Modelling matter as particles

Matter is any material that has mass and occupies volume. Chemists describe matter with a particulate model: a visible sample is imagined as a huge collection of tiny particles. This model explains why substances can have fixed compositions, why mixtures can be separated, and why chemical reactions rearrange particles instead of making matter appear or disappear.

An atom is the smallest particle of an element that retains that element’s chemical identity. Structure 1.2 opens up the atom properly. For now, the key idea is straightforward: atoms of different elements can be rearranged and bonded in different ways.

Elements, compounds and mixtures

A pure substance is a material made of only one chemical substance and therefore has a fixed composition throughout. Elements and compounds are both pure substances.

An element is a pure substance composed of only one type of atom and cannot be chemically decomposed into simpler substances. Copper, oxygen and sulfur are elements. Heating, filtering or dissolving an element will not split it into simpler chemical substances. In chemical symbols, an element is shown by one or two letters, such as C, Na or Fe.

A compound is a pure substance composed of atoms of two or more different elements chemically bonded in a fixed ratio. Water contains hydrogen and oxygen in a fixed particle ratio; sodium chloride contains sodium and chloride ions in a fixed ratio. A compound has its own properties. They are not just an average of the properties of the elements it contains. For example, a compound containing a reactive metal and a poisonous gas may be a stable crystalline solid.

A chemical bond is an electrostatic attraction that holds atoms or ions together in a substance. Since compounds involve bonding, separating a compound into its elements needs a chemical change, such as thermal decomposition or electrolysis, not a simple physical technique.

A mixture is a material containing two or more elements or compounds that are not chemically bonded together and are present in no fixed ratio. The phrase to watch is “no fixed ratio”. One sample of salty water can be more concentrated than another and still count as salty water. The components of a mixture usually keep their own chemical identities, so physical methods can separate mixtures.

A quick comparison is often the easiest way to keep the categories straight:

Homogeneous and heterogeneous mixtures

A homogeneous mixture is a mixture with uniform composition and properties throughout a sample. Air is a good model example: any small sample is still mainly the same gases mixed together. A solution is homogeneous too, because the dissolved particles are spread through the solvent at particle level.

A heterogeneous mixture is a mixture with non-uniform composition, so different regions of the sample have different properties or composition. Muddy water, oil-and-water dressing and granite are heterogeneous. A sample taken from one part may not have the same composition as a sample taken from another.

Do not decide “compound or mixture” just by appearance. A clear colourless liquid might be pure water, a solution, or a mixture of miscible liquids. Chemists look for evidence: fixed composition, sharp melting or boiling behaviour, whether components retain properties, and whether physical separation works.

Purity and melting point

A melting point is the temperature at which a solid and its liquid are in equilibrium at a stated pressure. Pure substances usually melt over a very narrow temperature range. Impurities usually lower the melting point and broaden the melting range, which makes melting point determination a useful Tool 1 technique when checking whether a product has been purified successfully.

In a practical, pack the sample into a thin capillary tube and place it in a melting point apparatus. Record the temperature when melting first begins and when the last solid disappears. That range tells you more than a single neat number when the sample is impure.

Choosing a separation method

A physical separation works when the components differ in a physical property. The property you choose to use determines the technique:

• particle size → filtration;
• solubility in a chosen solvent → solvation followed by filtration or crystallization;
• volatility or boiling point → evaporation or distillation;
• attraction to a mobile phase and a stationary phase → chromatography;
• magnetism or density → useful in some mixtures, though not the main syllabus list here.

Practical issues matter as well: safety, toxicity, flammability, cost, scale of the sample, environmental disposal, whether heating might decompose a substance, and the purity required. In class I always ask: “What property is different enough for us to use?” If you cannot answer that, you have not chosen a separation method yet.

Solvation, filtration, recrystallization and evaporation

Solvation is a process in which solute particles become surrounded by solvent particles and disperse through the solvent. If one solid dissolves and another does not, solvation gives a tidy route to separation. Add a suitable solvent, stir, and let the soluble component enter solution while the insoluble component stays as a solid.

