Bonding and Structure

Electron configuration (first 20 elements)

The number of electrons in each shell is 2-8-8-2 for the first 20 elements. The general rule is 2n². Up to potassium the fourth shell begins filling (2.8.8.1 for potassium, atomic number 19). The further an electron is from the nucleus, the less attraction it feels, so it needs more energy to stay.

Atomic number and electron configuration

The atomic number equals the number of electrons in a neutral atom, which allows you to write its electron configuration (e.g. 2.8.8).

Group number and outer shell electrons

The number of electrons in the outer shell can be determined by the group number. Every group has the same outer valency and every period has the same number of electron shells (energy levels).

Ion definition

An ion is a charged atom or group of atoms.

Anions and cations

If atoms gain electrons they form an anion (negative). If atoms lose electrons they form a cation (positive). Losing electrons requires energy, gaining electrons releases energy.

Stable ions and noble gas configurations

Stable ions of the first 20 elements usually have noble gas electron configurations.

Metal vs non-metal ion formation

Metals lose electrons to form positive ions, while non-metals gain electrons to form negative ions (except hydrogen).

Ion charge and valency

The charge of an ion is the opposite of the valency.

Elements that rarely form ions

Carbon, silicon and the noble gases cannot easily gain or lose electrons because it requires too much energy.

Formation of ions

Ions form because as you move to the right on the periodic table there is stronger attraction to the nucleus, so atoms are more likely to gain electrons.

Electron transfer in metal–non-metal reactions

When metals react with non-metals, electrons are transferred from the metal outer shell to the non-metal outer shell so each ion achieves a full outer shell (noble gas configuration).

Ionic compounds

Ionic compounds are made up of ions.

Compound definition

A compound is a pure substance consisting of two or more elements chemically bonded in a fixed ratio by mass.

Example ionic compound

Sodium chloride is an ionic compound made up of ions. It is not a mixture of sodium and chlorine otherwise it would show properties of both elements.

Binary ionic compounds

Binary ionic compounds consist of oppositely charged ions of two elements and have no overall charge.

Polyatomic ion definition

A polyatomic ion is a group of two or more atoms that carries a charge.

Ammonium ion

NH₄⁺, combining power (charge) 1.

Hydroxide ion

OH⁻, combining power (charge) 1.

Ethanoate / acetate ion

CH₃COO⁻, combining power (charge) 1.

Nitrate ion

NO₃⁻, combining power (charge) 1.

Hydrogencarbonate ion

HCO₃⁻, combining power (charge) 1.

Carbonate ion

CO₃²⁻, combining power (charge) 2.

Sulfate ion

SO₄²⁻, combining power (charge) 2.

Phosphate ion

PO₄³⁻, combining power (charge) 3.

Fixed metal charges

Zn always forms 2⁺ and Ag always forms 1⁺ ions.

Binary compound naming

The suffix “-ide” indicates a binary compound (e.g. nitride, sulfide). The metal is written first and positive, the non-metal second and negative.

Variable charges in metals

Some metals including transition metals can have variable combining powers (charges).

Naming variable-charge metals

The charge is shown with Roman numerals, e.g. Copper (II) chloride represents CuCl₂.

Ionic bonding

Ionic bonding is the electrostatic attraction between positive and negative ions.

Ionic compound structure

Ionic compounds contain cations and anions arranged in a regular three-dimensional network.

Electrostatic forces in ionic compounds

The attraction between oppositely charged ions are electrostatic forces known as ionic bonds.

Ionic lattice

Ionic compounds form infinite lattice networks with no fixed number of ions, so they are not molecules.

Ionic lattice arrangement

Ionic compounds consist of alternating cations and anions forming a lattice that maximises attraction.

Melting point of ionic compounds

Ionic compounds are solid at room temperature and have high melting points because strong electrostatic attractions require a lot of energy to break.

Brittleness of ionic compounds

Ionic compounds are brittle and break easily when external force is applied.

Reason for brittleness

When layers shift, like charges line up and repel, causing the lattice to fracture.

Conductivity of ionic compounds

Ionic compounds conduct electricity when molten or dissolved in water because ions are free to move and transfer charge.

Solubility of ionic compounds

Most ionic compounds are soluble in water.

Conductivity observation

Solid ionic substances do not conduct electricity, but molten ionic substances do because ions can move.

Solubility rule: Group I and ammonium

All Group I and ammonium salts are soluble.

Solubility rule: nitrates and acetates

All nitrates and acetates are soluble.

Solubility rule: halides

All chlorides, bromides and iodides are soluble except Ag and Pb.

Solubility rule: sulfates

All sulfates are soluble except Ba, Pb, and slightly Ag and Ca.

