2.4.1 Electron Structure

Shells

• The arrangement of electrons in an atom is called the electron 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
    • So, the first shell which is the closest to the nucleus is n = 1

  • The higher the principal quantum number, the greater the energy of the shell and the further away from the nucleus

• Each principal quantum number has a fixed number of electrons it can hold, as follows:
  • n = 1 : up to 2 electrons
  • n = 2 : up to 8 electrons
  • n = 3 : up to 18 electrons
  • n = 4 : up to 32 electrons

Subshells

• 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 shell
  • The energy of the electrons in the subshells increases in the order s < p < d

• The order of subshells appear to 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

Orbitals

• 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 them
  • Each atomic orbital can be occupied by a maximum of two electrons

• This means that the number of orbitals in each subshell is as follows:
  • s : one orbital (1 x 2 = total of 2 electrons)
  • p : three orbitals ( 3 x 2 = total of 6 electrons)
  • d : five orbitals (5 x 2 = total of 10 electrons)
  • f : seven orbitals (7 x 2 = total of 14 electrons)

• The orbitals have specific 3-D shapes

s orbital shape

• 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)

p orbital shape

• The p orbitals have a dumbbell shape
• 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

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

An overview of the shells, subshells and orbitals in an atom

• Electrons can be imagined as small spinning charges which rotate around their own axis in either a clockwise or anticlockwise direction
  • The spin of the electron is represented by its direction
  • The spin creates a tiny magnetic field with N-S pole pointing up or down

Electrons can spin either in a clockwise or anticlockwise direction around their own axis

• Electrons with the same spin repel each other which is also called spin-pair repulsion
  • Therefore, electrons will occupy separate orbitals in the same subshell first to minimise this repulsion and have their spin in the same direction
  • They will then pair up, with a second electron being added to the first p orbital, with its spin in the opposite direction

• This is known as Hund's Rule
  • E.g. if there are three electrons in a p subshell, one electron will go into each p_x, p_y and p_z orbital

 

Electron configuration: three electrons in a p subshell

• The principal quantum number indicates the energy level of a particular shell but also indicates the energy of the electrons in that shell
  • A 2p electron is in the second shell and therefore has an energy corresponding to n = 2

• Even though there is repulsion between negatively charged electrons, they occupy the same region of space in orbitals
• An orbital can only hold two electrons and they must have opposite spin - the is known as the Pauli Exclusion Principle
• This is because the energy required to jump to a higher empty orbital is greater than the inter-electron repulsion
• For this reason, they pair up and occupy the lower energy levels first

Ground state

• 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)
• The order of the subshells in terms of increasing energy does not follow a regular pattern at n = 3 and higher

The ground state of an atom is achieved by filling the lowest energy subshells first

• The electron configuration can also be represented using the orbital spin diagrams
• Each box represents an atomic orbital
• The boxes are arranged in order of increasing energy from lower to higher (i.e. starting from closest to the nucleus)
• The electrons are represented by opposite arrows to show the spin of the electrons
  • E.g. the box notation for titanium is shown below

The electrons in titanium are arranged in their orbitals as shown. Electrons occupy the lowest energy levels first before filling those with higher energy

• Writing out the electronic configuration tells us how the electrons in an atom or ion are arranged in their shells, subshells and orbitals
• This can be done using the full electron configuration or the shorthand version
  • The full electron configuration describes the arrangement of all electrons from the 1s subshell up
  • The shorthand electron configuration includes using the symbol of the nearest preceding noble gas to account for however many electrons are in that noble gas, followed by the rest of the electron configuration

• Ions are formed when atoms lose or gain electrons
  • Negative ions are formed by adding electrons to the outer subshell
  • Positive ions are formed by removing electrons from the outer subshell
  • The transition metals fill the 4s subshell before the 3d subshell, but they also lose electrons from the 4s first rather than from the 3d subshell

• The Periodic Table is split up into four main blocks depending on their electronic configuration:
  • s block elements (valence electron(s) in s orbital)
  • p block elements (valence electron(s) in p orbital)
  • d block elements (valence electron(s) in d orbital)
  • f block elements (valence electron(s) in f orbital)

The elements can be divided into four blocks according to their outer shell electron configuration

Exceptions to the Aufbau Principle

• Chromium and copper have the following electron configurations:
  • Cr is [Ar] 3d⁵ 4s¹ not [Ar] 3d⁴ 4s²
  • Cu is [Ar] 3d¹⁰ 4s¹ not [Ar] 3d⁹ 4s²

• This is because the [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹ configurations are energetically favourable
• By promoting an electron from 4s to 3d, these atoms achieve a half full or full d-subshell, respectively

Worked Example

Answer:

Answer 1:

• • Potassium has 19 electrons so the full electronic configuration is:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹

• • The 4s orbital is lower in energy than the 3d subshell and is therefore filled first
  • The nearest preceding noble gas to potassium is argon which accounts for 18 electrons so the shorthand electron configuration is:

[Ar] 4s¹

Answer 2:

• • Calcium has 20 electrons so the full electronic configuration is:

1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

• • The 4s orbital is lower in energy than the 3d subshell and is therefore filled first
  • The shorthand version is [Ar] 4s² since argon is the nearest preceding noble gas to calcium which accounts for 18 electrons

Answer 3:

• • Gallium has 31 electrons so the full electronic configuration is:

 [Ar] 3d¹⁰ 4s² 4p¹

Answer 4:

• • If you ionise calcium and remove two of its outer electrons, the electronic configuration of the Ca²⁺ ion is identical to that of argon:

Ca²⁺ is 1s² 2s² 2p⁶ 3s² 3p⁶                                   

Ar is also 1s² 2s² 2p⁶ 3s² 3p⁶

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