Thermochemistry: Part 2 SCH4U

Heat and Physical Changes ΔHvap= - ΔHcond ΔHmelt = - Δhfre

Common Molar Enthalpies

Methods of Measuring ΔH Calorimetry (experimental) Hess’s Law: using Standard Enthalpy of Reaction (ΔHrxn) of a series of reaction steps (indirect method). Standard Enthalpy of Formation (ΔHf ) used with Hess’s Law (direct method) Bond Energies used with Hess’s Law

Specific Heat Capacity the amount of heat (Q) required to raise the temperature of one gram (1g) of the substance by one degree Celsius. ( 1C)

Heat (Q) absorbed or released Q=mcΔT Q= amount of heat transferred (Joules, J) m= mass of a substance (g) c= specific heat capacity (J/g• 0C) ΔT = Tfinal – Tinitial (change in temperature) ( 0C)

Heat Examples When a 1.25kg sample of water was heated in a kettle, its temperature increased from 16.40C to 98.90C. How much heat did the water absorb?

Heat Examples When a 1.25kg sample of water was heated in a kettle, its temperature increased from 16.40C to 98.90C. How much heat did the water absorb? Q= 1.25x103g x 4.18 J/g0C x (+82.50C ) = 432 093.75 J = 4.32x105 J

Heat Examples 2) How much heat must be added to 128.62 g of steam at 126.00C to increase its temperature to 189.50C? A solid substance has a mass of 250.0g. It is cooled by 25.000C and loses 4.937kJ of heat. What is the specific heat capacity of the substance?

Heat Examples 2) How much heat must be added to 128.62g of steam at 126.00C to increase its temperature to 189.50C? A solid substance has a mass of 250.0g. It is cooled by 25.000C and loses 4.937kJ of heat. What is the specific heat capacity of the substance? Q = 128.62 gx 2.02J/g0Cx (189.5-126.0) 0C = 1.65x104 J c= Q/m(∆T) c= -4.937x103J/(250.0 g (-25.000C)) c= 0.7899 J/g0C

Additional Examples How much heat is released when the temperature of 789g of liquid ammonia decreases from 82.70C to 25.00C? 2) On a warm day, how much solar energy does a 3.982kg piece of concrete absorb as heat if its temperature increases from 13.600C to 14.500C

Additional Examples 1) How much heat is released when the temperature of 789g of liquid ammonia decreases from 82.70C to 25.00C? Q= mc∆T = 789 g x 4.70 J/g0C x (25.0-82.7) = - 213968.91 J change to -2.14 x105 J 2) On a warm day, how much solar energy does a 3.982kg piece of concrete absorb as heat if its temperature increases from 13.600C to 14.500C? Q= 3982 g x 0.88 J/g0C x (14.50-13.60) = 3153.744 J change to 3.2 x103J

Measuring Q experimentally Calorimeter  a device that is used to measure enthalpy changes for chemical and physical reactions. Qreactions = - Q insulated system

Using Flame Calorimetry To Determine the Enthalpy of Combustion UNIT 3 Chapter 5: Energy Changes Section 5.2 Using Flame Calorimetry To Determine the Enthalpy of Combustion Flame calorimeters are flame-resistant and often made of metals cans A flame calorimeter is used for determining ΔHcomb absorbs a great deal of energy, which must be included in energy calculations is used for burning impure materials like food; ΔHcomb is reported in kJ/g Image source: MHR, Chemistry 12 © 2011. ISBN 07-106010-3; page 306

Using Bomb Calorimetry To Measure Enthalpy Changes during Combustion UNIT 3 Chapter 5: Energy Changes Section 5.2 Using Bomb Calorimetry To Measure Enthalpy Changes during Combustion A bomb calorimeter is used for more accurately determining ΔHcomb determines ΔHcomb at constant volume has a particular heat capacity, C Image source: MHR, Chemistry 12 © 2011. ISBN 07-106010-3; page 307 Q = CΔT is used for bomb calorimetry calculations Bomb calorimeters are much more sophisticated than flame calorimeters or simple calorimeters.

Using a Simple Calorimeter UNIT 3 Chapter 5: Energy Changes Section 5.2 Using a Simple Calorimeter Constant-Pressure Calorimeter No boundary between system and surroundings Reactants and products are system, and water in which they dissolve in, is part of the surroundings the change in temperature of the water is measured; the solution absorbs or releases energy the solution is dilute enough so that the specific heat capacity of water is used Thermometer Styrofoam cover cups Stirrer Image source: MHR, Chemistry 12 © 2011. ISBN 07-106010-3; page 301

Practice #1

Practice #2 240 g of water (initially at 20oC) are mixed with an unknown mass of iron (initially at 500oC). When thermal equilibrium is reached, the system has a temperature of 42oC. Find the mass of the iron.

