8. Reaction Rates and Chemical Equilibrium Flashcards
Chemical kinetics
Study of reaction rates
Rate of reaction
-Change in concentration of reactant (or product) per unit time
-measured in laboratory
-not constant over a long period of time
Initial rate
-at the beginning, in most reactions, the change in concentration is directly proportional to timee
-later, as the reactant is used up, rate decreases
Effective collision
-a collision between two molecules or ions that results in a chemical reaction
*typically kase, nagba-bounce apart lng ung molecule A and B without reacting
Why are some collisions effective whereas others are not?
- Activation energy must be achieved
-energy required to break bonds of reactant molecules
-energy comes from collision
-energy depends on relative speed of colliding objects and on their angle of approach - A reaction may not take place if the molecules are not oriented properly when they collide
-HCl(g) and H2O(l)
-molecules must collide such that H of HCl hits O of water; otherwise no reaction
Energy diagrams
-schematic representation of energy changes that yake place as reactants are converted to products
Activation energy
-minimum amount of energy needed to break bonds
⬆️Ea -slower reaction rate
⬇️Ea -faster reaction rate
Bonds in reaction
-BOND FORMATION- releases energy
-BOND DISSOCIATION- requires energy input
Phase diagram
ENDOTHERMIC REACTION
-energy product > energy reactant
(+) ∆H°: more energy needed to break bonds than is released in forming bonds
-“uphill” reactions are endothermic
Phase diagram
EXOTHERMIC REACTION
-energy product < energy reactant
(-) ∆H°: more energy is released in forming bonds than needed to break bonds
-“downhill” reaction are exothermic
Energy of reaction
-difference between energy of reactant and product
Transition state
-top of the hill on an energy diagram
-one or more original bonds are partially broken and one or more new bonds is in the process of formation
Speed of reaction and collision
⬆️ speed of rxn ⬆️ probability of effective collision
Limiting rate (overall rate)
-slower reaction
How to change the rate of chemical reaction
1. Nature of reactant
-ions in aqueous solution are extremely rapid; activation energy is low because usually no covalent bond must be broken
-reactions between covalent molecules is slower
How to change the rate of chemical reaction
2. Concentration
-in most cases
⬆️ Concentration of reactant/ product ⬆️Reaction rate
Rate constant, k
Proportionality constant between molar concentration of reactants and rate of reaction
How to change the rate of chemical reaction
3. Temperature
⬆️ Temperature ⬆️Reaction rate
Collision theory: effects of temperature
- ⬆️ Temperature, molecules move more rapidly, collide more frequently, higher reaction rate
- ⬆️ Temperature, not only the average speed (kinetic energy) of the molecules is greater, but there is also distribution of speeds.
Number of very fast molecules increases much more than the number with average speed, more effective collision
How to change the rate of chemical reaction
4. Presence of catalyst
-any substance the increases rate of reaction without itself being used up
-allows reaction to take different pathway, one with LOWER activation energy
-catalyst may provide a surface where reactants can meet
Heterogeneous catalyst
A catalyst in a separate phase from reactants (i.e. platinum Pt(s) in reaction between formaldehyde CH2O (g) and H2(g)
Homogeneous catalyst
-catalyst in same phas as reactants
Reversible reaction
-reaction can be made to go in either direction
Dynamic equilibrium
-a state in which the rate of the forward reaction equals the rate of the reverse reaction
-no change in concentration
Equilibrium constant
-value calculated from the equilibrium expression fo a given reaction indicating in which direction reaction goes
-at constant temperature, equilibrium constants remain the same no matter what concentrations we have
-value of Keq changes when temperature changes
-large K proceeds to completion (to the right)
-small K hardly goes forward at all
Keq = ([C]^c[D]^d) / ([A]^a[B]^b)
*Concentrations at equilibrium
-in dilute solution, omit concentration of solvent because the molarity of solvent is effectively constant
-rate of reaction (how long it takes to reach equilibrium) and Keq have no relationship
Le Chatelier’s Principle
-when a stress is applied to a system in chemical equilibrium, the position of the equilibrium shifts in the direction that will relieve the applied stress
-Henri Le Chatelier (1888)
Types of stress that can be put on chemical equilibria
1. Addition of a reaction component
-if reactant is added:
equilibrium shifts to formation of more products (para maconsume yung dinagdag na reactant)
-if product is added:
equilibrium shifts to formation of reactants
Types of stress that can be put on chemical equilibria
2. Removal of a reaction component (i.e. thru distillation)
-if reactant is removed (or decreased in concentration), equilibrium shifts to reactant (para mapalitan ung nabawas)
-if product is removed, equilibrium shifts to product side
Types of stress that can be put on chemical equilibria
3. Change in temperature
•For exothermic reaction
-increase temperature (heat added): equilibrium shifts to reactant (endothermic yung reverse, so maaabsorb ung added heat)
-decrease temperature (heat removed): equilibrium shifts to product (exothermic forward) para magrelease ng heat
•For endothermic reaction
-increase temperature (heat added): equilibrium shifts to product (endothermic forward para maabsorb ung heat added
-decrease temperature (heat removed): equilibrium shifts to reactant (exothermic reverse) para magrelease ng heat
Types of stress that can be put on chemical equilibria
4. Change in oressure
-influences equilibrium only if one or more components of reaction mixture are gases
-increase of pressure shifts the reaction toward side with fewer moles of gas
-decrease of pressure shifts the reaction toward side with more moles of gas
Types of stress that can be put on chemical equilibria
5. Effects of catalyst
-for reversible reaction, catalyst always increase rates of BOTH FORWARD AND REVERSE reactions
Haber process
-artificial fixing of atmospheric nitrogen (instead of using bacteria)
-early problem: conflict between reaction rate and equilibrium
•exothermic so ideal low temp to produce more product
•kaso low temp, mabagal
N2 + 3H2 <–> 2NH3 + 22 kcal
-Fritz Haber (1908): discovered catalyst that permits the reaction to take place at a convenient rate at 500C
-NH3 converted to fertilizer