Solubility rule: hydroxides

Most hydroxides are insoluble except Group I metals and Ca, Sr, Ba, Ra.

Solubility rule: carbonates

All carbonates are insoluble except Group I.

General solubility pattern

Sodium compounds are generally soluble while lead compounds are generally insoluble.

Precipitation reaction

A precipitation reaction occurs when an insoluble compound forms when two aqueous solutions react.

Steps to predict precipitation

Write the reactants, swap partners, determine solubility and balance the equation.

NVR meaning

NVR means “no visible reaction”, so no chemical equation is needed.

Word and balanced equations

Chemical reactions can be written as word equations or balanced chemical equations.

Net ionic equation

The net ionic equation shows only the particles directly involved in the reaction.

Command term: Identify

Recognise and name.

Command term: Describe

Provide characteristics and features.

Command term: Explain

Relate cause and effect and explain why/how (Structure → Bonding → Properties).

Example ionic explanation

Ionic lattice made of alternating positive and negative ions held by strong electrostatic attractions called ionic bonds, requiring large energy to melt.

Molecule definition

A molecule is a finite structure with a fixed number of non-metal atoms of specific elements.

Examples of molecules

H₂O, NH₃ (ammonia), CH₄ (methane).

Covalent bond definition

A covalent bond is a shared pair of electrons between two atoms.

Ionic vs covalent bonding

Ionic bonds transfer electrons while covalent bonds share electrons.

Covalent naming prefixes

Mono, di, tri, tetra, penta, hexa indicate number of atoms. “Mono” is usually omitted for the first element.

Covalent compound characteristics

Covalent compounds consist of two non-metals, do not form ions and have no charge.

Order in covalent formulas

The element further right and higher in the periodic table is usually written second.

Electrostatic attraction in covalent bonds

Covalent bonding involves attraction between the shared electrons and the nucleus of each atom.

Lewis dot diagrams

Lewis dot diagrams show only outer shell (valence) electrons using atomic symbols and dots.

Octet rule

Atoms tend to share electrons so each atom has eight electrons in its outer shell.

Molecular substances

Molecular substances consist of discrete particles called molecules.

Intermolecular forces

Forces between molecules are weak and called intermolecular forces.

Cause of intermolecular forces

They arise due to dipoles (charge differences) and are electrostatic but much weaker than ionic bonds.

Melting point of covalent compounds

Covalent molecular substances have low melting points because intermolecular forces are weak.

Conductivity of covalent compounds

Covalent substances have low conductivity because they have no mobile charged particles.

Mechanical properties of covalent compounds

Covalent molecular substances are soft and brittle due to weak intermolecular forces.

Metallic bonding structure

Metals consist of positive metal ions arranged in a lattice surrounded by a sea of delocalised electrons.

Delocalised electrons in metals

Delocalised electrons move freely throughout the lattice and are not tied to one atom.

Metallic bonding force

Metal ions are held together by electrostatic attraction between positive ions and delocalised electrons.

Melting point of metals

Metals have high melting points because strong electrostatic forces require large amounts of energy to break.

Electrical conductivity of metals

Metals conduct electricity because delocalised electrons move freely.

Malleability and ductility of metals

Metals are malleable and ductile because electrons act like glue allowing ion layers to slide without breaking.

Practical comparison of bonding types

Experiments can compare properties of ionic, covalent and metallic substances.

Metal example: aluminium

Used in drink cans, foil, aircraft and boats because of low density, conductivity, malleability, ductility and reflective properties.

Ionic example: sodium chloride

Used for food flavouring and road salting; prevents freezing and is soluble in water.

Covalent example: paraffin wax

Low melting point, white and odourless; used in candle making, lubrication, crayons, candy coating and chewing gum additive.

Ionic structure summary

Rigid 3D lattice of alternating cations and anions.

Covalent structure summary

Discrete molecules made of small groups of covalently bonded atoms.

Metallic structure summary

3D lattice of positive metal ions in a sea of delocalised electrons.

Bonding comparison

Ionic: strong ionic bonds. Covalent: strong covalent bonds with weak intermolecular forces. Metallic: electrostatic attraction between metal ions and electrons.

Melting point comparison

Ionic high, covalent low, metallic high.

Conductivity comparison

Ionic solid poor but molten/aqueous good; covalent poor; metallic good.

Malleability comparison

Ionic brittle; covalent soft and brittle; metallic malleable.

Protons and neutrons density

Protons and neutrons are very dense particles in the nucleus.

Isotopes definition

Isotopes are atoms of the same element with different numbers of neutrons and different mass numbers.

Neutral atom rule

If charge is zero, the number of electrons equals the number of protons.

Salt definition

A salt is an ionic compound formed from the neutralisation of an acid and a base.