Chapter 8 Chemical and Physical Change: Energy, Rate, and Equilibrium Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Jan 18, 2016
Chapter 8
Chemical and Physical Change: Energy, Rate, and Equilibrium
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
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.