Filtration is a separation technique in which a mixture passes through a porous barrier that retains insoluble solid particles while liquid or solution passes through. The solid left on the filter paper is the residue, a solid material retained after a separation process. The liquid that passes through is the filtrate, a liquid or solution collected after filtration.

Evaporation is a separation technique in which solvent is vaporized to leave behind a dissolved solid. It is useful if the dissolved solid is stable on heating and if you do not need to collect the solvent. Evaporation to dryness is simple, but strong heating may produce small, impure crystals.

Recrystallization is a purification technique in which an impure solid is dissolved in hot solvent and then forms purer crystals as the solution cools. The desired substance is more soluble in hot solvent than in cold solvent, while many impurities either stay dissolved or are removed by filtration. The usual sequence is: dissolve in minimum hot solvent, filter if insoluble impurities remain, cool slowly, filter the crystals, wash with a little cold solvent, and dry.

For products of a reaction, a common purification pattern is to remove excess insoluble reactant by filtration, concentrate the filtrate by gentle evaporation, allow crystals to form on cooling, then filter and dry the crystals. The details change with the reaction, but the logic stays the same.

Distillation and chromatography

Distillation is a separation technique in which a liquid is vaporized and then condensed, allowing components with different volatilities or boiling points to be separated. Simple distillation is often used to separate a solvent from a dissolved solid, or to separate liquids with sufficiently different boiling points. The thermometer, condenser direction, and collection flask all matter. Vapour should be cooled efficiently so the distillate is collected as a liquid.

Chromatography is a separation technique in which components of a mixture move at different rates because they have different relative attractions to a mobile phase and a stationary phase. In paper chromatography, the solvent is the mobile phase and the paper is the stationary phase. A component more strongly attracted to the solvent moves further; a component more strongly attracted to the paper moves less far.

Structure 2.2 starts to peep through the door here: intermolecular forces influence whether substances mix, dissolve or separate. If particles of two substances attract each other strongly enough, a homogeneous mixture may form. If not, the mixture may remain heterogeneous or separate into layers.

Alloys are a useful link to Structure 2.3. An alloy is a mixture containing a metal and one or more other elements, usually with metallic bonding between particles. Alloys are generally treated as mixtures because their composition can vary over a range and their properties change with composition; they do not usually have the single fixed ratio expected for a compound.

The kinetic molecular theory

The kinetic molecular theory is a model that describes matter as particles in constant motion with spaces and attractions between them. The model doesn’t try to show every particle exactly as it is. Instead, it helps explain the properties we can observe, such as shape, volume, compressibility, diffusion and changes of state.

In a solid, the particles are held close together and mainly vibrate about fixed positions. In a liquid, they are still close together, but they can move past one another. In a gas, the particles are far apart compared with their own size, and they move rapidly in random directions.

Solids, liquids and gases

A solid is a state of matter with a fixed shape and fixed volume because its particles are held in fixed positions. Solids are not easily compressed, since the particles are already close together.

A liquid is a state of matter with fixed volume but no fixed shape because its particles remain close together while moving past one another. A liquid takes the shape of the bottom of its container and is also not easily compressed.

A gas is a state of matter with no fixed shape and no fixed volume because its particles are widely spaced and move independently. Gases fill their container. They are compressible because there is much more empty space between particles.

Whether a substance is solid or fluid under standard conditions depends on the balance between particle attractions and particle motion. Stronger attractions tend to favour solids; weaker attractions and greater particle motion favour liquids or gases. Structure 2 develops this using bonding and intermolecular forces.

State symbols in equations

State symbols show the physical state of each substance in a chemical equation:

• (s) means solid;
• (l) means liquid;
• (g) means gas;
• (aq) means aqueous, meaning dissolved in water.

For example, $H_2O(s)$, $H_2O(l)$ and $H_2O(g)$ show the same compound in different physical states. $NaCl(s)$ is solid sodium chloride, while $NaCl(aq)$ means sodium chloride dissolved in water. The symbol (aq) is not a fourth state of matter; it is a solution in water.

State symbols carry chemical information. If a reaction is done in solution, if a gas is released, or if an insoluble solid forms, the state symbols help show what you would actually observe.

Changes of state

Melting is a change of state in which a solid becomes a liquid. Freezing is a change of state in which a liquid becomes a solid. These two processes are opposites.

