Unit 14 Thermochemistry

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1 Unit 14 Thermochemistry Name May 5 6 Unit 13 Acids and Bases Test Intro to Thermochemistry Videos (p.2-3) HW: p Thermochemistry Interpret graphs Heat of reaction & Specific Heat Heat of reaction (p.6-8) (p. 11 and 12) Heat of formation (p ) Lab HW: p (p. 13) (HW: p. 16) HW: Finish p.12* (HW: p. 14) Quiz Heat of fusion & Phase changes (p. 18) (HW: p. 19) Review Due Calorimetry (p ) Cheeto Lab Unit 14 Thermodynamics Test 1

2 Energy and Chemistry: Crash Course Chemistry #17 Everything is made of chemicals except sound, heat, light, etc. All of those things (including chemicals, sound, heat, light, etc.) is. Energy can be found in several forms including mass,, and among many other forms. is energy contained within a system because of its position. The First Law of Thermodynamics is also known as the. It states that energy cannot be created and it cannot be destroyed. Thermodynamics is the study of heat, energy, and the ability of energy to do work. Energy is the capacity to or. Heat is an energy transfer. The symbol for heat is. We can split the universe into 2 parts: the and the. Positive ΔE means that work is done on the system or heat is transferred the system. Chemical energy stored in bonds is a kind of. A reaction in which heat flows out of the system is considered an reaction. A reaction in which heat flows into the system is considered an reaction. 2

3 Enthalpy: Crash Course Chemistry #18 yaxl5i Any bond between two atoms contains. Heat is the energy transferred to the motion of and molecules. The amount of heat or work done depends on the you take. The change in energy is the same in either case. The change of energy is independent of the pathway. In chemistry, we call that a. The only things that matter for state functions are the starting state and the ending state. We are interested in energy being transferred in or out of system because of reactions. In many cases, we are only interested in the part of energy. Enthalpy is represented by the letter. ΔH = (Change in enthalpy equals heat gained or lost in the reaction) When a reaction takes place and changes, that heat is transferring actual chemical bonds. We measure enthalpy with c. Germain Henri Hess was a chemist who has a law named after him:. Hess s Law says that the total enthalpy for a reaction doesn t depend on the pathway it takes but only depends on its and states. So as long as you start with the same reactants and end with the same products, the enthalpy change is. Standard state is just a set of criteria so that chemists can study stuff under the same. Standard state is at T = 278K and P = 1 atm Standard enthalpy of formation (ΔH f ) is the lost or gained when mole of a compound is formed from its constituent elements. The enthalpy change for a reaction is to the sum of the enthalpy of formation of all the products the sum of the enthalpy of formation of all the reactants. Σ means the of 3

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6 Thermochemistry Thermochemistry focuses on the changes that occur during a reaction. Heat ( ) that transfers from one object to another because of a between them. The standard international unit (SI unit) of heat is the ( ). **Note: Heat always flows from a warmer object to a cooler object. Enthalpy ( ) the of a system at a constant. Energy the for doing or supplying. Kinetic energy Potential energy o due to o o due to o stored in the between atoms in Law of Conservation of Energy Energy is neither nor ; but it can be from one form to another. chemical reactions involve a or of heat. Exothermic process to its surroundings (Temperature ) Endothermic process from its surroundings (Temperature ) The following graph shows the energy change during the reaction A + B C. energy: the amount of energy which a has to have in order for a chemical change to take place. 1) Does the product have more or less energy than the reactants?, 2) Would the reaction be endothermic or exothermic?. Heat is. 6

7 A thermochemical equation is a for a reaction that also includes. *Thermochemical equations must be balanced. Heat of Reaction ( ) = the of heat for a reaction under constant. Also known as change in enthalpy. Direction of heat flow Sign of H Reaction Type Heat flows out of system (heat on the side) Heat flows into the system (heat on the side) Practice Given the following balanced thermochemical reactions, answer the questions. 1: 4Fe (s) + 3O2 (g) 2Fe2O3 (s) kj Does this reaction release heat or absorb heat? How much? What does kj mean? (measurement of ) Endothermic or exothermic? 2: C (s) + 2S (s) kj CS2 (l) Is heat released or absorbed in this chemical reaction? How much? Endothermic or exothermic? 7

8 Thermochemical equations treat heat change ( ) just like any or Chemistry problems involving H are similar to problems; they depend on the number of of reactants and products involved, for example: CaO(s) + H2O(l) Ca(OH)2 (s) kj and 2 CaO(s) + 2 H2O(l) 2 Ca(OH)2 (s) kj You must multiple the heat of reaction by the number of moles. Practice H2 (g) + F2 (g) 2HF(g) H = -536 kj Calculate the heat change (in kj) for the conversion of 2 moles of H2 gas to HF gas at constant pressure. 2Al (s) + Fe2O3 (s) Al2O3 (s) + 2Fe (s) H = -851 kj Calculate the heat change (in kj) for the thermite reaction of 4 moles of Fe2O3 into Al2O3 at constant pressure K2O (s) + H2O (l) 2KOH (aq) H = 215 kj What is the heat change for the above reaction, at constant temperature if you begin with ½ mole of K2O? 8

