Lecture outline: Chapter 5

Lecture outline: Chapter 5 1. The nature of energy Thermochemistryh 2. First law of thermodynamics 3. Enthalpies of reaction 4. Hess law 5. Enthalpies of formation 1

Chemical Reactivity (1) Does a chemical reaction occur? (2) How rapidly does the reaction occur? (3) How far does the reaction go towards product(s) when all change has apparently ceased? 2

Chemical Thermodynamics The study of energy and it s transformations Thermochemistry The energy changes that take place during chemical reactions 3

A working definition for energy as it relates to chemical reactions The capacity to do work or supply heat Energy = work + heat 4

Matter- The physical material of the universe. Anything that occupies space and has mass. Energy- Much more complicated with multiple levels l of definitions. iti Potential to perform work (Ch. 5) Potential energy (Ch. 5) Kinetic energy (Ch. 5) Heat (Ch. 5) Radiant energy (Ch. 6) Chemical energy (Ch. 5) Free energy (Ch. 19) Electrical energy (Ch. 20) Nuclear N l energy (Ch. 21) 5

Kinetic and potential energy K. E.: Energy of motion; P. E.: stored energy E = k 1 2 mv 2 6

Potential energy in water Credit: http://commons.wikimedia.org/wiki/image:hoover_dam_from_air.jpg, author, snakefisch 7

Chemical energy A form of potential energy where chemical bonds of a molecule act as the storage medium Chemicals release potential energy as heat, light, or work when they undergo a reaction to form more stable products with more stable bonds. 8

Thermal energy (Heat) Energy transferred from one object to another as a result of a temperature difference Temperature is a measure of the kinetic energy of molecular motion Temperature is measured by determining i the direction and magnitude of heat flow from one object to another 9

Units of energy? How much kinetic energy does a bicyclist with a weight of 110 lbs (mass =50 kg) and travelling at a speed of 22 miles/hour (10 m/s) possess? E = k 1 2 mv 2 10

The S.I. unit of energy is the Joule 1joule = Other units of energy: 1kg m 2 s 2 calorie: one calorie is the amount of energy required to raise the temperature of one gram of water by one degree celsius 1cal= 4.184 J Calorie: 1000 calories (1 kcal). The Calorie is what you see for food contents. 11

Potential energy in food 250 Cal = 250,000 000 cal = 1,045000 J = 1045 kj 12

Potential energy in food 70 Cal 70 kcal 290 kj 13

Law of conservation of mass: atoms are not created or destroyed during chemical reactions. They simply rearrange. Mass before = mass after Law of conservation of energy: energy is not created or destroyed during reactions. It is converted dfrom one form to another. energy before = energy after Yaaaa zees ist true on das macroscopic level..but remember, mass und energy can be interconverted according to das theory of relativity, und E = mc 2. nichts vergessen kinder!! Gesundheit! 14

Law of conservation of energy All energy is potential Mix of kinetic and Where does potential energy energy go? 15

The first law of thermodynamics The energy of the UNIVERSE is constant Yaaaa but remember E = mc 2 means energy and mass can be interconverted..so t maybe better to say combined E + mc 2 is constant in das universe. Ich bin very smart-very much smarter dan du!! Nichts vergessen das. 16

The system : the components of a chemical reaction of interest Reactants Products Internal energy Kinetic (thermal) energy Potential (chemical) energy The total internal energy (Σ Ek + Ep) of an isolated system is constant 17

System and surroundings ΔE = E final -E initial Energy out of system to surroundings: (-) sign Exothermic Energy into system from surroundings: (+) sign Endothermic 18

E initial (potential energy) Energy out of system to surroundings: (-) sign Exothermic E final ΔE = E final -E initial 19

2 H 2(g) + O 2(g) 2 H 2 O (g) + 483.6 kj 2 H 2(g) + O 2(g) E initial (chemical energy) Energy out of system to surroundings: (-) sign Exothermic 2 H 2 O (g) E final ΔE = E final -E initial 20

