Reaction Rates and Equilibrium
Examples of exothermic processes
- Condensation - Combustion - Dissociation of strong acids - Solidification of cement, concrete, epoxy - Thermite reaction
Effect of Concentration on Reaction Rate
- Higher reactant concentration leads to more frequent collisions - More frequent collisions increases the rate of reaction. **As the CONCENTRATION (not necessarily the acc amount) of reactants increases, reaction rate increases.
Examples of endothermic processes:
- Melting - Photosynthesis - Many decomposition and dehydration reactions - Dissociation of some salts
Conditions for a reaction to occur:
- Molecules or particles must collide. - Collisions must be effective
Effect of Temperature on Reaction Rate
- Temperature represents a distribution of kinetic energies - Only a small fraction of molecules have sufficient energy to react. - Increasing temperature increases the proportion of reacting molecules *** As temperature increases, reaction rate increases.
Reaction rates depend on:
- reactant concentrations - temperature - pressure
Equilibrium constant (Keq)
: the ratio of the equilibrium concentrations of products raised to their stoichiometric ratios to the concentrations of reactants raised to their stoichiometric ratios; a measure of the relative quantities of reactants and products in a system at equilibrium Aa+ bB ↔ cC + dD [C]c[D]d Keq = ------ [A]a[B]b
reversible reaction
A reaction in which neither the forward nor reverse reaction is highly favored. - Reactants refer to the substances on the left side of the arrow - Products refer to the substances on the right side of the arrow. Example: CO₂ (aq) + H₂O(l) ↔H₂CO₃ (aq)
spontaneous reaction
A reaction that will proceed without any outside energy - Reactants have more free energy than products ∆Grxn < 0
Activation Energy and Reaction Rate
Activation energy is the minimum energy that reactants must attain for a reaction to occur. • Higher activation energy → more difficult to reach • Higher activation energy → lower reaction rate
Chemical Equilibrium
At equilibrium, when there are more reactants than products, • reactants are favored. • equilibrium "lies to the LEFT." At equilibrium, when there are more products than reactants, • products are favored. • equilibrium "lies to the RIGHT."
How do catalysts work?
Catalysts increase reaction rate: - Hold reactant molecules in proper orientation to react - Form temporary reaction intermediates with reactant molecules to increase rate • Provide alternate reaction pathway • Release catalyst when consumed - Weaken or break reactant bonds , increasing reactivity
Uses of Catalysts
Catalytic converters in automobiles: • reduce toxic exhaust Industrial chemical processes: • reduce waste. • speed production. • lower energy requirements and consumption Fuel cells: • increase reaction rates.
Changes in temperature
Changing the temperature of a system depends on how heat flows. *** ONLY stressor that changes Keq (the constant)!!! • EXOthermic reaction: reactants → products + heat • Adding heat causes shift to LEFT . • Removing heat causes shift to RIGHT. • Endothermic reaction: reactants + heat → products • Adding heat causes shift to RIGHT. • Removing heat causes shift to the LEFT
chemical bonds
Chemical reactions rearrange the atoms in molecules by making and/or breaking bonds. - Bonds are stable configurations of electrons - Electrons need enough energy to leave their current bonds to form new or different bonds.
Conditions for effective collisions:
Collision must be energetic enough. - Energy comes from the kinetic energy of random collisions Collision should be in the correct ORIENTATION
rate determining step
Consider the decomposition of ozone: Step 1: O₃(g) → O₂(g) + O(g) (fast) Step 2: O₃(g) + O(g) → 2O₂(g) (slow) Overall reaction: 2O₃(g) → 3O₂(g) ∙ The slower step 2 is the rate-determining step in a reaction mechanism that has the slowest rate and that therefore controls the rate of the overall reaction.
activation energy
EA: minimum energy needed for a chemical reaction to proceed - is a barrier to reaction - must be overcome before reaction will proceed. **the greater the activation energy for a reaction, the less likely the reaction is to proceed, even if it is spontaneous. speed of reaction and activation energy are inversely proportional.
Activation Energy calculation
Ea = Hactivatedcomplex - Hreactants
Endothermic Reaction Pathway
Endothermic reactions absorb energy - Reactant energy is LOWER than product energy - Magnitude of difference indicates magnitude of ΔHrxn.
*
Equilibrium can exist in a system even if change is occurring in the system.
