Week 8.1   thermodynamic and equilibria
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Week 8.1 thermodynamic and equilibria






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Week 8.1   thermodynamic and equilibria Week 8.1 thermodynamic and equilibria Presentation Transcript

  • Chapter 4 Thermodynamic and Equilibria First Law of Thermodynamics Second Law of Thermodynamics Physical Transformation of Pure Substances Prepared by: Mrs Faraziehan Senusi PA-A11-7C Simple Mixtures Chemical Equilibrium Reference: Chemistry: the Molecular Nature of Matter and Change, 6th ed, 2011, Martin S. Silberberg, McGraw-Hill
  • Thermochemistry • Thermodynamics ~ energy changes that accompany physical and chemical processes. Usually these energy changes involve heat. • Thermochemistry ~ is the study of heat change in chemical reactions.  concerned with how we observe, measure, and predict energy changes for both physical changes and chemical reactions  use energy changes to tell whether or not a given process can occur under specified conditions to give predominantly products (or reactants) and how to make a process more (or less) favorable.
  • Basic concepts • System : specific part of the universe that is of interest to us. The substances involved in the chemical and physical changes that we are studying • Surroundings: the rest of the universe boundary surrounding system
  • • Open system: can exchange mass and energy (heat) with surrounding • system where mass and energy can cross the boundary matter surroundings system energy • Closed system: allows transfer of heat but not mass • consists of a fixed amount of mass and no mass can cross its boundary • Energy in the form of heat or work can cross the boundary • Isolated system: does not allow transfer either mass or energy • A system where no mass, heat and work can cross the boundary. matter surroundings system energy matter system energy surroundings
  • THE FIRST LAW OF THERMODYNAMICS • The first law of thermodynamics (the conservation of energy principle) provides a sound basis for studying the relationships among the various forms of energy and energy interactions. The first law states that energy can be neither created nor destroyed; it can only change forms. When a rock falls, the decrease in potential energy is equals to the increase in kinetic energy. The increase in the energy of a potato in an oven is equal to the amount of heat transferred to it.
  • INTERNAL ENERGY, E Each particle in a system has potential and kinetic energy; the sum of these energies for all particles in a system is the internal energy, E. In a chemical reaction: when reactants are converted to products, E changes (DE). DE = Efinal - Einitial = Eproducts - Ereactants
  • Energy diagrams for the transfer of internal energy (E) between a system and its surroundings
  • Heat and Work • Energy transfer outward from the system or inward from the surroundings can appear in two forms, heat and work . DE = q + w where q = heat and w = work • Heat (or thermal energy, symbol q) is the energy transferred between a system and its surroundings as a result of a difference in their temperatures only. • All other forms of energy transfer (mechanical, electrical, and so on) involve some type of work (w), the energy transferred when an object is moved by a force.
  • Sign Conventions for q, w and DE • • The numerical values of q and w can be either positive or negative, depending on the change the system undergoes. Energy coming into the system is positive; energy going out from the system is negative. q + w = DE + + + + - depends on magnitudes of q and w - + depends on magnitudes of q and w - - - For q: (+) means system gains heat, (-) means system loses heat. For w: (+) means work done on system (compression), (-) means work done by system,(expansion).
  • Law of Conservation of Energy (First Law of Thermodynamics) The energy of the system plus the energy of the surroundings remains constant: energy is conserved. DEuniverse = DEsystem + DEsurroundings = 0 Units of Energy joule (J) 1 J = 1 kg m2/s2 calorie (cal) 1 cal = 4.184 J British Thermal Unit 1 Btu = 1055 J
  • Thermodynamic state of a system • The properties of a system—such as P, V, T—are called state functions • The value of a state function depends only on the state of the system and not on the way in which the system came to be in that state. • A change in a state function describes a difference between the two states. It is independent of the process or pathway by which the change.
  • • The most important use of state functions in thermodynamics is to describe changes. ΔX = ΔXfinal – ΔXinitial • When X increases, the final value is greater than the initial value, so ΔX is positive; a decrease in X makes ΔX a negative value.
  • Calorimetry • We can determine the energy change associated with a chemical or physical process by using an experimental technique called calorimetry. • This technique is based on observing the temperature change when a system absorbs or releases energy in the form of heat. • The experiment is carried out in a device called a calorimeter, in which the temperature change of a known amount of substance (often water) of known specific heat is measured. • The temperature change is caused by the absorption or release of heat by the chemical or physical process under study.
