Electronic Configuration-Static Model, Chemistry tutorial

Introduction:

The Dalton's theory stated that atoms are indivisible units of matter. The atom is the smallest unit of matter which can take part in a chemical change. Dalton's atomic theory satisfactorily described the laws of chemical combination however could not describe why substances react the way they do. Why is oxygen capable to react with the maximum of two atoms of hydrogen as in water? Why do several elements exist only as diatomic molecules? Why are a number of elements very reactive and a few inert? Dalton's law couldn't illustrate electrolysis neither could it describe the different masses of atoms of the similar element.

Nowadays we suppose that the atom consists of a substructure of its own. The atom comprises of much smaller particles which we call protons, neutrons and electrons.

What are the proofs in support of this new depiction of the atom? How many of such particles are present in the atom of elements? How are such particles positioned in the atom?

The Electrical Nature of the Atom:

Historical evidences:

A very early proof for the electrical nature of the atom came from the Faraday. The outcome of Faraday's experiment on electrolysis illustrated that the chemical change could be caused due to the passage of electricity via aqueous solutions of chemical compounds. This proof was closely followed by the discharge tube experiment. A heated metal cathode emitted negatively charged particles. This beam of particles is termed as cathode rays and the particles are the electrons.

J.J. Thompson worked on cathode rays and verified that they are negatively charged. Their charge mass ratio was found out and found to be - 1.76 x coulomb/g-1. A. Millikan found out the electronic charge in his famous oil drop experiment in the year 1908. The charge of electron is -1.602 x 1019 coulomb. The mass of electron was computed. It is 9.11 x 10-28 g.

2458_Millikan oil drop experiment.jpg

Fig: Millikan oil drop experiment

The atom is electrically neutral; therefore it is reasonable to suppose positively charged particles in the atom. The existence of positively charged particles was confirmed in a modified version of the set up employed by J. J. Thompson by the other scientist by name Goldstein. The charge/mass ratio of the positive particles is much smaller than for the electrons. The largest charge/mass was for hydrogen ion (H+) that now symbolizes the basic particle of positive charge in the atom. The proton is around 1835 times as heavy as the electron and carry a charge equivalent however of opposite sign to that of the electron. Protons and electrons are present in each and every atom.

The proof of radioactivity by H. Becquerel further illustrated the existence of subatomic particles. Becquerel noticed that some substances spontaneously emit radiations. The most significant of these radiations are the alpha, beta and gamma radiations. Chadwick later verified the existence of a neutral particle in the atom and termed this a neutron. This neutron consists of a mass approximately equivalent to that of the proton.

Table: Properties of subatomic particles

Particles

Mass

Charge

 

grams

amu

coulombs

electron charge

Proton    

1.67 x 10-24  

1.007274

+1.602 x 10-19

+1

Neutron   

1.68 x 10-24  

1.008665   

0

 

Electron  

9.11 x 10-28  

0.000549 

-1.602 x 10-19

-1

Atomic and mass numbers:

The atoms of different elements have varying numbers of protons, electrons and neutrons. The atomic number is the number of protons in an atom of the element and for the neutral atom; the atomic number is as well the number of electrons. The sum of the number of protons and neutrons is the mass number.

Isotopes are the atoms of similar element having different mass numbers. We are familiar that the chemical property of an element based on the number and also the arrangement of the atomic electrons. This describes why atoms of the similar element having different masses having the similar chemical properties 6C12 and 6C14 These are isotopes of carbon, mass number 12 and 14 and neutron numbers 6 and 8 correspondingly.

Atomic Models:

The nuclear atom:

The existence of subatomic particles posed challenge to early scientists. This was essential to state a model for the atom. J. J. Thompson proposed that the atom could be observed as positive matter in which the electrons are uniformly distributed to make it neutral at each and every point. This view was dropped due to the findings of two other scientists, Rutherford and Marsden. They bombarded a thin gold foil having fast moving alpha particles. They found that most of the alpha particles pass via the foil undeflected. A few were deflected as large angles as some were sent back on their paths. Alpha particles are positively charged and most of them passing via the foil undeflected recommended that most of the gold foil was vacant space.

763_Scattering and deflection of α-particles.jpg

Fig: Scattering and deflection of α-particles

Rutherford, employing the findings in the above experiment stated a model for the atom. He stated that the atom comprised of a tiny positively charged nucleus. The nucleus is centrally positioned in the atom and the electrons surround it. The extremely small number of deflections of alpha particles recommends that the nucleus takes a very small part of the atom. For heavy particles like the alpha particle to be so deflected proposes that the nucleus is a centre of heavy mass and positive charge. The protons and neutrons engage the nucleus whereas the electrons are arranged around the nucleus and move in orbits around it, as planets around the sun. This is Rutherford's nuclear model of the atom.

