Revision notes · Bonding, structure, and the properties of matter
How bonding and structure are related to the properties of
The three states of matter4.2.2.1
Definition: The three states of matter are solid, liquid and gas — the state a substance is in depends on the arrangement, movement and energy of its particles.
| Solid | Liquid | Gas | |
|---|---|---|---|
| Arrangement | Fixed, regular pattern | Close together, random | Far apart, random |
| Movement | Vibrate in place | Move around each other | Move quickly in all directions |
| Energy (relative) | Lowest | Middle | Highest |
- •Changes of state (melting, freezing, boiling, condensing) are physical changes — no new substance is made, and they are reversible.
- •The amount of energy needed to change state depends on the strength of the forces between the particles — stronger forces need more energy.
- •This simple particle model has limitations: it shows particles as solid, inelastic spheres with no forces between them — in reality there are forces between particles, and they aren't solid spheres.
⚠️ Common mistake: Thinking a change of state is a chemical change — it's always physical: the particles themselves don't change, just their arrangement and energy.
🧠 Remember: State changes are physical and reversible — same substance, different arrangement.
State symbols4.2.2.2
Definition: State symbols show the physical state of each substance in a chemical equation: (s) solid, (l) liquid, (g) gas, (aq) aqueous (dissolved in water).
- •State symbols are written immediately after each formula in a balanced equation.
- •(aq) specifically means the substance is dissolved in water, not just in liquid form.
⚠️ Common mistake: Using (l) for a substance dissolved in water — dissolved substances use (aq), not (l), even though the solution itself is liquid.
🧠 Remember: (s) solid, (l) liquid, (g) gas, (aq) = dissolved in water.
Properties of ionic compounds4.2.2.3
The properties of an ionic compound come directly from its giant ionic lattice structure.
- •High melting and boiling points — a lot of energy is needed to break the many strong electrostatic forces throughout the lattice.
- •Conduct electricity when molten or dissolved, because the ions are then free to move and carry charge; do not conduct as a solid.
- •Often soluble in water, since water molecules can separate the charged ions from the lattice.
⚠️ Common mistake: Saying ionic compounds don't conduct electricity at all — they do, but only once the ions are free to move (molten or dissolved), not as a solid.
🧠 Remember: Ionic: high melting point, conducts only when molten or dissolved.
Properties of small molecules4.2.2.4
Small covalent molecules like water and carbon dioxide are held together internally by strong covalent bonds, but attract each other only weakly.
- •Low melting and boiling points, because only the weak intermolecular forces between molecules need to be overcome — the strong covalent bonds within each molecule stay intact.
- •Intermolecular forces get stronger as molecule size increases, so larger molecules generally have higher melting and boiling points.
- •Small molecules don't conduct electricity, because they don't have an overall electric charge and have no free ions or electrons.
⚠️ Common mistake: Thinking a low melting point means covalent bonds are weak — the covalent bonds are strong; it is the much weaker forces between molecules that break when a substance melts or boils.
🧠 Remember: Small molecules melt easily because the forces between them (not within them) are weak.
Polymers4.2.2.5
Definition: Polymers are large molecules made from many small, repeating units called monomers, joined together by covalent bonds.
- •Polymers have much larger molecules than simple covalent substances, so the intermolecular forces between polymer chains are stronger.
- •Because of these stronger intermolecular forces, polymers are typically solid at room temperature.
⚠️ Common mistake: Thinking the covalent bonds along the polymer chain are what determine melting point — it is actually the intermolecular forces between chains that matter most, same as for any covalent substance.
🧠 Remember: Polymers = big molecules → strong intermolecular forces → solid.
Giant covalent structures4.2.2.6
Definition: Giant covalent structures are substances in which very large numbers of atoms are joined by strong covalent bonds in a giant lattice (e.g. diamond, silicon dioxide).
- •Very high melting and boiling points, because a huge number of strong covalent bonds must be broken.
- •Usually do not conduct electricity, since (with the exception of graphite) there are no free electrons or ions.
- •Are hard, because of the rigid network of strong bonds throughout the structure.
⚠️ Common mistake: Applying the 'weak forces between molecules, strong bonds within' rule from small molecules to giant covalent structures — there are no separate molecules here; the whole structure is one giant network of strong bonds, which is why the melting point is so much higher.
🧠 Remember: Giant covalent = one huge molecule of strong bonds — very high melting point.
Properties of metals and alloys4.2.2.7
Definition: An alloy is a mixture of a metal with one or more other elements, usually other metals.
Pure metals and alloys differ in how easily their layers of atoms can slide past each other.
- •Pure metals are usually too soft for many uses, because their identical-sized atoms are arranged in layers that slide over each other easily.
- •Alloys contain atoms of different sizes, which distorts the regular layers and makes it harder for them to slide, making alloys harder than the pure metal.
- •Both pure metals and alloys share typical metallic properties: good conductors, malleable, and generally high melting points.
⚠️ Common mistake: Thinking an alloy is a compound — an alloy is a mixture, not chemically bonded in fixed proportions, even though the different-sized atoms are what makes it harder.
🧠 Remember: Alloys are harder than pure metals because different-sized atoms disrupt the layers.
Metals as conductors4.2.2.8
Metals are excellent conductors of both electricity and heat, because of their delocalised electrons.
- •The delocalised electrons in metallic bonding are free to move throughout the structure, carrying electrical current.
- •These same delocalised electrons can also transfer thermal energy quickly, making metals good thermal conductors too.
⚠️ Common mistake: Thinking metals conduct heat and electricity for different reasons — both properties come from the same delocalised electrons.
🧠 Remember: One cause, two properties: delocalised electrons → electrical AND thermal conductor.
Feel like you’ve got it?
Practise it now →