  • • Specific heat, c : amount required to raise temperature of one gram of the substance by one degree Celsius (J/g.oC) • Heat capacity or calorimeter constant, C : amount of heat required to raise the temperature of a given quantity of the substance by one degree Celsius (J/oC) C = mc (where m is the mass) • If specific heat and amount of substance is known, change in sample temperature ∆T will tell us the amount of heat, q, that has been released/ absorbed in particular process. q = mc∆T • q is positive (endothermic) and negative for exothermic process
  • • Calorimeter can be used to measure the amount of heat absorbed or released when a reaction takes place in aqueous solution.
  • Example 1 We add 3.358 kJ of heat to a calorimeter that contains 50.00 g of water. The temperature of the water and the calorimeter, originally at 22.34°C, increases to 36.74°C. Calculate the heat capacity of the calorimeter in J/°C. The specific heat of water is 4.184 J/g.°C.  Calculate the amount of heat gained by the water in the calorimeter.  The rest of the heat must have been gained by the calorimeter.  Determine the heat capacity of the calorimeter.
  • Example 1 amount of heat gained by the water amount of heat gained by the calorimeter
  • Example 2 A 50.0 mL sample of 0.400 M copper(II) sulfate solution at 23.35°C is mixed with 50.0 mL of 0.600M sodium hydroxide solution, also at 23.35°C, in the coffee-cup calorimeter. Heat capacity of calorimeter is 24.0 J/ C. After the reaction occurs, the temperature of the resulting mixture is measured to be 25.23°C. The density of the final solution is 1.02 g/mL. Calculate the amount of heat evolved. Assume that the specific heat of the solution is the same as that of pure water, 4.184 J/g.°C.  The amount of heat released by the reaction is absorbed by the calorimeter and by the solution.  To find the amount of heat absorbed by the solution, we must know the mass of solution; to find that, we assume that the volume of the reaction mixture is the sum of volumes of the original solutions.
  • Example 2
  • The practice of calorimetry • Two common types are :  constant-pressure calorimeters A "coffee-cup" calorimeter is often used to measure the heat transferred (qp) in processes open to the atmosphere.  constant-volume calorimeters One type of constant-volume apparatus is the bomb calorimeter, designed to measure very precisely the heat released in a combustion reaction.
  • Constant-pressure Calorimetry • One common use is to find the specific heat capacity of a solid that does not react with or dissolve in water. • The solid (system) is weighed, heated to some known temperature, and added to a sample of water (surroundings) of known temperature and mass in the calorimeter. • With stirring, the final water temperature, which is also the final temperature of the solid, is measured. • The heat lost by the system (-qsys, or -qsolid) is equal in magnitude but opposite in sign to the heat gained by the surroundings (+qsurn or +qH2O): - qsolid = q H2O Or, - (c solid x mass solid x ΔT solid ) = c H2O x mass H2O x ΔTH2O
  • Example 3 Determining the Specific Heat Capacity of a Solid PROBLEM: A 25.64 g sample of a solid was heated in a test tube to 100.00 oC in boiling water and carefully added to a coffee-cup calorimeter containing 50.00 g of water. The water temperature increased from 25.10 oC to 28.49 oC. What is the specific heat capacity of the solid? (Assume all the heat is gained by the water) PLAN: It is helpful to use a table to summarize the data given. Then work the problem realizing that heat lost by the system must be equal to that gained by the surroundings. mass (g) c (J/g.K) Tinitial Tfinal DT 25.64 ? 100.00 28.49 -71.51 50.00 4.184 25.10 28.49 3.39 solid H2 O SOLUTION: 25.64 g x c x csolid = - 50.00 g x -71.51 K 4.184 J/g.K x 3.39 K 25.64 g x -71.51 K 4.184 J/g.K x 3.39 K = - 50.00 g x = 0.387 J/g.K
  • Constant-volume Calorimetry • • • Figure 6.8 depicts the preweighed combustible sample in a metal-walled chamber (the bomb), which is filled with oxygen gas and immersed in an insulated water bath fitted with motorized stirrer and thermometer. A heating coil connected to an electrical source ignites the sample, and the heat evolved raises the temperature of the bomb, water, and other calorimeter parts. Because we know the mass of the sample and the heat capacity of the entire calorimeter, we can use the measured ΔT to calculate the heat released.
  • Example 4 Calculating the Heat of Combustion PROBLEM: PLAN: SOLUTION: A manufacturer claims that its new diet dessert has ―fewer than 10 Calories (10 kcal) per serving‖. To test the claim, a chemist at the Department of Consumer Affairs places one serving in a bomb calorimeter and burns it in O2 (the heat capacity of the calorimeter = 8.151 kJ/K). The temperature increases by 4.937 oC. Is the manufacturer’s claim correct? - qsample = qcalorimeter qcalorimeter = heat capacity x DT = 8.151 kJ/K x 4.937 K = 40.24 kJ 40.24 kJ x kcal = 9.62 kcal < 10 Calories = 10 kcal 4.184 kJ The manufacturer’s claim is correct.
