Rates, Equilibrium and Redox

6.1 Physical and chemical changes

Signs of a chemical change: a new substance forms, there is a colour change, a gas is given off, a precipitate forms, or there is a change in temperature or light.

6.2 Rate of reaction

The rate of reaction is how fast reactants are used up or products are formed. It can be measured as the change in the amount of a reactant or product in a given time.

Collision theory

A reaction happens only when particles collide with enough energy. The collision theory says that a reaction occurs when:

1. the particles collide with each other, and
2. the collision has at least the activation energy, $E_\text{a}$ (the collision is "successful").

The rate depends on the frequency of successful collisions.

Factors that change the rate

The powder has a larger surface area, so more particles are exposed to the acid, and there are more frequent collisions between the acid and the calcium carbonate particles.

Investigating the rate

A reaction that gives off a gas can be followed in two ways, for example $\ce{CaCO3 + 2HCl -> CaCl2 + H2O + CO2}$.

Keep everything else the same (control variables): the same amounts, the same temperature and the same size of pieces, changing only one factor at a time.

Interpreting rate graphs

• The gradient (steepness) of the curve is the rate. The curve is steepest at the start, when the reactants are at their highest concentration, and it gets less steep as reactants are used up.
• The curve becomes flat (horizontal) when the reaction has finished and one reactant has run out.
• A steeper curve means a faster reaction. If the same amount of the limiting reactant is used, both curves end at the same final volume, but the faster one gets there sooner.
• If more of the limiting reactant is used, the curve ends at a higher level.

To find the rate at a particular time, draw a tangent to the curve and find its gradient, or compare the volumes made in equal times.

6.3 Reversible reactions and equilibrium

Some reactions are reversible: they can go in both directions. This is shown by the symbol $\rightleftharpoons$.

Hydrated and anhydrous compounds

Heating a hydrated compound drives off the water of crystallisation. Adding water to the anhydrous compound reverses the change.

$\ce{CuSO4.5H2O(s) <=> CuSO4(s) + 5H2O(l)}$

The forward reaction (heating) is endothermic and the reverse reaction (adding water) is exothermic: the anhydrous solid becomes hot as it turns blue again.

Dynamic equilibrium

At equilibrium both reactions are still happening, so it is dynamic. The amounts of reactants and products stay constant, though they need not be equal.

Changing the position of equilibrium

The position of equilibrium says how much reactant and product there is at equilibrium. If a condition is changed, the system opposes the change.

There is 1 mole of gas on the left and 2 moles on the right. Increasing the pressure moves the equilibrium to the side with fewer moles, the left, so less $\ce{NO2}$. Increasing the temperature favours the endothermic (forward) reaction, so it moves to the right and more $\ce{NO2}$ forms.

The Haber process

The Haber process makes ammonia, used for fertilisers.

$\ce{N2(g) + 3H2(g) <=> 2NH3(g)} \qquad \text{(forward reaction exothermic)}$

• Nitrogen comes from the air and hydrogen from methane (natural gas).
• Conditions: $450\,^\circ\text{C}$, $20\,000\ \text{kPa}$ ($200\ \text{atm}$) and an iron catalyst.
• The mixture is cooled so that ammonia liquefies and is removed. The unreacted nitrogen and hydrogen are recycled.

The Contact process

The Contact process makes sulfur trioxide, which is used to make sulfuric acid.

$\ce{2SO2(g) + O2(g) <=> 2SO3(g)} \qquad \text{(forward reaction exothermic)}$

• Sulfur dioxide comes from burning sulfur or roasting sulfide ores, and oxygen from the air.
• Conditions: $450\,^\circ\text{C}$, $200\ \text{kPa}$ ($2\ \text{atm}$) and a vanadium(V) oxide, $\ce{V2O5}$, catalyst.

6.4 Redox

Remember OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons).

Oxidation numbers

The oxidation number shows the charge an atom has, or would have, if the compound were made of ions.

• The oxidation number of an uncombined element is $0$ (for example $\ce{Mg}$, $\ce{O2}$, $\ce{Cl2}$).
• For a monatomic ion, it is the charge on the ion ($\ce{Na+}$ is $+1$, $\ce{Cl-}$ is $-1$).
• The sum of the oxidation numbers in a compound is $0$.
• The sum of the oxidation numbers in an ion is equal to the charge on the ion.
• Common values: hydrogen is $+1$ and oxygen is $-2$ in most compounds.

A Roman numeral in a name gives the oxidation number: iron(II) is $\ce{Fe^2+}$ with oxidation number $+2$, and iron(III) is $+3$. In manganate(VII), manganese is $+7$.

(a) Oxygen is $-2$, so $x + 2(-2) = 0$ and $x = +4$. (b) $2x + 7(-2) = -2$, so $2x = +12$ and $x = +6$. (c) $x + 3(-2) = -1$, so $x = +5$.

Identifying redox

Copper(II) oxide loses oxygen, so it is reduced (copper goes from $+2$ to $0$). Magnesium gains oxygen, so it is oxidised (from $0$ to $+2$). The oxidising agent is copper(II) oxide, and the reducing agent is magnesium.

Reactions of ions can be redox reactions too. In $\ce{Zn + Cu^2+ -> Zn^2+ + Cu}$, zinc loses two electrons (oxidised, so it is the reducing agent), and $\ce{Cu^2+}$ gains two electrons (reduced, so it is the oxidising agent). In $\ce{Cl2 + 2Br- -> 2Cl- + Br2}$, chlorine is the oxidising agent and bromide ions are oxidised.

Colour changes that show a redox reaction