Chapter 11. Thermochemistry. 1. Let s begin by previewing the chapter (Page 292). 2. We will partner read Pages

Transcription

1 Chapter 11 Thermochemistry 1. Let s begin by previewing the chapter (Page 292). 2. We will partner read Pages

2 The Flow of energy - heat Thermochemistry concerned with the heat changes that occur during chemical reactions. Energy the capacity to do work or supplying heat.

3 Chemical Potential Energy the energy stored within the structural units of chemical substances. Energy stored in bonds that is released when bonds are broken during chemical change Heat Represented by the letter q A type of energy that transfers from one object to another because of a temperature difference between them.

4 Endothermic vs. Exothermic System the part of the universe on which you focus your attention. Surroundings the part of the universe that includes everything else in the universe. Universe = system + surroundings Our focus is on how chemical reactions affect their surroundings.

5 Law of Conservation of Energy During any chemical or physical process, energy is neither created nor destroyed. All energy involved in a process must be accounted for as work, stored energy, or heat.

6 Endothermic Process A process that absorbs heat from the surroundings. The system gains heat from the surroundings, the thermometer will show a DECREASE in temperature in the surroundings. THE THERMOMETER IS IN THE SURROUNDINGS, WE STUDY THE SYSTEM, THE CHEMICAL BONDS.

7 Exothermic Process A process that releases heat to its surroundings. The system loses heat to the surroundings, the thermometer will RISE!

8 Heat Capacity and Specific Heat A calorie is the quantity of heat needed to raise the temperature of 1 gram of pure water by 1C. A Joule is the SI unit of heat and energy, named after the English physicist James Prescott Joule. One calorie = Joules

9 Heat capacity The amount of heat needed to increase the temperatur of an object exactly by one degree Celsius. A cup of water would have a much bigger heat capacity than a drop of water.

10 Specific Heat The amount of heat needed to raise the temperature of one gram of the substance by one degree Celsius. Water has a very high specific heat value: 1.00 cal/g-c or J/g-C.

11 Calculating heat, q We use this formula to calculate the amount of heat energy exchanged between the system and its surroundings: q = (mass in grams)x (specific heat) x (change in temperature) q = m c DT Let s try the example problems together.

12 Calorimetry The accurate and precise measurement of heat change for a chemical and physical process. It uses an instrument known as a calorimeter.

13 Coffee Cup Calorimeter for Constant Pressure Processes q = m c DT

14 Bomb Calorimeter for Constant Volume Combustions

15 Enthalpy The heat content of a system, denoted with the symbol H. When pressure is held constant, as in the case of our experiments in the lab, the Enthalpy of a reaction is equal to the heat. q = H at constant pressure Exothermic Reactions: DH = Negative Endothermic Reactions: DH = Positive

16 Relationship DH = q = m c DT is used to solve calorimetry problems. Let s try some example problems.

17 Thermochemical Equations An equation that includes the heat change associated with a chemical process. Let s try some examples together.

18 Enthalpy diagrams Show the exothermic or endothermic process as a function of heat content between reactants and products. Let s draw some together.

19 Heat in Changes of State Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

20 Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

21 Phase Diagrams The AB line is the liquid-vapor interface. It starts at the triple point (A), the point at which all three states are in equilibrium.

22 Phase Diagrams It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

23 Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

24 Phase Diagrams The AD line is the interface between liquid and solid. The melting point at each pressure can be found along this line.

25 Phase Diagrams Below A the substance cannot exist in the liquid state. Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

26 Phase Diagram of Water Note the high critical temperature and critical pressure: These are due to the strong van der Waals forces between water molecules.

27 Phase Diagram of Water The slope of the solid liquid line is negative. This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

28 Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures.

29 Phase Changes

30 Energy Changes Associated with Changes of State Heat of Fusion: Energy required to change a solid at its melting point to a liquid. DHfus

31 Energy Changes Associated with Changes of State Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas. DHvap Notice that the heat of vaporization is always larger than its heat of fusion.

32 Water The heat of fusion, or enthalpy of fusion, for ice is 6.01 kj/mol. The heat of vaporization, or enthalpy of vaporization, for water is 40.7 kj/mol. The heat of sublimation is the sum of heats of vaporization and fusion. For water = approx 47 kj/mol

33 Energy Changes Associated with Changes of State The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during the phase change.

34 Calculating DH for Temperature and Phase Changes Calculate the enthalpy change upon converting 1 mol of ice at -25 o C to water vapor (steam) at 125 o C under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.09 J/g-K, 4.18 J/g-K, and 1.84 J/g-K respectively. For H2O, DHfus = 6.01 kj/mol and DHvap = kj/mol.

35 Practice Exercise What is the enthalpy change during the process in which 100 g of water at 50.0 o C is cooled to ice at -30 o C?

36 Heats of Solution Heat changes can also occur when a solute dissolves in a solvent. The heat change caused by dissolution of one mole of substance is the Molar Heat of Solution, DHsol. Let s write some together, and work the practice problems.

37 Hess s Law of Heat Summation If you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction. Two rules If the reaction is reversed the sign of DH is changed If the reaction is multiplied or divided, so is DH

38 Using Hess s Law to Calculate DH The following information is known: C(s) + O2(g) CO2(g) DH1 = kj CO(g) + ½ O2 (g) CO2 (g) DH2 = kj Using these data, calculate the enthalpy for: C(s) + ½ O2(g) CO(g)

39 More Practice with Hess s Law Calculate DH for the reaction 2C(s) + H2(g) C2H2(g) Given the following chemical equations and their respective DH. C2H2(g) + 5/2O2 2CO2(g) + H2O (l) DH = kj C(s) + O2(g) CO2(g) H2(g) + ½ O2(g) H2O(l) DH = kj DH = kj

40 You Try It Calculate DH for the reaction NO(g) + O(g) NO2 (g) Given the following information: NO(g) +O3(g) NO2(g) + O2(g) DH = kJ O3(g) 3/2 O2(g) DH = kj O2(g) 2 O (g) DH = kj

41 Remember H is a state function, so for a particular set of reactants and products, DH is the same whether the reaction takes place in one step or in a series of steps.