Consider combustion reactions, incorporating energy in the balanced chemical equation 2 H 2(g) + O 2(g) 2 H 2 O (g) + 483.6 kj CH 4(g) + 2 O 2(g) 2 H 2 O (l) + CO 2(g) + 890 kj 21

What determines how much internal energy a system has? Chemical identity Sample size Temperature Pressure Physical state (s, l, g) The amount of an internal energy a system has is dependent on it s present condition (state). It is not dependent on the system s past history 22

State function A function or property whose value depends only on the present state of the system, not the path used to arrive at that condition Age Weight Location For any state function, the overall change is zero if the system returns to its original condition 23

State functions resulting from a change in a system A function or property whose value depends only on the present state of the system, not the path used to arrive at that condition Age Weight Location For any state function, the overall change is zero if the system returns to its original condition 24

Box Elder peak Wellsville cone Mendon peak Rattlesnake canyon E initial ΔE = E final E initial ~ 0 E final Deep canyon 25

State functions and energy changes Internal energy of a nonreacting system is a state function The energy change that occurs during a chemical reaction is a state function ΔE = E final -E initial The contributions of work and heat to ΔE will vary depending on the experimental conditions ΔE = q + w 2 H 2(g) + O 2(g) 2 H 2 O (g) + 483.6 kj 2(g) 2(g) 2 (g) 26

Electron flow with no conservation of energy e - e - http://commons.wikimedia.org/wiki/file:rjukanfossen07.jpg, licensed under the Creative Commons Attribution-Share Alike 3.0 Unported license. Attribution: MB. Gustavsen (user:hau-maggus) 27

Electron flow with conservation of energy e - e - Heggmoen kraftverk in Bodø, Norway. licensed under the Creative Commons Attribution 2.5 Generic license., Author: Røed 28

Energy fundamentals: common forms of energy Energy associated with physical objects Potential t energy: e E mass, gravity and height Kinetic energy: E mass and velocity 2 Energy associated with atoms and molecules Chemical potential energy: stored in bonds Thermal energy: kinetic energy of molecules Manifestations of energy changes in physical processes and chemical reactions: Heat: radiant energy Work: some form of movement 29

Energy fundamentals: energy changes Energy changes are referenced relative to the system The internal energy of the system before the reaction is E initial The internal energy of the system after the reaction is E final The change in energy (ΔE) from the perspective of the system is: ΔE = E final E initialiti If ΔE is negative, the system lost energy to the surroundings If ΔE is positive, the system gained energy from the surroundings E and ΔE are extensive properties ΔE = (-) n kj ΔE = (+) n kj 30

Energy is a state function A function or property whose value depends only on the present state of the system, not the path used to arrive at that condition E initial is a defined value for a given system E final is a defined d value for a given system ΔE = E final E initial ΔE is a state function ΔE = q + w q and w are not state functions, but their sum must equal the state function ΔE For any state function, the overall change is zero if the system returns to its original condition 31

Simplifying the components of energy ΔE = E final E initial = q + w A difference in pressure is necessary for work to occur in a chemical reaction Conduct the chemical reaction in a vessel open to the atmosphere so no pressure change can occur ΔE p = E final E initial = q p + 0 Enthalpy = ΔH = q p -ΔH: heat flows out of system (exothermic) +ΔH: heat flows into system (endothermic) 32

Stoichiometric relationships between coefficients and ΔH in a balanced chemical equation 2 H 2(g) + O 2(g) 2 H 2 O (g) + 483.6 kj 2 H 2(g) + O 2(g) H initial 2HO 2 (g) Hfinal Enthalpy (heat) out of system to surroundings: (-) sign Exothermic 33

What is the enthalpy change associated with the combustion in air of a balloon containing 3 liters (0.13 mols) of hydrogen gas? 2 H 2(g) + O 2(g) 2 H 2 O (g) + 483.6 kj

