Chapter 8

Chemical and Physical Change: Energy, Rate, and Equilibrium

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8.1 Thermodynamics

• Thermodynamics - the study of energy, work, and heat.– applied to chemical change

– applied to physical change

• The laws of thermodynamics help us to understand why some chemical reactions occur and others do not.

• As the bonds are broken and new bonds are formed, energy is required or released.

• We can measure the change in energy during these changes.

• System - the process under study

– Usually the chemical reaction or physical change of interest.

• Surroundings - the rest of the universe.

• We will be able to measure the change in energy in the form of heat as the temperature changes.

The Chemical Reaction and Energy

• Important points to kinetic molecular theory

– molecules and atoms in a reaction mixture are in constant, random motion;

– these molecules and atoms frequently collide with each other;

– only some collisions, those with sufficient energy, will break bonds in molecules; and

– when reactant bonds are broken, new bonds may be formed and products result.

Exothermic and Endothermic Reactions

1• The first law of thermodynamics – the energy of the universe is constant, E cannot be created nor destroyed.

• Where does the energy come from that is released and where does the energy go when it is absorbed?

• The chemical bond is stored chemical energy.

• If the energy required to form new bonds > the energy released when the old bonds are broken, there will need to be an external supply of energy…Endothermic reaction.

A-B + C-D A-D + C-B

These bonds must be broken.

This releases energy.

These bonds are formed.

This requires energy

• Enthalpy - represents heat energy.

• Change in Enthalpy (Ho) - energy difference between the products and reactants.

• Energy released (exothermic), enthalpy change is negative (energy diagram).

• Energy absorbed (endothoermic), enthalpy change is positive (energy diagram).

Entropy

• The second law of thermodynamics - the universe spontaneously tends toward increasing disorder or randomness.

• Entropy (So) - a measure of the randomness or disorder of a chemical system.

• High entropy - highly disordered system

• Low entropy - well organized system

• No such thing as negative entropy.

So of a reaction = So(products) - So(reactants)

• A positive So means an increase in disorder for the reaction.

• A negative So means a decrease in disorder for the reaction.

Spontaneous and Nonspontaneous Reactions

• Spontaneous reaction - occurs without any artificial external energy input.

• Often, but not always, exothermic reactions are spontaneous.

• Thermodynamics is used to help predict if a reaction will occur.

• If exothermic and positive ΔSo = Spontaneous

• If endothermic and negative ΔSo = Nonspontaneous

• For other situations, it depends on the relative size of ΔHo and ΔSo.

• Free energy (Go) - represents the combined contribution of the enthalpy and entropy values for a chemical reaction.

• Predicts spontaneity

• Negative Go…Spontaneous

• Positive Go…Nonspontaneous

Go = Ho - TSo

T in Kelvins

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8.2 Experimental Determination of Energy Change in Reactions

• Calorimetry - the measurement of heat energy changes in a chemical reaction.

• The change in temperature is used to measure the heat loss or gain.

• Calorie – the quantity of heat required to raise 1 g of water 1 °C.

• Nutritional Calorie (large Calorie) = 1kilocalorie (1kcal) or 1000 calorie.

• Specific heat (S.H.) - the number of calories of heat needed to raise the temperature of 1 g of the substance 1 oC.

• S.H. for water is 1.0 cal/goC

• To determine heat released or absorbed, need:– specific heat– total number of grams of solution– temperature change (increase or decrease)

S.H.T mQ

8.3 Kinetics

• Thermodynamics determines if a reaction will occur but tells us nothing about the time it will take.

• Kinetics - the study of the rate of chemical reactions.– Also gives the mechanism - step-by-step description of how

reactants become products.

• We will look at:– disappearance of reactants and

– appearance of products.

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Let’s consider the following reaction:

CH4(g) + 2O2(g) CO2(g) +2H2O(g) + 211 kcal

• C-H and O=O bonds must be broken and C=O and O-H bonds must be formed

• Energy is required to break the bonds.

– Comes from the collision of the molecules.

– Effective collision - one that leads to a chemical reaction.

• Activation energy - the minimum amount of energy required to produce a chemical reaction.

• Activated complex - extremely unstable complex, the formation of which requires energy.

Factors That Affect Reaction Rate

• Structure of the reacting molecules,

• Concentration of reactants,

• Temperature of reactants,physical state of reactants, and

• Presence of a catalyst

Structure of Reacting Molecules

• Oppositely charged species react more rapidly

• Ions with the same charge do not react.

