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Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle

Chemistry

Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle

Overview

Overview

🌳 Concept tree

  • Lesson 1 Chemical Reaction Rates and Le Chatelier's Principle
  • ├── The equilibrium system
  • ├── Types of chemical reactions
  • ├── Chemical equilibrium
  • ├── Rate of reaction
  • ├── Factors affecting the rate of a reaction
  • ├── Nature of the reactants
  • ├── Concentration and the law of mass action
  • ├── Temperature — collision theory
  • ├── Pressure
  • ├── Le Chatelier's principle
  • ├── Catalysts
  • └── Light
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 1 / 14
The equilibrium system

The equilibrium system

  • An equilibrium system looks stationary but is really dynamic.
  • Physical example. Heat water in a closed vessel. Two opposite processes occur — evaporation and condensation. At first evaporation predominates and the vapour pressure rises (vapour pressure = pressure due to water vapour in the air at a given temperature). It keeps rising until it equals the saturated vapour pressure (the maximum water-vapour pressure that can exist in air at that temperature). Then the number of molecules that evaporate per second equals the number that condense per second — equilibrium:
  • water (liquid) ⇌ water (vapour)
  • The macroscopic properties (the level of the liquid, the pressure) stop changing, but the two processes continue at equal rates.
  • That is what "dynamic" means.
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 2 / 14
Types of chemical reactions

Types of chemical reactions

Comparison aspectMain ideaKeywords
Complete (irreversible) reactions
  • go practically in one direction
  • because a product leaves the system as a gas or a precipitate
  • so the products cannot recombine under the same conditions. NaCl(aq) + AgNO₃(aq) → NaNO₃(aq) + AgCl↓ (white precipitate) Mg + 2HCl → MgCl₂ + H₂↑ (hydrogen escapes) Neutralization: NaOH + HCl → NaCl + H₂O
practicallydirectionbecauseproduct
Reversible reactions
  • proceed in both the forward and backward directions
  • so reactants and products are always present together in the reaction medium
proceedbothforwardbackward
Example: esterification.→CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O. According to the equation, mixing 1 mol acid with 1 mol alcohol should give 1 mol ester and 1 mol water. But if the reaction mixture is tested with blue litmus paper, it turns red, although the ester and water are neutral. The acid has not been fully consumed, because the reverse reaction (hydrolysis of the ester) takes place at the same time.
Comparison aspectMain ideaKeywords
ForwardCH₃COOH + C₂H₅OH → CH₃COOC₂H₅ + H₂OCH₃COOHC₂H₅OHCH₃COOC₂H₅
BackwardCH₃COOC₂H₅ + H₂O → CH₃COOH + C₂H₅OHCH₃COOC₂H₅CH₃COOHC₂H₅OH
•
  • Because both reactants and products are always present
  • the mixture is acidic
Becausebothreactantsproducts
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 3 / 14
Chemical equilibrium

Chemical equilibrium

  • A dynamic state reached in a reversible reaction when the rate of the forward reaction equals the rate of the backward reaction, and the concentrations of all reactants and products stay constant.
  • All substances remain in the reaction medium (no gas escapes, no precipitate forms).
  • The position of equilibrium does not change as long as the conditions (temperature, pressure, and so on) are unchanged.
  • Equal rates do not mean equal concentrations.
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 4 / 14
Rate of reaction

Rate of reaction

  • The rate of a chemical reaction is the change in the concentration of reactants (or products) per unit time.
  • Concentration is in mol/L and time in seconds or minutes, so the unit of rate is mol/L·s (or mol/L·min).
  • In a complete reaction the concentration of reactants falls until almost nil and that of products rises (Figure 3-2).
  • In a reversible reaction the reactant concentration falls and the product concentration rises until equilibrium, after which both curves become horizontal (Figure 3-1). The reactant concentration does not fall to zero.
  • Reactions differ in speed instantaneous (AgNO₃ + NaCl, which gives the precipitate of sparingly soluble AgCl as soon as the solutions are mixed), moderate (saponification of vegetable oil with caustic soda giving soap and glycerol), very slow (rusting of iron, taking months).
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 5 / 14
Factors affecting the rate of a reaction

Factors affecting the rate of a reaction

Six factors→(1) nature of the reactants, (2) concentration, (3) temperature, (4) pressure, (5) catalysts, (6) light.
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 6 / 14
Nature of the reactants

Nature of the reactants

Comparison aspectMain ideaExampleKeywords
Type of bondingionic reactants react extremely fast (ions combine as soon as the solutions are mixed
  • NaCl + AgNO₃)
  • covalent reactants react slowly (e.g. most organic reactions)
ionicreactantsreactextremely
Exposed surface area
  • equal masses of zinc as powder and as a block reacting separately with equal volumes of dilute HCl — the powder finishes sooner. Larger exposed surface → faster reaction
  • because more particles are in contact with the other reactant
—equalmasseszincpowder
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 7 / 14
Concentration and the law of mass action

