Atomic Structure

Atomic structure consists of a nucleus, protons, neutrons and electrons. Atoms are basic building blocks of matter.

This course introduces the concept of atomic structure and covers the following topics:

• matter
• atoms, subatomic particles
• atomic number and mass number
• isotopes
• various atomic theories

Matter

Matter is anything that has mass and takes up space (has volume). Everything around us — the food we eat, the air we breathe, clouds, stars, plants, animals, water, dust– is made up of matter.

Characteristics of particles of matter are the following:

• Particles of matter are very small and are normally not visible to the naked eye.
• Matter particles keep moving continuously .
• The kinetic energy associated with the continuous motion of the particles is directly proportional to their temperature.
• The particles of matter attract each other.
• The attracting force of the particles keeps the particles together; however, the strength of the attracting force varies from one kind of matter to another.

Atoms

Atoms are the building blocks of matter. The word ‘atom’ has been derived from the Greek word ‘a-tomio’ which means ‘non-divisible’.

The diversity of chemical behavior of different elements is due to the differences in the internal structure of atoms of these elements.

Atoms are composed of three type of particles: protons, neutrons, and electron (Figure 1). At the center of an atom is a nucleus, which is made up of protons and neutrons.

• Protons are positively charged particles.
• Neutrons are about the size of protons but have no charge.
• Electrons are negatively charged particles that orbit the nucleus.

Table showing Differences in Electron, Proton and Neutron

	Electron	Proton	Neutron
Definition	Negatively charged sub-atomic particle found in an atom	Positively charged sub-atomic particle found in an atom	Neutral sub-atomic particle found in an atom
Symbols	e	p	n
Location in the atom	Outside nucleus	Nucleus	Nucleus
Electric charge	1 –	1 +	0
Reactions	Take part in both chemical and nuclear reactions	Take part in nuclear reaction	Only gets exposed to nuclear reaction
Relative Mass	0 Atomic mass unit (amu)	1  amu	Very close to 1 amu
Mass (Kg)	9.109 x 10-31	1.673 x 10-27	1.675 x 10-27
Discovery	J.J Thompson	Ernest Rutherford	James Chadwick

Atomic Number

The positive charge on nucleus is due to protons. The atomic number (Z) is the number of protons present in the nucleus. For example, number of protons in hydrogen atom is 1 and therefore its atomic number is 1.

Mass Number

The mass number of an atom is the total number of protons plus neutrons in its nucleus (Figure 2). The mass number (A) is the sum of the atomic number (Z), which is the number of protons, and the number of neutrons (N) in the atomic nucleus of an isotope of an element, i.e. A=Z+N

Isotopes

Atoms with same atomic number (number of protons), but different mass numbers (number of protons and neutrons) are called isotopes. They occur naturally or can be produced artificially.

Isotopes are separated through mass spectrometry.

For example, the simplest and commonest form of hydrogen (Figure 3) has a nucleus that consists of a single proton; it is the only atom with no neutrons: its mass number is 1.

A rarer form of hydrogen known as deuterium has one proton and one neutron: its mass number is 2.

A third form of hydrogen known as tritium has one proton and two neutrons: its mass number is 3.

The stability of a nucleus depends on the ratio of protons to neutrons in it. Atoms with unstable nucleus are called radioactive atoms.

They spontaneously decay, emitting alpha, beta, or gamma rays until they reach a stability. For example, Uranium has three naturally occurring isotopes.

These are uranium-234, uranium-235, and uranium-238. Since each atom of uranium has 92 protons, the isotopes must have 142, 143 and 146 neutrons respectively.

Atomic theory

Around 500 BC, an Indian Philosopher Maharishi Kanad, postulated the concept of indivisible part of matter and named it ‘pramanu.’

Atomic theory has been revised over the years as scientists discovered more information about the atoms (Figure 4).

Atomic isotopes and the inter-conversion of mass and energy has been added. In addition, the discovery of subatomic particles indicated that atoms can be divided into smaller parts.

Dalton’s Atomic Theory

John Dalton (1766-1844) is the scientist credited for proposing the atomic theory in 1808. The atomic mass unit is designated as Dalton.

Dalton developed the law of multiple proportions based on the works of Antoine Lavoisier and Joseph Proust.

Dalton’s atomic theory states the following:

• All matter, whether an element, a compound, or a mixture is composed of small particles called atoms. Atoms are indivisible and indestructible.
• All atoms of a given element are identical in mass and properties
• Compounds are formed by a combination of two or more different kinds of atoms. In each compound, the relative number and kinds of atoms are constant.
• A chemical reaction is a rearrangement of atoms.

Discovery of Electrons

Electron was discovered by J. J. Thomson in 1897, when he was studying the properties of cathode ray.

He constructed a glass tube which was partially evacuated i.e. air was pumped out of the tube (Figure 5). Then he applied a high electrical voltage between two electrodes at either end of the tube.

He detected that a stream of particle (ray) was coming out from the negatively charged electrode (cathode) to positively charged electrode (anode).

It is called cathode ray and the tube is called cathode ray tube.

Following are the properties of cathode ray particle:

• They travel in straight lines
• They are independent of the material composition of the cathode

When electric field is applied in the path of cathode ray, it deflects the ray towards positively charged plate. Hence, cathode ray consists of negatively charged particles.

Discovery of Nucleus

Ernest Rutherford discovered the nucleus of the atom in 1911. His gold foil model contradicted Thomson’s atomic model (Figure 6).

Rutherford, in his experiment, directed high energy streams of α-particles from a radioactive source at a thin sheet (100 nm thickness) of gold.

To study the deflection caused due to α-particles, he placed a fluorescent zinc sulphide screen around the thin gold foil.

He observed that the major fraction of the α-particles bombarded towards the gold sheet passed through it without any deflection, some of them were deflected by the gold sheet by very small angles, while some were deflected back at 180⁰ (Figure 7).

Rutherford proposed the atomic structure of elements based on his observations. According to the Rutherford atomic model:

• The positively charged particles and most of the mass of an atom was concentrated in an extremely small volume called nucleus.
• Rutherford model proposed that the negatively charged electrons surround the nucleus of an atom. He also claimed that the electrons surrounding the nucleus revolve around it with very high speed in circular paths called orbits.
• Electrons and nucleus are held together by a strong electrostatic force of attraction.

Bohr’s Atomic Model

In 1913, Neils Bohr proposed his quantized shell model of the atom. His theory explained how electrons can have stable orbits around the nucleus.

According to this “planetary” model, electrons encircle the nucleus of the atom in orbits (Figure 8). Each orbit has a definite energy and is called an energy shell or energy level.

When the electron is in one of these orbits, its energy is fixed. Orbits further from the nucleus exist at higher energy levels and vice-versa.

When an electron jumps from a higher energy level to lower energy level, it emits energy, and when an electron absorbs sufficient energy it jumps from a lower energy level to a higher energy level.

External References:

https://phys.org/tags/atomic+structure/

https://www.pbslearningmedia.org/resource/lsps07.sci.phys.matter.theatom/the-atom/