Energy. Different types of energy exist (heat, potential, kinetic, chemical, nuclear etc.)

Transcription

1 Change in Energy

2 Energy Different types of energy exist (heat, potential, kinetic, chemical, nuclear etc.) Heat - the energy transferred between objects that are at different temperatures. Unit of heat is joules J. According to the Law of Conservation of Energy, energy cannot be created or destroyed. It can only by transferred or transformed. The total amount of energy in a system remains constant.

3 Potential Energy Diagram

4 Potential Energy Diagram Potential Energy Diagrams show changes in energy during chemical reactions. CHEMICAL EQUATIONS A + B ---> C + D Reactant Products In chemical reactions, reactants react to form products with an associated release or absorption of energy.

5 Energy Potential Energy Diagram Potential Energy in KJ O C O N O O C O N O CO + NO 2 REACTANTS Reaction Coordinate O C O N O No unit, fancy of saying start to finish PRODUCTS

6 Energy Potential Energy Diagram Heat of Activated Complex CO + NO 2 C REACTANTS A Heat of Reactants D Reaction Coordinate PRODUCTS Heat of Products

7 Energy Potential Energy Diagram O C O N O O C O N O B Activation Energy Amount of energy needed to start the reaction CO + NO 2 REACTANTS ΔH O C O N O Reaction Coordinate ΔH: ENTHALPY PRODUCTS Net energy of the reaction

8 Potential Energy Potential Energy Diagrams A catalyst is a substance that changes the speed of reactions by changing the activation energy. Catalysts speed up chemical reaction by lowering the activation energy Catalysts are not consumed during the reaction and can be recovered Activation Energy with no catalyst Activation Energy with a catalyst Enthalpy of Reaction Reaction Coordinate

9 H products < H reactants DH < 0 H products > H reactants DH > 0 Enthalpy (H) is used to quantify the heat flow into or out of a system in a process that occurs at constant pressure. DH = H (products) H (reactants) DH = heat given off or absorbed during a reaction at constant pressure

10 Change in Heat Energy 1. Exothermic Reaction The change in heat ( H) is NEGATIVE. Heat is released into the surroundings 2. Endothermic Reaction The change in heat ( H) is POSITIVE. Heat is absorbed by the system

11 Click Below for the Video Lectures Activation Energy Reaction Coordinate Catalyst

12 Exothermic and Endothermic

13 Exothermic Reaction A + B C + D + ENERGY a) Heat is released as a product. b) This heat is released into the environment surrounding the reaction, causing the temperature of the surroundings to increase. c) The products have less energy than the reactants after release. d) Since the products have less energy stored in their chemical bonds, they are more stable than the reactants were. Burning paper is exothermic. The ash formed by the burning is not flammable and will not burn.

14 General Exothermic Reaction A + B C + 70 kj ΔH A = 60 kj ΔH B = 40 kj ΔH c = 30 kj 100 KJ Energy Released as product Exothermic Law of conservation of energy states that the total energy of an isolated system cannot change.

15 Exothermic Reaction C (g) + O 2 (g) CO 2 (g) ΔH = kj The minus sign in front of the enthalpy kj indicates that this reaction is exothermic, that energy was released. This energy can be placed on the products side, as follows: C (g) + O 2 (g) CO 2 (g) kj

16 Synthesis C (g) + O 2 (g) CO 2 (g) ΔH = kj Decomposition 1 mol of CO 2 If we were to synthesize 2.3 moles of CO 2, how many kj would be released? 2.3 mol CO kj 1 mol CO 2 = kj If ΔH synthesis = kj/mole, then Δ H decomposition = kj/mole

17 Energy Potential Energy Diagram - Exothermic A + B C + D + 40 kj A + B REACTANTS Reaction Coordinate Activation Energy = KJ ΔH = kj, released C + D PRODUCTS

18 Endothermic Reaction A + B + ENERGY C + D a) Heat is absorbed by the reactants. b) This heat is absorbed from the environment surrounding the reaction, and the temperature of the surroundings decreases. c) The products have more energy than the reactants after absorption. d) This energy is stored in the chemical bonds of the products

19 General Endothermic Reaction A + B + 50 kj C ΔH A = 40 kj ΔH B = 20 kj ΔH c = 110 kj 60 KJ Energy absorbed as product Endothermic Law of conservation of energy states that the total energy of an isolated system cannot change.