Concentration and the law of mass action

  • More molecules per unit volume means more collisions, hence a faster reaction.
  • The Norwegian scientists Waage and Guldberg stated the law of mass action:
  • At constant temperature, the rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants, each raised to a power equal to the number of molecules (or ions) of it in the balanced equation.
  • Demonstration. Adding pale-yellow FeCl₃ solution gradually to colourless NH₄SCN solution gives a blood-red colour of iron(III) thiocyanate:

FeCl₃ + 3NH₄SCN ⇌ Fe(SCN)₃ + 3NH₄Cl

Adding excess FeCl₃ deepens the red colour→more product forms.
Comparison aspectMain ideaKeywords
The equilibrium constant.
  • For the reaction above
  • forward rate r₁ = k₁[FeCl₃][NH₄SCN]3 and backward rate r₂ = k₂[Fe(SCN)₃][NH₄Cl]3. At equilibrium r₁ = r₂
  • so k₁[FeCl₃][NH₄SCN]3 = k₂[Fe(SCN)₃][NH₄Cl]3 and
reactionaboveforwardrate
•Kc = k₁/k₂ = [Fe(SCN)₃][NH₄Cl]3 / ([FeCl₃][NH₄SCN]3)NH₄ClFeCl₃NH₄SCN
•Kc is the equilibrium constant (concentrations in mol/L; k₁, k₂ are the rate constants of the forward and backward reactions).equilibriumconstantconcentrationsrate
General form
  • for aA + bB ⇌ cC + dD
  • Kc = [C]c [D]d / ([A]a [B]b)
Rate and the law of mass action — worked examples.
  • For the reaction 2A + B → products
  • the rate is r = k[A]2[B]. If the concentration of A is doubled (B constant) the rate becomes 22 = 4 times
  • if B is doubled it becomes 2 times
  • if both are doubled
  • 4 × 2 = 8 times
  • if the volume of the (gaseous) mixture is halved
  • every concentration doubles
  • so the rate becomes 8 times
reactionproductsconcentrationdoubled
Comparison aspectMain ideaKeywords
Example 1.
  • For H₂ + I₂ ⇌ 2HI the equilibrium concentrations are [H₂] = 0.5
  • [I₂] = 0.5 and [HI] = 3.0 mol/L
equilibriumconcentrations
•Kc = (3.0)2/(0.5 × 0.5) = 9/0.25 = 36.
Example 2 (finding Kc from the initial amounts).1 mol H₂ and 1 mol I₂ are placed in a 1 L vessel; at equilibrium 1.6 mol of HI is present. Forming 1.6 mol HI used 0.8 mol of each reactant, so [H₂] = [I₂] = 1 − 0.8 = 0.2 mol/L. Kc = (1.6)2/(0.2 × 0.2) = 2.56/0.04 = 64.placedvesselequilibriumpresent
Example 3 (degree of dissociation).0.10 mol of N₂O₄ is placed in a 1 L flask where it dissociates 40 % according to N₂O₄ ⇌ 2NO₂. Dissociated: 0.04 mol; remaining [N₂O₄] = 0.06 mol/L; [NO₂] = 2 × 0.04 = 0.08 mol/L. Kc = (0.08)2/0.06 = 0.0064/0.06 = 0.107.N₂O₄placedflaskdissociates
  • Remarks on Kc
  1. A small Kc (< 1) means the product of the product concentrations (the numerator) is smaller than the product of the reactant concentrations (the denominator): the reaction hardly proceeds towards the products and the reverse reaction is effective. Example: AgCl(s) ⇌ Ag+ + Cl− has K ≈ 1.7 × 10−10, which shows that AgCl is sparingly soluble in water.
  2. A large Kc (> 1) means the reaction goes almost to completion (forward reaction predominates), e.g. H₂ + Cl₂ ⇌ 2HCl, whose constant is enormous (about 1033 at room temperature).
  3. Pure solids, precipitates and pure water (as solvent) are not written in the Kc expression, because their concentrations stay constant whatever their quantity. Examples: for CaCO₃(s) ⇌ CaO(s) + CO₂(g), Kc = [CO₂]; for Fe₃O₄(s) + 4H₂(g) ⇌ 3Fe(s) + 4H₂O(g), Kc = [H₂O]4/[H₂]4.
  4. At a given temperature Kc does not change with concentration, pressure or catalyst; it changes only with temperature.
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 8 / 14
Temperature — collision theory

Temperature — collision theory

  • For a reaction to occur, molecules of the reactants must collide, but only molecules of very high speed react, because their kinetic energy is high enough to break the bonds inside the molecules.
  • The minimum energy a colliding molecule must have in order to react is the activation energy.
  • Activated molecules are those whose kinetic energy equals or exceeds the activation energy.
  • Raising the temperature increases the fraction (proportion) of activated molecules, so the rate increases.
  • For many reactions the rate doubles for a 10 °C rise.
  • For instance, if the rate doubles for every 10 °C, heating from 20 °C to 50 °C multiplies the rate by 23 = 8.
  • Effect of temperature on equilibrium (NO₂/N₂O₄).

