Electrons & Ions — IAL Edexcel Modular (Chemistry) Revision Notes & Topical Questions

Section 1 of 18Ionisation energy

1First ionisation energy

First ionisation energy the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

• EquationX(g) → X⁺(g) + e^-, for example Na(g) → Na⁺(g) + e^- (+496 kJ mol^-1).
• Units are kJ mol^-1, measured under standard conditions.

2Second ionisation energy

Second ionisation energy the energy required to remove one mole of electrons from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions

• EquationX⁺(g) → X²⁺(g) + e^-.

3Later ionisation energies

• Later ionisation energies the nth IE removes electrons from gaseous (n−1)+ ions
• EquationX⁽ⁿ⁻¹⁾⁺(g) → Xⁿ⁺(g) + e^-, for example Al²⁺(g) → Al³⁺(g) + e^-.

4Always endothermic

• Always endothermic energy must be supplied to overcome the attraction between the electron and the nucleus
• All ionisation energies are therefore positive.

Section 2 of 18Factors affecting ionisation energy

1Nuclear charge

• Nuclear charge more protons attract the electrons more strongly, so ionisation energy increases

2Distance from the nucleus

• Distance from the nucleus attraction falls off rapidly with distance, so outer electrons further away are removed more easily

3Shielding

• Shielding more inner shells and sub-shells shield the outer electron from the nucleus, so it is removed more easily

4Spin-pair repulsion

• Spin-pair repulsion two electrons in one orbital repel each other, so a paired electron is removed more easily

5Down a group the first IE decreases

• Down a group the first IE decreases larger atomic radius and more shielding outweigh the extra nuclear charge
• Group 1 metals have low ionisation energies and the noble gases have very high ones.

Section 3 of 18Successive ionisation energies

Successive ionisation energies

Successive IEs always increase

• Successive IEs always increase each electron is removed from an increasingly positive ion

Large jumps show a change of shell

• Large jumps show a change of shell the next electron comes from a shell closer to the nucleus
• Smaller jumps within a shell show a change of sub-shell, which supports the sub-shell model.

Deducing the group

• Deducing the group the position of the first large jump gives the number of outer electrons
• A huge jump after the first IE places an element in group 1, as in sodium (1s² 2s² 2p⁶ 3s¹).
• A huge jump after the second IE places an element in group 2, as in magnesium.
• A huge jump after the third IE places an element in group 3, as in aluminium.

Use of the data

• Use of the data confirms the electronic configuration and the number of outer-shell electrons

Section 4 of 18Shells, sub-shells and orbitals

• Shells
  • numbered n = 1, 2, 3 and so on, with the energy rising as n increases
  • Shell n holds up to 2n² electrons.

Sub-shell	Orbitals	Maximum electrons
s	1	2
p	3	6
d	5	10
f	7	14

Orbital

Orbital a region of space that holds up to two electrons of opposite spin

• This is the Pauli exclusion principle.

Occurrence

• Occurrence the first shell has only a 1s orbital, p orbitals begin at n = 2 and d orbitals at n = 3

Degenerate orbitals

• Degenerate orbitals orbitals in the same sub-shell, such as 2p_x, 2p_y and 2p_z, have equal energy

Section 5 of 18Shapes of orbitals

s orbitals

• s orbitals spherical
• They get larger as the shell number increases, so 3s is bigger than 1s.

p orbitals

• p orbitals dumbbell-shaped with two lobes
• The three p orbitals lie along the x, y and z axes at right angles to each other.
• The lobes become larger and longer as the shell number increases.

d orbitals

• d orbitals their shapes are not required

Section 6 of 18Filling orbitals

Energy order

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p

Aufbau principle

• Aufbau principle sub-shells are filled in order of increasing energy
• The 4s sub-shell fills before the 3d sub-shell.

Hund's rule

• Hund's rule orbitals of equal energy are filled singly with parallel spins before any pairing
• Three electrons in a p sub-shell occupy p_x, p_y and p_z singly, so nitrogen has three unpaired 2p electrons.

Electrons-in-boxes

• Electrons-in-boxes each box is an orbital and each electron is an arrow, with opposite arrows for opposite spins
• The boxes are arranged in order of increasing energy.

Section 7 of 18Writing electronic configurations

1

Find the number of electrons from the atomic number and adjust for any charge.

2

Fill the sub-shells in order from 1s, putting at most two electrons in each orbital.

3

Fill degenerate orbitals singly first.

4

Write the sub-shells in numerical order, with 3d before 4s.

5

For positive ions, remove electrons from the outermost shell, and from 4s before 3d.

Full configuration

• Full configuration lists every sub-shell from 1s
• Potassium is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.
• Germanium is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p².

