Ionic & Metallic Bonding & Structure — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

Section 1 of 20Giant ionic lattices and their properties

Giant ionic lattices and their properties

Ionic bond

• Ionic bond the electrostatic attraction between a positive metal ion (cation) and a negative non-metal ion (anion), acting in all directions.

Lattice

• Lattice ions form a giant regular repeating lattice of alternating positive and negative ions.
• The lattice type depends on the relative sizes of the ions, with NaCl and MgO cubic.
• Electron density maps show the density falls to zero between ions.

Melting and boiling points

• Melting and boiling points high, because a lot of energy is needed to overcome the strong electrostatic attraction in the lattice.
• Higher ionic charge gives a higher melting point: Na₂O melts at 1405 K and CaO at 2845 K.
• Size is not the factor here, because Na⁺ and Ca²⁺ are similar in size.

Electrical conductivity

• Electrical conductivity conduct when molten or in solution because ions are free to move.
• Solids do not conduct because the ions are fixed in the lattice.

Solubility

• Solubility many dissolve in polar solvents such as water.
• Polar molecules disrupt the lattice and surround each ion, with the + end towards anions and the - end towards cations.
• Ion-dipole attractions hydrate the ions.
• Solubility depends on the lattice attraction versus the ion-solvent attraction.
• Greater ionic charge generally means lower solubility, as in NaCl (356.9 g dm^-3) and CaCl₂ (74.4 g dm^-3), but there are many exceptions.

Section 2 of 20Evidence for the existence of ions

Electrolysis

• Electrolysis positive ions move to the negative electrode and negative ions to the positive electrode.

Copper(II) chromate(VI)

• Copper(II) chromate(VI) the olive green solution contains blue Cu²⁺ and yellow CrO₄^2- ions.
• On electrolysis blue colour appears at the negative electrode.
• Yellow colour appears at the positive electrode.

Section 3 of 20Forming ions and ionic dot-and-cross diagrams

Electron transfer

• Electron transfer metals lose valence electrons to form cations and non-metals (right) gain electrons to form anions.
  • (left of the Periodic Table)
• The ions usually have a full outer shell like a noble gas, as K⁺ and Cl^- are both [2,8,8] like argon.

Ionic bond

• Ionic bond the electrostatic attraction between the ions acts in all directions and is strong.
  • (non-directional)
• The lattice is regular, with positive and negative charges cancelling, so the whole lattice is electrically neutral.

Dot-and-cross diagram

• Dot-and-cross diagram shows only outer electrons, with dots for one atom and crosses for the other.
• Show the transfer so each ion has a full outer shell.
• Draw each ion in square brackets with the charge at the top right.
• Use enough atoms to make the total charge zero.

Section 4 of 20Examples of ionic formulae

Calcium fluoride

• Ca loses 2 electrons (Ca²⁺) and each F gains 1 (F^-), so CaF₂ needs two fluoride ions.

Lithium nitride

• Li loses 1 (Li⁺) and N gains 3 (N^3-), so Li₃N has three lithium ions.

Aluminium oxide

• Al loses 3 (Al³⁺) and O gains 2 (O^2-), so Al₂O₃ has two Al³⁺ and three O^2-.

Section 5 of 20Ionic radius, charge density and bond strength

Atomic radius

• Atomic radius the distance from the nucleus to the outermost electron, found by halving the distance between the nuclei of two atoms of the same element.
• It decreases across a period because the nuclear charge increases while shielding stays the same.

Ionic radius

• Ionic radius cations are smaller than their atoms and anions are larger.
• Positive ionsfewer electrons feel the same nuclear charge, so a greater charge gives a smaller ion.
• Negative ionsextra electrons are held more weakly, so a greater negative charge gives a larger ion.

Isoelectronic ions

• N^3-, O^2-, F^-, Na⁺, Mg²⁺ and Al³⁺ all have 1s² 2s² 2p⁶.
• More protons pull the electrons closer, so radii fall from N^3- (0.171 nm) to Al³⁺ (0.054 nm).