NO₂ (reddish brown) ⇌ N₂O₄ (colourless) + heat

  • A flask of NO₂ placed in a freezing mixture loses its brown colour gradually (the forward exothermic reaction is favoured and N₂O₄ forms); in hot water the brown colour deepens.
  • When the flask is brought back to room temperature (25 °C) the colour returns to its first state.
  • So cooling an exothermic reaction at equilibrium shifts it forward (to liberate heat); heating shifts it backward.
  • For an endothermic reaction the reverse is true.
  • Temperature is the only factor that changes the numerical value of K.
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 9 / 14
Pressure

Pressure

  • Concentrations of substances in solution are given as molarity in square brackets [].
  • For reactants or products in the gaseous state, concentration is expressed by partial pressure, and the equilibrium constant is written Kp.
  • For example ammonia is manufactured from its elements:

N₂ + 3H₂ ⇌ 2NH₃, ΔH = −92 kJ

Kp = (PNH₃)2 / (PN₂ × (PH₂)3)

  • Four moles of gas become two moles, so the number of moles and the volume decrease.
  • Increasing pressure (and cooling) increases the rate of ammonia formation: raising the pressure on a gaseous equilibrium shifts it toward the side with fewer gas moles (smaller volume).
  • If the number of moles of gas is equal on both sides (e.g.
  • H₂ + I₂ ⇌ 2HI), pressure has no effect on the position.
  • Kp does not change with the partial pressures at constant temperature; the total pressure is the sum of the partial pressures of the gases and depends on the number of moles of each gas.
Comparison aspectMain ideaKeywords
Example 4.For 2SO₂ + O₂ ⇌ 2SO₃ the equilibrium partial pressures are P(SO₂) = 0.5 atm, P(O₂) = 0.25 atm and P(SO₃) = 2 atm: Kp = (2)2/((0.5)2 × 0.25) = 4/0.0625 = 64.2SO₂2SO₃equilibriumpartial
Example 5.For N₂ + 3H₂ ⇌ 2NH₃, if P(N₂) = 4 atm, P(H₂) = 2 atm and P(NH₃) = 8 atm, then Kp = 82/(4 × 23) = 64/32 = 2 atm−2.2NH₃then
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 10 / 14
Le Chatelier’s principle

Le Chatelier’s principle

  • Formulated by the French scientist Le Chatelier, it predicts the effect of factors such as concentration, temperature and pressure on a system at equilibrium:
  • If a system at equilibrium is subjected to a change in one of the factors affecting it (concentration, pressure, temperature), the equilibrium shifts in the direction that opposes (cancels) the change.
Change Shift
Add a reactant forward (consumes it)
Add a product (or remove a reactant) backward
Remove a product forward
Increase temperature toward the endothermic direction
Decrease temperature toward the exothermic direction
Increase pressure / decrease volume toward fewer gas moles
Catalyst no shift (only speeds up reaching equilibrium)
Example 6.→Consider N₂ + O₂ ⇌ 2NO, ΔH = +180 kJ (endothermic). Effect on the amount of NO:

adding more N₂ — forward, NO increases;

removing NO as it forms — forward, NO increases;

  • raising the temperature — toward the endothermic (forward) direction, NO increases, and K increases;
  • raising the pressure — no effect (2 moles of gas on each side);
  • adding a catalyst — no change in the amount of NO at equilibrium, but equilibrium is reached faster.
  • Example 7. For 2SO₂ + O₂ ⇌ 2SO₃, ΔH = −198 kJ: the yield of SO₃ is increased by high pressure (3 moles → 2 moles), by adding excess O₂ or removing SO₃, and by a lower temperature (the reaction is exothermic); in practice a moderate temperature and V₂O₅ catalyst are used so that the rate is acceptable.
Example 8 (temperature and dissociation).→For which reaction does heating increase the percentage dissociation? 2NH₃ ⇌ N₂ + 3H₂ with ΔH = +92 kJ (endothermic in the direction of dissociation): yes, heating shifts to the endothermic side so dissociation rises. For an exothermic dissociation heating would reduce it.
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 11 / 14
Catalysts

