14.S: Chemical Kinetics (Summary)

A general chemistry Libretexts Textmap organized around the textbook
 Chemistry: The Central Science
by Brown, LeMay, Bursten, Murphy, and Woodward

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These is a summary of key concepts of the chapter in the Textmap created for "Chemistry: The Central Science" by Brown et al.

14.1: Factors that Affect Reaction Rates

chemical kinetics – area of chemistry dealing with speeds/rates of reactions

• rates of reactions affected by four factors
  1. concentrations of reactants
  2. temperature at which reaction occurs
  3. presence of a catalyst
  4. surface area of solid or liquid reactants and/or catalysts

14.2: Reaction Rates

• reaction rate – speed of a chemical reaction
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14.1.1 Rates in Terms of Concentrations

• rate calculated in units of M/s
• brackets around a substance indicate the concentration
• instantaneous rate – rate at a particular time
• instantaneous rate obtained from the straight line tangent that touches the curve at a specific point
• slopes give instantaneous rates
• instantaneous rate also referred to as the rate

14.1.2 Reaction Rates and Stoichiometry

• for the reaction
• 

14.3: Concentration and Rate

• equation used only if C and D only substances formed
• Rate = k[A][B]
• Rate law – expression that shows that rate depends on concentrations of reactants
• K = rate constant

14.2.1 Reaction Order

• Rate = k[reactant 1]^m[reactant 2]ⁿ
• m, n are called reaction orders
• m+n, overall reaction order
• reaction orders do not have to correspond with coefficients in balanced equation
• values of reaction order determined experimentally
• reaction order can be fractional or negative

14.2.2 Units of Rates Constants

• units of rate constant depend on overall reaction order of rate law
• for reaction of second order overall
• units of rate = (units of rate constant)(units of concentration)²
• units of rate constant = M^-1s¹

14.2.3 Using Initial Rates to Determine Rate Laws

• zero order – no change in rate when concentration changed
• first order – proportional changes in rate
• second order – increase rate by 2² or 3³, etc…
• rate constant does not depend on concentration

14.4: The Change of Concentration with Time

• rate laws can be converted into equations that give concentrations of reactants or products

14.3.1 First-Order Reactions

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• corresponds to y = mx + b
• equations used to determine:
• 1) concentration of reactant remaining at any time
• 2) time required for given fraction of sample to react
• 3) time required for reactant concentration to reach a certain level

14.3.2 Half-Life

• half-life of first order reaction
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• half-life – time required for concentration of reactant to drop by one half initial value
• t_1/2 of first order independent of initial concentrations
• half-life same at any given time of reaction
• in first order reaction – concentrations of reactant decreases by ½ in each series of regularly spaced time intervals

14.3.3 Second-Order Reactions

• rate depends on reactant concentration raised to second power or concentrations of two different reactant each raised to first power
• Rate = k[A]²
• 
• 
• half life dependent on initial concentration of reactant

14.5: Temperature and Rate

• chemiluminescent reaction – reaction that produces light
• rate constant must increase with increasing temperature

14.4.1 The Collision Model

• collision model – molecules must collide to react
• greater number of collisions the greater the reaction rate
• for most reactions only small amount of collisions lead to a reaction

14.4.2 Activation Energy

• Svante Arrhenius
• Molecules must have a minimum amount of energy to react
• Energy comes from kinetic energy of collisions
• Kinetic energy used to break bonds
• Activation energy, E_a – minimum energy required to initiate a chemical reaction
• Activated complex or transition state – atoms at the top of the energy barrier
• Rate depends on E_a
• Lower E_a means faster reaction
• Reactions occur with collisions and orientation of molecules

14.4.3 The Arrhenius Equation

• reaction rate data:
• (Arrhenius Equation)
• k = rate constant, E_a = activation energy, R = gas constant (8.314 J/mol K), T = absolute temperature, A = frequency factor
• A relates to frequency of collisions, favorable orientations
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• ln k vs 1/t graph has slope –E_a/R and y-intercept ln A
• for two temperatures:
• 
• used to calculate rate constant, k₁ and T₁

14.6: Reaction Mechanisms

• reaction mechanism – process by which a reaction occurs

14.5.1 Elementary Steps

• elementary steps – each step in a reaction
• molecularity – if only one molecule involved in step
• unimolecular – if only one molecule involved in step
• bimolecular – elementary step involving collision of two reactant molecules
• termolecular – elementary step involving simultaneous collision of three molecules
• elementary steps in multistep mechanism must always add to give chemical equation of overall process
• intermediate – product formed in one step and consumed in a later step

14.5.2 Rate Laws of Elementary Steps

• if reaction is known to be an elementary step then the rate law is known
• rate of unimolecular step is first order (Rate = k[A])
• rate of bimolecular steps is second order (Rate = k[A][B]
• first order in [A] and [B]
• if double [A] than number of collisions of A and B will double

14.5.3 Rate laws of multistep mechanisms

• rate-determining step – slowest elementary step
• determines rate law of overall reaction

14.5.4 Mechanisms with an Initial First Step

• intermediates usually unstable, low and unknown concentrations
• whenever a fast step precedes a slow one, solve for concentration of intermediate by assuming that equilibrium is established in fast step

14.7: Catalysis

• catalyst – substance that changes speed of chemical reaction without undergoing a permanent chemical change

14.6.1 Homogeneous Catalysis

• homogeneous catalyst – catalyst that is present in same phase as reacting molecule
• catalysts alter E_a or A
• generally catalysts lowers overall E_a for chemical reaction
• catalysts provides a different mechanism for reaction

14.6.2 Heterogeneous Catalysis

• exists in different phase from reactants
• initial step in heterogeneous catalyst is adsorption
• adsorption – binding of molecules to surface
• adsorption occurs because ions/atoms at surface of solid extremely reactive

14.6.3 Enzymes

• biological catalysts
• large protein molecules with molecular weights 10,000 – 1 million amu
• catalase – enzyme in blood and liver that decomposes hydrogen peroxide into water and oxygen
• substrates – substances that undergo reaction at the active site
• lock-and-key model – substrate molecules bind specifically to the active site
• enzyme-substrate complex – combination of enzyme and substrate
• binding between enzyme and substrate involves intermolecular forces (dipole-dipole, hydrogen bonding, and London dispersion forces)
• product from reaction leaves enzyme allowing for another substrate to enter enzyme
• enzyme inhibitors – molecules that bind strongly to enzymes
• turnover number – number of catalyzed reactions occurring at a particular active site
• large turnover numbers = low activation energies