Thermochemistry Chapter 5 BLB 12th
Expectations • Heat & enthalpy – same or different? • Heat calculations: • Temp. change • Phase change (11.4, p. 438) • Reactions • Enthalpy calculations • Read the chapter, study, and apply!
5.1 The Nature of Energy Chemistry ⇐ ? ⇒ Energy • Energy – capacity to do work or transfer heat • Potential – stored energy; chemical • Kinetic – released energy; energy of motion; thermal • Electrostatic potential – interaction between charged particles
Energy, cont. Units of energy: • Joule (J) – SI unit of energy; • calorie (cal) • amount of energy required to raise the temperature of exactly 1 gram of pure water by 1°C (from 14.5°C to 15.5°C) • 1 cal = 4.184 J (exactly) • Calorie (dietary calorie),Cal • 1 Cal = 1000 cal = 1 kcal
Energy, cont. System and Surroundings • System – component(s) of interest • Open – matter and energy can be exchanged between system and surroundings • Closed – can exchange energy but not matter • Isolated – neither energy nor matter can be exchanged • Surroundings – everything outside of the system
Energy, cont. Transferring Energy • Work (w) – energy used to move an object against a force; w = F x d • Heat (q) – energy transferred from a hotter object to a cooler one • Energy – capacity to do work or transfer heat; ΔE = q + w
Combustion heat & work
5.2 The First Law of Thermodynamics Energy can be neither created nor destroyed. • Energy is conserved. • Internal energy, E – sum of all the kinetic and potential energy of the system’s components What kinds of energy are in here? What changes could occur?
5.2 The First Law of Thermodynamics • More interested in the change in energy: ΔE = Efinal – Einitial • Need to give number, units, and sign for all thermodynamic quantities. • ΔE > 0 - Efinal> Einitial, system has gained energy; endergonic • ΔE < 0 - Efinal< Einitial,system has lost energy; exergonic Note: Opposite change occurs with respect to the surroundings.
Energy, heat & work • ΔE = q + w • Sign of ΔE depends upon sign and magnitude of q and w.
Exothermic • Efinal < Einitial
Sample Exercise 5.2 A(g) + B(g) → C(s) • System loses 1150 J of heat to the surroundings. • The piston move downwards doing 480 J of work on the system. • ΔE = ?
Calculate ΔE (in J); exothermic or endothermic? • Balloon heating by adding 900 J of heat and expands doing 422 J of work on atmosphere. • 50 g of H2O cooled from 30°C to 15°C losing 3140 J of heat. • Reaction releases 8.65 kJ of heat, no work done.
Heat or Thermal Energy (q) • Exothermic: system → surroundings • Heat energy released to surroundings • q < 0 • e.g. combustion reaction, crystallization • Surroundings get warmer • Endothermic: system ← surroundings • Heat energy flows into the system • q > 0 • e.g. melting, boiling, dissolution of NH4NO3 • Surroundings get colder
Heat, cont. • Evidenced by a change in temperature • Spontaneously transferred from the hotter to the cooler object • Atoms or molecules with more energy move faster • Temperature-dependent • Extensive property (depends on amount) • Total energy of system is the sum of the individual energies of all the atoms and molecules of the system.
Work (w) • Force acting over a distance w = F x d = −PΔV • Compression: work ← surroundings • Work is done on the system. • ΔV < 0 • w > 0 • Expansion: work → surroundings • Work is done on the surroundings. • ΔV > 0 • w < 0
Heat & Work, cont. • Work and heat are pathways by which energy can be transferred. • State function – depends only on the system’s present state; independent of the pathway; internal energy, P, V, ΔE, ΔH, ΔS are state functions • Energy is a state function, as is enthalpy.
5.3 Enthalpy • Enthalpy – heat flow at constant pressure; from Gr. enthalpien – to warm • Enthalpy change (ΔH) – energy transferred as heat at constant pressure; ΔH = Hproducts – Hreactants • H = E + PV • For a change @ constant pressure: ΔH = ΔE + PΔV ΔH = ΔE − w =qP
5.3 Enthalpy • ΔH < 0 exothermic: reactants → products + heat • ΔH > 0 endothermic: reactants + heat → products • ΔH = q/mol • Enthalpy (or heat) of reaction,ΔHrxn – enthalpy change that accompanies a reaction
5.4 Enthalpies of Reaction, ΔHrxn • ΔH is an extensive property; value depends upon the BALANCED equation. 2 H2(g) + O2(g) → 2 H2O(g) ΔH = +483.6 kJ per 2 moles of H2O H2(g) + ½ O2(g) → H2O(g) ΔH = −241.8 kJ per mole of H2O
5.4 Enthalpies of Reaction • For reverse reactions: ΔH values are equal in magnitude, but opposite in sign. For water: ΔHvap = +44.0 kJ/mol ΔHcond = −44.0 kJ/mol
5.4 Enthalpies of Reaction • ΔH is dependent upon physical state. ΔHf values: C6H6(g) = 82.9 kJ/mol C6H6(l) = 49.0 kJ/mol H2O(l) → H2O(g) ΔH = +44 kJ
CH3OH(g) → CO(g) + 2 H2(g) ΔHrxn= +90.7 kJ • Exothermic or endothermic? • Heat transferred for 1.60 kg CH3OH? • If 64.7 kJ of heat were used, how many grams of H2 would be produced?