• Bond strength plays a role.

• Magnitude of the activation energy is related to bond strength

• Size and shape influence the rate.

• Large molecules may block the reactive part of the molecule.

The Concentration of Reactants

• Rate will generally increase as concentration increases.

• Caused by a greater number of collisions

The Temperature of Reactants

• Rate increases as the temperature increases.

• Higher temp. means higher K.E.

• Higher K.E. means higher percentage of these collisions will result in product formation.

Physical State of Reactants

• Solid state:

• atoms, ions or molecules are close but restrictive in motion.

• Gaseous state:

• particles are free to move but are far apart causing collisions to be relatively infrequent.

• Liquid state:

• particles are free to move and are close together.

• The typical order of rate per state of reactant?

• Liquid > gas> solid

Presence of a Catalyst

• Catalyst - a substance that increases the reaction rate.

• Catalysts interact with the reactants to create an alternative pathway for product formation by lowering the activation energy.

• Enzyme - a biological catalyst that controls and speeds up thousands of essential biochemical reactions.

8.4 Equilibrium

Rate and Reversibility of Reactions

• Equilibrium reactions - chemical reactions that do not go to completion (incomplete reactions).

• After no further obvious change, measurable quantities of reactants and products remain.

• Reversible reaction - a process that can occur in both directions.– Use the double arrow symbol

• Dynamic equilibrium - the rate of the forward process is exactly balanced by the rate of the reverse process.

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Equilibrium• Example: Sugar in Water

– If put 2-3 g of sugar in 100 mL water, all will dissolve.

– sugar (s) sugar (aq)

• If dissolving 100 g in 100 mL of water, not all of it will dissolve.

– Over time, you observe no further change in the amount of dissolved sugar.

– Individual sugar molecules are constantly going into and out of solution and happens at the same rate.

• The double arrow serves as– an indicator of a reversible process– an indicator of an equilibrium process, and– a reminder of the dynamic nature of the process.

• Equilibrium constant (Keq)- ratio of the two rate constants.

sugar(s) sugar(aq)

The Generalized Equilibrium-Constant Expression for a Chemical Reaction

ba

dc

B][[A]

D][C][Keq

9aA + bB cC + dD

Writing Equilibrium-Constant Expressions

• Each chemical reaction has a unique equilibrium constant value at a specified temperature.

• The brackets represent molar concentration.

• All equilibrium constants are shown as unitless.

• Only the concentration of gases and substances in solution are shown.

• concentration for pure liquids and solids are not shown.

Interpreting Equilibrium Constants

• The value of Equil. constant tells us the extent to which reactants have converted to products.

1. Keq greater than 1 x 102.

• Large value of Keq: numerator > denominator.

• At equilibrium mostly product present.

2. Keq less than 1 x 10-2.

• Small value of Keq: numerator < denominator.

• At equilibrium mostly reactant present.

3. Keq between 1 x 10-2 and 1 x 102.

• Equilibrium mixture contains significant concentration of both reactants and products.

• LeChateleir’s Principle - if a stress is placed on a system at equilibrium, the system will respond by altering the equilibrium composition in such a way as to minimize the stress.

We will examine the following “stresses.”1. Effect of Concentration2. Effect of Heat3. Effect of Pressure4. Effect of Catalyst

1. Effect of Concentration

• Adding or removing reactants and products at a fixed volume.

• Addition of N2 or H2. To minimize the stress, which way will the reaction shift?

• To the right. Forming more products.

• If NH3 is put in the reaction vessel?

• Equilibrium shifts to the left, forming more reactants.

N2(g) + 3H2(g) 2NH3(g)

• Addition of heat is similar to increasing the amount of product. If heat is generated, which way will the equilibrium shift? To the left.

2. Effect of Heat

• Exothermic reactions: treat heat as a product

N2(g) + 3H2(g) 2NH3(g) + 22 kcal

• Endothermic Reaction - treat heat as a reactant.

39 kcal + 2N2(g) + O2(g) 2NH3(g)

• Which way will this reaction shift if the reaction is heated? To the right.

3. Effect of Pressure

• Pressure affects the equilibrium only if one or more substances in the reaction are gases

• Relative number of gas moles on reactant and product side must differ.

• When pressure goes up…shift to side with less moles of gas. When pressure goes downs…shifts to side with more moles of gas.

4. Effect of a Catalyst

• A catalyst has no effect on the equilibrium composition. It increases the rate of both the forward and reverse reaction to the same extent.