Understanding Chemical Kinetics

• Chemical Kinetics helps understanding how chemical reactions occur.
• Chemistry deals with change; substances transform into others with different properties via chemical reactions.
• Chemists aim to determine the feasibility, extent and speed of chemical reactions.
• Thermodynamics predicts reaction feasibility (∆G < 0 at constant temperature and pressure.)
• Chemical equilibrium determines the extent a reaction will proceed.
• Chemical kinetics focuses on the rate of reaction (speed) and factors influencing it.
• Understanding reaction rates is crucial in applications like food spoilage, dental material setting, and fuel burning in engines.
• Chemical kinetics originates from 'kinesis', a Greek term for movement.
• Thermodynamics indicates diamond transforms to graphite, but the reaction rate is imperceptible.
• Kinetic studies determine reaction speed and conditions to alter reaction rates, including concentration, temperature, pressure, and catalysts.
• At the macroscopic level, the interests are about amounts reacted or formed and their rates.
• At the molecular level, reaction mechanisms involve molecular orientation, energy during collisions.

Rate of a Chemical Reaction

• Reactions like silver chloride precipitation occur rapidly.
• Rusting of iron is slow.
• Inversion of cane sugar and starch hydrolysis proceed at moderate speeds.
• Reaction speed is the change in reactant or product concentration per unit time
• Reaction rate is expressed as decrease in reactant concentration or increase in product concentration.
• For a reaction R → P, rate is the change in concentration of R or P over time, assuming constant volume.

Average Rate

• Average rate relies on change in reactant/product concentration and time taken.
• Average rate is expressed as:
  • Rate of disappearance of R = -Δ[R]/Δt
  • Rate of appearance of P = +Δ[P]/Δt,
• Δ[R] is negative, rate is multiplied by -1 to yield a positive value.

Instantaneous Rate

• Instantaneous rate is the rate at a specific moment, found with the average rate at the smallest time interval (Δt approaches zero).
• Mathematically:
  • rinst = -d[R]/dt = d[P]/dt
• Graphically, instantaneous rate at time t is the tangent's slope on concentration vs. time curve.

Reaction Rate Units

• Reaction rate units are concentration/time, e.g., mol L⁻¹ s⁻¹.
• For gases, rate is expressed with partial pressures, units are atm s⁻¹.
• The unit of the rate constant depends on the order of the reaction

Rate Expression

• The rate of a chemical reaction at a given temperature depends the concentration of reactants or products.
• Rate law expresses reaction rate regarding reactant concentrations.
• Rate law is also known as rate equation or rate expression.

Rate Constant

• Rate constant is the proportionality constant in the differtial rate equation.
• Table 3.1 illustrates a reaction's rate decreases with time as reactant concentrations decline, so increasing reactant concentrations generally increase rates.

Rate Law

• Rate law: Reaction rate is given in molar concentration with each term raised to some power.
• Powers may or may not equal to stoichiometric coefficients.
• For a reaction aA + bB → cC + dD: Rate ∝ [A]ˣ [B]ʸ
• Rate = k [A]ˣ [B]ʸ.
• Where:
  • k = Rate constant
  • x and y shows how sensitive the rate is to the presence of concentrations of A and B
• Rate expression has reaction rate in terms of reactant molar concentration, each raised to a power possibly unequal to its stoichiometric coefficient.
• aA + bB → cC + dD. Rate = k [A]ˣ [B] ʸ.

Differential Rate Equation

• Rate = k [A]ˣ [B] ʸ is the differential rate equation, where k is the rate constant.

Order

• Order = x + y
• Cannot be predicted by a balanced equation, has to be theoretically determined

Instantaneous Rate Example

• Instantaneous rate for is given by
• Rate of reaction where stoichiometric coefficients of reactants and products are the same is:
• Rate=−Δ[𝐻𝑔]/Δ𝑡=− Δ[𝐶𝑙2]/Δ𝑡=Δ[𝐻𝑔𝐶𝑙2]/Δ𝑡*

