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Energy, Rate, and Equilibrium
Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Chapter 7 Energy, Rate, and Equilibrium Denniston Topping Caret 6th Edition
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7.1 Thermodynamics Thermodynamics – the study of energy, work, and heat applied to chemical change Calculate the quantity of heat obtained from combustions of one gallon of fuel oil applied to physical change Determine the energy released by boiling water The laws of thermodynamics help us to understand why some chemical reactions occur and others do not
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The Chemical Reaction and Energy
Basic Concepts – from Kinetic Molecular Theory Molecules and atoms in a reaction mixture are in constant, random motion Molecules and atoms frequently collide with each other Only some collisions, those with sufficient energy, will break bonds in molecules When reactant bonds are broken, new bonds may be formed and products result
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Change in Energy and Surroundings
Absolute value for energy stored in a chemical system cannot be measured Can measure the change in energy during these chemical changes System: contains the process under study Surroundings: the rest of the universe
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Changes in the System Energy can be lost from the system to the surroundings Energy may be gained by the system at the expense of the surroundings This energy change is usually in the form of heat This change can be measured
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Law of Conservation of Energy
First Law of Thermodynamics – energy of the universe is constant This law is also called the Law of Conservation of Energy Where does the reaction energy come from that is released and where does the energy go when it is absorbed?
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Changes in Chemical Energy
A-B + C-D A-D + C-B Consider the reaction converting AB and CD to AD and CB Each chemical bond is stored chemical energy If a reaction will occur Bonds must break Breaking bonds requires energy
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Exothermic Reactions If the energy required to break the bonds is less than the energy released when the bonds are formed, there is a net release of energy This is called an Exothermic reaction Energy is a product in this reaction A-B + C-D A-D + C-B These bonds must be broken in the reaction, requiring energy These bonds are formed, releasing energy
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Endothermic Reactions
If the energy required to break the bonds is larger than the energy released when the bonds are formed, there will need to be an external supply of energy This is called an Endothermic reaction A-B + C-D A-D + C-B These bonds must be broken in the reaction, requiring energy These bonds are formed, releasing energy
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Exothermic Reaction Combustion CH4(g) + 2O2(g) CO2(g) + 2H2O(g) kcal Endothermic Reaction Decomposition 22 kcal + 2NH3(g) N2(g) + 3H2(g)
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Enthalpy Enthalpy – represents heat energy
Change in Enthalpy (DHo) – energy difference between the products and reactants of a chemical reaction Energy released, exothermic reaction, enthalpy change is negative In the combustion of CH4, DHo = -211 kcal Energy absorbed, endothermic, enthalpy change is positive. In the decomposition of NH3, DHo=+22 kcal
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Spontaneous and Nonspontaneous Reactions
Spontaneous reaction - occurs without any external energy input Most, but not all, exothermic reactions are spontaneous Thermodynamics is used to help predict if a reaction will occur Another factor is needed, Entropy
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Spontaneous and Nonspontaneous Reactions
DSo is positive DSo is negative
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Entropy The second law of thermodynamics – the universe spontaneously tends toward increasing disorder or randomness Entropy (So) – a measure of the randomness of a chemical system High entropy – highly disordered system, the absence of a regular, repeating pattern Low entropy – well organized system such as a crystalline structure
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Entropy of Reactions DSo of a reaction = So(products) - So(reactants)
A positive DSo means an increase in disorder for the reaction A negative DSo means a decrease in disorder for the reaction
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Processes Having Positive Entropy
Phase change Melting Vaporization Dissolution All of these processes have a positive DSo
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Entropy and Reaction Spontaneity
If exothermic and positive DSo… SPONTANEOUS If endothermic and negative DSo… NONSPONTANEOUS For any other situations, it depends on the relative size of DHo and DSo
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Free Energy Free energy (DGo) – represents the combined contribution of the enthalpy and entropy values for a chemical reaction Free energy predicts spontaneity of chemical reactions Negative DGo…Always Spontaneous Positive DGo…Never Spontaneous DGo = DHo - TDSo T in Kelvin
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Free Energy and Reaction Spontaneity
Need to know both DH and DS to predict the sign of DG, making a statement on reaction spontaneity Temperature also may determine direction of spontaneity DH +, DS - : DG always +, regardless of T DH -, DS + : DG always -, regardless of T DH +, DS + : DG sign depends on T DH -, DS - : DG sign depends on T
