IB Chemistry SL Structure 1.3
Electron configuration
The electromagnetic spectrum is a range of frequencies that covers all electromagnetic radiation and their respective wavelengths and energy
It is divided into bands or regions, and is very important in analytical chemistry.
The spectrum shows the relationship between frequency, wavelength and energy
Frequency is how many waves pass per second, and wavelength is the distance between two consecutive peaks on the wave
Gamma rays, X-rays and UV radiation are all dangerous - you can see from that end of the spectrum that it is high frequency and high energy, which can be very damaging to your health
The electromagnetic spectrum spans a broad spectrum from very long radio waves to very short gamma rays
All light waves travel at the same speed; what distinguishes them is their different frequencies
The speed of light (symbol ‘c’) is constant and has a value of 3.00 x 108 ms-1
As you can see from the spectrum, frequency (symbol ‘f') is inversely proportional to wavelength (symbol ‘λ')
In other words, the higher the frequency, the shorter the wavelength
The equation that links them is c = fλ
Since c is constant you can use the formula to calculate the frequency of radiation given the wavelength, and vice versa
A continuous spectrum in the visible region contains all the colours of the spectrum
This is what you are seeing in a rainbow, which is formed by the refraction of white light through a prism or water droplets in rain
A continuous spectrum shows all frequencies of light
However, a line spectrum only shows certain frequencies
The line spectrum of helium which shows only certain frequencies of light
This tells us that the emitted light from atoms can only be certain fixed frequencies - it is quantised (quanta means 'little packet')
Electrons can only possess certain amounts of energy - they cannot have any energy value
The formula that relates frequency and wavelength is printed in Section 1 of the IB Chemistry Data Booklet so you don’t need to learn it
You will also find the speed of light and other useful constants in Section 2
Electrons move rapidly around the nucleus in energy shells
If their energy is increased, then they can jump to a higher energy level
The process is reversible, so electrons can return to their original energy levels
When this happens, they emit energy
The frequency of energy is exactly the same, it is just being emitted rather than absorbed:
The difference between absorption and emission depends on whether electrons are jumping from lower to higher energy levels or the other way around
The energy they emit is a mixture of different frequencies
This is thought to correspond to the many possibilities of electron jumps between energy shells
If the emitted energy is in the visible region, it can be analysed by passing it through a diffraction grating
The result is a line emission spectrum
The line emission (visible) spectrum of hydrogen
Each line is a specific energy value
This suggests that electrons can only possess a limited choice of allowed energies
These packets of energy are called 'quanta' (plural quantum)
What you should notice about this spectrum is that the lines get closer together towards the blue end of the spectrum
This is called convergence and the set of lines is converging towards the higher energy end, so the electron is reaching a maximum amount of energy
This maximum corresponds to the ionisation energy of the electron
These lines were first observed by the Swiss school teacher Johannes Balmer, and they are named after him
We now know that these lines correspond to the electron jumping from higher levels down to the second or n = 2 energy level
A larger version of the hydrogen spectrum from the infrared to the ultraviolet region looks like this
The full hydrogen spectrum
In the spectrum, we can see sets or families of lines
Balmer could not explain why the lines were formed - an explanation had to wait until the arrival of Planck's Quantum Theory in 1900
Niels Bohr applied the Quantum Theory to electrons in 1913, and proposed that electrons could only exist in fixed energy levels
The line emission spectrum of hydrogen provided evidence of these energy levels and it was deduced that the families of lines corresponded to electrons jumping from higher levels to lower levels
Electron jumps in the hydrogen spectrum
The findings helped scientists to understand how electrons work and provided the backbone to our knowledge of energy levels, sublevels and orbitals
The jumps can be summarised as follows:
Jumps | Region | Energy |
n∞→ n3 | Infrared | Low |
n∞ → n2 | Visible | ↓ |
n∞ → n1 | Ultraviolet | High |
Which electron transition in the hydrogen atom emits visible light?
