Chem Unit 6: Reaction Rates

Reaction rates consider how

fast or slow a reaction takes place

Rate of Reaction

the rate of disappearance of the reactants OR the rate of appearance of the products depending on what is easiest to observe

• units: M/s

Reaction Rate Causes

1. number of collisions

2. orientation of collisions

3. sufficient energy of collision

Effective collisions

when reaction rate conditions are met

For an “effective collision” to occur, the reactant molecules must be oriented in space correctly to facilitate the

breaking and forming of bonds and the rearrangement of atoms that result in the formation of product molecules

Potential Energy diagrams represent

the energy pathway of a reaction

The activated complex is formed at

the top of the curve (or energy hill)

• at this point in the reaction process bonds are breaking and new bonds are forming

The activation energy

the energy needed to form the activated complex

• if the reaction does not have enough the reactants will not react

• if sufficient energy is available, then products are formed

*shown as the numerical distance of potential energy from the highest peak to the reactant*

The heat of the reaction (ΔH)

shown on the graph as the difference of potential energy between the products and the reactants (can be negative or positive depending exo/endo)

• final-initial

Reaction Rate Factors

1. Nature of the reactants

2. Temperature

3. Surface Area/ Stirring

4. Presence of catalysts

5. Concentration of reactants

Nature of Reactants

rates depend on reactivity of the reactant

• ionic reaction: happen almost instantly

• covalent reaction: slower because of atoms and electrons needing to rearrange

Temperature

increased temperature increases the energy of the collisions resulting in more effective collisions

• it takes energy to break bonds (endothermic)

• energy is released when new bonds form (exothermic)

Surface Area/Stirring

increased surface area (smaller pieces)causes increased collisions

• stirring has the same effect by removing reacted particles quickly

Presence of Catalysts

A catalyst is a substance that increases the rate of the chemical reaction without being permanently changed or used in the reaction

• a catalyst works by providing an alternate pathway that has a lower activation energy than the original reaction

• in our bodies catalysts are called enzymes

• alternatively, an inhibitor is the opposite and forms a new higher activation energy or pathway

Concentration of Reactants

increased concentrations of reactants will cause more collisions

• for gases increasing the pressure will cause an increase in concentration because more particles per volume increases the collisions

The Rate Law

a mathematical expression that related the concentration of the reactants to the reaction rate

• found experimentally and is based on the slowest step in a reaction mechanism

Rate Law Equation

A + B → C

Rate= k[A]^x [B]^y

• where [] indicates concentration

• x and y are the order of the reactants

• k is the rate law constant

• products do not appear in the rate law

Zero Order

the concentration of the reactant doesnt affect the rate of production of product

• the concentration of reactant is doubled, the production of the product doesnt change

First Order

the concentration of the reactant does affect the rate of production of product

• the concentration of reactant is doubled, the production of the product double

Second Order

the concentration of the reactant does affect the rate of production of product

• the concentration of reactant is doubled, the production of the product quadruples

Third Order

the concentration of the reactant does affect the rate of production of product

• the concentration of the reactant is doubled, the production of the product increases by a factor of 8

Overall Order

the sum of all orders

an elementary reaction is

a reaction in which reactants are converted to products in a single step (a single step reaction)

• mA +nB → C (m and n are moles in the balanced equation)

• Rate=k[A]^m [B]^n

• for single step reactions only, the balanced equation can be used to determine the rate law

• peaks

Reaction Mechanism for Multi-Step Reactions

• chemical equations describe reactions, but do not show the reaction mechanism

• most reactions are not single step but require a series of steps, we call this the reaction mechanism, which can be used to determine the rate law if the slow or rate determining step is known

• the reaction mechanism must agree with the experimentally determined rate law if the rate law is known

In a multistep chemical reaction, the steps do not all progress

at the same rate

a reaction can never proceed faster than

its slowest step

Intermediates

substances that are produced by one reaction and later consumed by another reaction

• starts as a product, ends as a reactant

• not included in the rate law or the chemical equation

• valleys

catalyst

substances that starts a reaction and later produced by another reaction

• starts as reactant, ends as a product

• not included in the rate law (for our class)

• must be completely canceled out

  • good: NO2 → NO2

  • bad: NO2 + NO2 → NO2

on a multi-step reaction energy diagram, each peak represents

an elementary reaction or a step in the reaction

on a multi-step reaction energy diagram, each valley is

an intermediate

• always one less than the elementary

on a multi-step reaction energy diagram, the highest peak is

the activated complex

When looking at a diagram, the overall activation energy is the

the distance between the starting value and the highest peak

When looking at a diagram, the rate-determining step for a reaction is the

highest peak

On a diagram, exothermic looks like

lower at products

On a diagram, endothermic looks like

higher at products

Gibbs Free Energy (G)

the energy associated with a chemical reaction that is able to do work

By spontaneous it means

that the reaction will go on its own, if given the necessary energy to get over the energy of activation “barrier”

-ΔG

reaction is spontaneous

+ΔG

reaction is not spontaneous

Entropy is the measure of

disorder in a system

The Second Law of Thermodynamics states that

the entropy of the universe must always increase

Things that increase the entropy of the universe

• a system going from ordered → disordered

• heat being released into the surroundings

Entropy Examples +ΔS

• your room getting messy

• a wall crumbling

• going from s→l→g

• breaking into pieces (ex: decomposition reaction, or a solid dissolving)

• any reaction that has more particles as products than as reactants

• increasing temperature

*-ΔS is the opposite of all the above*

Gibbs Free Energy Equation

ΔG = ΔH -TΔS

• ΔG= Gibbs Free Energy

• T= Temperature (in K)

• ΔH= Enthalpy (+ if endo, - is exo)

• ΔS= Entropy (+ if increasing disorder, - if decreasing disorder)

-ΔH / +TΔS / -ΔG

yes, always spontaneous

+ΔH / +TΔS / ?ΔG (not on test)

yes, spontaneous, at high temps

no, not spontaneous, at low temps

-ΔH / -TΔS / ?ΔG (not on test)

no, not spontaneous, at high temps

yes, spontaneous, at low temps

+ΔH / -TΔS / +ΔG

nope, never spontaneous