Vaporization is a change of state in which a liquid becomes a gas. It includes evaporation, a surface process that can occur below the boiling point, and boiling, a process throughout the liquid when bubbles of vapour form within the liquid. Condensation is a change of state in which a gas becomes a liquid.

Sublimation is a change of state in which a solid becomes a gas without first becoming a liquid. Deposition is a change of state in which a gas becomes a solid without first becoming a liquid. Solid carbon dioxide subliming is one useful mental picture; frost forming from water vapour is another.

These changes are physical changes, not chemical changes. The particles themselves do not turn into different substances. Their arrangement, spacing and motion change.

Energy changes during changes of state

A move from a more ordered or condensed state to a less condensed state needs an input of energy. Melting, vaporization and sublimation are endothermic processes, meaning processes that absorb energy from the surroundings.

A move from a less condensed state to a more condensed state releases energy. Freezing, condensation and deposition are exothermic processes, meaning processes that transfer energy to the surroundings.

This links to Reactivity 1.2: energy is not destroyed during a change of state; it is transferred. During melting or boiling, energy is used mainly to overcome attractions between particles rather than to increase the temperature of the sample. During freezing or condensation, attractions form more strongly and energy is released.

Temperature and particle motion

The temperature $T$ means thermodynamic temperature, measured on the Kelvin scale ($K$). The kinetic energy $E_k$ is the energy a particle has because it is moving, where $E_k$ is kinetic energy (J). Temperature measures the average kinetic energy of the particles in a sample.

That word “average” matters. In a real sample, particles don’t all have exactly the same kinetic energy. Some move faster and others slower. At a higher temperature, the distribution shifts, so the average kinetic energy is greater. In Reactivity 2.2 you will meet this as a graphical distribution of kinetic energies.

For chemical reactions, the link is direct: particles must collide, and successful collisions need enough kinetic energy and a suitable orientation. So not every collision at a given temperature reacts. Temperature changes the fraction of particles able to react.

Heating and cooling curves

When a single substance is heated within one state, its temperature usually rises because the particles’ average kinetic energy increases. In a solid, the particles vibrate more vigorously; in a liquid or gas, they move faster on average.

During a change of state of a pure substance at constant pressure, temperature stays constant even though energy is still being transferred. This catches students out every year, so be precise: during melting or boiling, the added energy changes the arrangement and separation of particles. It does not increase their average kinetic energy.

A heating curve therefore has sloping regions and flat regions. Sloping regions show one state warming up. Flat regions show a change of state, such as melting or boiling, where two states are present together. During cooling, the same idea runs in reverse: condensation and freezing occur at constant temperature for a pure substance at constant pressure.

Observable changes match the curve. A solid may soften and then coexist with liquid during melting. A boiling liquid forms vapour bubbles throughout the liquid. During condensation, gas becomes visible liquid droplets if the liquid is normally transparent. During deposition, solid appears directly from gas on a cold surface.

Kelvin and Celsius

The kelvin is the SI unit of thermodynamic temperature. It is an absolute temperature scale: $0\ \text{K}$ is absolute zero, the lowest possible temperature on the thermodynamic scale. At absolute zero, particles have minimum possible thermal motion; a substance cannot be cooled below this.

The Celsius and Kelvin scales use the same size interval. A rise of $1\ \text{K}$ is the same temperature change as a rise of $1\ ^\circ\text{C}$. They differ in their zero point: Celsius is based historically around water’s freezing and boiling points, while Kelvin starts at absolute zero.

Use the conversion:

Rearranged:

$\theta = T - 273.15$.

So a temperature of $25.00\ ^\circ\text{C}$ is $298.15\ \text{K}$, and a temperature change of $20\ ^\circ\text{C}$ is a temperature change of $20\ \text{K}$. Do not put a degree sign with kelvin: write $\text{K}$, not $\degree\text{K}$.

Kelvin temperature is proportional to average kinetic energy, but Celsius temperature is not. The reason is the zero: $0\ \text{K}$ corresponds to zero on the thermodynamic scale, while $0\ ^\circ\text{C}$ is simply a convenient reference point for water under ordinary pressure, not zero particle kinetic energy.