  • Enthalpy • The quantity of heat transferred into or out of a system as it undergoes a chemical or physical change at constant pressure. • Extensive property: magnitude depends on amount of substance present • Impossible to determine enthalpy of substance • Measure change in enthalpy, ∆H
  • • Enthalpy of reaction, ∆H ∆H = H (product) – H (reactant) • Exothermic : negative • Endothermic: positive • EXOTHERMIC PROCESS – a process that releases energy in the form of heat into its surroundings. (Ex: combustion reaction) • ENDOTHERMIC PROCESS – a process that absorbs energy from its surroundings
  • Enthalpy diagrams for exothermic and endothermic processes CH4(g) + 2O2(g) CO2(g) + 2H2O(g) + heat heat + H2O(s) Heat is released; enthalpy decreases. H2O(l) Heat is absorbed; enthalpy increases.
  • Some Important Types of Enthalpy Change heat of combustion (DHcomb) 1C4H10(l) + 13/2O2(g) 4CO2(g) + 5H2O(g) heat of formation (DHf) K(s) + 1/2Br2(l) 1KBr(s) heat of fusion (DHfus) 1NaCl(s) NaCl(l) heat of vaporization (DHvap) 1C6H6(l) C6H6(g) Standard quantity of either reactant or product: 1 mol
  • Thermochemical equations • A balanced chemical equation, together with its value of ΔH • The ΔHrxn value shown refers to the amounts (moles) of substances and their states of matter in that specific equation. Combustion of methane : CH4 (g) + 2O2 (g)  CO2 (g) + 2H2O(l) ∆H= -890.4kJ
  • • 1367 kJ of heat is released when one mole of C2H5OH(l) reacts with three moles of O2(g) to give two moles of CO2(g) and three moles of H2O(l). • We can refer to this amount of reaction as one mole of reaction, which we abbreviate ―mol rxn.‖ • We can also write the thermochemical equation as • We always interpret ΔH as the enthalpy change for the reaction as written; as (enthalpy change)/(mole of reaction), where the denominator means ―for the number of moles of each substance shown in the balanced equation.‖
  • • Guidelines for writing thermochemical equations and interpreting – Stoichiometric coefficient refer to number of moles of substance – Reverse equation, magnitude of ∆H same but sign changes – Multiply equation by factor of n then ∆H must also change by the same factor – Must always specify the physical state
  • AMOUNT (mol) Summary of the relationship between amount (mol) of substance and the heat (kJ) transferred during a reaction of compound A AMOUNT (mol) molar ratio from balanced equation of compound B HEAT (kJ) DHrxn (kJ/mol) gained or lost
  • Example PROBLEM: Using the Heat of Reaction (DHrxn) to Find Amounts The major source of aluminum in the world is bauxite (mostly aluminum oxide). Its thermal decomposition can be represented by: Al2O3(s) 2Al(s) + 3/2O2(g) DHrxn = 1676 kJ If aluminum is produced this way, how many grams of aluminum can form when 1.000 x 103 kJ of heat is transferred? PLAN: SOLUTION: heat (kJ) 1.000 x 103 kJ x 2 mol Al 1676 kJ 1676 kJ = 2 mol Al mol of Al x M (g/mol) g of Al = 32.20 g Al x 26.98 g Al 1 mol Al
  • Standard heat of reaction Specifying Standard States For a gas, the standard state is 1 atm; ideal gas behavior is assumed. For a substance in aqueous solution, the standard state is 1 M concentration (1 mol/liter solution). For a pure substance (element or compound), the standard state is usually the most stable form of the substance at 1 atm and the temperature of interest (usually 25 oC (298 K). DHorxn = standard heat of reaction (enthalpy change determined with all substances in their standard states)
  • Standard enthalpy of formation • Standard enthalpy of formation (∆Hof) : heat change that results when one mole of a compound is formed from its elements at a pressure of 1 atm (standard state). • ∆Hof of any element in its most stable form is zero ∆Hof (O2) = 0 ∆ Hof (O3) ≠ 0
  • Direct method of measuring ∆Hof Formation Equations In a formation equation, 1 mol of a compound forms from its elements. The standard heat of formation (DHof) is the enthalpy change for the formation equation when all substances are in their standard states. C(graphite) + 2H2(g) C(graphite) + O2 (g)  CO2 (g) CH4(g) DHof = -74.9 kJ ∆Horxn= - 393.5kJ An element in its standard state is assigned a DHof of 0.
