Two Basic Drives (HSC SSCE Chemistry): Revision Notes

Two Basic Drives

Introduction: predicting whether reactions will occur

With countless chemical compounds available, chemists face an important question: will a proposed reaction actually happen? When starting materials are mixed in the laboratory, sometimes no observable reaction occurs. This could mean either:

• The reaction is very slow under current conditions but might proceed under different conditions, or
• The reaction simply will not occur under any conditions.

Being able to calculate whether a reaction is possible would save significant time and effort. Fortunately, such calculations can determine:

• The reaction will not occur under any conditions, or
• The reaction may occur.

This section explores the question: Is a reaction spontaneous or not?

The energy drive: nature's preference for lower energy

Our everyday experiences show that systems naturally move towards states of lower energy:

• Stones roll down hills rather than up
• Water flows from clifftops to the ground below
• Compressed springs unwind and relax

In each case, the system moves from high potential energy to lower potential energy.

Chemical reactions follow the same principle. Exothermic reactions agree with this concept—chemicals move from higher chemical energy to lower chemical energy, releasing energy (usually as heat) in the process. This is why exothermic reactions are much more common than endothermic ones.

The puzzle: spontaneous endothermic reactions

However, some endothermic reactions occur spontaneously, which seems contradictory. Consider these examples:

Endothermic chemical reactions:

1. Magnesium chloride and sodium carbonate solutions react to form magnesium carbonate precipitate:

$$MgCl_2(aq) + Na_2CO_3(aq) \rightarrow MgCO_3(s) + 2NaCl(aq)$$ $$\Delta H = +48 \ \text{kJ mol}^{-1} \ \text{(16.1)}$$

2. Crystalline barium hydroxide reacts with crystalline ammonium thiocyanate:

$$Ba(OH)_2 \cdot 8H_2O(s) + 2NH_4SCN(s) \rightarrow 2NH_3(g) + 10H_2O(l) + Ba(SCN)_2(aq)$$ $$\Delta H = +2176 \ \text{kJ mol}^{-1} \ \text{(16.2)}$$

3. Sodium hydrogen carbonate reacts with citric acid solution:

$$C_6H_8O_7(aq) + 3NaHCO_3(s) \rightarrow Na_3C_6H_5O_7(aq) + 3CO_2(g) + 3H_2O(l)$$ $$\Delta H = +80 \ \text{kJ mol}^{-1} \ \text{(16.3)}$$

Endothermic physical changes:

1. Liquids like ethanol evaporate:

$$C_2H_5OH(l) \rightarrow C_2H_5OH(g)$$ $$\Delta H = +43 \ \text{kJ mol}^{-1} \ \text{(16.4)}$$

2. Sodium thiosulfate dissolves readily in water:

$$Na_2S_2O_3 \cdot 5H_2O(s) \rightarrow Na_2S_2O_3(aq) + 5H_2O(l)$$ $$\Delta H = +46 \ \text{kJ mol}^{-1} \ \text{(16.5)}$$

All these processes absorb heat as they occur, meaning the systems move naturally towards states of higher energy. This suggests another natural tendency must exist alongside the drive towards lower energy.

The entropy drive: nature's preference for chaos

In nature, there is a second fundamental drive: the drive towards increased chaos or increased randomness.

Observable examples of the entropy drive

Think about these everyday observations that demonstrate the tendency towards disorder:

• Mixed coloured lollies never spontaneously separate into groups by colour
• When gases mix together, they never spontaneously separate
• A mixture of alcohol and water never separates on its own
• A tidy room becomes messy over time (never the reverse!)

Defining entropy

The physical quantity entropy (symbol: $S$) measures the amount of randomness, chaos, or lack of ordered structure in a substance.

In the five endothermic processes described above (equations 16.1–16.5), there was a change from a fairly ordered state to a much less ordered one. Therefore, $\Delta S$ is positive for each process. The drive towards greater entropy (greater chaos) is stronger than the drive towards lower energy, so these reactions occur as written despite being endothermic.

Competition between the two drives

In chemical reactions, usually the energy drive dominates, which explains why exothermic reactions are more common. However, in some reactions the entropy drive is stronger than the energy drive, allowing endothermic reactions to be spontaneous.

The direction of a reaction depends on which drive is larger.

The relationship between drives and changes

Investigation 16.1: modelling entropy changes in reactions

This investigation helps students understand entropy by constructing models of substances undergoing physical or chemical changes. The key question is: Does randomness increase or decrease?

Aim

To model entropy changes in reactions and understand the concept of disorder.

Key parts of the investigation

Important safety considerations

The teacher demonstration in Part F involves hazardous chemicals:

Hazard	Safety measures
Ammonia gas is released	Teacher demonstration only. Perform in fume cupboard or well-ventilated area. Check if students have asthma or respiratory conditions.
Barium hydroxide is very corrosive to skin and eyes	Wear safety glasses. Perform in fume cupboard. Use plastic spoon to transfer chemicals.

Generalizations about randomness and entropy

Investigation 16.1 demonstrates several important principles about how entropy relates to the physical state and behavior of matter.

1. Entropy increases with state changes

Let's understand why this ordering exists:

• Solids: Particles are close together, vibrating about fixed positions. This is an orderly, low-entropy arrangement.
• Liquids: Particles are further apart, moving from place to place as well as vibrating. This is less orderly—higher entropy than solids.
• Gases: Particles are much further apart, moving rapidly and randomly. This is the most chaotic state—highest entropy.

2. Similar states have similar entropies

• Different solids have fairly similar entropy values per mole
• Different liquids have fairly similar entropy values per mole
• Different gases have fairly similar entropy values per mole

3. Dissolution increases entropy

When a solid dissolves in a solvent, there is an increase in entropy. The ordered crystal structure breaks down, and ions or molecules become dispersed throughout the solution.

Using these generalizations

To estimate qualitatively the entropy change ($\Delta S$) for a chemical reaction, consider:

1. Count the moles of gas on each side of the equation. If gas moles increase, $\Delta S$ is positive.
2. Look for solids dissolving to form solutions. This increases entropy.
3. Compare number of particles in different states.

Predicting reaction spontaneity: which drive wins?

The table below summarizes how the relative magnitudes of energy and entropy drives determine whether a reaction occurs as written:

Direction of energy drive	Direction of entropy drive	Which drive is larger?	Does reaction go as written?
Forward	Forward	Does not matter	Yes
Forward	Backward	Energy	Yes
Forward	Backward	Entropy	No
Backward	Forward	Energy	No
Backward	Forward	Entropy	Yes
Backward	Backward	Does not matter	No

Key insights from the table

Worked example: predicting entropy changes

Practice questions

Try predicting whether $\Delta S$ is positive or negative for these reactions:

a) $\text{N}_2(\text{g}) + 3\text{H}_2(\text{g}) \rightarrow 2\text{NH}_3(\text{g})$ ($\Delta H$ is negative)

b) $\text{CuSO}_4\text{.5H}_2\text{O}(\text{s}) \rightarrow \text{CuSO}_4(\text{s}) + 5\text{H}_2\text{O}(\text{g})$ ($\Delta H$ is positive)

Then deduce which drive is larger, given that reaction (a) goes as written while reaction (b) goes in the reverse direction.

Summary