20 Energy Potential Energy Diagram - Endothermic 150 A + B + 40 kj C + D 70 C + D PRODUCTS 30 A + B REACTANTS ΔH = kj, Absorbed Activation Energy = KJ Reaction Coordinate

21 Click Below for the Video Lectures Exothermic and Endothermic

22 Calorimetry

23 Calorimetry q = m C ΔT units Joules (J) Grams (g) J/g o C o C What Each Variable Means q is the quantity of heat that is absorbed or released by a physical or chemical change m is the mass of water in the calorimeter cup that absorbs heat from the change or releases heat to the change. C is the specific heat of water, the rate at which water gains or loses heat if energy is absorbed or removed from it. T is the Temperature change of the water in the calorimeter cup as a result of the physical or chemical change.

24 C Specific Heat The amount of heat required to raise the temperature of one gram of substance by one degree Celsius. cal/g C J/g C (or J/gK) water aluminum copper silver gold

25 Use this equation on the back of periodic table Specific Heat of water. Use the one for correct unit! They are NOT equations!!!

26 1. Circle the numbers, underline what you are looking for. 2. Make a list of number you circled using variables. 3. Write down the formula 4. Derive the formula to isolate the variable you are looking for. 5. Plug in the numbers How many joules are absorbed by grams of water if the temperature is increased from 35.0 o C to 50.0 o C? m = T i= T f= q= C= g 35.0 o C 50.0 o C? 4.18 j/g o C q = mcδt q = (100.0 g)(4.18 j/g o C)(50.0 o C o C) q = 6270 J

27 1. Circle the numbers, underline what you are looking for. 2. Make a list of number you circled using variables. 3. Write down the formula 4. Derive the formula to isolate the variable you are looking for. 5. Plug in the numbers q = m = ΔT = C= 300 J 50g? 300. J of energy is absorbed by a 50. g sample of water in a calorimeter. How much will the temperature change by? q = mcδt ΔT = q / mc ΔT = 300 J/ (50g)(4.18 j/g o C) 4.18 j/g o C ΔT = = o C ΔT = = 1.4 o C

28 Checking for understanding 1. How many joules are required to raise the temperature of 100. grams of water from 20.0 o C to 30.0 o C? 2. How many grams of water can be heated from 20.0 o C to 75.0 o C using J?

29 We use an insulated device called a calorimeter to measure this heat transfer. A typical device is a coffee cup calorimeter. Calorimetry

30 To measure ΔH for a reaction 1.dissolve the reacting chemicals in known volumes of water 2.measure the initial temperatures of the solutions 3.mix the solutions Calorimetry 4.measure the final temperature of the mixed solution

31 Calorimetry The heat generated by the reactants is absorbed by the water. We know the mass of the water, mwater. We know the change in temperature, Twater. We also know that water has a specific heat of Cwater = 4.18 J/ C-g. We can calculate the heat of reaction by: qsys = H = qsurr = -mwater Cwater Twater

32 When 25.0 ml of water containing mol of HCl at 25.0 C is added to 25.0 ml of water containing mol of NaOH at 25.0 C in a coffee cup calorimeter, a reaction occurs. Calculate H (in kj) during this reaction if the highest temperature observed is 32.0 C. Assume the densities of the solutions are 1.00 g/ml. Knowns: Vfinal = VHCl + VNaOH = ( ) ml = 50.0 ml Dwater = 1.00 g/ml Twater = Tfinal Tinitial = 32.0 C 25.0 C = +7.0 C Cwater = 4.18 J/ C-g Calculation: mwater = 50.0 g H = m C T = (50.0 g)(4.18 J/ C-g)(7.0 C) = 1463 J = J

33 Phase Change Heating and Cooling Curve

34 Freeze Melt Sublime Condense Evaporate Solid Liquid Gas Deposit 34 4/3/2017

35 Phase Change During a phase change, there is a change in heat energy but no change in temperature. At the melting point, heat energy is being used to break down the crystalline lattice. At the boiling point, heat energy is being used to convert the liquid to a gas. During a phase change, heat is being used to change phase, not increase the kinetic energy of the particles. It is for this reason that Temperature remains constant in a phase change while heat energy changes.

36 Temperature LIQUID TO GAS BP EVAPORATING D E GAS SOLID to LIQUID LIQUID MP B MELTING C A SOLID Energy

37 Phase Change During phase change, temperature stays constant but energy increases.

38 Phase Change Energy needed to melt is called Latent Heat of Fusion ( ΔH fus ) Energy need to vaporize is called Latent Heat of Vaporization (ΔH vaporization )

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40 Triple point the point on a phase diagram at which the three states of matter: gas, liquid, and solid coexist Critical point the point on a phase diagram at which the substance is indistinguishable between liquid and gaseous states

41 q = mcδt q = mcδt q = mcδt

42 q = m Δ H fus q = m Δ H vap

43 How many joules does it take to melt 100. grams of water at its melting point? q = m Δ H fus = 100. grams X 334 J/gram = Joules

