Unit 1 Kinetics and Equilibrium Chemistry 3202

3:11 PM

Part 1: Reaction Kinetics (Chp. 12) Reaction Kinetics is the study of the rate of a chemical reaction Qualitative: Reactions may be described as being FAST or SLOW 

Fast – burning, explosions, precipitation  Slow – rusting, fermentation 

3:11 PM

Reaction Kinetics Quantitative: The rate of a reaction measures how fast products are formed or how fast reactants are consumed POSSIBLE UNITS ??

Rate = Change in quantity Change in time 3:11 PM

Measuring Reaction Rate The method used to determine reaction rate will depend on the reaction being studied. (p. 466) Methods: 1. monitor pH if there is an acid or base in the equation 2. record gas volume or changes in pressure if there is a gas in the reaction 

3:11 PM

Measuring Reaction Rate 

Methods: (cont’d) 3. record changes in mass if solids are present 4. monitor absorption of light if there is a color change 5. changes in electrical conductivity indicate changes in ion concentration 3:11 PM

MC: What could we use to measure the rate of this reaction? Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq)

a) pressure b) pH

c) d)

gas volume mass

Answer: d) because a solid is present 3:11 PM

MC: What could we use to measure the rate of this reaction? SO3(g) + H2O(l)  H2SO4(aq) a) pressure c) gas volume b) pH d) mass Answers: a) and c) because a gas is present b) because an acid is being produced 3:11 PM

p. 468; # 4

3:11 PM

What determines RATE?? All chemical reactions are bond breaking/bond forming events  The rate of a reaction depends on how quickly bonds are broken and how rapidly new bonds form.  KMT and Collision Theory are used to explain reaction rates. 

3:11 PM

Kinetic Molecular Theory (KMT) Matter is made of particles (atoms, ions, or molecules) in continuous motion  An increase in temperature:  increases the speed of particles  reduces the forces of attraction between particles 

3:11 PM

Kinetic Molecular Theory (KMT) 

KMT is supported by:

Diffusion – particles of a gas spread to fill their container (‘perfume in a room’) - solids dissolve uniformly in liquids over time. Pressure – a balloon remains inflated because gas particles are continuously hitting the sides of the balloon 3:11 PM

MC: Which observation best supports the Kinetic Molecular Theory? (A) Acetic acid odour is detected from across the room. (B) Liquid water freezes at 0°C under standard conditions. (C) Nitrogen dioxide gas is dark brown in colour. (D) When burned, butane produces more heat per mole than propane.

3:11 PM

Energy Distribution of Particles 25 °C

# of particles

200 °C

Kinetic Energy 3:11 PM

Collision Theory reactant particles must collide with one another for a chemical reaction to occur  particles must collide with proper orientation  collisions must have enough intensity to break old bonds and allow new bonds to form 

3:11 PM

Collision Theory 

to increase reaction rate you must increase the number of successful collisions between reactant molecules

VIDEO (VHS): Reaction Rates  LASERDISK: 3 VIDEO CLIPS 

3:11 PM

MC: Which must occur for a chemical reaction to take place? (A) addition of a catalyst (B) addition of energy (C) collisions between reacting particles (D) formation of a reaction intermediate

3:11 PM

Factors Affecting Reaction Rate Only works for (aq) or (g) reactants

Concentration – an increase in the concentration of a reactant usually increases the rate of a chemical reaction - the rate increases because there are: - more particles resulting in - more collisions between particles & - more successful collisions. 1.

3:11 PM

Factors Affectingmore Reaction Rate successful collisions Temperature - an increase in the temperature increases the rate of a chemical reaction - the higher temperature results in: faster rate - faster moving particles - more collisions between particles - more intense collisions NOTE: A temperature increase of 10 ºC usually causes reaction rate to double. 2.

3:11 PM

Factors Affecting Reaction Rate 3.

Nature of Reactants

– compounds with fewer bonds to break will react more rapidly than compounds with many bonds eg. propane (C3H8) burns faster than candlewax (C25H52) because it has fewer bonds

3:11 PM

Factors Affecting Reaction Rate Nature of Reactants - compounds with weak bonds react more rapidly than compounds with strong bonds – ions will react more rapidly than atoms and molecules 3.

3:11 PM

Factors Affecting Reaction Rate Surface Area - crushing a solid to produce a powder, or changing a substance to the gas phase, exposes more particles for collision if more particles are available for collision there will be: faster rate - more collisions - more successful collisions 4.

3:11 PM

3:11 PM

MC: Which factor explains why coal dust is explosive? (A) concentration (B) pressure (C) surface area (D) temperature

3:11 PM

PRACTICE: p. 484; #’s 1 & 2 p. 486; #’s 1,2, 4, 6, & 7 Kinetics & Equilibrium #1

3:11 PM

Factors Affecting Reaction Rate Catalysts - a catalyst increases the reaction rate by providing a different reaction pathway or mechanism with a lower activation energy - a catalyst IS NOT consumed by a chemical reaction. 5.