The Rate Expression Equation

• Expressing:
• 2𝐻𝐼(𝑔)→𝐻_2(𝑔)+𝐼_2(𝑔)*
• Rate=−1/2 Δ[𝐻𝐼]/Δ𝑡=Δ[𝐻2]/Δ𝑡=Δ[𝐼2]/Δ𝑡*
• 5𝐵𝑟^−(𝑎𝑞)+𝐵𝑟𝑂_3^−(𝑎𝑞)+6𝐻^+(𝑎𝑞)→3𝐵𝑟_2 (𝑎𝑞)+3𝐻_2 𝑂(𝑙)*
• Rate=−1/5 Δ[𝐵𝑟^−]/Δ𝑡=−Δ[𝐵𝑟𝑂_3^−]/Δ𝑡=−1/6 Δ[𝐻^+]/Δ𝑡=1/3 Δ[𝐵𝑟_2]/Δ𝑡=1/3 Δ[𝐻_2 𝑂]/Δ𝑡*
• In gaseous reactions, concentration is directly proportional to partial pressure. Rate can be expressed as change in partial pressure.

Order of a Reaction

• Rate = k [A]ˣ [B] ʸ, x and y show how sensitive the rate is to changes in A and B concentration.
• Overall order is x + y.
• x is order re: A and y is order re: B.
• Sum of reactant concentration powers in the rate law expresses reaction order.
• Rate can be 0, 1, 2, 3, or fraction.
• Zero order: Rate is independent of reactant concentration.

Elementary Reactions

• Balanced chemical equation does not show how a reaction takes place, a reaction rarely finishes in one step.
• Elementary reactions take place in one step.
• Complex reactions involve steps (mechanism).

Consecutive Reactions

• These may be consecutive reactions, for example, when Oxidation of ethane to CO2 and H2O and passes through intermediate stage(s) in which alcohol, aldehyde and acid are formed.

Rate Constant

• k = Rate/[A]*[B]
• Rate constant unit depends order, measured in mol L-1 and time(s)

Molecularity

• This helps to understand mechanism.
• Reacting species numbers (atoms, ions or molecules) involved in an elementary reaction, which collide simultaneously is the molecularity of a reaction.
• Molecularly can be:
• Unimolecular
• Bimolecular reactions
• Trimolecular. or termolecular reactions'
• Complex reactions: order given the slowest step + molecularity of slowest step is same order as reaction.

Integrated Rate Equations

• Zero-order reaction: One that means that the reaction is proportional to zero power of reactants concentrations
• R-->P
• This gives Rate =-d[R]/dt =k[R]0
• This also means d[R] = k x 1

Integrating sides of a zero-order reaction

• [R] = -kt + I
• When At t=0, the reaction R =[R], this is the initial concentration the equations = [R] = - k * 0 + 1, resulting to solving I
• You can find the rate constant, which is k =([R]-[R])\t*

First Order Chemical Reactions

• the rate reaction is proportional with the power of the concentration which also depends with proportional of reaction R for example R-->P
• Rate= −d[R]/dt=k[R] or d[R]/ [R] =-kdt, then it gives with integrationIn [R] = −kt+ I, then you can easily find In[R] = -Kt =In [R]

Catalysis

• Catalyst increases the rate of a reaction without any permanent chemical change.
• Catalyst can assist in a reaction, so the substance can be called inhibitor.
• Catalytic theory: a catalyst forms temporary bonds with the reactants as an intermediate complex. The transitory assists in creating product + catalyst.
• It is believed catalyst causes alternate pathway to increase and causes lowering energy path and hence potential energy can be achieved fast.
• Small amount of catalyst large amounts the rate, and does not affect AG of action... Also found that a catalyst can affect equilibrium constant.

Arrhenius Equation

• According to this Equation:
• k=𝐴𝑒^(−𝐸_𝑎/𝑅𝑇)*
• Where temperature is raised, the energy will move to maximum height, and broadens at large such creates larger reactions will occur at high speeds and under the must be at all times. Increasing :
• In k=−𝐸_𝑎/𝑅 [1/𝑇]+In 𝐴* it means the increasing a decrease and cause and and the temperature.

Collision Theory

• This theory mentions that the reactant molecules are likely hard spheres and the reaction will occur in the the volume's measurements, The # os all per second can is called Frequency.
• Rate=𝑍_(𝐴𝐵) 𝑒^(−𝐸_𝑎/𝑅𝑇)* = It has accurate constants
• You can take to the account collisions/