Charge density

• Charge density charge divided by size.
• Higher charges and smaller radii give higher charge density, stronger ionic bonds and higher melting points.

Section 6 of 20Polarisation of ions

1Perfect ionic model

• Perfect ionic model assumes spherical ions and purely electrostatic attraction.

2Polarisation

• Polarisation a cation attracts electrons from the anion and distorts its electron density.
• It introduces covalent character, so the measured lattice energy differs from the theoretical value.

3Polarising power of cations

• Polarising power of cations increases with charge density.
• K⁺ is weaker than Li⁺ because of its larger radius.
• Na⁺ is weaker than Mg²⁺ because of its smaller charge.
• Small, highly charged cations are the most polarising.

4Polarisability of anions

• Polarisability of anions increases with ionic radius.
• Br^- is more polarisable than Cl^-.

Section 7 of 20Metallic bonding and the lattice

Lattice

• Lattice metal atoms are tightly packed in a giant lattice, often hexagonal layers or a cubic arrangement.

Delocalised electrons

• Delocalised electrons outer-shell electrons are free to move throughout the structure and are not bound to one atom.

Metal ions

• Metal ions the atoms become positive ions suspended in a sea of delocalised electrons.

Metallic bond

• Metallic bond the electrostatic attraction between the positive ions and the delocalised electrons.
• It acts throughout the lattice and is non-directional.
• It keeps the lattice stable although the positive ions repel one another.

Diagram

• Diagram draw layers of close-packed positive ions in a regular arrangement and label the metal ions and delocalised electrons.

Section 8 of 20Strength of metallic bonding

1Stronger metallic bonding

• Stronger metallic bonding increases with each of these.
• More delocalised electrons per atom.
• A higher positive charge on the metal ions.
• Smaller metal ions.

2Example

• Example melting point rises Na < Mg < Al.

Section 9 of 20Properties of metals

Property	Explanation
High melting and boiling points	Strong attraction between the ions and delocalised electrons in a giant lattice needs much energy to overcome
Good electrical conductors	Delocalised electrons move and carry charge, while the ions stay fixed
Good thermal conductors	Mobile electrons transfer kinetic energy, and vibrating closely packed ions also pass on energy
Malleable and ductile	Layers of ions slide over each other and the delocalised electrons move with them, keeping the attraction so the bonding is not broken

• Giant structure
  • melting a metal overcomes metallic bonding, not separate molecules.

Section 10 of 20Alloys

Alloy

• Alloy formed when atoms of another element, such as carbon in iron, are added to a metal lattice.

Properties

• Properties the different-sized atoms stop the layers of metal ions sliding over each other easily, so alloys are stronger and harder than pure metals.

Section 11 of 20Forming ions

Forming ions

Metals lie on the left of the Periodic Table and non-metals on the right.

Ionic bonding involves transfer of electrons from a metal atom to a non-metal atom.

• Metals lose valence electrons to form positive cations.
• Non-metals gain electrons to form negative anions.
• Both ions usually end with a full outer shell, the same as a noble gas.

K⁺ and Cl⁻ both have the argon configuration [2,8,8].

Section 12 of 20The ionic bond and lattice

Ionic bond

• Ionic bond the electrostatic attraction between oppositely charged ions, acting in all directions.
  • (non-directional)

Each ion is attracted to all the oppositely charged ions around it.

• The ions form a regular repeating lattice in which positive and negative charges cancel.
• The lattice is overall electrically neutral.

The attraction is very strong, so a lot of energy is needed to overcome it and ionic compounds have high melting points.

Section 13 of 20Ionic dot-and-cross diagrams

1

Draw only the outer-shell electrons of each atom, as dots for one atom and crosses for the other.

2

Show the electrons transferred so that each ion has a full outer shell.

3

Draw each ion in square brackets with its charge at the top right.

4

Use enough atoms of each element so the total charge is zero.

CaF₂

• CaF₂ calcium (Group 2) loses 2 electrons to form Ca²⁺, and each fluorine (Group 7) gains 1 electron to form F⁻, so 2 fluoride ions are needed.

