SectionC5_BB - Tagged
Page 1: Introduction to Chemical Kinetics
Course Title: CHEM10021 Chemistry for Bioscientists I
Section: C5 Chemical kinetics, Rate laws
Year: 1824
Institution: The University of Manchester
Page 2: Overview of Chemical Kinetics
Study Focus: Rates of chemical reactions
Key Measurements:
Consumption of reactants
Formation of products
Influencing Factors:
Pressure, temperature, catalysts
Goals:
Gain insights into the mechanisms of reactions
Identify the sequence of elementary steps
Page 3: Stoichiometry
Definition: Study of the numbers of reactant and product molecules
Example Reaction:
1 C6H12O6 + 6 O2 -> 6 CO2 + 6 H2O
Concentration Monitoring:
Concentrations change with time
Temperature (T) held constant during reaction monitoring
Page 4: Measuring Concentrations
Biological Reactions Timescales:
Ranges from 10^-15 to approximately 10^7 seconds
Required Experimental Techniques:
Pressure measurement (for gaseous species)
Spectrophotometry (very fast)
Conductivity measurement (ionic solutions)
pH measurement (for H+/OH- reactions)
Polarimetry (for chiral species)
Page 5: Reaction Rate
Definition: Rate of reaction is defined as the number of moles of reaction per unit time
Example Reaction:
1 C6H12O6 + 6 O2 -> 6 CO2 + 6 H2O
Calculation Method:
Rate of loss of C6H12O6 or 1/6 rate of production of CO2
Units: mol dm^-3 s^-1
Monitoring: Concentration changes of a single reactant or product
Page 6: General Rate Expression
General Reaction Form:
aA + bB -> cC + dD
Rate Expressions:
Rate = -1/a * (d[A]/dt) = -1/b * (d[B]/dt) = 1/c * (d[C]/dt) = 1/d * (d[D]/dt)
Estimation: Rate can be estimated by measuring concentration at two time points and dividing by the elapsed time.
Page 7: Instantaneous Rate
Definition: Initial rates can be determined
Focus: Comparison of rates at later times (t1, t2) against initial time (t)
Page 8: Rate Laws and Rate Constants
Form of Rate Law:
rate of reaction = k[A][B]
Components Explained:
k is the rate constant; unique to each reaction and temperature-dependent
Units of k: dm³ mol^-1 s^-1
Relationship: rate (mol dm^-3 s^-1) = k (dm³ mol^-1 s^-1) * [A] (mol dm^-3) * [B] (mol dm^-3)
Page 9: Order of Reaction
Definition: Order of a reaction refers to the power to which the concentration is raised in the rate law
Examples:
rate = k[A][B]: first order in both A and B
rate = k[A]²: second order in A
Overall Order: Sum of the individual orders
Page 10: Types of Reaction Orders
Zero Order: Rate is independent of concentration
Example: rate = k (when [A]° = 0)
Fractional Order: Example rate = k[A]^(1/2)[B]
Negative Order: Example: rate = k[C]/[A][B]
Note: If the rate law does not fit conventional forms, the reaction does not have an order.
Page 11: Finding the Order of Reaction: Isolation Method
Important Note: Rate laws are deduced experimentally and cannot solely depend on the stoichiometric equation.
Relation: Rate law is often influenced by the reaction mechanism.
Page 12: Isolation Method Explained
Concept: All reactants except one are kept constant
Approximation Example:
True rate law might be: rate = k[A][B]²
Approximating [B] by its initial value [B]₀ leads to a pseudo first-order reaction.
Page 13: Initial Rate Method
Methodology: Measurements taken before concentrations have varied significantly.
Finding Order: By plotting log(initial rate) (ro) against log(concentration).
Page 14: Logarithmic Representation
Derivation for Second Order Reaction:
If rate = k[B]² and ro is the initial rate:
log(ro) = log(k) + 2log([B])
Mathematical Representation: y = mx + c, where slope (m) is 2.
Page 15: Conclusion
Section: C5 Chemical kinetics, Rate laws
Year: 1824
Institution: The University of Manchester