Shorthand configuration

• Shorthand configuration the core is replaced by the preceding noble gas in brackets
• Potassium is [Ar] 4s¹.

Exceptions

• Exceptions chromium is [Ar] 3d⁵ 4s¹ and copper is [Ar] 3d¹⁰ 4s¹
• A half-full or full 3d sub-shell is energetically stable.
• Iron is [Ar] 3d⁶ 4s² and Fe³⁺ is [Ar] 3d⁵.

Section 8 of 18Blocks of the Periodic Table and chemical properties

Blocks of the Periodic Table and chemical properties

Blocks

• Blocks an element's block is the sub-shell that holds its outermost electron
• s-block elements are in groups 1 and 2, p-block elements are in groups 3 to 0, and d-block elements are the transition metals.

Periods and groups

• Periods and groups elements in a period have the same number of shells and elements in a group have the same number of outer electrons

Reactions

• Reactions atoms lose or gain electrons to reach a stable noble gas configuration
• Mg²⁺ and O^2- both have the configuration [Ne].

Noble gases

• Noble gases full outer sub-shells and high ionisation energies make them very unreactive

Section 9 of 18More from the specification

Electronic excitation between energy levels

• Electronic excitation between energy levels removing valence electrons requires overcoming electrostatic attraction, promoting charges across quantized electronic energy levels.

Ionisation energy

• Ionisation energy the energy to remove one mole of electrons from one mole of gaseous atoms;
  • it is endothermic.

Subshells

• Subshells s, p, d and f subshells fill in order of energy, giving electronic structures such as calcium 1s2 2s2 2p6 3s2 3p6 4s2.

Principal quantum number

• Principal quantum number shows the main energy level, with higher shells further from the nucleus.

Valence electron

• Valence electron an outer-shell electron involved in bonding, which is lost first in ionisation.

Trends

• Trends successive ionisation energies jump when an electron comes from an inner shell closer to the nucleus.

Subshell

• A set of orbitals of the same type within a shell: s, p, d or f.

Quantum

• A packet of energy.

Valence electron

• An electron in the outermost shell of an atom.

Energy level

• A main shell of fixed energy labelled by principal quantum number n.

Section 10 of 18Ionisation energy definitions

1

First ionisation energy (IE1)

First ionisation energy (IE1) the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

• It is measured in kJ mol^-1 under standard conditions of 298 K and 101 kPa.
• Ca(g) → Ca⁺(g) + e^-, IE1 = +590 kJ mol^-1.

2

Successive ionisation energies

• Successive ionisation energies each further electron is removed from a more positive ion
• Ca⁺(g) → Ca²⁺(g) + e^-, IE2 = 1145 kJ mol^-1.
• Ca²⁺(g) → Ca³⁺(g) + e^-, IE3 = 4912 kJ mol^-1.

Section 11 of 18Factors affecting ionisation energy

Nuclear charge

• Nuclear charge more protons attract electrons more strongly

Distance

• Distance attraction falls rapidly with distance from the nucleus

Shielding

• Shielding inner electrons shield outer electrons, so more shells make removal easier

Spin-pair repulsion

• Spin-pair repulsion two electrons in one orbital repel, so a paired electron is removed more easily

Successive ionisation energies always increase

• Big jumps mark a change of shell and small jumps mark a change of sub-shell.

Section 12 of 18Using successive ionisation energies

Using successive ionisation energies

The largest jump shows the number of outer electrons, and so the group

• Sodium has a huge jump between IE1 and IE2, so it is in group 1.
• Magnesium has a huge jump between IE2 and IE3, so it is in group 2.

Period 2 element

• Period 2 element a largest jump between IE6 and IE7 means six outer electrons
• The element is in group 6, which is oxygen.

Group 1 and noble gases

• Group 1 and noble gases group 1 metals have low IE1 and noble gases have very high IE1,
  • because helium, neon and argon have full shells that make them very unreactive

Section 13 of 18Shells, sub-shells and orbitals

Sub-shell	Orbitals	Maximum electrons	Shape
s	1	2	Spherical, larger with higher n
p	3 (px, py, pz)	6	Dumbbell, at right angles
d	5	10	Not required
f	7	14	Not required

Orbital

• Orbital holds a maximum of two electrons with opposite spins
• This is the Pauli exclusion principle.

Degenerate orbitals

• Degenerate orbitals orbitals in one sub-shell have equal energy
• Electrons occupy separate orbitals with parallel spins first, because same-spin electrons repel less.
• They pair up only when no empty orbital is left.

