The Bohr Model of the Atom

The Bohr model of the atom

One of the simplest working models of the atom is that developed by Niels Bohr. In this, a central nucleus containing protons and neutrons is orbited by electrons. The electrons are arranged in shells or energy levels. Electrons in shells nearer the nucleus have lower energy than those farther away. These orbits had different energy values. Electrons could gain or lose energy by jumping from one orbit to another. The Bohr model gave rise to the idea that electrons could be found in shells or energy levels.

The shell closest to the nucleus is lowest in energy. As we move further from the nucleus the electrons at the higher levels have more energy.

Electronic Configurations

The electronic configuration of an atom describes the arrangement of the electrons in the shells.

The electrons are placed in the lowest available energy level first. For instance hydrogen has one electron and this will be found in the first energy level, which is sited close to the centre of the atom. Higher levels are found further from the atom.

Each shell can hold a maximum number of electrons and this is shown in the table below:-

Principal Quantum Number n= / Shell / Maximum Electrons

Electrons are placed into shells, starting with the lowest energy level first. Each shell must be full before the next shell starts to fill.

Sub Shells

The newer models of electronic structure depend on a complex understanding of Mathematics, Schrodinger’s equation and Heisenberg’s uncertainty principle. Explanations of these can be found in many textbooks but a knowledge of these is well beyond that required at A/S level.

It is now thought that the following hold:-

· The energy levels (shells) of principal quantum numbers n=1,2,3,4 etc do not have precise energy values. Instead they each contain a set of sub shells, which contain orbitals with different energy values.

· The sub shells are of different types labeled s,p,d and f. An s-subshell contains one orbital; a p-subshell contains three orbitals and a d subshell contains five orbitals. An f-subshell contains seven orbitals but this is beyond the scope of our treatment.

Orbitals

A region of space within which there is a high probability of finding an electron is called an orbital. There are two types of orbitals you must know at this level:-

The s-orbital

The p-orbitals

px orbital py orbital pz orbital

Student activity 1

Write true (T) or false (F) next to each statement.

a) An orbital is a region of space in which an electron can be found.

b) The p-orbitals can each hold 6 electrons.

c) The third shell can hold 8 electrons.

d) d-orbitals have higher energy electrons than s orbitals of the same principal quantum number.

e) The outer shell electrons have a higher energy than the other electrons in an atom.

The Production of Electronic Configurations

In order to write electronic configurations of the atoms we have to use some specific rules.

First it is important to know how many orbitals there are at each principal quantum number and the order in which the orbitals are filled.

Types of orbital at each level

Principal Level n = / 1 / 2 / 3 / 4
Sub Levels / s / s and p / s, p and d / s,p,d and f

Maximum Occupancy of the Orbitals

s orbitals can hold a maximum of ….. electrons at a particular energy level

p orbitals can hold a maximum of ….. electrons at a particular energy level

d orbitals can hold a maximum of ….. electrons at a particular energy level

Rule 1 : The Aufbau Rule

The word “Aufbau” is German for building up. This rule states that the sub shells are filled in order or energy. The lowest energy levels are filled first. This can be shown in the diagram below:

______

Energy ______

___

______

___

______

___

Rule 2 : The Pauli Exclusion Principle

This states that an orbital cannot contain more than two electrons and then only if they have opposite spins.

All electrons spin on an axis. They can spin either clockwise or anti-clockwise.

Rule 3: Hund’s rule of maximum multiplicity

This states that the orbitals of a sub-shell must be occupied singly and with parallel spin before being occupied in pairs. Example: Nitrogen (Atomic Number =7)

This has the electronic configuration

It can also be represented in electron in boxes format

Electrons in box format can also be used to represent electronic configurations, as it is the only way of showing electrons with spin. When using this notation each box represents one orbital and so an s orbital is represented by one box, three boxes represent p sub shells and five boxes represent d sub shells. Conventionally however the electronic configurations are usually written in s, p, d and f notation.

The electronic configurations of the elements

At AS you have to be able to write the electronic configuration of the elements from hydrogen to krypton.

Student activity 2

a) Write the order of the sublevels (without any electrons in them- no superscripts) up to and including 4p

…………………………………………………………………………………………………..

b) Write the full electronic configurations for the following atoms:

H………….. C…………….. F…………………. Ne…………….

Na………………. Mg……………………. Al…………………….

Cl………………. and K………………….

Electron Configuration Shorthand (noble gas shorthand):

A further refinement of electronic configurations is the “noble gas core” notation. This is a shortening tool, which simplifies the configurations of the elements. Consider Calcium with 20 electrons. Its electronic configuration is

1s2 2s2 2p6 3s2 3p6 4s2

The first eighteen electrons are the same as the eighteen electrons in argon and so we can shorten this by using the symbol [Ar], which is a shorthand notation for 18 electrons. The configuration of the element now becomes [Ar] 4s2.This has obvious benefits the further down the periodic table one goes. Similar short hand notation is available for eight electrons

[Ne] = 10 electrons

[Ar] = 18 electrons

[Kr] = 36 electrons

Student activity 3

1. Give the electronic configuration for the following atoms giving full s,p,d notation

S …………………….. Sc …………………… Br…………………….