Find the mass of the iron. 240 g of water (initially at 20oC) are mixed with an unknown mass of iron (initially at 500oC). When thermal equilibrium is reached, the system has a temperature of 42oC. Find the mass of the iron. Fe T = 500oC mass = ? grams T = 20oC mass = 240 g - LOSE heat = GAIN heat - [(Cp,Fe) (mass) (DT)] = (Cp,H2O) (mass) (DT) - [(0.4495 J/goC) (X g) (42oC - 500oC)] = (4.184 J/goC) (240 g) (42oC - 20oC)] Drop Units: - [(0.4495) (X) (-458)] = (4.184) (240 g) (22) 205.9 X = 22091 X = 107.3 g Fe Calorimetry Problems 2 question #5

Practice #3 A 97 g sample of gold at 785oC is dropped into 323 g of water, which has an initial temperature of 15oC. If gold has a specific heat of 0.129 J/goC, what is the final temperature of the mixture? Assume that the gold experiences no change in state of matter.

A 97 g sample of gold at 785oC is dropped into 323 g of water, which has an initial temperature of 15oC. If gold has a specific heat of 0.129 J/goC, what is the final temperature of the mixture? Assume that the gold experiences no change in state of matter. Au T = 785oC mass = 97 g T = 15oC mass = 323 g LOSE heat = GAIN heat - - [(Cp,Au) (mass) (DT)] = (Cp,H2O) (mass) (DT) Drop Units: - [(0.129 J/goC) (97 g) (Tf - 785oC)] = (4.184 J/goC) (323 g) (Tf - 15oC)] - [(12.5) (Tf - 785oC)] = (1.35 x 103) (Tf - 15oC)] -12.5 Tf + 9.82 x 103 = 1.35 x 103 Tf - 2.02 x 104 3 x 104 = 1.36 x 103 Tf Tf = 22.1oC Calorimetry Problems 2 question #8

Practice #4 If 59 g of water at 13oC are mixed with 87 g of water at 72oC, find the final temperature of the system.

If 59 g of water at 13oC are mixed with 87 g of water at 72oC, find the final temperature of the system. T = 13oC mass = 59 g T = 72oC mass = 87 g LOSE heat = GAIN heat - - [(Cp,H2O) (mass) (DT)] = (Cp,H2O) (mass) (DT) Drop Units: - [(4.184 J/goC) (87 g) (Tf - 72oC)] = (4.184 J/goC) (59 g) (Tf - 13oC) - [(364.0) (Tf - 72oC)] = (246.8) (Tf - 13oC) -364 Tf + 26208 = 246.8 Tf - 3208 29416 = 610.8 Tf Tf = 48.2oC Calorimetry Problems 2 question #9

Practice #5 A 322 g sample of lead (specific heat = 0.138 J/goC) is placed into 264 g of water at 25oC. If the system's final temperature is 46oC, what was the initial temperature of the lead?

A 322 g sample of lead (specific heat = 0 A 322 g sample of lead (specific heat = 0.138 J/goC) is placed into 264 g of water at 25oC. If the system's final temperature is 46oC, what was the initial temperature of the lead? T = ? oC Pb mass = 322 g Ti = 25oC mass = 264 g Pb Tf = 46oC - LOSE heat = GAIN heat - [(Cp,Pb) (mass) (DT)] = (Cp,H2O) (mass) (DT) Drop Units: - [(0.138 J/goC) (322 g) (46oC - Ti)] = (4.184 J/goC) (264 g) (46oC- 25oC)] - [(44.44) (46oC - Ti)] = (1104.6) (21oC)] - 2044 + 44.44 Ti = 23197 44.44 Ti = 25241 Ti = 568oC Calorimetry Problem

Molar Enthalpy Change ΔHn: Molar Enthalpy Change (ΔHn): enthalpy change associated with any change (physical, chemical or nuclear) in one mole of a substance; J/mol or kJ/mol ΔHrxn = nΔHn

Example: Ethanol is used to disinfect skin before receiving a flu shot. How does this feel on your skin? The enthalpy of vaporization (ΔHvap)of ethanol is 38.6 kJ/mol. What is the enthalpy change (ΔHrxn) of this reaction if 1.0g of ethanol is rubbed on your skin?