Catalysts

  • Heating speeds slow reactions but is expensive in industry, since it adds fuel costs to the cost of the product.
  • Chemists found that slow reactions can be accelerated by using certain substances which increase the rate without raising the temperature.
  • A catalyst is a substance, used in a small amount, that changes the rate of a reaction without itself being changed and without changing the position of equilibrium.
  • It lowers the activation energy and increases the rate of the forward and backward reactions equally and at the same time, so equilibrium is reached sooner but K is unchanged.
  • Catalysts are used in over 90 % of industrial processes (food, petrochemicals) and in car catalytic converters that turn the harmful gaseous combustion products which pollute the air into harmless ones.
  • Catalysts may be metals, metal oxides or other compounds.
  • Enzymes are protein molecules made in living cells that act as biological catalysts.
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 12 / 14
Light

Light

  • Some reactions are affected by light:
  • Photosynthesis — chlorophyll in plants absorbs light to form carbohydrates from CO₂ and water.
  • Photographic film — silver bromide in a gelatin layer light causes the positive silver ion to accept an electron from the bromide ion, Ag+ + e− → Ag; the bromine is absorbed in the gelatin. More light gives more silver.
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Part 14

Part 14

  • 📖

🔍 In-Depth Explanations, Deductions & Multi-Book Insights:

Chemistry Thanaweya Amma (2026/2027)

  • Lesson 1 – Chemical Reaction Rates and Le Chatelier's Principle
  • Detailed Academic Lesson Overview:
  • Chemical reactions are categorized into complete (irreversible) and reversible reactions.
  • When the rates of forward and backward reactions equalize, dynamic chemical equilibrium is established.
  • Le Chatelier's principle predicts how changing concentration, temperature, or pressure alters equilibrium.

🌳 Hierarchy Classification & Conceptual Roadmap

Chemical Equilibrium Principles

├── 1. Reaction Types & Rate (v)

  • │ ├── Irreversible Escaping gas or precipitate leaves reaction medium
  • │ └── Reversible Reactants and products coexist dynamically at equilibrium (rf = rb)
  • ├── 2. Law of Mass Action & Equilibrium Constant (Kc & Kp)
  • │ ├── Kc = [Products]p / [Reactants]r (Pure solids and liquids omitted)
  • │ └── Kp = (Pproducts)p / (Preactants)r (For gas phase reactions)

└── 3. Le Chatelier's Principle (Stress Factors)

Comparison aspectMain ideaKeywords
├── Concentration
  • Adding reactant shifts equilibrium forward
  • removing shifts backward
Addingreactantshiftsequilibrium
├── Pressure
  • Affects only unequal gaseous moles
  • increased pressure shifts toward fewer moles
Affectsunequalgaseousmoles
├── TemperatureONLY factor that changes the numerical value of Kc / Kpfactorchangesnumericalvalue
│ ├── Exothermic (ΔH < 0)Heating decreases Kc and shifts backwardHeatingdecreasesshiftsbackward
│ └── Endothermic (ΔH > 0)Heating increases Kc and shifts forwardHeatingincreasesshiftsforward
└── CatalystsIncrease rates of both forward and reverse equally without shifting positionIncreaseratesbothforward
•📋 Master Teacher's Blackboard Rules & Exam SecretsMasterTeacherBlackboardRules
RULE 1ONLY temperature can change the numerical magnitude of the equilibrium constant (Kc or Kp)temperaturechangenumericalmagnitude
RULE 2In exothermic reactions (Heat on right side): Raising temperature shifts backward and Kc DECREASESexothermicreactionsHeatright
RULE 3
  • Pressure has NO effect on reactions where total gaseous moles on reactants equal total gaseous moles on products (e.g
  • H2(g) + I2(g) ⇌ 2 HI(g))
Pressureeffectreactionstotal
EXAM TRAP
  • Catalysts decrease activation energy and accelerate the time required to reach equilibrium
  • but DO NOT shift the position of equilibrium and DO NOT alter Kc
Catalystsdecreaseactivationenergy
📊 Detailed Comparison MatrixEffect of External Stresses on Reversible System: N2(g) + 3 H2(g) ⇌ 2 NH3(g) (ΔH = -92 kJ)EffectExternalStressesReversible
External Stress Applied Direction of Equilibrium Shift Effect on [NH3] Effect on Kc Value
Increase Pressure Forward (towards fewer moles, 4 → 2) Increases Remains constant
Decrease Temperature (Cooling) Forward (produces heat) Increases Increases
Increase [N2] or [H2] Forward Increases Remains constant
Add Iron Catalyst No shift (rf and rb increase equally) No change Remains constant
Increase Temperature (Heating) Backward (absorbs heat) Decreases Decreases
Lesson 1: Chemical Reaction Rates and Le Chatelier’s Principle · 14 / 14