By counting the number of α-particles deflected in different directions, Rutherford was capable to exhibit that the diameter of a nucleus is around 1/100,000 times the diameter of the atom.

Electronic energy levels:

The main problem of Rutherford's model of the atom is that the electron (negatively charged) rotating around the nucleus (positively charged) will lose energy constantly due to the electrostatic force of attraction of the nucleus on the electrons. This is not noticed in practice. Intact energy absorption and emission through elements is discontinuous.

Have you ever observe the rainbow in the sky? The colors you spot range from violet to red having no sharp line separating one color from the other. This is a constant spread of colors and is termed as a continuous spectrum. The various colors are constituent colors of light. In the laboratory the separation of light to its component colors as well occurs whenever light passes via a glass prism.

The light from vapor of an element doesn't provide a continuous spectrum. Each and every element consists of its own characteristic bright lines in specific positions. This is a line spectrum recommending that the light energy absorption or emissions through elements is only at specific energies characteristic of the element.

On the base of the above examinations Niel Bohr introduced a model for the atom in which the electrons move round the nucleus only in allowed orbits numbered in sequence. The orbit closest to the nucleus is assigned number 1 and is the orbit of lowest energy. Allowed transitions are transitions from one orbit to the other and will lead to emission or absorption of the energy difference between the orbits.

In Bohr's model light, there are electronic energy levels in atoms corresponding to various orbits of electron motion that are allowed. Such energy levels are at times termed to as electron shells and designated K, L, M, N and so on corresponding to orbit numbers 1, 2, 3, 4 and so on correspondingly.

Bohr's model provides satisfactory illustration of the hydrogen spectrum. The theory is limited though in its description for multi electron atoms. The wave mechanical treatment of the atom overcomes this restriction of Bohr's theory.

Electronic configuration:

Electronic configuration provides the arrangement of electrons in the energy levels in the atom. In the ground state (that is, most stable state of lowest energy) of the atom, electron assignment to energy levels is according to the given rules:

a) The order of filling is K → L → M → N and so on. The first shell is filled first prior to filling the second.

b) The K shell can contain maximum of two (2) electrons.

c) The L shell can contain maximum of eight (8) electrons.

d) The M shell contains eighteen (18) electrons.

Whenever 8 electrons are accommodated to the M shell though, there is additional stability and the next 2 electrons go to the N shell. Consequently any electron goes to the M shell till it accommodates the maximum of 18 electrons.

e) There are higher energy levels which accommodate a larger number of electrons than the K, L, M shells.

Table: Electron arrangement in a few elements

Element   

Symbol 

Atomic number

Arrangement of electrons

Hydrogen

H

1

1

Helium

He

2

2

Lithium

Li

3

2,1

Beryllium

Be

4

2,2

Nitrogen

N

7

2,5

Oxygen

O

8

2,6

Fluorine

F

9

2,7

Neon

Ne

10

2,8

Sodium

Na

11

2,8,1

Magnesium

Mg

12

2,8,2

Aluminium

Al

13

2,8,3

Silicon

Si

14

2,8,4

Sulphur

S

16

2,8,6

Chlorine

Cl

17

2,8,7

Argon

Ar

18

2,8,8

Calcium

Ca

20

2,8,8,2

The electronic structure 2, 8, 1 represents 2 electrons in the K shell, 8 in the L shell and 1 in the M shell. From the above table we note the given facts:

i) He and Ne; each consists of full shell of electrons and are both inert gases.

ii) F and Cl; each consists of one electron short of a full electron arrangement.

iii) H and Na; each consist of one electron outside a full shell electron arrangement. These pair of elements having similar electron arrangements as well consists of similar chemical properties.

Ion formation:

An ion is made up if an atom gains or losses electron example: H+, F- and Na+. The proton number is not influenced by ion formation however the electron number increases or decreases.

Ion

Electron number

Electron arrangement

Li+

2

2

F-

10

2, 8

02-

10

2, 8

Mg2+

10

2, 8

Cl-

18

2, 8, 8

Ca2+

18

2, 8, 8

You will notice that these ions have closed or full shell electron arrangements and are stable ions. This examination and the full shell arrangements for the inert gases, recommend that a full shell electron arrangement is the stable electron arrangement.

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