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What do you think could increase the rate or speed of a chemical reaction? Define activation energy. Give a real-life example. Define a catalyst. How does it work?. IQ #1. Chapter 18: “Reaction Rates and Equilibrium”. Section 18.1 Rates of Reaction. OBJECTIVES

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What do you think could increase the rate or speed of a chemical reaction?

Define activation energy. Give a real-life example.

Define a catalyst. How does it work?

IQ #1


Chapter 18 reaction rates and equilibrium

Chapter 18: chemical reaction?“Reaction Rates and Equilibrium”


Section 18 1 rates of reaction
Section 18.1 Rates of Reaction chemical reaction?

  • OBJECTIVES

    • Describe how to express the rate of a chemical reaction.

    • Identify four factors that influence the rate of a chemical reaction.


Collision theory
Collision Theory chemical reaction?

  • Reactions can occur:

    • Very fast – such as a firecracker

    • Very slow – such as the time it took for dead plants to make coal

  • A “rate” is a measure of the speed of any change that occurs within an interval of time

  • In chemistry, reaction rate is expressed as the amount of reactant changing per unit time


Collision model
Collision Model chemical reaction?

Key Idea: Molecules must collide to react.

However, only a small fraction of collisions produces a reaction. Why?


Collision model1
Collision Model chemical reaction?

  • Collisions must have enough energy to produce the reaction (must equal or exceed the activation energy– the minimum energy needed to react).

  • Think of clay clumps thrown together gently – they don’t stick, but if thrown together, they stick together!



Collision model2
Collision Model chemical reaction?

  • An “activated complex” is an unstable arrangement of atoms that forms momentarily (typically about 10-13 seconds) at the peak of the activation-energy barrier.

    • This is sometimes called the transition state

  • Results in either:

    a) forming products

    b) reformation of reactants

    • Both outcomes are equally likely


- Page 543 chemical reaction?


Collision model3
Collision Model chemical reaction?

  • The collision theory explains why some naturally occurring reactions are very slow

    • Carbon and oxygen react when charcoal burns, but this has a very high activation energy

    • At room temperature, the collisions between carbon and oxygen are not enough to cause a reaction


Factors affecting rate
Factors Affecting Rate chemical reaction?

1) Temperature

-Increasing temperature always increases the rate of a reaction.

2) Surface Area

-Increasing surface area increases the # of collisions/rate of a reaction

3) Concentration

-Increasing concentration USUALLY increases the rate of a reaction

4) Presence of Catalysts


Catalysts
Catalysts chemical reaction?

  • Catalyst: A substance that speeds up a reaction, without being consumed itself in the reaction

  • Enzyme: A large molecule (usually a protein) that catalyzes biological reactions.

    • Human body temperature = 37 oC, much too low for digestion reactions without catalysts.

  • Inhibitors – interfere with the action of a catalyst; reactions slow or even stop


Endothermic reaction with a catalyst
Endothermic Reaction with chemical reaction?a Catalyst


Catalysts increase the number of effective collisions
Catalysts Increase the Number of chemical reaction?Effective Collisions


Section 18 2 reversible reactions and equilibrium
Section 18.2 Reversible Reactions and Equilibrium chemical reaction?

  • OBJECTIVES

    • Describe how the amounts of reactants and products change in a chemical system at equilibrium.

    • Identify three stresses that can change the equilibrium position of a chemical system.

    • Explain what the value of Keq indicates about the position of equilibrium.


Reversible reactions
Reversible Reactions chemical reaction?

  • Some reactions do not go to completion as we have assumed

    • They may be reversible– a reaction in which the conversion of reactants to products and the conversion of products to reactants occur simultaneously

  • Forward: 2SO2(g) + O2(g)→ 2SO3(g)

  • Reverse: 2SO2(g) + O2(g)← 2SO3(g)


Reversible reactions1
Reversible Reactions chemical reaction?

  • The two equations can be combined into one, by using a double arrow, which tells us that it is a reversible reaction:

    2SO2(g) + O2(g)↔ 2SO3(g)

  • A chemical equilibrium occurs, and no net changeoccurs in the actual amounts of the components of the system.


List the four factors that affect reaction rates. How does each do so?

What is an activated complex?

What does it mean when we say that a reaction is “reversible”?

IQ #2


Reversible reactions2
Reversible Reactions each do so?