Examples

• Examples three electrons in a p sub-shell are one in each of p_x, p_y and p_z, and a fourth electron pairs up in one orbital

Section 14 of 18Writing configurations of atoms and ions

Order of filling

• 1s 2s 2p 3s 3p 4s 3d 4p
• The 4s sub-shell fills before the 3d sub-shell.

Potassium

• Potassium full 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ and shorthand [Ar] 4s¹
• Argon accounts for 18 electrons.

Calcium and gallium

• Ca is [Ar] 4s² and Ga is [Ar] 3d¹⁰ 4s² 4p¹, written in numerical order

Exceptions

• Cr is [Ar] 3d⁵ 4s¹ and Cu is [Ar] 3d¹⁰ 4s¹
• Half-full and full d sub-shells are more stable.

Ions

• Ions positive ions lose electrons from the outer sub-shell and negative ions gain electrons there
• Transition metals lose 4s electrons first.
• Ca²⁺ has the same configuration as argon.

Boxes notation

• Boxes notation each box is an orbital, ordered by increasing energy upwards, with arrows for spin

Section 15 of 18Atoms, ions and isotopes

Atoms, ions and isotopes

Atom

• Atom neutral, so the number of protons equals the number of electrons, and the nuclear charge equals the number of protons

Relative isotopic mass

Relative isotopic mass the mass of an isotope relative to 1/12 of a carbon-12 atom

Fragment ions in mass spectrometry

• Fragment ions in mass spectrometry only positively charged fragments are detected
• Propanone, CH₃COCH₃, gives M⁺ at m/z 58, CH₃CO⁺ at 43 and CH₃⁺ at 15.

Section 16 of 18More from the specification

Quantum electron configurations and subshells

• Electrons occupy principal energy levels subdivided into s, p, d,
• and f subshells, filling orbitals in order of increasing energy according to the Aufbau principle.

Hund's rule and Pauli exclusion principle

• Orbitals in a subshell are singly occupied with parallel electron spins before pairing occurs, with a maximum of two electrons per orbital having opposite spins.

Valence electrons and ionization

• The outermost valence electrons dictate chemical reactivity;
  • transition metals lose 4s electrons before 3d electrons during cation formation.

Principal quantum shells and subshells

• Principal quantum shells and subshells electrons occupy discrete principal quantum shells subdivided into s, p, d and f subshells possessing distinct shapes.

Valence electron shielding and ionisation

• Valence electron shielding and ionisation outer valence electron removal requires less energy as inner electron subshells increase electrostatic shielding from the nucleus.

Orbital spin-pair repulsion

• Orbital spin-pair repulsion two electrons occupying the same atomic orbital spin in opposite directions, represented by antiparallel arrows in boxes.

Subshell

• A set of orbitals of the same type within a shell: s, p, d or f.

Quantum

• A packet of energy.

Property

• A feature that describes what a substance is like and what it does.

Valence electron

• An electron in the outermost shell of an atom.

Chemical property

• How a substance reacts; observable only when identity changes.

Energy level

• A main shell of fixed energy labelled by principal quantum number n.

Section 17 of 18Key terms

Term	Meaning
First ionisation energy	The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Orbital	A region of space around the nucleus that holds up to two electrons of opposite spin.
Sub-shell	A group of orbitals of the same type and energy within a shell.
Principal quantum number	The number n of the main energy level (shell) an electron occupies.
Degenerate orbitals	Orbitals of equal energy in the same sub-shell.
Hund's rule	Orbitals of equal energy are filled singly before electrons pair up.
Shielding	Repulsion of outer electrons by inner electrons, reducing the nuclear attraction.

Term	Meaning
First ionisation energy	The energy required to remove one mole of electrons from one mole of gaseous atoms to form 1+ ions.
Shielding	Reduction of the nuclear attraction on outer electrons by inner electrons.
Orbital	A region of space holding up to two electrons of opposite spin.
Pauli exclusion principle	An orbital holds at most two electrons, with opposite spins.
Relative isotopic mass	The mass of an isotope relative to 1/12 of carbon-12.

Section 18 of 18Exam technique

1

Writing an ionisation energy as exothermic — all ionisation energies are endothermic and positive

2

Forgetting state symbols in ionisation equations — the atoms and ions must be gaseous

3

Removing 3d electrons before 4s when forming transition metal ions — 4s electrons are lost first

4

Giving chromium as [Ar] 3d⁴ 4s² or copper as [Ar] 3d⁹ 4s² — they are [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹

5

Pairing electrons in a p sub-shell before each orbital has one — Hund's rule requires single occupancy first

1

Saying 4s electrons are removed last from transition metal ions — they are removed first

2

Saying paired electrons have the same spin — they have opposite spins

3

Forgetting that ionisation energies refer to gaseous atoms and ions

4

Quoting units as kJ instead of kJ mol⁻¹