Structure (Edexcel A Level Chemistry): Flashcards

A temporary dipole in one molecule can induce a dipole in a neighbouring molecule, creating an attractive force between them.

The three types of intermolecular forces are:

1. Induced dipole-dipole forces (London dispersion forces)

2. Permanent dipole-dipole forces

3. Hydrogen bonding

True.

Intramolecular forces (such as covalent bonds) are significantly stronger than intermolecular forces. For example, a hydrogen bond is approximately one tenth the strength of a covalent bond.

Permanent dipole-dipole forces are intermolecular attractions between neighbouring polar molecules, where the δ⁺ end of one molecule attracts the δ⁻ end of another.

Propanone is a polar molecule with permanent dipole-dipole forces, while butane is non-polar and only has induced dipole-dipole forces. Permanent dipole forces are stronger, so more energy is needed to separate propanone molecules.

Hydrogen bonding requires a hydrogen atom covalently bonded to a highly electronegative atom such as oxygen, nitrogen or fluorine.

False.

Induced dipole-dipole forces exist between all atoms and molecules, whether polar or non-polar. They arise from the movement of electron clouds, which occurs in all atoms.

Induced dipole-dipole forces are weak intermolecular forces that arise from temporary dipoles caused by the constantly moving electron cloud. They exist between all atoms and molecules, including non-polar ones.

Hydrogen bonding requires: (1) a hydrogen atom covalently bonded to a highly electronegative atom (O, N or F), and (2) a lone pair on an O, N or F atom in a neighbouring molecule available to accept the bond.

Alcohols, amines and carboxylic acids can all form hydrogen bonds because each contains a bond between hydrogen and a highly electronegative atom.

A hydrogen bond is a strong type of permanent dipole-dipole interaction that forms between a δ⁺ hydrogen atom (bonded to O, N or F) and a lone pair on an O, N or F atom in a neighbouring molecule.

False.

Water can form up to four hydrogen bonds per molecule — two as a hydrogen bond donor (via its two O-H bonds) and two as a hydrogen bond acceptor (via its two lone pairs). Ammonia has only one lone pair, limiting its ability to accept hydrogen bonds, so it forms fewer hydrogen bonds per molecule than water overall.

Water has an anomalously high boiling point because of the strong hydrogen bonds between molecules. A large amount of energy is needed to break these bonds when boiling, far more than expected for a molecule of such small size.

True.

In ice, water molecules form an open lattice structure held by hydrogen bonds, placing them slightly further apart than in liquid water. This makes ice approximately 9% less dense than liquid water.

Ice floats because it is less dense than liquid water. In ice, water molecules are arranged in an open lattice held by hydrogen bonds, spacing them further apart than in the liquid state and giving ice a lower density.

In ice, water molecules form an open lattice structure held together by hydrogen bonds, making it less dense than liquid water.

Surface tension is the ability of a liquid surface to resist external forces. In water, it arises from hydrogen bonds pulling surface molecules inward and downward, compressing the surface layer and making it harder to break.

Branching reduces the surface area of a molecule, which decreases London dispersion forces between molecules. This lowers the boiling point compared to unbranched isomers with the same number of electrons.

Branching reduces the surface area of a molecule, which weakens the strength of London dispersion forces and lowers the boiling point.

London dispersion forces are weak intermolecular attractions caused by temporary dipoles that arise from the uneven distribution of electrons. Their strength increases with the number of electrons in the molecule.

True.

More electrons increase the likelihood of temporary dipoles forming and increase their magnitude. This strengthens London dispersion forces and raises the melting and boiling points of the substance.

Alcohols can form hydrogen bonds between molecules due to their O-H groups, which are much stronger than the London dispersion forces found in alkanes. More energy is needed to separate alcohol molecules, raising their boiling points.

False.

HF has a higher boiling point than HCl due to hydrogen bonding between HF molecules, not because of more electrons. HCl actually has more electrons than HF but lacks the strongly polar H-F bond needed for hydrogen bonding.

'Like dissolves like' means that polar substances dissolve best in polar solvents and non-polar substances dissolve best in non-polar solvents. Compatible intermolecular forces between solute and solvent are required for dissolution.

Polar covalent substances generally dissolve in polar solvents because they can form dipole-dipole interactions or hydrogen bonds with the solvent.

Ethanol is small enough that its polar O-H group can form hydrogen bonds with water, enabling dissolution. In hexanol, the large non-polar hydrocarbon chain outweighs the polar O-H group, reducing compatibility with the polar water solvent.

Metallic bonding is the strong electrostatic attraction between a lattice of positive metal cations and a surrounding 'sea' of delocalised electrons.

Metal atoms become positively charged ions because their outer-shell electrons become delocalised, leaving the atoms with a net positive charge.

In a metallic lattice, the outer-shell electrons are free to move and are described as delocalised. The metal atoms that remain are positively charged.

False.