Exothermic Reaction Pathway
Exothermic reactions RELEASE energy - Reactant energy higher than product energy - Magnitude of difference indicates magnitude of ΔHrxn.
stresses
External changes called stresses can disrupt a chemical system at equilibrium. Stresses include: • change in concentration. (i.e. adding/subtracting more reactant/product) • change in pressure. • change in temperature.
Forward and Reverse Reaction Rates
Forward and reverse rates may be equal or unequal.
G
Free Energy
Le Chatelier's principle
If a stress is applied to a chemical system at equilibrium, the system will respond by shifting in a direction to counteract the stress, and a new equilibrium will be established.
which description explains the role of activation energy in a chemical reaction??
It provides reactants with sufficient energy for bonds to break and reform.
activation complex
It's the peak on the graph.
According to Le Chatelier's Principle, what happens to K when the concentrations of reactants is doubled.
K remains the same. ** Remember that TEMP is the ONLY stressor that actually changes K!!!!
Interpreting K
Keq < 1: equilibrium lies to the LEFT Keq > 1: equilibrium lies to the RIGHT Keq =1:neither reactants nor products favored • The value of Keq is different for each reaction system. • Keq depends on pressure, temperature , and volume of system. • Changes in conditions can change the position of equilibrium.
Using the Elementary step, determine the Molecularity and Rate Law: A +A --> B + C
Molecularity: Bimolecular - rate is dependent on the concentration of both reactants (2 molecules of A) Rate Law: rate = k[A][A] or rate = k[A]²
Using the Elementary step, determine the Molecularity and Rate Law: A --> B + C
Molecularity: Unimolecular - rate is dependent only on the concentration of A Rate Law: rate = k[A]
Using the Elementary step, determine the Molecularity and Rate Law: A + B → C + D
Molecularity: Bimolecular - rate is dependent on the concentration of both reactants (A and B) Rate Law: rate = k[A][B]
Kinetic theory
Molecules at a given temperature have different speeds
consider the balanced equation for the decomposition of ozone: 2O3(g)--> 3O2(g) Can the rate law for this chemical reaction be deduced from its balanced equation?
No; the rate law must be determined experimentally.
What is the rate of the reaction under these conditions? R = k [NO]^2 [O2], k = 7,000 M-2 • s-1 [NO] = 0.040 M, [O2] = 0.060 M
R = 0.672 M • s-1 How? R = 7,000 * [0.04]^2 * [0.06] = 0.672 M
Consider the following chemical reaction of bromothymol blue indicator. It appears yellow in undissociated form and blue in its dissociated aqueous solution. HC2H3O2(aq) <--> H+(aq) + C2H3O2-(aq) yellow blue What will be the color of the solution if a large amount of H2CO3 is added? The solution will remain yellow. The solution will turn blue. The solution will turn pink. The solution will turn green.
REmain yellow Carbonic acid dissociation in water: H₂CO₃ ⇄ H⁺(aq) + HCO₃⁻(aq). According to Le Chatelier's Principle the position of equilibrium moves to counteract the change.
Using Rate Laws
Rate laws - find new rates when the concentration changes • Zero order: rate does not change • First order: proportional to concentration • Doubling concentration results in doubling the rate • Second order: two types • Having all concentrations results in quartering the rate
Forward reaction rate:
Rate of reaction from left to right A + B → C + D
Reverse reaction rate:
Rate of reaction from right to left C + D → A + B
Effect of Surface Area on Reaction Rate
Reactions occur at the surface. - Increasing the surface area increases particles available for reaction. - Decreasing the particle size increases the surface area. **As the surface area of a ______reactant increases, reaction rate increases.
Factors Affecting Reaction Rate
Reactions require collisions. • Factors that increase frequency or energy of collisions will increase reaction rate. • Factors that decrease frequency or energy of collisions will decrease reaction rate.
How will a DECREASE in pressure affect the following equation? 2H2 (g) + O2 (g) <--> 2H2O (g) + heat
Shift towards REACTANTS.