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7.2 Experimental Determination of Energy Change in Reactions
Calorimetry – the measurement of heat energy changes in a chemical reaction Calorimeter – device which measures heat changes in calories The change in temperature is used to measure the loss or gain of heat
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Heat Energy in Reactions
Change in temperature of a solution, caused by a chemical reaction, can be used to calculate the gain or loss of heat energy for the reaction Exothermic reaction – heat released is absorbed Endothermic reaction – reactants absorb heat from the solution Specific heat (SH) – the number of calories of heat needed to raise the temperature of 1 g of the substance 1 oC
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Heat Energy in Reactions
Specific heat of the solution along with the total number of g solution and the temperature change permits calculation of heat released or absorbed during the reaction 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)
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Calculation of Heat Energy in Reactions
Q is the product ms is the mass of solution in the calorimeter DTs is the change in temperature of the solution from initial to final state SHs is the specific heat of the solution Calculate with this equation Units are: calories = gram ºC calories/gram ºC
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Calculating Energy Involved in Calorimeter Reactions
If 0.10 mol HCl is mixed with 0.10 mol KOH in a “coffee cup” calorimeter, the temperature of 1.50 x 102 g of the solution increases from 25.0oC to 29.4oC. If the specific heat of the solution is 1.00 cal/goC, calculate the quantity of energy evolved in the reaction DTs = 29.4oC oC = 4.4oC Q = ms x DTs x SHs = 1.50 x 102 g solution x 4.4oC x 1.00 cal/goC = 6.6 x 102 cal
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Calculating Energy Involved in Calorimeter Reactions
Is the reaction endothermic or exothermic 0.66 kcal of heat energy was released to the surroundings, the solution The reaction is exothermic What would be the energy evolved for each mole of HCl reacted? 0.10 mol HCl used in the original reaction [6.6 x 102 cal / 0.10 mol HCl] x 10 = 6.6 kcal
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Bomb Calorimeter and Measurement of Calories in Foods
Nutritional Calorie (large “C” Calorie) = 1 kilocalorie (1kcal) 1000 calories the fuel value of food Bomb Calorimeter is used to measure nutritional Calories
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Calculating the Fuel Value of Foods
1 g of glucose was burned in a bomb calorimeter x 103 g H2O was warmed from 24.5oC to 30.6oC. Calculate the fuel value of the glucose (in Kcal/g). DTs = 30.6oC oC = 6.1oC Surroundings of calorimeter is water with specific heat capacity = 1.00 cal/g H2O oC Fuel Value = = 1.25 x 103 g H2O x 6.1oC x 1.00 cal/g H2O oC = 7.6 x 103 cal 7.6 x 103 cal x 1 Calorie / 103 cal = 7.6 nutritional Calories
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7.3 Kinetics Thermodynamics determines if a reaction will occur spontaneously but tells us nothing about the amount of time the reaction will take Kinetics – the study of the rate (or speed) of chemical reactions Also supplies an indication of the mechanism: step-by-step description of how reactants become products
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Kinetic Information Kinetic information represents changes over time, seen here: disappearance of reactant, A appearance of product, B
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Kinetic Data Assessed by Color Change
Change in color over time can be used to monitor the progress of a chemical reaction The rate of color change can aid in calculating the rate of the chemical reaction
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CH4(g) + 2O2(g) CO2(g) +2H2O(g) + 211 kcal
The Chemical Reaction CH4(g) + 2O2(g) CO2(g) +2H2O(g) kcal C-H and O=O bonds must be broken C=O and O-H bonds must be formed Energy is required to break the bonds This energy comes from the collision of the molecules If sufficient energy available, bonds break and atoms recombine in a lower energy arrangement Effective collision is one that produces product molecules leads to a chemical reaction
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Activation Energy and the Activated Complex
Activation energy: the minimum amount of energy required to initiate a chemical reaction Picture a chemical reaction in terms of changes in potential energy occurring during the reaction Activated complex: an extremely unstable, short-lived intermediate complex Formation of this activated complex requires energy (Ea) to overcome the energy barrier to start the reaction
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Activation Energy and the Activated Complex
Reaction proceeds from reactants to products via the activated complex Activated complex – can’t be isolated from the reaction mixture Activation energy (Ea) is the difference between the energy of the reactants and that of the activated complex To be an Exothermic reaction requires a net release of energy (DHo)
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Activation Energy in the Endothermic Reaction
This figure diagrams an endothermic reaction Reaction takes place slowly due to the large activation energy required The energy of the products is greater than that of the reactants
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Factors That Affect Reaction Rate
Structure of the reacting species Concentration of reactants Temperature of reactants Physical state of reactants Presence of a catalyst