Delocalised electrons are not bound to any individual atom. They are free to move throughout the entire metallic lattice.

Delocalised electrons are outer-shell electrons that are no longer bound to a specific atom and are free to move throughout the metallic lattice.

The positive metal ions are held in their regular arrangement by the strong electrostatic attraction from the surrounding 'sea' of negatively charged delocalised electrons.

Metal atoms are arranged in a lattice structure. The strong forces between the positive metal centres and the sea of delocalised electrons are called metallic bonding.

True.

Metal atoms are tightly packed in a lattice structure, with the positive metal ions held in place by the surrounding sea of delocalised electrons.

The metallic lattice is held together by the strong electrostatic attraction between the positively charged metal cations and the sea of negatively charged delocalised electrons.

A lattice is a regular, repeating three-dimensional arrangement of ions or atoms extending throughout a solid.

The three main types of giant lattice structure are giant ionic, giant metallic and giant covalent lattices.

Giant ionic lattice is a regular, repeating three-dimensional arrangement of cations and anions held together by electrostatic attraction acting in all directions.

In sodium chloride, each Na⁺ ion is surrounded by 6 Cl^- ions and each Cl^- ion is surrounded by 6 Na⁺ ions. This is called 6:6 coordination.

False.

A giant ionic lattice is overall electrically neutral because the positive and negative charges of the cations and anions cancel each other out.

Giant covalent lattice is an extended three-dimensional network of atoms bonded throughout by strong covalent bonds, with no discrete molecules present.

A giant metallic lattice contains positive metal cations arranged in regular close-packed layers, surrounded by a sea of delocalised electrons. The bonding is the strong electrostatic attraction between the cations and the delocalised electrons.

Giant lattice structures do not contain discrete molecules. For such substances, use relative formula mass rather than relative molecular mass.

True.

Diamond, graphite and silicon(IV) oxide (SiO₂) are all giant covalent lattices in which atoms are bonded throughout by strong covalent bonds.

A greater charge on the metal cation increases the strength of metallic bonding, because the electrostatic attraction between the cations and the sea of delocalised electrons is stronger.

Simple molecules are small covalent structures in which atoms are bonded by shared electrons, held together in the bulk by weak intermolecular forces rather than strong covalent bonds between molecules.

Examples of simple molecular compounds include H₂O, NH₃, CH₄ and HCl. Other examples are H₂, Cl₂, O₂ and N₂.

Iodine exists as a simple molecule but forms a crystalline solid held together by weak London dispersion forces.

Allotropes are different structural forms of the same element. Diamond, graphite and buckminsterfullerene are all allotropes of carbon.

False.

Diamond cannot conduct electricity because all four outer electrons on every carbon atom are used in covalent bonds, leaving no freely moving electrons available to carry charge.

In graphite, each carbon atom forms three covalent bonds, leaving one free electron per atom. These delocalised electrons can move along the carbon layers when a voltage is applied, allowing graphite to conduct electricity.

In graphite, carbon layers are held together by weak intermolecular forces, which means layers can slide over each other, making graphite soft and slippery.

Each carbon atom in diamond forms four covalent bonds with neighbouring carbon atoms, creating a tetrahedral arrangement throughout the giant covalent lattice.

True.

Graphene consists of a single layer of carbon atoms covalently bonded in a continuous hexagonal arrangement, making it essentially a two-dimensional molecule.

A giant covalent lattice is an extended three-dimensional network of atoms bonded throughout by strong covalent bonds, requiring a large amount of energy to break and resulting in very high melting and boiling points.

A giant ionic structure typically shows a high melting point, solubility in water and the ability to conduct electricity when molten or dissolved, but not in the solid state.

To predict the structure and bonding of a substance means to use its physical properties (such as melting point, electrical conductivity and solubility) to determine whether it has a giant ionic, giant metallic, giant covalent or simple molecular structure.

False.

A low melting point and low electrical conductivity are characteristic of a simple molecular structure, where only weak intermolecular forces must be overcome to melt the substance.

A giant metallic structure has a very high melting point and conducts electricity in both the solid and liquid state, due to the presence of delocalised electrons that are free to move throughout the lattice.

A substance with a very high melting point, poor conductivity as both a solid and a liquid, and insolubility in water is most likely a giant covalent structure.

Electrical conductivity is the ability of a substance to allow the flow of electric charge. It is determined by the presence of delocalised electrons (in metals and graphite) or mobile ions (in molten or dissolved ionic compounds).

In the solid state, the ions are held in fixed positions in the lattice and cannot move. When molten, the ions are free to move and carry charge, allowing the compound to conduct electricity.

True.

All giant structures (ionic, metallic and covalent) have high melting points because a large amount of energy is required to overcome the strong bonds or electrostatic attractions holding the lattice together.

Both have very high melting points, but a giant metallic structure conducts electricity in both solid and liquid states due to delocalised electrons, whereas a giant covalent structure (except graphite) does not conduct electricity in either state.