Structure of the Atom — Study Notes

1. Introduction

In the previous chapter we learned that matter is made of atoms, and that atoms of elements take part in chemical reactions. But what is an atom made of? The answer to this question was discovered over many years through the experiments of scientists such as J.J. Thomson, Ernest Rutherford and Niels Bohr. In this chapter we study the discovery of subatomic particles, the models proposed to explain the structure of the atom, and the concepts of atomic number, mass number and isotopes.

An atom is electrically neutral overall, yet it contains positively charged protons and negatively charged electrons. How are these particles arranged inside the atom? How is the mass of the atom distributed? These questions led to the development of atomic models, each improving upon the earlier ones.

We will also study how electrons are distributed in different shells, how the chemical behaviour of an atom is related to its electrons, and why atoms of the same element can differ in mass.

2. Charged Particles in Matter

When certain materials are rubbed, they acquire a charge. For example, rubbing a glass rod with silk makes it positively charged, and rubbing an ebonite rod with fur makes it negatively charged. This shows that matter contains charged particles.

Experiments with the discharge of electricity through gases led to the discovery of the electron. When a high voltage is applied across a discharge tube, cathode rays are produced which travel from the cathode to the anode. In 1897, J.J. Thomson showed that these cathode rays are made up of negatively charged particles, which he called electrons. The mass of an electron is 9.1 x 10^-31 kg, which is about 1/2000 times the mass of a hydrogen atom. The charge on an electron is -1.6 x 10^-19 coulomb.

The positively charged particle inside the atom is called the proton. It was discovered by E. Goldstein through canal rays (also called anode rays) in a discharge tube. The charge on a proton is +1.6 x 10^-19 coulomb, equal in magnitude but opposite in sign to the electron's charge. The mass of a proton is 1.672 x 10^-27 kg.

There is a third subatomic particle called the neutron, discovered by James Chadwick in 1932. The neutron has no charge and a mass slightly greater than that of a proton. Neutrons are present in the nucleus of all atoms except the common form of hydrogen.

3. Thomson's Model of the Atom

J.J. Thomson proposed the first model of the atom in 1904. According to Thomson, the atom is a sphere of positive charge in which negatively charged electrons are embedded, much like the seeds in a watermelon or the plums in a plum pudding. This model is therefore called the plum pudding model.

The positive charge and the electrons balance each other, making the atom electrically neutral. Thomson's model could explain that the atom is neutral, but it could not explain the results of later experiments on the scattering of alpha particles.

4. Rutherford's Alpha Particle Scattering Experiment

Ernest Rutherford bombarded a very thin gold foil with fast-moving alpha particles. Most of the alpha particles passed straight through the foil, a small number were deflected by small angles, and a very few (about 1 in 20000) bounced back at large angles.

From this experiment Rutherford made the following observations and conclusions:

1. Most of the space inside the atom is empty, because most alpha particles passed straight through.
2. A few particles were deflected, showing that the positive charge of the atom is concentrated in a very small volume called the nucleus.
3. Very few particles bounced back, showing that the nucleus is very dense and rigid.
4. The electrons revolve around the nucleus in circular orbits.

Rutherford's nuclear model of the atom: The atom has a central, positively charged nucleus which contains almost the entire mass of the atom. The electrons revolve around the nucleus in circular paths. The number of electrons is equal to the number of protons, so the atom is neutral.

Drawbacks of Rutherford's model: The revolution of electrons in circular orbits was expected to make them emit energy continuously and eventually fall into the nucleus, causing the atom to collapse. But atoms are stable. Rutherford's model could not explain the stability of the atom.

5. Bohr's Model of the Atom

To overcome the drawbacks of Rutherford's model, Niels Bohr proposed his model in 1913. The main postulates are:

1. Electrons revolve around the nucleus in fixed circular paths called orbits or shells or energy levels.
2. These shells are denoted by the letters K, L, M, N or the numbers 1, 2, 3, 4.
3. Each shell is associated with a fixed amount of energy; hence the shells are called energy levels.
4. As long as an electron stays in a particular shell, it neither gains nor loses energy. The energy of the electron increases as the shell is farther from the nucleus.
5. The energy of the electron changes only when it jumps from one shell to another.

The maximum number of electrons that can be accommodated in a shell is given by the formula 2n^2, where n is the shell number. For the first shell (K, n=1), the maximum is 2; for the second (L, n=2), it is 8; for the third (M, n=3), it is 18; and for the fourth (N, n=4), it is 32.

6. Electronic Configuration and Valency

The distribution of electrons in the various shells of an atom is called its electronic configuration. Some important rules for writing the configuration are:

1. The maximum number of electrons in a shell is 2n^2.
2. The outermost shell of an atom cannot have more than 8 electrons, and the shell next to the outermost shell cannot have more than 18 electrons.
3. Electrons fill the inner shells first.

For example, the electronic configuration of sodium (atomic number 11) is 2, 8, 1. That of oxygen (atomic number 8) is 2, 6, and of magnesium (atomic number 12) is 2, 8, 2.

The valency of an element is the number of electrons lost, gained or shared by its atom to complete the octet (8 electrons) in the outermost shell. Sodium has 1 electron in its outermost shell, so its valency is 1. Oxygen has 6 electrons in its outermost shell, so it needs 2 more to complete the octet, and its valency is 2. Noble gases like helium and neon have complete outermost shells, so their valency is zero.

7. Atomic Number and Mass Number

Atomic number (Z): The atomic number of an element is the number of protons present in the nucleus of its atom. Since the atom is electrically neutral, the atomic number is also equal to the number of electrons in the atom. For example, the atomic number of carbon is 6, which means a carbon atom has 6 protons and 6 electrons.

Mass number (A): The mass number of an atom is the total number of protons and neutrons present in its nucleus. For example, the mass number of a carbon atom with 6 protons and 6 neutrons is 12.

Mass number = Number of protons + Number of neutrons

The number of neutrons can be found by subtracting the atomic number from the mass number: Number of neutrons = Mass number - Atomic number.

We represent the atomic number and mass number of an element as follows: the element is written with the mass number as a superscript and the atomic number as a subscript. For example, carbon-12 is written as 12C with 12 as the mass number.

8. Isotopes and Isobars

Isotopes: Isotopes are atoms of the same element which have the same atomic number but different mass numbers. This means they have the same number of protons but different numbers of neutrons. For example, carbon has three isotopes: carbon-12 (6 protons, 6 neutrons), carbon-13 (6 protons, 7 neutrons) and carbon-14 (6 protons, 8 neutrons). Hydrogen has three isotopes: protium (1 proton, 0 neutrons), deuterium (1 proton, 1 neutron) and tritium (1 proton, 2 neutrons).

Isotopes of an element have the same chemical properties but different physical properties. Some applications of isotopes are: - An isotope of uranium (U-235) is used as a fuel in nuclear reactors. - The isotope of cobalt (Co-60) is used in the treatment of cancer. - An isotope of iodine (I-131) is used in the treatment of goitre.

Isobars: Isobars are atoms of different elements which have the same mass number but different atomic numbers. For example, calcium (atomic number 20, mass number 40) and argon (atomic number 18, mass number 40) are isobars.

Quick Revision Tables

Mind Map

Important Diagrams (SVG)

Diagram 1: Rutherford's Gold Foil Experiment

Diagram 2: Bohr's Model - Electronic Configuration of Sodium

Common Mistakes

1. Believing that the proton and electron have equal masses; the electron is about 2000 times lighter than a proton.
2. Thinking that electrons are found in the nucleus; electrons revolve around the nucleus in shells, while protons and neutrons are inside the nucleus.
3. Forgetting that the outermost shell of an atom cannot have more than 8 electrons even if 2n^2 gives a larger number.
4. Writing the electronic configuration without filling the inner shells first, such as writing the configuration of sodium incorrectly.
5. Confusing isotopes and isobars; isotopes have the same atomic number, isobars have the same mass number.
6. Believing that all elements have the same number of neutrons as protons; the number of neutrons = mass number - atomic number.
7. Saying that Rutherford's model explained the stability of the atom; in fact his model could not explain it, and Bohr's model did.

Exam Tips

1. Memorise the scientists and their contributions: Thomson (electron), Goldstein (proton), Chadwick (neutron), Rutherford (nucleus), Bohr (shells).
2. Learn the observations and conclusions of Rutherford's gold foil experiment, including why most alpha particles passed straight through.
3. Remember the maximum electron capacity formula 2n^2 and the limitation that the outermost shell holds at most 8 electrons.
4. Practise writing electronic configurations and finding valency from the outermost shell electrons.
5. Be able to compute the number of neutrons using mass number minus atomic number.
6. Define isotopes and isobars with examples: isotopes of carbon (C-12, C-14) and isobars like calcium and argon (both mass number 40).
7. Know one application each of isotopes: U-235 in nuclear reactors, Co-60 in cancer treatment, I-131 in treatment of goitre.

Conclusion

In this chapter we traced the discovery of the electron, proton and neutron, and saw how different models of the atom were proposed to explain experimental results. Thomson's plum pudding model gave way to Rutherford's nuclear model, which placed the mass and positive charge in a tiny nucleus, and finally to Bohr's model, which introduced fixed energy shells for electrons. We learned how to write electronic configurations, calculate valency, and define atomic number and mass number. The concepts of isotopes and isobars helped us understand why atoms of the same element can have different masses and how isotopes are used in medicine and nuclear energy. These ideas form the basis of the modern understanding of atomic structure and the periodic table.