Unit 1: Reaction Kinetics

Reaction kinetics

• Study of the rates of reactions and the factors that affect the rates.

• Rate of reaction = speed at which a reaction occurs

• Expressed in terms of the change in amount of a reactant (consumption/ decrease) or product (production/ increase) in a certain interval of time.

Mass change

• Open system, the NO₂ (g) escapes, total mass of system decreases.

• Electronic balance and stopwatch.

• Rate = Δmass/Δtime

Colour change

• Spectrophotometer and stopwatch

• Colour intensity increases as more Cu(NO₃)₂ (aq) is produced.

• Rate = Δcolour intensity/Δtime

Pressure change

• Closed (sealed) container, a pressure gauge and stopwatch.

• To measure pressure change, there must be a change in the total number of moles of gas in the reaction.

• Rate = Δpressure/Δtime

Temperature change

• Exothermic = temperature increases

• Endothermic = temperature decreases

• Thermometer and stopwatch

• Rate = Δtemperature/Δtime

pH change

• pH meter and stopwatch

• HNO₃ (aq) is used up, acidity decreases and pH of the solution increases

• Rate = ΔpH/Δtime

Concentration-Time graph

• Rates of reaction do not typically remain constant for the entire duration of a reaction.

• Initially, rates are fast because [reactants] are high.

• Rates decrease as reaction proceeds since [reactants] decrease.

• The exact rate at any particular time can be obtained by determining the slope of a line that is tangent to the concentration-time curve at that point.

Some reactions may show an increase in rate after a slow start

• exothermic reactions

• changes in surface area

• coatings on the reactant surface

• autocatalysis

TEMPERATURE affects reaction rates

• When temperature is increased, the time required for a reaction decreases.

• Rate = Δamount of reactant or product/Δtime

• Increasing temperature increases the rate.

• In general, for many SLOW reactions, a 10°C increase in temperature doubles the reaction rate.

CONCENTRATION affects reaction rates

Increasing reactant concentration increases rate.

PRESSURE affects reaction rates

• Partial pressure of a gas is proportional to the moles of a gas when temperature is constant.

• Increasing partial pressure of a gas is equivalent to increasing concentration; therefore, rate increases.

NATURE OF REACTANTS affects reaction rates

• Some reactions are naturally faster than others.

• Reactions that involve breaking weak bonds or transferring electrons that are weakly held are faster than those in which bonds are strong and electrons are held strongly.

• These are fundamental differences in the chemical reactivity of different substances which we have no control.

SURFACE AREA affects reaction rates

• Heterogeneous reaction = reactants are in different phases (i.e., solid and liquid or solution)

• Homogeneous reaction = reactants in the same phase (e.g., two gases).

• Increasing exposed area (surface area) increases the rate.

  • Crushing, powdering, grinding, etc.

• Only affects the rate of heterogeneous reactions.

CATALYSTS AND INHIBITORS affect the reaction rates

• Catalyst = chemical which reduces reaction rate but is regenerated in its original form at the end of a reaction.

• Inhibitor = chemical substance that reduces reaction rate by combining with a catalyst or one of the reactants in such a way that it prevents the reaction from occurring.

PHASE OF REACTANTS affects the reaction rates

• Oppositely charged aqueous ionic reactants lead to very fast reaction rates.

• Few bonds or weak bonds between reactants have faster rates than many and strong ones.

• Homogeneous reactants have faster rates than heterogeneous phases.

• Undergoing 2-particle collisions is faster than those involving 3 or more particles.

• Order of reaction rate (fastest → slowest): Aqueous ions > Gases or Liquids > Solids

  • Particles in a solid do not have free movement.

  • Particles in a solution are close and have free movement.

Collision Theory (Kinetic Molecular Theory)

• Reactions depend on collisions between reactant molecules; however, not all collisions lead to a reaction.

• Successful or effective collisions lead to the formation of products, while those that don’t are ineffective or unsuccessful.

Orientation or geometry

Colliding reactant molecules must be oriented in a favourable position to allow bonds to break and atoms to rearrange.

Activation energy

For a reaction to occur, molecules need to collide with sufficient energy to break bonds so that atoms can rearrange and form new bonds.

• Activation energy = minimum amount of energy needed for a reaction to occur.