Spontaneous Reactions and Activation Energy
Spontaneous reactions are self-sustaining. • - Energy released during the reaction provides energy to additional reactant molecules to overcome their activation energy - Initial energy input may be required to overcome activation energy
Determining the Rate Law if the First Step Is a Fast Step Consider the following reaction mechanism:
Step 1: NO(g) + Cl2(g) ↔ NOCl2(g) (Fast) Step 2: NOCl2(g) + NO(g) → 2NOCl(g) (slow) Overall reaction: 2NO(g) + Cl2(g) → 2NOCl(g) • Since Step 2 is the slow step, you would expect the rate law for the reaction to be the same as the rate law for Step 2: rate = k₂[NOCl₂][NO]. • But NOCl₂ is an intermediate, so it cannot appear in the rate law.
Determining Rate Law from Molecularity of Elemental Steps Consider the following reaction mechanism
Step 1: NO2(g) + NO2(g) → NO3(g) + NO(g) (slow) Step 2: NO3(g) + CO(g) → NO2(g) + CO2(g) (fast) Overall reaction: NO2(g) + CO(g) → NO(g) + CO2(g) • Rate law for Step 1: rate = k₁ [NO₂]² • Rate law for Step 2: rate = k₂ [NO₃][CO] • Rate law for overall reaction determined by rate law of the slow (rate-determining) step • Therefore, rate law for overall reaction: rate = k₁ [NO₂]²
Consider the decomposition of Ozone
Step 1: O₃ (g) → O₂ (g) + O (g) Unimolecular Step 2: O₃ (g) + O (g) → 2O₂ (g) Bimolecular Sum: 2O₃(g) + O (g) → 3O₂ (g) + O (g) Cut O(g) from each side. Overall reaction: 2O₃ (g) → 3O₂ (g)
Enzyme function (cont,)
Structure or reactivity changes when body conditions change. • The structure and reactivity of an enzyme are controlled by feedback mechanisms in the body. • These changes allow for precise regulation of body conditions.
catalyst
Substance that increases the rate of a reaction, but is not used up. - Provides an alternative reaction pathway with lower activation energy - Is not a reactant or product
what is the most likely reason that refrigerating most foods reduces the rate at which they spoil?
The lower temperature reduces molecule speeds, reducing the number of effective collisions.
Enzyme function
The structure of enzymes relates to the substrate it binds to. • An enzyme and its substrate work like a "lock-and-key" mechanism. • Only correct substrate(s) can bind to an enzyme
Enthalpy
The total heat content of a system at a constant pressure, commonly denoted as H.
Effect of Pressure on Reaction Rate
To make a GAS reaction go faster: - Add reactants - Decrease volume For reactions involving gases: - Increasing pressure increases concentration, which increases reaction rate. *** Increasing pressure increases reaction rate if the reactants contain more moles of gas than the products.
Heterogeneous Catalyst
a catalyst that is in a different phase than the reactants in a chemical reaction' - Easier separation of reactants and catalyst
Homogeneous catalyst
a catalyst that is in the same phase as the reactants in a chemical reaction Examples include: • Cl in atmospheric decomposition of O₃ • Carbonic anhydrase in conversion of CO2 to H₂CO₃ in blood
Dynamic equilibrium:
a condition in a chemical system in which the rates of forward and reverse reactions are equal At first only A and B are present. So the forward reaction rate is very high, whereas the reverse reaction rate is 0 because we don't have any C or D. Once C and D are present, the reverse reaction rate starts increasing while the forward reaction rate is decreasing. We have a smaller and smaller concentration of A and B. • The reaction still occurs. • Concentrations of reactants and products do not change. Eventually, the two reaction rates are going to be ______. At that point, the concentration of the reactants and the products does not change, but the reaction is still occurring.
Reaction pathway graph
a diagram indicating the change in energy between reactants and products - Is sometimes called potential energy diagram - Has energy on vertical axis - Has reaction progression on horizontal axis - Has reactants to the left - Has products to the right
Collision Theory
a model for chemical reactions that requires particles to collide in order to react - • Every reaction begins with collision of molecules or particles.
Elementary reaction
a single step in a reaction mechanism • The elementary steps always add up to give the balanced chemical equation of the overall reaction process. Step 1: NO(g) + Cl₂(g) → (g) Step 2: NOCl₂(g) + NO(g) → 2NOCl(g) Sum: 2NO(g) + Cl₂(g) + NOCl₂(g) → 2NOCl(g) + NOCl₂(g) Simplified: 2NO(g) + Cl₂(g) → 2NOCl(g)
Reaction intermediate
a substance that is formed in one step of a reaction mechanism and consumed in a subsequent step • The reaction intermediate does not appear in the balanced equation for the overall reaction.