A. n = 1 to n = 2
B. n = 2 to n = 3
C. n = 2 to n = 1
D. n = 3 to n = 2
Answer
Option D is correct
Emission in the visible region occurs for an electron jumping from any higher energy level to n = 2
The arrangement of electrons in an atom is called the electronic configuration
Electrons are arranged around the nucleus in principal energy levels or principal quantum shells
Principal quantum numbers (n) are used to number the energy levels or quantum shells
The lower the principal quantum number, the closer the shell is to the nucleus
The higher the principal quantum number, the greater the energy of the electron within that shell
Each principal quantum number has a fixed number of electrons it can hold
n = 1 : up to 2 electrons
n = 2 : up to 8 electrons
n = 3 : up to 18 electrons
n = 4 : up to 32 electrons
There is a pattern here - the mathematical relationship between the number of electrons and the principal energy level is 2n2
So for example, in the third shell n = 3 and the number of electrons is 2 x (32) = 18
Electrons are arranged in principal quantum shells, which are numbered by principal quantum numbers
The principal quantum shells are split into subshells which are given the letters s, p and d
Elements with more than 57 electrons also have an f subshell
The energy of the electrons in the subshells increases in the order s < p < d
The order of subshells overlap for the higher principal quantum shells as seen in the diagram below:
Electrons are arranged in principal quantum shells, which are numbered by principal quantum numbers
The subshells contain one or more atomic orbitals
Orbitals exist at specific energy levels and electrons can only be found at these specific levels, not in between
Each atomic orbital can be occupied by a maximum of two electrons
The orbitals have specific 3D shapes:
Representation of orbitals (the dot represents the nucleus of the atom) showing spherical s orbitals (a), p orbitals containing ‘lobes’ along the x, y and z axis
Note that the shape of the d orbitals is not required for IB Chemistry
An overview of the shells, subshells and orbitals in an atom
The ground state is the most stable electronic configuration of an atom which has the lowest amount of energy
This is achieved by filling the subshells of energy with the lowest energy first (1s)
This is called the Aufbau Principle
The order of the subshells in terms of increasing energy does not follow a regular pattern at n = 3 and higher
The Aufbau Principle - following the arrows gives you the filling order
The principal quantum shells increase in energy with increasing principal quantum number
E.g. n = 4 is higher in energy than n = 2
The subshells increase in energy as follows: s < p < d < f
The only exception to these rules is the 3d orbital which has slightly higher energy than the 4s orbital, so the 4s orbital is filled before the 3d orbital
Relative energies of the shells and subshells
Each shell can be divided further into subshells, labelled s, p, d and f
Each subshell can hold a specific number of orbitals:
s subshell : 1 orbital
p subshell : 3 orbitals labelled px, py and pz
d subshell : 5 orbitals
f subshell : 7 orbitals
Each orbital can hold a maximum number of 2 electrons so the maximum number of electrons in each subshell is as follows:
s : 1 x 2 = total of 2 electrons
p : 3 x 2 = total of 6 electrons
d : 5 x 2 = total of 10 electrons
f : 7 x 2 = total of 14 electrons
In the ground state, orbitals in the same subshell have the same energy and are said to be degenerate, so the energy of a px orbital is the same as a py orbital
Shells are divided into subshells which are further divided into orbitals
Principle quantum number, n (shell) | Subshells possible (s, p, d, f) | Orbitals per subshell | Orbitals per principle quantum number | Electrons per subshell | Electrons per shell |
1 | s | 1 | 1 | 2 | 2 |
2 | s | 1 | 4 | 2 | 8 |
p | 3 | 6 | |||
3 | s | 1 | 9 | 2 | 18 |
p | 3 | 6 | |||
d | 5 | 10 | |||
4 | s | 1 | 16 | 2 | 32 |
p | 3 | 6 | |||
d | 5 | 10 | |||
f | 7 | 14 |
The s orbitals are spherical in shape
The size of the s orbitals increases with increasing shell number
E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum shell (n = 1)
The s orbitals become larger with increasing principal quantum number
The p orbitals are dumbbell-shaped
Every shell has three p orbitals except for the first one (n = 1)
The p orbitals occupy the x, y and z axes and point at right angles to each other, so are oriented perpendicular to one another
The lobes of the p orbitals become larger and longer with increasing shell number
The p orbitals become larger and longer with increasing principal quantum number
Electron configuration
The electromagnetic spectrum is a range of frequencies that covers all electromagnetic radiation and their respective wavelengths and energy
It is divided into bands or regions, and is very important in analytical chemistry.