CH3OH(g) → CO(g) + 2 H2(g) ΔHrxn= +90.7 kJ • ΔH of reverse reaction? Heat (in kJ) released when 32.0 g of CO(g) reacts completely?
5.5 Calorimetry • Calorimetry – science of measuring heat flow • Calorimeter – a device used to measure heat flow Coffee-cup calorimeter ⇒
Heat Capacity and Specific Heat • Heat capacity (C) - amount of heat required for a 1°C temperature change: J/°C = J/K • extensive property • ΔTin K = ΔT in °C
Heat Capacity and Specific Heat • Specific heat capacity (Cs) – heat capacity for 1 g; J/g·°C or J/g·K • Molar heat capacity – heat capacity for 1 mole; J/mol·°C or J/mol·K
Heat Capacity and Specific Heat Specific heat values (more on p. 176): Fe 0.45 J/g·K glass 0.84 J/g·K water 4.18 J/g·K(highest of all liquids and solids except ammonia)
Calculating heat (q) • To calculate the quantity of heat transferred: q = Csx m xΔT q – heat (J) Cs – specific heat (J/g∙K) m – mass (g) ΔT – change in temp. (K or°C)
Calculate the heat (in J) required to raise the temperature of 62.0 g toluene from 16.3°C to 38.8°C. The specific heat of toluene is 1.13 J/g·K.
Calculate the specific heat of lead if 78.2 J of heat were required to raise the temperature of a 45.6-g block of lead by 13.3°C.
Constant-Pressure Calorimetry • Constant-pressure, ΔH = qP and ΔE = qP + w • Assume no heat is lost to surroundings. • Usually exothermic (qrxn < 0) • Applications: • Heat transfer between objects • Reactions in aqueous solutions • Use specific heat of water (4.18 J/g·K). • Use mass (or moles) of solution.
Solution Calorimetry • heat lost by reaction = heat gained by solution −qrxn = qsoln qrxn =−(Cs,solnxmsolnxΔT) • Enthalpy of reaction (ΔHrxn) per mole ΔHrxn = qrxn/mol of specified reactant
A 19.6-g piece of metal was heated to 61.67°C. When the metal was placed into 26.7 g water, the temperature of the water increased from 25.00 to 35.00°C. Calculate specific heat of the metal.
A 15.0-g piece of nickel at 100.0°C is dropped into a coffee-cup calorimeter containing 55.0 g H2O at 23.0°C. What is the final temperature of the water and nickel after reaching thermal equilibrium? The specific heat capacity of nickel is 0.444 J/g·K and of water is 4.18 J/g·K.
In a coffee-cup calorimeter, 2.50 g of MgOwas reacted with 125 mL of 1.0 M HCl. The temperature increased by 9.6°C. Calculate the enthalpy of reaction per mole of MgO for the following reaction. Mg2+(aq) + H2O(l) →MgO(s) + 2 H+(aq)
Constant-Volume Calorimetry (p. 178) • Bomb calorimetry • No work is done (ΔV = 0), so ΔE = qV • Used for combustion reactions • The bomb components absorb the heat lost by the reaction. • Heat capacity of the bomb (Ccal) needed to calculate the heat of combustion (reaction) qrxn =−(CcalxΔT)
A 1.320-g sample of a new organic substance is combusted in a bomb calorimeter (Ccal = 8.74 kJ/K). The temperature of the bomb increased from 22.14°C to 26.82°C. What is the heat of combustion per gram of the substance?
Phase Changes (Fig. 11.20, p. 439) (or crystallization)
Phase Changes(p. 440) Endothermic → Cs= 1.84 J/g·K ΔHvap = 40.67 kJ/mol ← Exothermic Cs= 4.18 J/g·K ΔHfus = 6.01 kJ/mol Cs= 2.03 J/g·K
Heat transfer – phase changes • To calculate the quantity of heat transferred during a change of state: q = ΔHprocessx m or q = ΔHprocessxmol • No change in temperature, so no ΔT. • For a complete process, add together the heat transferred for each segment. • See Sample Exercise 11.3, p. 441.