3:11 PM

Ea with catalyst

# of particles

Ea without catalyst

Kinetic Energy 3:11 PM

Potential Energy Diagrams (p. 473) PE diagrams show changes in potential energy (stored chemical energy) during chemical reactions  Exothermic reactions release more energy than they absorb (eg. burning)  Endothermic reactions absorb more energy than they release (eg. photosynthesis) 

3:11 PM

Potential Energy Diagrams ∆H represents the heat of reaction or enthalpy of reaction  ∆H is the difference between the PE of the reactants and the PE of the products  the minimum energy needed for a chemical reaction to occur is the activation energy 

3:11 PM

Potential Energy Diagrams 

the activated complex for a reaction is a temporary, unstable, intermediate species that quickly decomposes to products eg. H2 + I2 → H2I2 → 2 HI

activated complex

3:11 PM

ENDOTHERMIC site of the activated complex

Products

Eareverse

PE ∆H (positive)

Reactants

activation energy (Ea forward)

Reaction Progress 3:11 PM

EXOTHERMIC PE

site of AC Ea reverse

Reactants

∆H (negative)

Ea forward

Products

Reaction Progress 3:11 PM

ΔH & equations 

The energy term may be included in a chemical equation,

eg. CO(g) + 2 H2(g) → CH3OH(g) + 65 kJ Energy is PRODUCED

EXOTHERMIC

OR written as ΔH to the right of the equation. eg. CO(g) + 2 H2(g) → CH3OH(g) 3:11 PM

ΔH = - 65 kJ

ΔH & equations Another eg. eg. N2(g) + O2(g) + 90 kJ → 2 NO(g) Energy is REQUIRED

ENDOTHERMIC

OR N2(g) + O2(g) → 2 NO(g) ΔH = + 90 kJ

3:11 PM

Formula: (OPTIONAL) Eaforward - Eareverse = ΔH This formula is NOT necessary if you prefer using the PE diagram.

Animation 3:11 PM

Ea fwd

Ea rev

25

ΔH

Endothermic or Exothermic

-30

50

20

150

250 65

28

Sketch a PE diagram for each reaction

Photosynthesis

Earev C6H12O6 + O2

PE Eafwd CO2 + H2O

Reaction Progress 3:11 PM

∆H

Respiration Earev

Eafwd

PE

C6H12O6 + O2

∆H CO2 + H2O

Reaction Progress 3:11 PM

∆H

p. 474 3:11 PM

∆H

3:11 PM

∆H

p. 475 3:11 PM

Effect of a catalyst

PE catalyzed no catalyst Reaction Progress 3:11 PM

EXOTHERMIC

no catalyst

PE

Reaction Progress 3:11 PM

MC:

3:11 PM

MC: Carbon monoxide, CO(g), reacts with nitrogen dioxide, NO2(g) according to the reaction below. Which describes the reaction if Ea (forward) = 134 kJ? CO(g) + NO2(g) → CO2(g) + NO(g) + 226 kJ Ea (reverse) (A) 92 (B) 92 (C) 360 (D) 360

3:11 PM

Reaction type endothermic exothermic endothermic exothermic

 Sample

problem: p. 475  Questions: p. 476; #’s 1 – 4 p. 484; #’s 3 & 4

3:11 PM

Reaction Mechanisms (pp. 477 – 485) reaction mechanism – the steps that occur in a chemical reaction elementary reaction - each step in a reaction mechanism reaction intermediate – a molecule, atom or ion formed in one step and consumed in a later step NOTE: reaction intermediates are NOT included in the overall equation 3:11 PM

Reaction Mechanisms eg. #1 Step #1 Step #2

NO(g) + O2(g)  NO3(g) NO3(g) + NO(g)  2 NO2(g)

Overall Equation:

2 NO(g) + O2(g)  2 NO2(g)

3:11 PM

HBr + O2 → HOOBr fast HOOBr + HBr → 2 HOBr slow 2 HOBr + 2 HBr → 2 H2O + 2 Br2 fast

p. 478 #’s 5 – 8 3:11 PM

Reaction Mechanisms rate-determining step (RDS) - the RDS is the slowest step in a reaction mechanism - to increase the rate of a reaction you must speed up the RDS - increasing the concentration of a reactant will increase the rate ONLY IF the reactant is in the RDS 3:11 PM

Reaction Mechanisms PE diagrams - every step in a reaction mechanism has an activation energy which can be drawn on a PE diagram

3:11 PM

Reaction Mechanisms 3-step mechanism #2

PE

#1

RDS ?? #3



Reaction Progress 3:11 PM

Reaction Mechanisms eg: Step #1

H2CO2 + H+  H2CO2H+

fast

Step#2

H2CO2H+  HCO+ + H2O

slow

Step #3

HCO+  CO + H+

fast

3:11 PM

Reaction Mechanisms eg: Overall

H2CO2  H2O + CO

Omit H+

- catalyst

Omit H2CO2H+ & HCO+ - reaction intermediates

3:11 PM

Reaction Mechanisms PE

H2CO2 + H+

Reaction Progress 3:11 PM

CO + H+

Practice #5 p. 484 #’s 5 – 9 p. 485 #’s 10, 12 p. 486 #’s 8, 10, 11 p. 487 #’s 14, 17 p. 829 #’s 128,129, 131, 132

3:11 PM

p. 829 # 128 Step 1

H2(g) + NO(g) → H2O(g) + N(g)

Step 2 Step 3

H2(g) + O(g) → H2O(g)

2H2(g) + 2NO(g) → N2(g) + 2H2O(g)