The spectrum shows the relationship between frequency, wavelength and energy
Frequency is how many waves pass per second, and wavelength is the distance between two consecutive peaks on the wave
Gamma rays, X-rays and UV radiation are all dangerous - you can see from that end of the spectrum that it is high frequency and high energy, which can be very damaging to your health
The electromagnetic spectrum spans a broad spectrum from very long radio waves to very short gamma rays
All light waves travel at the same speed; what distinguishes them is their different frequencies
The speed of light (symbol ‘c’) is constant and has a value of 3.00 x 108 ms-1
As you can see from the spectrum, frequency (symbol ‘f') is inversely proportional to wavelength (symbol ‘λ')
In other words, the higher the frequency, the shorter the wavelength
The equation that links them is c = fλ
Since c is constant you can use the formula to calculate the frequency of radiation given the wavelength, and vice versa
A continuous spectrum in the visible region contains all the colours of the spectrum
This is what you are seeing in a rainbow, which is formed by the refraction of white light through a prism or water droplets in rain
A continuous spectrum shows all frequencies of light
However, a line spectrum only shows certain frequencies
The line spectrum of helium which shows only certain frequencies of light
This tells us that the emitted light from atoms can only be certain fixed frequencies - it is quantised (quanta means 'little packet')
Electrons can only possess certain amounts of energy - they cannot have any energy value
The formula that relates frequency and wavelength is printed in Section 1 of the IB Chemistry Data Booklet so you don’t need to learn it
You will also find the speed of light and other useful constants in Section 2
Electrons move rapidly around the nucleus in energy shells
If their energy is increased, then they can jump to a higher energy level
The process is reversible, so electrons can return to their original energy levels
When this happens, they emit energy
The frequency of energy is exactly the same, it is just being emitted rather than absorbed:
The difference between absorption and emission depends on whether electrons are jumping from lower to higher energy levels or the other way around
The energy they emit is a mixture of different frequencies
This is thought to correspond to the many possibilities of electron jumps between energy shells
If the emitted energy is in the visible region, it can be analysed by passing it through a diffraction grating
The result is a line emission spectrum
The line emission (visible) spectrum of hydrogen
Each line is a specific energy value
This suggests that electrons can only possess a limited choice of allowed energies
These packets of energy are called 'quanta' (plural quantum)
What you should notice about this spectrum is that the lines get closer together towards the blue end of the spectrum
This is called convergence and the set of lines is converging towards the higher energy end, so the electron is reaching a maximum amount of energy
This maximum corresponds to the ionisation energy of the electron
These lines were first observed by the Swiss school teacher Johannes Balmer, and they are named after him
We now know that these lines correspond to the electron jumping from higher levels down to the second or n = 2 energy level
A larger version of the hydrogen spectrum from the infrared to the ultraviolet region looks like this
The full hydrogen spectrum
In the spectrum, we can see sets or families of lines
Balmer could not explain why the lines were formed - an explanation had to wait until the arrival of Planck's Quantum Theory in 1900
Niels Bohr applied the Quantum Theory to electrons in 1913, and proposed that electrons could only exist in fixed energy levels
The line emission spectrum of hydrogen provided evidence of these energy levels and it was deduced that the families of lines corresponded to electrons jumping from higher levels to lower levels
Electron jumps in the hydrogen spectrum
The findings helped scientists to understand how electrons work and provided the backbone to our knowledge of energy levels, sublevels and orbitals
The jumps can be summarised as follows:
Jumps | Region | Energy |
n∞→ n3 | Infrared | Low |
n∞ → n2 | Visible | ↓ |
n∞ → n1 | Ultraviolet | High |
Which electron transition in the hydrogen atom emits visible light?