  • Even though the rates of the forward and reverse are equal, the concentrations of components on both sides may not be equal

    • An equlibrium position may be shown:

      A B or A B

      1% 99% 99% 1%

  • It depends on which side is favored; almost all reactions are reversible to some extent


Le chatelier s principle
Le Chatelier’s Principle each do so?

  • The French chemist Henri Le Chatelier (1850-1936) studied how the equilibrium position shiftsas a result of changing conditions

  • Le Chatelier’s principle:If stress is applied to a system in equilibrium, the system changes in a way that relieves the stress


Le chatelier s principle1
Le Chatelier’s Principle each do so?

  • What items did he consider to be stresson the equilibrium?

    • Concentration

    • Temperature

    • Pressure

  • Concentration– adding more reactant produces more product, and removing the product as it forms will produce more product

Each of these will now be discussed in detail


Le chatelier s principle2
Le Chatelier’s Principle each do so?

  • Temperature– increasing the temperature causes the equilibrium position to shift in the direction that absorbs heat

    • If heat is one of the products (just like a chemical), it is part of the equilibrium

    • so cooling an exothermic reaction will produce more product, and heating it would shift the reaction to the reactant side of the equilibrium


Le chatelier s principle3
Le Chatelier’s Principle each do so?

  • Pressure– changes in pressure will only effect gaseous equilibria

    • Increasing the pressure will usually favor the direction that has fewer molecules

      N2(g) + 3H2(g)→ 2NH3(g)

  • For every two molecules of ammonia made, four molecules of reactant are used up – the equilibrium shifts to the right with an increase in pressure


Lechatelier example 1
LeChatelier Example #1 each do so?

A closed container of ice and water is at equilibrium. Then, the temperature is raised.

Ice + Energy  Water

The system temporarily shifts to the _______ to restore equilibrium.

right


Lechatelier example 2
LeChatelier Example #2 each do so?

A closed container of N2O4 and NO2 is at equilibrium. NO2 is added to the container.

N2O4 (g) + Energy  2 NO2(g)

The system temporarily shifts to the _______ to restore equilibrium.

left


Lechatelier example 3
LeChatelier Example #3 each do so?

A closed container of water and its vapor is at equilibrium. Vapor is removed from the system.

water + Energy  vapor

The system temporarily shifts to the _______ to restore equilibrium.

right


Lechatelier example 4
LeChatelier Example #4 each do so?

A closed container of N2O4 and NO2 is at equilibrium. The pressure is increased.

N2O4 (g) + Energy  2 NO2(g)

The system temporarily shifts to the _______ to restore equilibrium, because there are fewer moles of gas on that side of the equation.

left


Applications of Le Chatelier’s Principle each do so?

1. Glasses/Sunglasses

Reaction:

Stress: Add Sunlight

Stress: Remove Sunlight

AgCl + energy Ag + Cl

clear

dark

Shifts right, glasses become dark

Shifts left, glasses become clear


Biochemical Equilibrium System (Respiration) each do so?

Reactions:

STRESS: Exercise vigorously (____________________)

STRESS: Breathe faster (_____________________)

CO2 + H2O H2CO3

H2CO3 H+ + HCO3-

More CO2 produced

Both reactions shift right

[H+] increases (acid), so the brain sends a message to exhale more rapidly.

Exhale CO2

Both reactions shift left

pH returns to normal level


Equilibrium constants
Equilibrium Constants each do so?

  • Chemists generally express the position of equilibrium in terms of numerical values

    • These values relate to the amountsof reactants and products at equilibrium

    • This is called the equilibrium constant, and abbreviated Keq


Equilibrium constants1
Equilibrium Constants each do so?

  • consider this reaction:

    aA + bB cC + dD

    • The equilibrium constant (Keq) is the ratio of product concentration to the reactant concentration at equilibrium, with each concentration raised to a power (= the balancing coefficient)


Equilibrium constants2
Equilibrium Constants each do so?

  • consider this reaction:

    aA + bB cC + dD

    • Thus, the “equilibrium constantexpression” has the general form:

      [C]c x [D]d

      [A]a x [B]b

      (brackets: [ ] = molarity concentration)

Keq =


Equilibrium constants3
Equilibrium Constants each do so?

  • the equilibrium constants provide valuable information, such as whether products or reactants are favored:

    Keq > 1, products favored at equilibrium

    Keq < 1, reactants favored at equilibrium


- Page 557 each do so?


Section 18 4 entropy and free energy
Section 18.4 each do so?Entropy and Free Energy

  • OBJECTIVES

    • Identify two characteristics of spontaneous reactions.

