Rate of Reaction and Chemical Equilibrium
FoundationReaction rate is measured from how gas volume, mass or colour depth changes with time; concentration, temperature, particle size and catalysts speed reactions up by increasing effective collisions. A reversible reaction in a closed system reaches dynamic equilibrium; changes in concentration or temperature shift the equilibrium position, while a catalyst only hastens its arrival.
- The gradient of a rate curve is the rate; follow progress by gas volume, mass change or colour depth.
- Collision theory: concentration, temperature and particle-size effects are all explained by collision frequency and the proportion of effective collisions.
- A catalyst offers a lower-energy route, is unchanged itself, and never shifts the equilibrium position.
文體小知識：科學知識說明文——以大理石與鹽酸等實例說明影響反應速率的因素與碰撞理論，再由可逆反應引入動態平衡與平衡移動的判斷。
§Text Passage
1Iron rusts over years, yet the explosion of hydrogen with oxygen is over in an instant. The rate of reaction describes how fast a reaction proceeds, usually as the amount of reactant consumed or product formed per unit time. In practice we follow the progress by measuring the volume of gas, the mass of the mixture or the depth of colour, plotting the data against time: the gradient of the curve is the rate — the steeper the curve, the faster the reaction; as the curve flattens, reactants are running out.
2Why do changing conditions affect the rate? Collision theory says particles must collide to react, and only collisions with enough energy are effective. Raising the concentration puts more particles into each unit volume, making collisions more frequent; raising the temperature makes particles move faster and increases the proportion of sufficiently energetic collisions, so the rate rises markedly; grinding a solid into powder enlarges the contact surface; and a catalyst provides an alternative route of lower energy demand.
3Some reactions run both ways: the forward reaction turns reactants into products while the reverse reaction simultaneously turns products back into reactants — a reversible reaction. In a closed system the forward and reverse rates eventually become equal and the concentrations of all species stop changing: dynamic equilibrium. The word dynamic reminds us that the reactions have not stopped; the two opposing changes cancel out, gaining and losing equal amounts in the same time.
4Equilibrium is disturbed when conditions change. Adding a reactant shifts the equilibrium towards the products to consume the excess; raising the temperature shifts it in the endothermic direction. In short, after a change of conditions the equilibrium moves in whichever direction lessens the effect of the change — the principle for predicting shifts. At constant temperature the equilibrium constant is a fixed value, unaffected by concentration changes; only a temperature change alters it.
5This knowledge of equilibrium directly supports industry: plants making ammonia and sulfuric acid control variables such as pressure, concentration and temperature to secure the best combination of yield and rate. In daily life, the rush of bubbles when a soft drink is opened and the changing oxygen levels of mountain streams are examples of equilibrium shifts. Mastering rate and equilibrium together lets you explain and control chemical changes around you.
1Summary 重點摘要
Reaction rate is measured from how gas volume, mass or colour depth changes with time; concentration, temperature, particle size and catalysts speed reactions up by increasing effective collisions. A reversible reaction in a closed system reaches dynamic equilibrium; changes in concentration or temperature shift the equilibrium position, while a catalyst only hastens its arrival.
2Vocabulary cards Vocabulary
The amount of reactant consumed or product formed per unit time, describing how fast a reaction goes.
The theory that particles must collide with sufficient energy to react; it explains the factors affecting rate.
A graph of reaction progress against time; its gradient is the rate, which decreases as reactants run out.
A substance that speeds up a reaction by providing a lower-energy route, itself unchanged in mass and chemical nature.
The area of a solid reactant exposed to others; grinding the solid enlarges it and speeds up the reaction.
A reaction that proceeds in both directions, shown with the double arrow ⇌.
The state in a closed system where forward and reverse rates are equal and all concentrations stay constant.
When conditions such as concentration or temperature change, the equilibrium shifts in the direction that lessens the change.
The relative amounts of reactants and products at equilibrium; a change of conditions shifts it.
The constant Kc relating the equilibrium concentrations at a given temperature; only a temperature change alters it.
3Step by step Steps
- Step 1: How to measure the rate of reactionThe method depends on the reaction, but the common approach is to follow an observable change: collecting and measuring the volume of gas produced, tracking the fall in total mass, or following the depth of colour. For example, marble reacts with dilute hydrochloric acid giving off carbon dioxide; plotting the gas volume against time, the steepest gradient comes at the start, when the rate is greatest, and the curve flattens as reactants run out. Reading rate curves is a must-skill in Paper 1: gradient and final height each carry their own meaning and must not be confused.
- Step 2: Concentration and pressureCollision theory explains that reacting requires collisions, and only collisions with sufficient energy are effective. Raising the concentration (or the pressure of a gas) puts more particles into each unit volume, so collisions become more frequent and the number of effective collisions per unit time increases — the rate rises. For example, marble fizzes far more vigorously in 2 mol dm⁻³ hydrochloric acid than in 0.5 mol dm⁻³; steelworks blast in high-pressure air for the same reason. In explanations, write 'the frequency of collisions increases' — merely 'more collisions' is not precise enough.