44 How many joules does it take to condense 50.0 grams of water at its boiling point? q = m Δ H vap = 50.0 grams X 2260 J/gram = Joules

45 Calculating Enthalpy Change The enthalpy change for a reaction is the enthalpy of the products minus the enthalpy of the reactants DH = H (products) H (reactants)

46 Hess s Law The overall enthalpy change in a reaction is equal to the sum of enthalpy changes for the individual steps in the process Ex: Elemental carbon can be found as graphite and diamond at 25 o C C(diamond) C(graphite) Takes millions of years to go from diamond to graphite Reaction is too slow to measure the heat change

47 Hess s Law 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 ΔH to give the final heat of reaction Example: Conversion of diamond to graphite C(diamond) C(graphite) C(s, diamond) + O 2 (g) CO 2 (g) ΔH = kJ C(s, graphite) + O 2 (g) CO 2 (g) ΔH = kJ REVERSE the above equation, so that we can show diamond being converted into graphite CO 2 (g) C(s, graphite) + O 2 (g) ΔH = 393.5kJ

48 Add the equations together! C(s, diamond) O CO 2 (g) (g) CO C(s, graphite) 2 (g) C(s, diamond) C(s, graphite) kj ΔH = -1.9 kj The overall ΔH is negative, so you can see this is an exothermic process! O 2 2 (g) DH DH kj

49 Standard Enthalpy (Heat) of Formation o DH f - standard heat of formation the change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states at 298 K (25 o C) Free elements at 298 K & 1 atm have an enthalpy of zero. Examples, H O N Cl Br I F Other ΔHº f values have been calculated ΔHº reaction can be calculated as follows: DH o reaction DH o f products DH o f reactants

50 Let s do an example! What is the standard heat of reaction (ΔHº) for the following equation: O 2 (g) + 2CO (g) 2CO 2 From your text: ΔHº f O 2 (g) = 0 (free element) ΔHº f CO (g) = kj/mol ΔHº f CO 2 (g) = kj/mol Calculate the ΔHº f of the reactants: 2 mol CO, 1 mol O 2 mol of CO 2 mol x kj/mol = kj 1 mol O 2 = 0 KJ (free element) TOTAL: 0 kj kj = -221 kj (reactants) Calculate the ΔHº f of the products: 2 mol CO 2 2 mol CO 2 2 x kj/mol = kj Use o o to find o DH DH DH the total! reaction f products f reactants ΔHº= kj (-221.0) kj ΔHº = kj

51 Try a Practice Problem 2NO (g) + O 2 (g) 2NO 2 (g) ΔHº f NO = ΔHº f O 2 = 0.0 ΔHº f NO 2 = (33.85) [2(90.37) + 0] = kj

52 Spontaneous Reactions Changes in enthalpy and entropy determine whether a process is spontaneous Spontaneous Process: any physical or chemical change that once begun, occurs without any outside intervention. Examples: Rusting of iron Melting of ice Vaporization of water

53 Entropy Entropy = a measure of the disorder or randomness of the particles that make-up a system The Second Law of Thermodynamics Spontaneous processes always proceed in a way that the entropy of the universe increases

54 Water Entropy Melting S solid < S liquid Evaporation S liquid < S vapor Dissolving S solute + S solvent < S solution

55 Energy Calculations Standard entropy (S o ) is measured at 25 o C. Units on S o = J/(K*mol) DS 0 rxn = S 0 products - S 0 reactants Predicting if S is positive or negative: 1. State changes which would cause S to increase? Decrease? 2. Dissolving a gas in a solvent? Ex O 2 (g) O 2 (aq) 3. Number of gaseous particles present? Ex 2SO 3 (g) 2SO 2 (g) + O 2 (g) 4. Dissolving? 5. Temperature Changes?

56 Energy Calculations Which is more important if they are working against each other: Enthalpy or Entropy? Gibbs Free Energy (ΔG): the enthalpy of the system minus the product of the temperature times the entropy of the system. DG o DH o - TDS o T in Kelvins In a spontaneous reaction, ΔG is always negative!

57 Reaction Spontaneity and the Signs of DH 0, DS 0, and DG 0 DH 0 DS 0 -TDS 0 DG 0 Description or or - Spontaneous at all T Nonspontaneous at all T Spontaneous at higher T; nonspontaneous at lower T Spontaneous at lower T; nonspontaneous at higher T

58 Practice Problems Using the values for ΔH and Sº, determine whether the reaction is spontaneous at 25ºC: C (s, graphite) + O 2 (g) CO 2 (g) Substance ΔH (kj/mol) Sº (J/Kxmol) C (s, graphite) O CO