2. Give ‘electron in box’ notation for the following

3. Use ‘Noble gas core’ notation for the following

Al …………………….. Li …………………….. P ……………………..

The Electronic Configuration of the Chromium and Copper

These two elements have different electronic configurations to that predicted by the “Aufbau” principle. In each case the 4s orbital is half filled, containing only one electron.

Cr 1s22s22p63s23p64s13d5

Cu 1s22s22p63s23p64s13d10

The reason for having only one electron in the 4s orbital is that by placing an electron into the 3d orbital this creates either a full or half filled 3d orbital which gives the atom increased stability.

THIS ONLY OCCURS FOR THE TWO ATOMS STATED ABOVE.

The Electronic Configuration of Ions

When ions are formed electrons are either lost or gained. Electrons will be lost from the outer electron shells when positive ions are formed or will be added to the outer energy levels when negative ions are formed.

7N 3- (Note: 3 electrons added)

11Na+ (Note: 1 electron removed)

In transition metal ions the 4s electrons are always lost before the 3d electrons. This will be explained later when we consider the chemistry of the transition metals.

Student activity 4

Write the electronic configuration of the following ions.

a) K+…………………….. b) O2-………………… c) Mg2+………………….

d) Cl-……………… e) H-………………

Ionisation is the removal of an electron from an atom to make a positive ion. Once the electron has been removed we say that the atom has been ionized. Energy is needed to remove electrons from the atom and this is generally called the ionisation energy.

Definition The first ionisation energy is the energy required to remove one electron from each atom in one mole of gaseous atoms to make one mole of gaseous unipositive ions.

Example equation:

Definition The second ionisation energy is the energy required to remove one electron from each unipositive ion in one mole of gaseous unipositive ions to make one mole of gaseous dipositive ions.

Example equation:

Both of the above equations are endothermic because energy is needed to overcome the electrostatic force of attraction between the negative electron and the positive nucleus.

Student activity 5

Write an equation to represent each of the following:-

a) The second ionisation energy of calcium

……………………………………………………………………………………………………………….

b) The third ionisation energy of aluminium.

……………………………………………………………………………………………………………….

Factors Influencing ionisation energy

There are three factors which influence the magnitude of ionisation energy and these are:-

1. Nuclear Charge

Example : Hydrogen and Helium

The first ionisation energy of helium is higher than that of hydrogen because helium has more protons in its nucleus than hydrogen. It therefore has a greater nuclear charge and exerts a greater attractive force on the outer electrons.

2. Distance from the Nucleus

Attraction is significantly reduced with increasing distance. Example: Lithium and Sodium

Lithium has a higher first ionisation energy than sodium because the outer electron in sodium is further from the nucleus. The further the electron is from the nucleus the easier it is to remove.

3. Shielding Effect

Any inner shells between the electrons and the nucleus repel the outer electrons. This repelling effect reduces the attraction experienced by the outer electrons from the nucleus. The electrons in the inner shells shield the outer electrons from the attractive force of the nucleus.

Student activity 6

State and explain which element in the following pairs of elements has the highest first ionisation energy

(a) Carbon and Nitrogen

……………………………………………………………………………………………………………….

(b) Potassium and Lithium

……………………………………………………………………………………………………………….

………………………………………………………………………………………………………………..

Successive ionisation energies

Electronic configurations of the elements can be predicted from successive ionisation energy data. Successive ionisation energies provide evidence for different energy levels. An element has as many ionisation energies as it has electrons.

Successive ionisation energy values are very large and so we plot log10 (ionisation energy) against number of electrons removed to fit the graph on a suitable scale.

Student activity 7

Using the data given in the table below calculate log10 (ionisation energy) for each ionisation energy.

No of Electrons Removed / Ionisation Energy (kJmol-1) / log10(Ionisation energy)
1 / 490
2 / 4560
3 / 6490
4 / 9540
5 / 13400
6 / 16600
7 / 20100
8 / 25500
9 / 28900
10 / 141000
11 / 158700

Give an explanation of each of the following statements:

1. Each time an electron is removed the ionisation energy increases.

2. The second ionisation energy is much greater than the first.

3. The second to ninth successive ionisation energies are of a similar order of magnitude.

4. There is a large difference in value between the ninth and tenth ionisation energy.

5. Sketch graphs to show the successive ionisation energies for:-

(a) Neon (10 electrons) (b) Phosphorus (15 electrons)

6. Using your knowledge of successive ionisation energy data identify the group in the periodic table that the following elements come from.

1st I.E. / 2nd I.E. / 3rd I.E. / 4th I.E.
Element X / 494 / 4560 / 6900 / 9500
Element Y / 527 / 1800 / 2740 / 11600
Element Z / 736 / 1450 / 7400 / 10500