- Step 3: The effect of temperatureRaising the temperature raises the average kinetic energy of the particles. Collisions become more frequent, but crucially the proportion of collisions energetic enough to react rises sharply, so the rate increases markedly. For example, refrigerators slow both the chemical spoilage of food and the micro-organisms at work; in the classic experiment timing how long sodium thiosulfate takes to cloud dilute sulfuric acid, each ten-degree rise shortens the time noticeably. As a rule of thumb, a ten-degree rise roughly doubles to triples the rate — the same reason a pressure cooker cooks faster.
- Step 4: Particle size and catalystsOnly the surface of a solid can collide, so grinding it into powder vastly enlarges the contact surface and multiplies the chances of collision. For example, the same mass of powdered marble reacts with 2 mol dm⁻³ hydrochloric acid far faster than marble chips, giving a much steeper initial curve; flour mills ban naked flames because airborne flour dust, with its huge surface, can explode violently. A catalyst offers an alternative route with a lower energy threshold, so more collisions succeed. For example, hydrogen peroxide barely decomposes on its own, yet a pinch of manganese(IV) oxide releases oxygen briskly — and the powder's mass and nature are unchanged afterwards.
- Step 5: Reversible reactions and dynamic equilibriumSome reactions run both ways, marked with the double arrow ⇌. In a closed system the forward reaction starts fast and the reverse slow; as products accumulate the reverse speeds up, until both rates are equal and all concentrations stay constant — dynamic equilibrium. For example, in a sealed bottle of water the rates of evaporation and condensation become equal: the water level and vapour concentration no longer change, yet molecules keep crossing between phases. Remember the two signatures of equilibrium — equal forward and reverse rates, constant concentrations — and the requirement of a closed system; all three are favourite exam checks.
- Step 6: Predicting shifts of equilibriumA change of conditions upsets the equilibrium, and the system adjusts until a new one forms. The guiding rule: the equilibrium shifts in the direction that lessens the change. Adding a reactant shifts equilibrium towards the products; raising the pressure of a gas mixture shifts it towards the side with fewer gas molecules; raising the temperature shifts it in the endothermic direction, lowering it in the exothermic direction. For example, reddish-brown nitrogen dioxide and colourless dinitrogen tetroxide coexist as N₂O₄ ⇌ 2NO₂; on warming the colour deepens, showing the equilibrium has shifted towards NO₂ — so that direction is the endothermic one. A catalyst speeds up both forward and reverse reactions equally, shortening only the time to reach equilibrium; it never shifts the position nor changes the yield.
- Step 7: Industrial wisdom of rate and equilibriumIndustry cares about rate and yield at once: a slow reaction earns nothing, and a badly placed equilibrium gives a poor yield. For example, ammonia synthesis runs at a moderate temperature (balancing rate against equilibrium), high pressure (favouring the ammonia side) and with a catalyst (hastening equilibrium), while the ammonia is liquefied and removed as it forms so the equilibrium keeps replenishing it. This 'react and remove' strategy is the most valuable application of equilibrium. Soft drinks are another everyday case: with the cap on, high pressure keeps carbon dioxide dissolved; once opened, the pressure plunges, the equilibrium shifts to release gas, and the foam surges out — a live demonstration of equilibrium shifting.
4Diagrams Diagrams
5Worked example Example
Experiment (a): equal masses of powdered marble and marble chips each react with excess 1 mol dm⁻³ hydrochloric acid, giving two curves of carbon dioxide volume against time. Experiment (b): a sealed syringe holds an N₂O₄ ⇌ 2NO₂ mixture; on warming, the reddish-brown colour deepens. Compare the curves with explanation, then deduce the direction of the equilibrium shift and the heat change of that direction.
- Experiment (a): the powder's curve starts with a much steeper gradient — a faster initial rate — but both curves level off at the same final volume.
- Explanation: the powder has a larger contact surface area, so collisions per unit time are more frequent and the rate rises; with acid in excess and equal masses of marble, the total carbon dioxide is the same.
- Experiment (b): the deepening colour on warming means the concentration of NO₂ has risen, so the equilibrium has shifted from left to right, towards nitrogen dioxide.
- By the general rule, on heating the equilibrium moves in the endothermic direction, so the forward change from N₂O₄ to NO₂ is endothermic.
- Answer: Powdered marble reacts faster at first (larger surface area, more frequent collisions) but gives the same final gas volume; warming shifts N₂O₄ ⇌ 2NO₂ to the right (more NO₂, deeper colour), and that direction is endothermic.
6Common mistakes Mistakes
7Quick check Quick check
Q1Which observable changes can be followed to measure the rate of a reaction?
Q2Use collision theory to explain why raising the temperature speeds up a reaction.
Q3How does a catalyst speed up a reaction, and does it shift the equilibrium position?
Q4When the N₂O₄ ⇌ 2NO₂ mixture is warmed, the colour deepens. Which way does the equilibrium shift, and is that direction endothermic or exothermic?
Q5What are the two features of a system at dynamic equilibrium, and what condition is required?
8Recap Recap
- The gradient of a rate curve is the rate; follow progress by gas volume, mass change or colour depth.
- Collision theory: concentration, temperature and particle-size effects are all explained by collision frequency and the proportion of effective collisions.
- A catalyst offers a lower-energy route, is unchanged itself, and never shifts the equilibrium position.
- Dynamic equilibrium: equal forward and reverse rates, constant concentrations, in a closed system.
- After a change of conditions the equilibrium shifts to lessen the change; Kc is constant only at constant temperature.