Reaction quotient (Q):
a value calculated by applying actual concentrations of components in a chemical reaction to the equation for that reaction's equilibrium constant *** Still works for NON-equilibrium concentrations!!! ** can be used to predict the direction of change
When Q < K
abundance of REACTANT shift to PRODUCT
When Q > K
abundance of products shift to REACTANTS
On a reaction pathway graph, the energy of the activated complex is ___________ than that of the reactants or products. a. always higher b. always lower c. sometimes higher
always higher
Unimolecular
an elementary reaction that involves a single reactant molecule A→ product
Bimolecular
an elementary reaction that involves two reactant molecules A + A → product OR A + B → product
common ion
an ion that is present in a system at equilibrium and an external ionic compound that can be added to the system. EX: If HCl is added to the system at equilibrium below, the common ion, Cl⁻ , is added.
The synthesis of water proceeds according to this equation: 2H2(g) + O2(g) --> 2H2O(g) How can the rate of this reaction be increased?
by increasing the temperature of the reactants
What is the best definition of a rate-determining step?
elementary step that LIMITS HOW FAST overall reaction proceeds
rate constant
is the proportionally constant in a rate law, k ***The rate constant does not depend on concentration.
How does changing the concentration of the reactants change the parts of a rate law?
it changes the rate, R
An exothermic reaction a. requires energy b. produces energy c. is energy-neutral.
produces energy
rate law
rate=k[A]ⁿ[B]ⁿ A + B → C + D an equation that shows the relationship between the concentration of reactants and the reaction rate.
nonspontaneous reaction
reaction that does not favor the formation of products - Products have more free energy than Reactants - ∆Grxn > 0
Look at the representation of a reaction mechanism. Step 1: Unimolecular (fast) Step 2: Bimolecular (slow) Step 3: Unimolecular (fast) Which conclusion about the rate law for the reaction above is reasonable?
second order overall because step 2 is bimolecular.
Molecularity
the number of molecules or atoms that participate as reactants in an elementary step - Unimolecular - Bimolecular
order of reaction
the power to which a reactant is raised in a rate law. *** Order of reaction cannot be determined from the chemical equation. It is determined experimentally.
reaction rate
the rate at which reactants are converted into products Reaction rates depend on activation energy - High activation energy barriers result in low rates. - Low activation energy barriers result in high rates. Reaction rate is not determined by ∆Grxn
Reaction mechanism:
the sequence of molecular events, or reaction steps, that define the pathway from reactants to products. 2NO(g) + Cl₂(g) → 2NOCl(g) • To happen in one step, 3 molecules would have to collide with enough energy and in the correct orientation • More likely, it happens in more than one step: Step 1: NO(g) + Cl₂(g) → (g) Step 2: NOCl₂(g) + NO(g) → 2NOCl(g)
Activated complex
the short-lived high-energy intermediate between reactants and products
changes in pressure
when pressure INCREASES, the equilibrium will shift in the direction that REDUCES the total number of gas molecules. when pressure DECREASES, eq will shift in the direction that INCREASES the total # of gas molecules ** (pretty sure that pressure is only relevant to solids and liquids)
Enzymes
• An enzyme is any of numerous complex proteins that are produced by living cells and CATALYZE specific biochemical reactions at body temperatures. • Most enzymes are specific to particular reactions or substances The functioning of your body and of all other living things depends on chemical reactions. • Most biochemical reactions proceed slowly at body temperature.
These always appear in equilibrium expression:
• Gases • Aqeuous species
Maxwell-Boltzmann distribution
• Graph shows fraction of particles (vertical axis) and the particle's speed ( horizontal axis). - A graph with a lower y-value has a higher temp. - A graph with a higher y-value has a lower temp
Examples of Heterogeneous catalysts
• Platinum in catalytic converter (solid catalyst, gaseous reactants) • Iron in production of ammonia (solid catalyst, gaseous reactants)
These do not appear in equilibrium expression:
• Pure solids • Pure liquids
Calculating ΔHrxn from a Reaction Pathway Graph
∆H rxn = ∆Hf,products - ∆Hf,reactants = -877.14 - (-74.8)
Activation energy (Ea)
− the MINIMUM amount of energy needed to initiate a chemical reaction -