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Structure of Reacting Species
Oppositely charged species react more rapidly Dissociated ions in solution whose bonds are already broken have a very low activation energy Ions with the same charge do not react Bond strength plays a role Covalent molecules bonds must be broken with the activation energy before new bonds can be formed Magnitude of the activation energy is related to bond strength Size and shape influence the rate Large molecules may obstruct the reactive part of the molecule Only molecular collisions with correct orientation lead to product formation
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The Concentration of Reactants
Rate is related to the concentration of one or more of the reacting substances Rate will generally increase as concentration increases Higher concentration means more reactant molecules per unit volume More reactant molecules means more collisions per unit time
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The Temperature of Reactants
Rate increases as the temperature increases Increased temperature relates directly to increased average kinetic energy Greater kinetic energy increases the speed of particles Faster particles increases likelihood of collision Higher Kinetic Energy means a higher percentage of these collisions will result in product formation
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The Physical State of Reactants
Reactions occur when reactants can collide frequently with sufficient energy to react Solid state: atoms, ions, compounds are close together but restricted in motion Gaseous state: particles are free to move but often are far apart causing collisions to be relatively infrequent Liquid state: particles are free to move and are in close proximity Reactions to be fastest in the liquid state and slowest in the solid state Liquid > Gas> Solid
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The Presence of a Catalyst
Catalyst – a substance that increases the reaction rate Undergoes no net change (i.e., is not consumed) Does not alter the equilibrium Interacts with the reactants to create an alternative pathway for product production
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Use of a Solid Phase Catalyst
N2+3H2 2NH3 Haber Process is a synthesis of ammonia facilitated by a solid phase catalyst Diatomic gases bind to the surface Bonds are weakened Dissociation of diatomic gases and reformation of NH3 Newly formed NH3 leaves the solid surface with catalyst unchanged
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Mathematical Representation of Reaction Rate
Consider a decomposition reaction with the following balanced chemical equation: When heated N2O5 decomposes to 2 products – NO2 and O2 When holding all factors constant, except concentration, rate of reaction is proportional to the concentration
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Mathematical Representation of Reaction Rate
Reaction rate is proportional to reactant concentration: Concentration of N2O5 is denoted as [N2O5] Replace proportionality symbol with = and proportionality constant k k is called the rate constant
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Rate Equation For a reaction Aproducts we write the equation:
rate = k[A]n This is called the rate equation (or rate law) The exponent n is the order of the reaction If n=1, first order If n=2, second order n must be determined experimentally This exponent is not the same as the coefficient of the reactant in the balanced equation
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Rate Equation For the equation A + B products the rate equation is:
rate = k[A]n[B]m What would be the general form of the rate equation for the reaction: CH4+2O2CO2+2H2O Rate = k[CH4]n[O2]m Knowing the rate equation and the order helps industrial chemists determine the optimum conditions for preparing a product
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Writing Rate Equations
Write the form of the rate equation for the oxidation of ethanol (C2H5OH) The reaction has been experimentally determined to be first order in ethanol and third order in oxygen Rate expression involves only the reactants Concentrations: [C2H5OH][O2] Raise each to exponent corresponding to its order rate = k [C2H5OH][O2]3 Remember that 1 as an exponent is understood and NOT written Knowing the rate equation and the order helps industrial chemists determine the optimum conditions for preparing a product
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7.4 Equilibrium Rate and Reversibility of Reactions
Equilibrium reactions: chemical reactions that do not go to completion Completion – all reactants have been converted to products Equilibrium reactions are also called Incomplete reactions Seen with both physical and chemical processes After no further observable change, measurable quantities of reactants and products remain
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Physical Equilibrium Physical equilibria are reversible reactions
Dissolved oxygen in lake water Sugar dissolved in water Reversible reaction – a process that can occur in both directions Use the double arrow symbol Dynamic equilibrium – the rate of the forward process in a reversible reaction is exactly balanced by the rate of the reverse process
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Sugar in Water If add 2-3 g of sugar into 100 mL water
All will dissolve with stirring in a short time No residual solid sugar, sugar dissolved completely Sugar (s) Sugar (aq) If add 100 g of sugar in 100 mL of water Not all of it will dissolve even with much stirring Over time, you observe no further change in the amount of dissolved sugar Appears nothing is happening – Incorrect!