    • Describe the role of entropy in chemical reactions.

    • Identify two factors that determine the spontaneity of a reaction.

    • Define Gibbs free-energy change.


Free energy and spontaneous reactions
Free Energy and Spontaneous Reactions each do so?

  • Many chemical and physical processes release energy, and that can be used to bring about other changes

    • The energy in a chemical reaction can be harnessed to do work, such as moving the pistons in your car’s engine

  • Free energyis energy that is available to do work

    • That does not mean it can be used efficiently


Free energy and spontaneous reactions1
Free Energy and Spontaneous Reactions each do so?

  • Your car’s engine is only about 30 % efficient, and this is used to propel it

    • The remaining 70 % is lost as friction and waste heat

  • No process can be made 100 % efficient

    • Even living things, which are among the most efficient users of free energy, are seldom more than 70 % efficient


Free energy and spontaneous reactions2
Free Energy and Spontaneous Reactions each do so?

  • We can only get energy from a reaction that actually occurs, not just theoretically:

    • CO2(g)→ C(s) + O2(g)

  • this is a balanced equation, and is the reverseof combustion

  • Experience tells us this does not tend to occur, but instead happens in the reverse direction


Free energy and spontaneous reactions3
Free Energy and Spontaneous Reactions each do so?

  • The world of balanced chemical equations is divided into two groups:

    • Equations representing reactions that actually occur

    • Equations representing reactions that do not tend to occur, or at least not efficiently


Free energy and spontaneous reactions4
Free Energy and Spontaneous Reactions each do so?

  • The first, (those that actually occur) and more important group involves processes that are spontaneous

    • A spontaneous reaction occurs naturally, and favors the formation of products at the specified conditions

    • They produce substantial amounts of product at equilibrium, and release free energy

    • Example: a fireworks display – page 567


Free energy and spontaneous reactions5
Free Energy and Spontaneous Reactions each do so?

  • In contrast, a nonspontaneous reaction is a reaction that does not favor the formation of products at the specified conditions

    • These do not give substantial amounts of product at equilibrium

  • Think of soda pop bubbling the CO2 out: this is spontaneous, whereas the CO2 going back into solution happens very little, and is non-spontaneous


Spontaneous reactions
Spontaneous Reactions each do so?

  • Do not confuse the words spontaneous and instantaneous. Spontaneous just simply means that it will work by itself, but does not say anything about how fast the reaction will take place – it may take 20 years to react, but it will eventually react.

    • Some spontaneous reactions are very slow: sugar + oxygen → carbon dioxide and water, but a bowl of sugar appears to be doing nothing (it is reacting, but would take thousands of years)

    • At room temperature, it is very slow; apply heat and the reaction is fast; thus changing the conditions (temp. or pressure) may determine whether or not it is spontaneous


Entropy
Entropy each do so?

  • Entropy is a measure of disorder, and is measured in units of J/mol.K; there are no negative values of entropy

  • The law of disorder states the natural tendency is for systems to move to the direction of maximum disorder, not vice-versa

    • Your room NEVER cleans itself (disorder to order?)

  • An increase in entropy favors the spontaneous chemical reaction

  • A decrease in entropy favors the nonspontaneous reaction


- Page 570 each do so?


Enthalpy and entropy
Enthalpy and Entropy each do so?

  • Reactions tend to proceed in the direction that lowers the energy of the system (H, enthalpy).

and,

  • Reactions tend to proceed in the direction that increases the disorder of the system (S, entropy).


Enthalpy and entropy1
Enthalpy and Entropy each do so?

  • These are the two “drivers” to every equation.

    • If they both AGREE the reaction should be spontaneous, IT WILL be spontaneous at all temperatures, and you will not be able to stop the reaction without separating the reactants

    • If they both AGREE that the reaction should NOT be spontaneous, it will NOT work at ANY temperature, no matter how much you heat it, add pressure, or anything else!


Enthalpy and entropy2
Enthalpy and Entropy each do so?

  • The size and direction of enthalpy and entropy changes together determine whether a reaction is spontaneous

  • If the two drivers disagree on whether or not it should be spontaneous, a third party (Gibb’s free energy) is called in to act as the “judge” about what temperatures it will be spontaneous, and what the temp. is.

    • But, it WILL work and be spontaneous at some temperature!


Spontaneity of reactions
Spontaneity of Reactions each do so?