ΔH forward reaction = - ΔH reverse reaction 2 H 2(g) + O 2(g) 2 H 2(g) + O 2(g) H initial H initial H out H in 2 H 2 O (g) H final 2 H 2 O (g) H final 35

ΔH depends on the physical states of reactants and products 36

Determination of ΔH for a reaction: Determine experimentally: calorimetry Calculate from enthalpy data for other reactions 37

Experimental determination of ΔH for a reaction q Heat capacity = Δ T 38

Specific heat q specific heat = mass ΔT What is the specific heat of water, if it takes 2090 J of energy to raise the temperature of 100 ml of water from 25 C to 30 C? (d H = 1.00 g/ml) 2O 39

Experimental determination of ΔH for a reaction A + B C + D The system: A, B, C, D The surroundings: the water inside id the calorimeter. Exothermic reaction: heat flows from system to water (T of water goes up) endothermic reaction: heat flows from water to system (T of water goes down By knowing the mass of the water, the temperature t change, and the specific heat of the surroundings (water), you can determine ΔH for the reaction 40

Experimental determination of ΔH for a reaction Carry out reaction Measure temperature t change of reaction Know specific heat of calorimeter Know mass of heat absorbing material in the calorimeter 41

Determine ΔH for the neutralization of H 2 SO 4 with NaOH, given the following data: 50 ml 1.0 M NaOH 25 ml 1.0 M H 2 SO 4 T i = 25 C T f = 33.9 C SH solution =418J/gK 4.18 J/g K D solution = 1.00 g/ml specific heat = q mass ΔT Carry out reaction Measure temperature change of reaction Know specific heat of calorimeter Know mass of heat absorbing material in the calorimeter 42

Determine ΔH for the neutralization of H 2 SO 4 with NaOH, given the following data: q 50 ml 1.0 M NaOH 25 ml 1.0 M H 2 SO 4 T i =25 C T f = 33.9 C SH solution = 4.18 J/g K D solution = 1.00 g/ml specific c heat = mass ΔT 43

ΔH values should be reported relative to the stoichiometry of the balanced chemical equation 2NaOH (aq) + H 2 SO 4(aq) Na 2 SO 4(aq) + 2H 2 O (l) 44

50.0 ml of 0.100 M AgNO 3 are mixed with 50.0 ml of 0.100 M HCl in a calorimeter. The temperature increases from 22.30 C to 23.11 C. What reaction occurred? What is ΔH for the reaction? 45

Energy fundamentals: energy changes Energy changes are referenced relative to the system The internal energy of the system before the reaction is E initial The internal energy of the system after the reaction is E final The change in energy (ΔE) from the perspective of the system is: ΔE = E final E initialiti If ΔE is negative, the system lost energy to the surroundings If ΔE is positive, the system gained energy from the surroundings E and ΔE are extensive properties ΔE = (-) n kj ΔE = (+) n kj 46

50.0 ml of 0.100 M AgNO 3 are mixed with 50.0 ml of 0.100 M HCl in a calorimeter. The temperature increases from 22.30 C to 23.11 C. What reaction occurred? What is ΔH for the reaction? 47

Hess Law ΔH for a given reaction is the sum of the ΔH values for each individual step in the reaction, or the sum of ΔH values for a sequence of reactions that add together to give the reaction of interest 48

Compare the H values for the combustion of hydrogen gas to produce either water vapor or liquid water 49

Determine H for the reaction: C (s) + 2H 2(g) CH 4(g) Given: Gven C C(s) +O 2(g) CO 2(g) ΔH = -393.5 kj H 2(g) + ½ O 2(g) H 2 O (l) ΔH = -285.8 kj CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (l) ΔH = -890.3 kj 50

Determine H for the reaction: C (s) + 2H 2(g) CH 4(g) Given: Gven C C(s) +O 2(g) CO 2(g) ΔH = -393.5 kj H 2(g) + ½ O 2(g) H 2 O (l) ΔH = -285.8 kj CH 4(g) + 2O 2(g) CO 2(g) + 2H 2 O (l) ΔH = -890.3 kj 51