Li₃N

• Li₃N lithium (Group 1) loses 1 electron to form Li⁺, and nitrogen (Group 5) gains 3 electrons to form N³⁻, so 3 lithium ions are needed.

Al₂O₃

• Al₂O₃ aluminium (Group 3) forms Al³⁺ and oxygen (Group 6) forms O²⁻, so 2 Al³⁺ balance 3 O²⁻.

Section 14 of 20Atomic and ionic radius

Atomic radius

• Atomic radius half the distance between the nuclei of two atoms of the same element.
• It is called the metallic radius in metals and the covalent radius in non-metals.

Atomic radius decreases across a period.

• Nuclear charge and the number of electrons rise by one for each element.
• The number of shells, and so the shielding, stays the same.
• The nucleus therefore attracts the electrons more strongly and pulls them closer.

Positive ions are smaller than their atoms.

• Nuclear charge is unchanged but fewer electrons are attracted more strongly.
• The greater the positive charge, the smaller the ion.

Negative ions are larger than their atoms.

• The outer electrons are further from the nucleus and held more weakly.
• The greater the negative charge, the larger the ion.

Isoelectronic ions have the same configuration (1s2 2s2 2p6) but different proton numbers.

• The series N³⁻, O²⁻, F⁻, Na⁺, Mg²⁺, Al³⁺ has 7 to 13 protons.
• As the number of protons rises, the electrons are pulled closer, so the radius falls from 0.171 nm for N³⁻ to 0.054 nm for Al³⁺.

Section 15 of 20Ionic bond strength

Ionic bond strength depends on both the charges and the sizes of the ions.

Charge density

• Charge density ionic charge relative to ionic size.

Higher charge and smaller radius give a higher charge density and a stronger bond.

Stronger ionic bonds give higher melting and boiling points and lower volatility.

• Compounds of 2+/2− ions are harder and melt at higher temperatures than those of 1+/1− ions.

Section 16 of 20Polarisation of ions

1

The theoretical lattice energy assumes a perfect ionic model of 100% spherical ions with purely electrostatic attractions.

2

The measured lattice energy can differ from the theoretical value because of covalent character.

• The cation attracts electrons from the anion and distorts its electron density.
• This is polarisation of the anion.

3Polarising power of a cation

• Polarising power of a cation increases with charge density.
• Li⁺ has greater polarising power than K⁺ because it is smaller.
• Mg²⁺ has greater polarising power than Na⁺ because its charge is larger.
• A small, highly charged cation has the greatest polarising power.

4Polarisability of an anion

• Polarisability of an anion increases with ionic radius.
• Br⁻ is more polarisable than Cl⁻ because it is larger.

Section 17 of 20Metallic bonding and lattices

1

In a metal the atoms are tightly packed in a lattice, often hexagonal layers or a cubic arrangement.

2

The outer-shell electrons are delocalised: they move freely through the structure and belong to no single atom.

3

The atoms become positive ions, held in a sea of delocalised electrons by strong electrostatic attraction.

4

Metallic bonding is stronger with:

• more delocalised electrons per atom,
• a higher positive charge on the ions,
• smaller metal ions.

5High melting and boiling points

Properties

• High melting and boiling points
  • a lot of energy is needed to overcome the strong attraction between ions and electrons.
• Good electrical conductors
  • delocalised electrons move and carry charge.
• Good thermal conductors
  • electrons transfer kinetic energy and ions vibrate.
• Malleable and ductile
  • layers of ions slide without breaking the bonding, because the electrons are not tied to any one ion.

6Alloys

• Alloys adding atoms of another element stops the layers of ions sliding easily, so alloys are stronger than the pure metal.
  • (e.g. carbon in iron)

Section 18 of 20More from the specification

More from the specification

Giant Ionic Lattices and Electrostatic Forces

• Ionic compounds form giant three-dimensional crystalline lattices held together by strong omnidirectional electrostatic forces of attraction between oppositely charged cations and anions.

Trends in Ionic Radii

• Across an isoelectronic series, ionic radii decrease as nuclear charge increases, pulling electrons closer;
  • cations are smaller than their parent atoms,
• while anions are larger due to added electron-electron repulsion.