A. n = 1 to n = 2
B. n = 2 to n = 3
C. n = 2 to n = 1
D. n = 3 to n = 2
Answer
Option D is correct
Emission in the visible region occurs for an electron jumping from any higher energy level to n = 2
The arrangement of electrons in an atom is called the electronic configuration
Electrons are arranged around the nucleus in principal energy levels or principal quantum shells
Principal quantum numbers (n) are used to number the energy levels or quantum shells
The lower the principal quantum number, the closer the shell is to the nucleus
The higher the principal quantum number, the greater the energy of the electron within that shell
Each principal quantum number has a fixed number of electrons it can hold
n = 1 : up to 2 electrons
n = 2 : up to 8 electrons
n = 3 : up to 18 electrons
n = 4 : up to 32 electrons
There is a pattern here - the mathematical relationship between the number of electrons and the principal energy level is 2n2
So for example, in the third shell n = 3 and the number of electrons is 2 x (32) = 18
Electrons are arranged in principal quantum shells, which are numbered by principal quantum numbers
The principal quantum shells are split into subshells which are given the letters s, p and d
Elements with more than 57 electrons also have an f subshell
The energy of the electrons in the subshells increases in the order s < p < d
The order of subshells overlap for the higher principal quantum shells as seen in the diagram below:
Electrons are arranged in principal quantum shells, which are numbered by principal quantum numbers
The subshells contain one or more atomic orbitals
Orbitals exist at specific energy levels and electrons can only be found at these specific levels, not in between
Each atomic orbital can be occupied by a maximum of two electrons
The orbitals have specific 3D shapes:
Representation of orbitals (the dot represents the nucleus of the atom) showing spherical s orbitals (a), p orbitals containing ‘lobes’ along the x, y and z axis
Note that the shape of the d orbitals is not required for IB Chemistry
An overview of the shells, subshells and orbitals in an atom
The ground state is the most stable electronic configuration of an atom which has the lowest amount of energy
This is achieved by filling the subshells of energy with the lowest energy first (1s)
This is called the Aufbau Principle
The order of the subshells in terms of increasing energy does not follow a regular pattern at n = 3 and higher
The Aufbau Principle - following the arrows gives you the filling order
The principal quantum shells increase in energy with increasing principal quantum number
E.g. n = 4 is higher in energy than n = 2
The subshells increase in energy as follows: s < p < d < f
The only exception to these rules is the 3d orbital which has slightly higher energy than the 4s orbital, so the 4s orbital is filled before the 3d orbital
Relative energies of the shells and subshells
Each shell can be divided further into subshells, labelled s, p, d and f
Each subshell can hold a specific number of orbitals:
s subshell : 1 orbital
p subshell : 3 orbitals labelled px, py and pz
d subshell : 5 orbitals
f subshell : 7 orbitals
Each orbital can hold a maximum number of 2 electrons so the maximum number of electrons in each subshell is as follows:
s : 1 x 2 = total of 2 electrons
p : 3 x 2 = total of 6 electrons
d : 5 x 2 = total of 10 electrons
f : 7 x 2 = total of 14 electrons
In the ground state, orbitals in the same subshell have the same energy and are said to be degenerate, so the energy of a px orbital is the same as a py orbital
Shells are divided into subshells which are further divided into orbitals
Principle quantum number, n (shell) | Subshells possible (s, p, d, f) | Orbitals per subshell | Orbitals per principle quantum number | Electrons per subshell | Electrons per shell |
1 | s | 1 | 1 | 2 | 2 |
2 | s | 1 | 4 | 2 | 8 |
p | 3 | 6 | |||
3 | s | 1 | 9 | 2 | 18 |
p | 3 | 6 | |||
d | 5 | 10 | |||
4 | s | 1 | 16 | 2 | 32 |
p | 3 | 6 | |||
d | 5 | 10 | |||
f | 7 | 14 |
The s orbitals are spherical in shape
The size of the s orbitals increases with increasing shell number
E.g. the s orbital of the third quantum shell (n = 3) is bigger than the s orbital of the first quantum shell (n = 1)
The s orbitals become larger with increasing principal quantum number
The p orbitals are dumbbell-shaped
Every shell has three p orbitals except for the first one (n = 1)
The p orbitals occupy the x, y and z axes and point at right angles to each other, so are oriented perpendicular to one another
The lobes of the p orbitals become larger and longer with increasing shell number
The p orbitals become larger and longer with increasing principal quantum number