Enthalpy of formation (ΔH f ) The change in enthalpy (heat input or heat output) associated with the formation of one mol of a compound from its constituent elements in their standard state forms Standard state 25 C and one atmosphere of pressure The standard state form of an element is the form most stable at 25 C and one atmosphere of pressure 52

Standard state 25 C and one atmosphere of pressure Element H O N C Cl I Std. state form? H 2(g) (diatomic molecule) O 2(g) (diatomic molecule) N 2(g) (diatomic molecule) C (s)graphite (empirical) Cl 2(g) (diatomic molecule) I 2(s) (diatomic molecule) Hg Hg (l) (metallic solid) S S 8(s) (octatomic molecule) 53

Tabulated enthalpies of formation (ΔH f ) Standard enthalpies of formation for selected compounds at 298.15 K (25 ) Substance H f (kj/mol) Substance H f (kj/mol) CO 2 (g) -393.5 NH 3 (g) -46.19 CH 4 (g) -74.8 NaCl(s) -410.9 C 6 H 12 O 6 (s) -1273 C(s) diamond 1.88 CH 3 OH(l) -238.6 H 2 O(g) -241.8 C 2 H 2 (g) 226.7 H 2 O(l) -285.85 ΔH f for an element in it s standard state = 0 54

Using a table of enthalpies of formation to calculate the enthalpy change (ΔH ) for a chemical reaction of interest ΔH rxn = Σ nδh f (products) - Σ mδh f (reactants) Appropriate stoichiometric coefficients in the balanced chemical equation Standard enthalpies of formation for selected compounds at 298.15 K (25 ) Substance H f (kj/mol) Substance H f (kj/mol) CO 2 (g) -393.5 NH 3 (g) -46.19 CH 4 (g) -74.8 NaCl(s) -410.9 C 6 H 12 O 6 (s) -1273 C(s) diamond 1.88 CH 3 OH(l) -238.6 H 2 O(g) -241.8 C 2 H 2 (g) 226.7 H 2 O(l) -285.85 55

What enthalpy change is associated with the production of glucose and oxygen from carbon dioxide and water (the reaction of photosynthesis)? 6CO 2(g) + 6H 2 O (l) C 6 H 12 O 6(s) + 6O 2(g) 56

What enthalpy change is associated with the production of glucose and oxygen from carbon dioxide and water (the reaction of photosynthesis)? 6CO 2(g) + 6H 2 O (l) C 6 H 12 O 6(s) + 6O 2(g) Standard enthalpies of formation for selected compounds at 298.15 K (25 ) Substance H f (kj/mol) Substance H f (kj/mol) CO 2 (g) -393.5 NH 3 (g) -46.19 CH 4 (g) -74.8 NaCl(s) -410.9 C 6H 12O 6 6( (s) -1273 C(s) () diamond 1.88 CH 3 OH(l) -238.6 H 2 O(g) -241.8 C 2 H 2 (g) 226.7 H 2 O(l) -285.85 ΔH rxn = Σ nδh f (products) - Σ mδh f (reactants) 57

What enthalpy change is associated with the production of glucose and oxygen from carbon dioxide and water (the reaction of photosynthesis)? 6CO 2(g) + 6H 2 O (l) C 6 H 12 O 6(s) + 6O 2(g) Standard enthalpies of formation for selected compounds at 298.15 K (25 ) Substance H f (kj/mol) Substance H f (kj/mol) CO 2 (g) -393.5 NH 3 (g) -46.19 CH 4 (g) -74.8 NaCl(s) -410.9 C 6 H 12 O 6 (s) -1273 C(s) diamond 1.88 ΔH rxn = Σ nδh f (products) - Σ mδh f (reactants) CH3 OH(l) -238.6 H 2 O(g) -241.8 C 2 H 2 (g) 226.7 H 2 O(l) -285.85 58