actually lose energy. In other words, the energy of the electron is stationary. It therefore, revolves in the same orbit without losing any energy.

According to Bohr's theory:

• Electrons revolve around the nucleus in well-defined orbits or shells. Each shell have a definite amount of energy associated with the electrons in it. Therefore, these shells are also called energy levels.
• The energy associated with the electrons in an orbit increase as the radius of the orbit increase. These shells are also known as $\mathrm{K},\mathrm{L},\mathrm{M},\mathrm{N}$. starting from the one closest to the nucleus.
  
  Neils Bohr
• An electron in a shell can move to a higher or lower energy shell by absorbing or releasing a fixed amount of energy.
  
  Different energy levels

The order of the energy of these energy shell is: $\mathrm{K}<\mathrm{L}<\mathrm{M}<\mathrm{N}<\mathrm{O}<\dots ..$ or $1<2<3<4<5<$...... While revolving in an orbit, the electron is not in a position to either lose or gain energy. In other words, its energy remains stationary. Therefore, these energy states for the electrons are also known as stationary states.

Electrons have the negligible mass as compared to protons and neutrons.

5.0Distribution of electrons in different orbits (shells)

We know that atoms of different elements differ in their atomic numbers and hence have different number of electron. These are distributed in the various energy shells which were given by Bohr i.e. K, L, M, N .... etc. The distribution or arrangement of the electrons in the different shells of the atom is called the electronic configuration of the element. It is based on certain guidelines or rules given by Bohr and Bury. This is known as Bohr-Bury Scheme. According to this scheme, (i) The maximum number of electrons which can be present in a particular energy shell of an atom is given by $2n\mathbf{n}$. Here ' $n$ ' is the number of the energy shell or energy level. Hence the maximum number of electrons in different shells are as follows. First orbit (K-shell) $=2\times 1^2=2$ Second orbit (L-shell) $=2\times 2^2=8$ Third orbit (M-shell) $=2\times 3^2=18$ Fourth orbit ( N -shell) $=2\times 4^2=32$ (ii) The maximum number of electrons that can be accommodated in the outermost orbit is 8 . The outermost energy shell in an atom cannot have more than eight electrons even if it has a capacity to take up more electrons. (iii) Electrons do not enter into a new shell unless the inner shells are completely filled. In other words, the shells are filled in a step-wise manner.

6.0Discovery of neutrons ( n )

In 1920, Rutherford found that except for hydrogen atom, the atomic masses of no other elements could be explained by electrons and protons only.

According to ${}_1{\mathrm{H}}^1$ mass of protium will depend on protons but if we see ${}_2{\mathrm{He}}^4$ the mass of He is not equal to number of P. So, definitely an another particle is present, then neutron is discovered.

To solve this problem, Rutherford predicted the presence of some neutral particles in the nucleus of the atom. Its mass must be same as that of proton but it is expected to have no charge.

The material particles were actually discovered by Chadwick in 1932. He bombarded the nucleus of some light elements like beryllium and boron with fast moving $\alpha$-particles. He found that some neutral particles were ejected from the nucleus. Each of these particles carried no charge but had a mass equal to that of proton. This particle was named as neutron.

Charge and mass of subatomic particles

7.0Atomic number and mass number

Atomic number (Z)

Protons are present in the nucleus of an atom. It is the number of protons of an atom, which determines its atomic number. It is denoted by ' $Z$ '. Thus, atomic number $=$ No. of protons = No. of electrons

Therefore, the atomic number is defined as the total number of protons present in the nucleus of an atom.

For example, ${}_6{\mathrm{C}}^{12}$ means atomic number of carbon is 6 . Nucleus of carbon has 6 protons. Nucleus of carbon has 6unit positive charge. There are also 6 electrons, revolving around the nucleus of carbon.

For example, ${}_{12}^{24}\mathrm{Mg}$

Mass number is merely a number but the atomic mass is the relative mass of an atom and it has specific units. It is expressed in a.m.u. or $u$.

The mass number of magnesium is 24 . The total number of protons and neutrons in the nucleus of magnesium is 24 , number of protons is 12 . Number of neutrons is $=24-12=12$.

Mass number (A)

The mass number of an atom is defined as the sum of the total number of protons and neutrons present in the nucleus of an atom. For example, mass number of carbon is 12 because it has 6 proton and 6 neutron that is $6+6=12$ Mass number: Number of protons + number of neutrons. The notation for an atom: The atomic number, mass number and symbol of the element are to be written as follows.