Reactions proceed spontaneously in the direction that lowers their Gibb’s free energy, G.

G = H - TS(T is kelvin temp.)

If G is negative, the reaction is spontaneous. (system loses free energy)

If G is positive, the reaction is NOT spontaneous. (requires work be expended)


Spontaneity of reactions1
Spontaneity of Reactions each do so?

  • Therefore, if the enthalpy and entropy do not agree with each other as to what should happen:

    • Gibbs free-energy says that they are both correct, the reaction will occur

    • But the Gibbs free-energy will decide the conditions of temperature that it will happen

  • Figure 18.25, page 572


- Page 572 each do so?


Entropy1

6. Calculation of each do so?Srxn

**Use the Tables from Thermochemistry**

7. Example calculation: Calculate the Sorxn of the following reaction:

2 H2O(l)→ 2 H2(g) + O2(g)

Entropy

Sorxn = So Products - So Reactants

Predict: Favorable or unfavorable

Sorxn = So Products - So Reactants

Sorxn = [2(H2) + (O2)] - [2(H2O)]

Sorxn = [2 mol(130.58 J/mol·K) + (1 mol(205.0 J/mol·K)] – [2 mol(69.91 J/mol·K)]

Sorxn = + 326.3 J/K


C. each do so?Gibb’s Free Energy (___): Determines the ___________ of a reaction by combining the two driving forces: ________ and _______.

1. Units of Gibb’s Free Energy: ______.

2. ΔG = ___ → Reaction is ________________.

3. ΔG = ___ → Reaction is ________________.

4. Calculations of Δ Grxn

Combines BOTH driving forces.

spontaneity

G

enthalpy

entropy

kJ/mol

+

non-spontaneous

-

spontaneous

G = H - TS

T = Temperature of Kelvin

Standard Temp = 25 oC = 298 K


Example calculation: Given that the ∆H each do so?o is +571.6 KJ, calculate the ∆Go of the reaction at standard temperature.

2 H2O(l)→ 2 H2(g) + O2(g)

G = H - TS

G = + 571.6 kJ – (298 K)(0.3263 kJ/K)

G = + 474.4 kJ

 The rxn is non-spontaneous


Always each do so?

Spontaneous

+

or

Spontaneous

or Non-spon.

+

or

Spontaneous

or Non-spon.

Always

Non-spontaneous

H

S

G

Condition

+

Exothermic

Favorable

Spontaneous

(favorable)

H > TS

Exothermic

Unfavorable

(favorable)

+

+

TS > H

Endothermic

Favorable

(unfavorable)

+

+

Endothermic

Non-spontaneous

Unfavorable

(unfavorable)


6.(a) Calculate each do so?Δ Ho (exo or endo?), Δ So (favorable or unfavorable?), and Go (Spontaneous or non-spontaneous?) for the following reaction: Al2O3(s) + 2 Fe(s)→ 2 Al(s) + Fe2O3(s).

(b) Calculate the minimum temperature at which this reaction is spontaneous.

Horxn = [2(Al) + (Fe2O3)] – [(Al2O3) + 2(Fe)]

Horxn = [2 (0) + (-822.16)] – [(-1669.8) + 2 (0)]

Horxn = + 847.6 kJ

Endothermic

 Unfavorable

Sorxn = [2(28.32) + (89.96)] - [(51.00) + 2(27.15)]

Sorxn = + 41.30 J/K

 Favorable

Go = H - TS

Go = + 847.6 kJ – (298 K)(0.0413 kJ/K)

Go = + 835.3 kJ

 Non-spontaneous

= 20,200 K


7. For the following reaction: NH each do so?4NO3(s) → N2O(g) + 2 H2O(g)

∆H°f (kJ/mol)S° (J/mol K)

NH4NO3(s) -365.6 +151.1

N2O(g) +82.1 +219.7

H2O(g) -241.8 +188.7

Calculate ∆H° (exo or endo?), ∆S° (favorable or unfavorable?), and ∆G° (spontaneous or non-spontaneous?) @ standard temp.

(b) Calculate the minimum temperature at which this reaction is spontaneous.

Horxn = [(82.1) + 2(-241.8)] – [(-365.6)]

= - 35.9 kJ

Exothermic

 Favorable

Sorxn = [(219.7) + 2(188.7)] - [(151.1)]

 Favorable

= + 446.0 J/K

Go = - 35.9 kJ – (298 K)(0.446 kJ/K)

 Spontaneous

= - 169 kJ

Always


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