Lattice Enthalpy and Melting Point

• Lattice dissociation enthalpy increases with higher ionic charges and smaller ionic radii, yielding high melting points that require substantial thermal energy to overcome electrostatic bonds.

Electrical Conductivity of Ionic Compounds

• Solid ionic compounds cannot conduct electricity because ions are fixed in lattice positions;
  • once melted into a liquid or dissolved in aqueous solution, mobile ions act as mobile charge carriers.

Giant Ionic Lattices and Electrostatic Forces

• Ionic compounds form giant three-dimensional crystalline lattices held together by strong omnidirectional electrostatic forces of attraction between oppositely charged cations and anions.

Trends in Ionic Radii

• Across an isoelectronic series, ionic radii decrease as nuclear charge increases, pulling electrons closer;
  • cations are smaller than their parent atoms,
• while anions are larger due to added electron-electron repulsion.

Lattice Enthalpy and Melting Point

• Lattice dissociation enthalpy increases with higher ionic charges and smaller ionic radii, yielding high melting points that require substantial thermal energy to overcome electrostatic bonds.

Electrical Conductivity of Ionic Compounds

• Solid ionic compounds cannot conduct electricity because ions are fixed in lattice positions;
  • once melted into a liquid or dissolved in aqueous solution, mobile ions act as mobile charge carriers.

Giant Ionic Lattices and Electrostatic Forces

• Ionic compounds form giant three-dimensional crystalline lattices held together by strong omnidirectional electrostatic forces of attraction between oppositely charged cations and anions.

Trends in Ionic Radii

• Across an isoelectronic series, ionic radii decrease as nuclear charge increases, pulling electrons closer;
  • cations are smaller than their parent atoms,
• while anions are larger due to added electron-electron repulsion.

Lattice Enthalpy and Melting Point

• Lattice dissociation enthalpy increases with higher ionic charges and smaller ionic radii, yielding high melting points that require substantial thermal energy to overcome electrostatic bonds.

Electrical Conductivity of Ionic Compounds

• Solid ionic compounds cannot conduct electricity because ions are fixed in lattice positions;
  • once melted into a liquid or dissolved in aqueous solution, mobile ions act as mobile charge carriers.

Electrostatic force

• Force of attraction between opposite charges.

Section 19 of 20Key terms

Term	Meaning
Cation	A positively charged ion.
Anion	A negatively charged ion.
Charge density	The charge of an ion divided by its size.
Polarisation	Distortion of the electron cloud of an anion by a nearby cation.
Delocalised electrons	Outer electrons free to move throughout a metal lattice.
Metallic bond	The electrostatic attraction between positive metal ions and delocalised electrons.
Isoelectronic ions	Ions with the same electron configuration but different numbers of protons.
Alloy	A metal mixed with atoms of another element, making it stronger than the pure metal.

Term	Meaning
Cation	A positively charged ion.
Anion	A negatively charged ion.
Charge density	The charge of an ion relative to its size.
Isoelectronic ions	Ions that have the same electron configuration.
Polarisation	Distortion of the electron cloud of an anion by a nearby cation.
Delocalised electrons	Outer electrons free to move throughout a metal lattice.
Alloy	A mixture of a metal with other elements, which is stronger than the pure metal.

Section 20 of 20Exam technique

1

Drawing ionic compounds with shared electron pairs — electrons are transferred.

2

Omitting square brackets and charges on ions in dot-and-cross diagrams.

3

Saying metal ions carry the current — the delocalised electrons carry the charge.

4

Saying ionic solids conduct because ions are present — the ions must be free to move (molten or dissolved).

5

Describing melting of a metal as breaking molecules apart — metallic bonding in a giant lattice is overcome.

1

Drawing ionic compounds with shared electron pairs — electrons are transferred, not shared.

2

Omitting square brackets and charges on ions in dot-and-cross diagrams — each ion needs both.

3

Saying metal ions move in an electric current — the delocalised electrons carry the charge.

4

Thinking the size of the cation does not matter for bond strength — smaller ions give higher charge density and stronger bonds.