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Sugar in Water If add 100 g of sugar in 100 mL of water
Appears nothing is happening – Incorrect! Individual sugar molecules are constantly going into and out of solution Both happen at the same rate Over time the amount of sugar dissolved in the measured volume of water does not change An equilibrium situation has been established Some molecules dissolve and others return to the solid state; the rate of each process is equal Sugar(s) Sugar(aq)
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Dynamic Equilibrium sugar(s) sugar(aq)
The double arrow serves as an indicator of a reversible process an equilibrium process the dynamic nature of the process Continuous change is taking place without observable change in the amount of sugar in either the solid or the dissolved form
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Equilibrium Constant ratef = forward rate rater = reverse rate
at equilibrium: ratef = rater ratef = kf[sugar(s)] rater = kr[sugar(aq)] kf[sugar(s)]=kr[sugar(aq)] Equilibrium constant (Keq): ratio of the two rate constants
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Chemical Equilibrium The Reaction of N2 and H2 N2(g) + 3H2(g) 2NH3(g)
Mix components at elevated temperature Some molecules will collide with sufficient energy to break N-N and H-H bonds Rearrangement of the atoms will produce the product NH3
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Chemical Equilibrium N2(g) + 3H2(g) 2NH3(g)
Initially the forward reaction is rapid Reactant concentrations are high Product concentration negligible Forward reaction rate decreases with time Concentrations of reactants are decreasing Product concentration increasing Equilibrium occurs when the rate of reactant depletion is equal to the rate of product depletion Rates of forward and reverse reactions are Equal
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The Generalized Equilibrium-Constant Expression for a Chemical Reaction
aA + bB cC + dD A and B are reactants C and D are products a, b, c, and d are the coefficients of the balanced equation
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Writing Equilibrium-Constant Expressions
Equilibrium constant expressions can only be written after a correct, balanced chemical equation Each chemical reaction has a unique equilibrium constant value at a specified temperature The brackets represent molar concentration (except gases, they are in partial pressures) 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
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Writing an Equilibrium-Constant Expression
Write an equilibrium-constant expression for the reversible reaction: H2(g) + F2(g) HF(g) No solids or liquids are present All reactants and products appear in the expression Numerator term is the product term [HF]2 Denominator term is the reactants [H2] and [F2] Each term contains an exponent identical to the corresponding coefficient in the balanced equation Keq = [HF]2 [H2][F2]
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Writing an Equilibrium-Constant Expression
Write an equilibrium-constant expression for the reversible reaction: MnO2(s) + 4H+(aq) + 2Cl-(aq) Mn2+(aq) + Cl2(g) + H2O(l) MnO2 is a solid H2O(l) is a product, but negligible compared to solvent water Numerator term is the product terms [Mn2+] and [Cl2] Denominator term is the reactants [H+]4 and [Cl-]2 Each term contains an exponent identical to the corresponding coefficient in the balanced equation Keq = [Mn2+] [Cl2] [H+]4 [Cl-]2
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Interpreting Equilibrium Constants
Reversible arrow in chemical equation indicates equilibrium exists The numerical value of the equilibrium constant tells us the extent to which reactants have converted to products Keq greater than 1 x 102 Large value of Keq indicates numerator (product term) >>> denominator (reactant term) At equilibrium mostly product present
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Interpreting Equilibrium Constants
Keq less than 1 x 10-2 Small value of Keq indicates numerator (product term) <<< denominator (reactant term) At equilibrium mostly reactant present Keq between 1 x 10-2 and 1 x 102 Equilibrium mixture contains significant concentration of both reactants and products
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Calculating Equilibrium Constants
A reversible reaction is allowed to proceed until the system reaches equilibrium Amount of reactants and products no longer changes Analyze reaction mixture to determine molar concentrations of each product and reactant 2NO2(g) N2O4(g)
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Calculating an Equilibrium Constant
HI placed in a sealed container and comes to equilibrium; equilibrium reaction is: 2HI(g) H2(g) + I2(g) Equilibrium concentrations: [HI] = 0.54 torr [H2] = 1.72 torr [I2] = torr Substitute concentrations: Keq = [H2] [I2] [HI]2 Keq= [1.72] [1.72] = 10.1 [0.54]2
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LeChateleir’s Principle
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 If reactants and products are present in a fixed volume and more NH3 is added into the container, the system will be stressed Stressed = the equilibrium will be disturbed N2(g) + 3H2(g) NH3(g)
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LeChateleir’s Principle
Adding NH3 to the system causes stress To relieve stress, remove as much of added material as possible by converting it to reactants Adding N2 or H2 to the system causes stress also To relieve stress, remove as much of added material as possible by converting it to product N2(g) + 3H2(g) NH3(g) Equilibrium shifted Product introduced: Reactant introduced:
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Effect of Concentration
N2(g) + 3H2(g) NH3(g) Adding or removing either reactants or products at a fixed volume is saying that the concentration is changed Removing material decreases concentration System will react to this stress to return concentrations to the appropriate ratio
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Effect of Heat Exothermic reactions: treat heat as a product
N2(g) + 3H2(g) NH3(g) + 22 kcal Addition of heat is treated as increasing the amount of product More product shifts equilibrium to the left Increases amount of reactants Decreases amount of product
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Effect of Heat Endothermic Reaction – treat heat as a reactant
39 kcal + 2N2(g) + O2(g) N2O(g) This reaction shift will shift to the right if heat is added by increasing the temperature
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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
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Effect of Pressure 2HI(g) H2(g) + I2(g) N2(g) + 3H2(g) 2NH3(g)
If increase pressure, which way will the equilibrium shift? Increased pressure favors decreased volume with more product (2 moles) formed and less reactant (4 moles) 2HI(g) H2(g) + I2(g) If increase the pressure in this reaction, which way will the equilibrium shift? No shift in equilibrium as both reactant and product have 2 moles of gas
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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 While equilibrium composition and concentration do not change, equilibrium is reached in a shorter time
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