Collision Theory - Increasing Concentration

• Increasing the concentration of reactants (or partial pressure if gases) increases the frequency of possible collisions and hence increases the rate.

• The percentage or fraction of collisions that are effective remains the same.

Collision Theory - Increasing Temperature

• Increasing temperature increases the average kinetic energy of the molecules (i.e., increases the speed at which the molecules are moving).

• Two effects:

  • molecules collide more often

  • molecules collide with more energy

• Increasing temperature increases the percentage or fraction of effective collisions.

• Rate increase that accompanies an increase in temperature is primarily due to the increased energy of collisions.

Potential energy

Stored energy

• Related to the energy of the electrons in the chemical bonds, as well as the number and types of atoms in the molecule.

• Potential energy increases when bonds are broken and decreases when new bonds are formed.

  • To break a bond, energy must be put into the reaction

  • To form a bond, energy is released

Kinetic energy

Energy of motion

• Energy exists as a result of a movement of molecules within a system.

• Kinetic energy can be related to the temperature of the system.

Enthalpy (H)

Heat of reaction

• TOTAL kinetic and potential energy that exists in a system at constant pressure.

• ΔH = H_products - H_reactants

During a chemical reaction, the bonds of the reactant molecules are broken, the atoms are rearranged, and new bonds are formed.

• Exothermic reaction: H_products < H_reactants

  • ΔH < 0

  • Heat is released into the surroundings → Warmer temperature

• Endothermic reaction: H_products > H_reactants

  • ΔH > 0

  • Heat is absorbed into the system → Cooler temperature

Kinetic energy distribution (Maxwell-Boltzmann distribution)

• Room temperature and pressure: molecules undergo about 10¹⁰ collisions/second.
  → The lack of reactivity is NOT due to the lack of collisions.

• Temperature increases → Molecules have more energy → Reaction rate increases.

• Some molecules have high KE while others have low KE.

• Increasing temperature increases the average energy of the system.

• Only the molecules with KE ≥ minimum energy will react.

• The increased reaction rate due to an increase in temperature is PRIMARILY DUE to the increased number of molecules with sufficient energy to react.

Activation energies

• Activation energy (E_a) = the minimum amount of energy required for the reactants to form the activated complex.

• Reactant molecules approach each other → Slow down → KE is converted into PE.

• e^-repulsion → e^- gain PE by absorbing KE → slow down the reaction → reactant molecules gain enough energy (≥ E_a) → activated complex is formed → e^-repulsion → product molecules move away from each other → PE is converted into KE.

Activated complex

• High energy, unstable arrangement of atoms which occurs when reactants are in the process of rearranging to form new products.

• Total energy = PE + KE

• Molecules do not have ideal geometry → Reaction takes place with an increased activation energy.

• E_a = Activated Energy - H_reactant

• ΔH = H_products - H_reactants

Reversible reaction

• Reactant ⇆ Product

• E_a(fwd): activation energy for the forward reaction (reactant to activated complex).

• E_a(rev): activation energy for the reverse reaction (products to activated complex).

• E_a(fwd) = E_a(rev) + ΔH (ΔH > 0)

• E_a(rev) = E_a(fwd) + ΔH (ΔH < 0)

• The higher the activation energy, the slower the reaction rate, and vice versa.

  • High E_a = Low reaction rate

  • Low E_a = High reaction rate

Reaction mechanism

• Reactions often occur as a result of several elementary steps.

• Elementary process = individual step in a reaction mechanism.

• Reaction mechanism = sequence of steps that make up an overall reaction.

  • Each step has its own peak.

  • Overall activation energy: the difference between the reactants and the highest peak.

  • Activation energy for each step is the PE difference between the activated complex and the reactants involved in that step.

• Rate-determining step = the slowest step in a mechanism (highest E_a).

Effects of catalysts on E_a

• Catalysts do NOT change the ΔH.

• Catalysts lower the activation energy → More molecules have enough energy to react → Reaction rate increases.

• Catalysts decrease both the E_a(fwd) and E_a(rev).

• Catalysts work by providing an alternative reaction mechanism with a lower E_a.

Reaction intermediate vs. Catalyst

• Reaction intermediate: product → reactant

• Catalyst: reactant → product

• Both intermediate species and catalyst cancel out when the individual steps are added to get the overall reaction.