Atomic parameter of first 20 elements

8.0Valency of elements

The atoms of noble gases are regarded as stable atoms because they have maximum number of permissible electrons based on certain rules. What about the other elements? Their atoms have incomplete outermost shells and have urge or desire to complete these. For this, they approach or invite the atoms of other elements to combine with them. The purpose behind these combinations is to have either 2 or 8 electrons in their outermost shell. The combining capacity of an element with atoms of other elements in order to acquire 8 electrons ( 2 in some exceptional cases) is called the valency of an element. The valency of the element can be calculated as, (a) The atoms having 1, 2, 3 or 4 valence electrons normally lose these to the combining atoms and they show valencies of $1,2,3$ or 4 respectively. (b) The atoms having 5, 6 or 7 valence electrons generally take up 3, 2 or 1 electrons respectively from the atoms of the other elements and show valencies of 3,2 or 1 respectively.

Formation of ions

The species carrying charge are called ions. The charge on the ions is equal to the number of electrons lost or gained by the atom and represents the valency of an element.

Positive ions or cations:

Negative ions or anions:

9.0Isotopes

Isotopes are atoms of the same element, having the same atomic number but different mass number. For example, isotopes of hydrogen atom, namely protium $\left({1H}^1\right)$, deuterium $\left(1{\mathbf{H}}^2\right)$, tritium ( ${1H}^3$ ). The atomic number of each one is 1 , but the mass number is 1,2 , and 3 respectively.

10.0Examples of isotopes

Applications of Isotopes

Since the chemical properties of all isotopes of an element are the same but some isotopes have special properties which find them useful in various fields. Some of them are

• Uranium $\left({}_{92}^{235}\mathrm{U}\right)$ is used as a fuel in nuclear reactors.
• Cobalt $\left({}_{27}^{60}\mathrm{Co}\right)$ is used in the treatment of cancer.
• Iodine $\left({}_{53}^{131}\mathrm{I}\right)$ is used in the treatment of goitre.
• Sodium $\left({}_{11}^{24}\mathrm{Na}\right)$ is used for differentiating cancerous tissues from the normal tissues.
• Carbon $\left({}_6^{14}\mathrm{C}\right)$ is used in dating of fossil samples known as carbon dating.
• Phosphorus $\left({}_{15}^{32}\mathrm{P}\right)$ is used in the treatment of blood cancer.

11.0Calculation of average atomic mass of an element from the atomic mass of its isotopes

The atomic mass of an element is the weighted arithmetic mean of the atomic masses of its isotopes present in the sample of the element.

Fractional atomic masses indicate average mass of all the isotopes taken together with respect to their relative abundance in environment. Let us consider a sample of an element X containing its two isotopes ${\mathrm{X}}_1$ and ${\mathrm{X}}_2$.

Average atomic mass of the element

$\mathrm{X}=\frac{\%\text{ of }{\mathrm{X}}_1\times \text{ Atomic mass of }{\mathrm{X}}_1+\%\text{ of }{\mathrm{X}}_2\times \text{ Atomic mass of }{\mathrm{X}}_2}{100}$ For example, this method can be illustrated by taking the case of chlorine. The two isotopes of chlorine, ${}_{17}^{35}\mathrm{Cl}$ and ${}_{17}^{37}\mathrm{Cl}$ occur in the ratio $3:1$. Then, Atomic mass of chlorine $=\frac{(35u\times 3)+(37u\times 1)}{3+1}=\frac{105u+37u}{4}=\frac{142u}{4}=35.5u$

Isobars

Atoms of different elements with different atomic number which have the same mass number, are known as isobars. Chemical properties of isobars are different.

For example, (1) Calcium and argon $20{\mathrm{Ca}}^{40}{}_{18}{\mathrm{Ar}}^{40}$

The atomic number of argon is 18 , calcium is 20 but the mass number of these elements is same. (2) ${}_6^{14}\mathrm{C}$ and ${}_7^{14}\mathrm{N}$

12.0Basic terminology

• $\alpha$ particle - Doubly charged helium atoms $\left({\mathrm{He}}^{+2}\right)$.
• Atomic nucleus - It is the small, dense region consisting of protons and neutrons at the center of an atom.
• Shell (orbit) - The space available around the nucleus on which electrons revolve.
• Atomic number - Number of protons = number of electrons.
• Mass number - Number of protons + number of neutrons.
• Isotope - The atoms of the same element, having the same atomic number but different mass number.
• Isobars - Isobars are atoms of different elements which have a different atomic number but the same mass number.
• Valency - Combining capacity of the element.