  • Selected Standard Heats of Formation at 25 oC (298 K) Table 6.5 Formula DHof (kJ/mol) calcium Ca(s) CaO(s) CaCO3(s) carbon C(graphite) C(diamond) CO(g) CO2(g) CH4(g) CH3OH(l) HCN(g) CS2(l) chlorine Cl(g) 0 -635.1 -1206.9 0 1.9 -110.5 -393.5 -74.9 -238.6 135 87.9 121.0 Formula DHof (kJ/mol) Formula 0 -92.3 silver Ag(s) AgCl(s) hydrogen H(g) H2(g) 218 0 sodium nitrogen N2(g) NH3(g) NO(g) 0 -45.9 90.3 DHof (kJ/mol) Cl2(g) HCl(g) oxygen O2(g) O3(g) H2O(g) 0 143 -241.8 H2O(l) -285.8 Na(s) Na(g) NaCl(s) 0 -127.0 0 107.8 -411.1 sulfur S8(rhombic) 0 S8(monoclinic) 2 SO2(g) -296.8 SO3(g) -396.0
  • Indirect method measuring ∆Hof • Many compounds cannot be directly synthesized from their elements (proceed too slowly/ side reactions produce other substance than desired compound) • Hess’s law: when reactants are converted to products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps • General rule of applying: – Arrange series of chemical equation in a way that when added together all species will cancel except reactant and product that appear in the overall reaction – Multiply or reverse equation
  • Hess’s Law of Heat Summation The enthalpy change of an overall process is the sum of the enthalpy changes of its individual steps. Used to predict the enthalpy change (a) of an overall reaction that cannot be studied directly, and/or (b) of an overall reaction that can be separated into distinct reactions whose enthalpy changes can be measured individually.
  • Example PROBLEM: Using Hess’s Law to Calculate an Unknown DH Two gaseous pollutants that form auto exhaust are CO and NO. An environmental chemist is studying ways to convert them into less harmful gases through the following equation: CO(g) + NO(g) CO2(g) + 1/2N2(g) DH = ? Given the following information, calculate the unknown DH. Equation A: CO(g) + 1/2O2(g) 2NO(g) DHB = +180.6 kJ Equation B: N2(g) + O2(g) PLAN: CO2(g) DHA = -283.0 kJ Equations A and B have to be manipulated by reversal and/or multiplied by factors in order to sum to the target equation. SOLUTION: Multiply Equation B by 1/2 and reverse it. CO(g) + 1/2O2(g) NO(g) CO(g) + NO(g) CO2(g) 1/2N2(g) + 1/2O2(g) CO2(g) + 1/2N2(g) DHA = -283.0 kJ DHB = -90.3 kJ DHrxn = -373.3 kJ
  • • From ∆Hof values, the standard enthalpy of reaction ∆Horxn can be calculated aA + bB  cC + dD ∆Horxn = [c∆Hof (C) + d∆Hof (D)] – [a∆Hof (A) + b∆Hof (B)] ∆Horxn = ∑n∆Hof (Product) –∑m∆Hof(Reactant)
  • The general process for determining DHorxn from DHof values
  • Example Calculating the Standard Heat of Reaction from Standard Heats of Formation PROBLEM: Nitric acid, whose worldwide annual production is about 8 billion kg, is used to make many products, including fertilizers, dyes and explosives. The first step in the industrial production process is the oxidation of ammonia: 4NH3(g) + 5O2(g) 4NO(g) + 6H2O(g) Calculate DHorxn from DHof values. SOLUTION: DHorxn = S nDHof (products) - S mDHof (reactants) DHorxn = [4DHof NO(g) + 6DHof H2O(g)] - [4DHof NH3(g) + 5DHof O2(g)] = [(4 mol)(90.3 kJ/mol) + (6 mol)(-241.8 kJ/mol)] [(4 mol)(-45.9 kJ/mol) + (5 mol)(0 kJ/mol)] DHorxn = -906 kJ
  • Bond Energies • Chemical reactions involve the breaking and making of chemical bonds. • Energy is always required to break a chemical bond. • The bond energy (B.E.) is the amount of energy necessary to break one mole of bonds in a gaseous covalent substance to form products in the gaseous state at constant temperature and pressure. • The greater the bond energy, the more stable (stronger) the bond is, and the harder it is to break. Thus bond energy is a measure of bond strengths. • A special case of Hess’s Law involves the use of bond energies to estimate heats of reaction.
  • A schematic representation of the relationship between bond energies and ΔHrxn for gas phase reactions. (a) For a general reaction (exothermic) (b) For the gas phase reaction : H2 (g) + Br2 (g)  2HBr (g)
  • Example 1 Use the bond energies listed in Table 15-2 to estimate the heat of reaction at 298 K for the following reaction:
  • Example 2 Use the bond energies listed in Table 15-2 to estimate the heat of reaction at 298 K for the following reaction: