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Structure of the Atom

Chapter: Structure of the Atom

Definition of an Atom

Different kinds of matter exist because of the various atoms present in them. Atom is the smallest unit which defines the chemical elements and their respective isotopes. Every substance which can be felt or touched is made up of atoms. Any solid, liquid or gas is made up of atoms.

Atoms are tiny in size and the size of an atom is measured in Pico meters (trillionths of a meter). A single strand of human hair is about one million carbon atoms wide.

Structure of an Atom

Electrons and protons are the two fundamental particles discovered inside an atom. This discovery led to the failure of Dalton’s atomic theory. Then we felt the need to know the arrangement of electrons and protons inside an atom. Many scientists proposed various atomic models to describe this arrangement.

J.J. Thomson was the first one to propose a model for the structure of an atom.

 

 

1. Thomson's Model of an Atom

J.J. Thomson is a very famous physicist. In 1897, he worked on cathode rays and discovered the electron and he stated that electrons are a component of every atom. Thus he proved wrong the belief that atoms are the indivisible and ultimate particles of matter. He also postulated that the electrons are of low mass and negatively charged. They are distributed throughout the atom. The atom contains a uniform sea of positive charge. We can also think of a water melon as shown below, the positive charge in the atom is spread all over like the red edible part of the water melon, while the electrons are studded in the positively charged sphere like the seeds in the water melon.

Thomson’s model of an atom

Thomson proposed that:-

(i) An atom consists of a positively charged sphere and the electrons are embedded in it.

(ii) The negative and positive charges are equal in magnitude. So, the atom as a whole is electrically neutral.

Although Thomson’s model explained that atoms are electrically neutral, the results of experiments carried out by other scientists could not be explained by this model, as we will see below.

2. Rutherford's Model of an Atom

 

 

Rutherford beamed alpha particles (doubly charged helium ions) through gold foil and detected them as flashes of light or scintillations on a screen. When alpha particles collide on the screen, they scintillate. The gold foil was only 0.0004 centimeter thick, meaning a few hundreds of atoms thick. He expected the alpha particles to pass with relatively little deflection and strike the fluorescent screen directly behind the foil.

This experiment showed that-

  1. Most of the mass of an atom (and all of its positive charge) is concentrated in a very small region in the center of an atom known as nucleus.

  2. Most of the volume of an atom is empty space.

  3. The size of nucleus is very small as compared to the size of atom.

However, Rutherford also envisioned the electrons as moving around the nucleus similar to the way in which planets orbit the sun. Today, we know that the planetary motion is not a good model for the movement of electrons around the nucleus of an atom. Instead, we think of electrons as a cloud of negative charge surrounding the nucleus.

Some of the Drawbacks of Rutherford’s model of the atom:

The revolution of the electron in a circular orbit is not predictable to be stable. Any particle in a circular orbit would experience acceleration. Charged particles would radiate energy during the process of acceleration. So we can say that the electron which is revolving would lose energy and gradually fall into the nucleus. If this happens, atoms could become unstable and matter would not exist in the form that we know. But we know that atoms are pretty stable. So this model had to be modified.

3. Bohr's Model of an Atom

In order to overcome the objections raised against Rutherford’s model of the atom, Neils Bohr put forward the following postulates about the model of an atom:

  1. Only certain special orbits known as discrete orbits of electrons, are allowed inside the atom.

  2. The electrons do not radiate energy while revolving in discrete orbits.

  3. If additional energy is supplied to an electron in a given energy level it can move up to a higher unfilled energy level. This transition requires the energy equal to the difference in energy between two energy levels(Einitial and Efinal).

  4. When the electron drops down to a lower level it releases that energy.

These special orbits or shells are called energy levels. Energy levels in an atom are shown in the figure below:

 

Neutrons

In 1932, J. Chadwick discovered another subatomic particle that had no charge, but a mass almost equal to that of a proton. It was eventually named as neutron. Neutrons are present in the nucleus of all atoms, except hydrogen. In general, a neutron is represented as ‘n’. The mass of an atom is therefore given by the sum of the masses of protons and neutrons present in the nucleus.

Distribution of electrons in an orbits:

The distribution of electrons into different 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 number of electrons present in a shell is given by the formula 2n2, where ‘n’ is the orbit number or energy level index, 1,2,3,…. Hence the maximum number of electrons in different shells are as follows: first orbit or K-shell will be = 2 × 12 = 2, second orbit or L-shell will be = 2 × 22 = 8, third orbit or M-shell will be = 2 ×32 = 18, fourth orbit or N-shell will be= 2 × 42= 32, and so on.

  2. The maximum number of electrons that can be accommodated in the outermost orbit is 8.

  3. Electrons are not accommodated in a given shell, unless the inner shells are filled. That is, the shells are filled in a step-wise manner.

Valency

The electrons present in the outermost shell of an atom are known as the valence electrons. From the Bohr-Bury scheme, we also know that the outermost shell of an atom can accommodate a maximum of 8 electrons. It was observed that the atoms of elements, having a completely filled outermost shell show little chemical activity. In other words, their combining capacity or valency is zero.

 

Composition of Atoms of the First Eighteen Elements with Electron Distribution in Various Shells

 

Atomic Number and Mass Number

Atomic Number:

It is known that protons are present in the nucleus of an atom. The atomic number of a chemical element which is also known as its proton number is defined as the number of protons found in the nucleus of an atom of that element and  it is denoted by ‘Z’. All atoms of an element have the same atomic number Z. In fact, elements are defined by the number of protons they possess.

Mass Number:

The mass number (A), also called atomic mass number or nucleon number, is the total number of protons and neutrons (together known as nucleons) in an atomic nucleus. The mass number is different for each different isotope of a chemical element. This is not the same as the atomic number (Z) which denotes the number of protons in a nucleus, and thus uniquely identifies an element. Hence, the difference between the mass number and the atomic number gives the number of neutrons (N) in a given nucleus: N=A−Z

The mass number is written either after the element name or as a superscript to the left of an element's symbol. For example, the most common isotope of carbon is carbon-12, or 12C, which has 6 protons and 6 neutrons.

Isotopes:

Isotopes are variants of a particular chemical element such that, while all isotopes of a given element have the same number of protons in each atom, they differ in the number of neutrons. The number of nucleons (both protons and neutrons) in the nucleus is the atom's mass number, and each isotope of a given element has a different mass number.

For example, carbon-12, carbon-13 and carbon-14 are three isotopes of the element carbon with mass numbers 12, 13 and 14 respectively. The atomic number of carbon is 6, which means that every carbon atom has 6 protons, so that the neutron numbers of these isotopes are 6, 7 and 8 respectively.

Isobars:

Isobars are atoms (nuclides) of different chemical elements that have the same number of nucleons. Correspondingly, isobars differ in atomic number (or number of protons) but have the same mass number. Let us consider two elements — calcium with atomic number 20, and argon with atomic number 18. The number of electrons in these atoms is different, but the mass number of both these elements is 40. That is, the total number of nucleons is the same in the atoms of this pair of elements. Atoms of different elements with different atomic numbers, which have the same mass number, are known as isobars.