9 Homework: 1. What is the law of conservation of energy? 2. Explain how kinetic and potential energy are involved in a chemical reaction. 3. Explain from where the heat in an exothermic reaction comes. (Hint: look at the answers for the last two questions. 4. Classify these processes as exothermic or endothermic. (Think about whether the object is warming up/accepting heat or cooling down/releasing heat.) A. Burning alcohol ( exothermic / endothermic ) B. Baking a potato ( exothermic / endothermic ) C. Combustion of gasoline ( exothermic / endothermic ) 5. Indicate whether ΔH is positive (+) or negative (-) for the following: A. N 2 + O kcal 2 NO ΔH = B. 2 C 2H O 2 4 CO H 2O kj ΔH = C. the reactants contain more enthalpy than the products ΔH = D. the products contain more enthalpy than the reactants ΔH = E. the surroundings lose heat as a reaction occurs ΔH = F. the temperature increases as a reaction occurs ΔH = 9

10 Determine if the following reactions are endothermic or exothermic. 6. N 2 (g) + O 2 (g) kj 2 NO (g) ( exothermic / endothermic ) 7. 2 C 2H 6 (g) + O 2 (g) 4 CO 2 (g) + 6 H 2O (l) kj ( exothermic / endothermic ) 8. C 3H 8 (g) C 3H 8 (l) kj ( exothermic / endothermic ) Each of the following equations (with the ΔH provided) has been rewritten. Find the ΔH for the new equation. 9. Given: CuO (s) Cu (s) + ½ O 2 (g) ΔH = 37.1 kj 2 Cu (s) + O 2 (g) 2 CuO (s) ΔH = 10. Given: C 2H 2 (g) + 5/2 O 2 (g) 2 CO 2 (g) + H 2O (l) kj ΔH = 2 C 2H 2 (g) + 5 O 2 (g) 4 CO 2 (g) + 2 H 2O (l) ΔH = 11. Given: H 2O (g) H 2O (l) ΔH = kj 3 H 2O (g) 3 H 2O (l) ΔH = 12. Rewrite the following equations by expressing the energy change as a term in the equation: a) H 2O (g) H 2O (l) ΔH = kj b) 4 Al + 3 O 2 2 Al 2O 3 ΔH = kj c) 2 H 2SO 4 2 SO H 2O + O 2 ΔH = kj d) H 2O (g) H 2O (l) ΔH = kj 10

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13 Standard Heat of Formation ( ) The heat of the reaction can be calculated from standard heats of formation (see table at the end of the packet) when it is difficult to measure the heat change for a reaction. The standard heat of formation of a compound is 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 25 C. All free elements have a Hf 0 of Formula: Guided Practice: 1) What is the standard heat of reaction for the reaction 2CO(g) + O2(g) 2CO2(g) Hf 0 O2(g) = Hf 0 CO(g) = Hf 0 CO2 (g) = a. H = b. This reaction is endothermic/exothermic. c. Rewrite the equation with the change in enthalpy represented on the correct side of the thermochemical equation. 2) Calculate the standard heat of reaction for the reaction: CaCO3(s) CaO(s) + CO2(g) Hf 0 CaCO3 = Hf 0 CaO = Hf 0 CO2 = a. H = b. This reaction is endothermic/exothermic. c. Rewrite the equation with the change in enthalpy represented on the correct side of the thermochemical equation. 13

14 Standard Heats of Formation Homework: ( Use values from the chart on last page of packet!!!) 1. Use the Standard heats of formation to calculate the standard heats of the reaction for the reaction. Br2(g) Br2(l) H f 0 Br 2(g) = kj/mol H f 0 Br 2(l) = 0.0 kj/mol This reaction is endothermic/exothermic. 2. Use the standard heats of formation to calculate the standard heats of reaction for the reaction: 2NO(g) + O2(g) 2NO2(g) H f 0 NO (g) = H f 0 O 2(g) = H f 0 NO 2(g) = 3. Calcium carbonate decomposes at high temperature to form carbon dioxide and calcium oxide: Using the heat of formation table, determine the heat of reaction. CaCO3 CO2 + CaO Is this reaction endothermic or exothermic? 4. Determine the heat of reaction when carbon tetrachloride is formed by reacting chlorine with methane. CH4 + 2 Cl2 CCl4 + 2 H2 Is this process endothermic or exothermic? 5. When potassium chloride reacts with oxygen under the right conditions, potassium chlorate is formed. Determine the heat of reaction. 2 KCl + 3 O2 2KClO3 Is this process endothermic or exothermic? 6. The following is known as the thermite reaction: 2Al(s) + Fe 2O 3(s) Al 2O 3(s) + 2Fe(s) a. Find the heat of reaction (ΔH ) for the thermite reaction. b. The thermite reaction is highly exothermic. Does your answer support this piece of information? 14

15 Refresher Specific Heat The specific heat capacity ( ), or simply the of a substance is the amount of heat it takes to raise the temperature of of the substance. Formula: Cp = Q = heat energy in Joules, m = mass, ΔT = change in temperature Therefore the units for specific heat are J/(g C) or cal/(g C) ΔT = T final T initial To solve for heat energy, rearrange the equation for q: Q= Some substances, such as metals, have specific heats. This means it doesn t take a lot of energy to cause a change. Other substances, such as water, have high specific heats. It takes more energy to cause a temperature change. On a summer day, why does the concrete desk around a swimming pool become hot, while the water stays much cooler? Refresher Part 2: Phase Changes Phase changes are changes in the of a substance. Phase changes can be exothermic or endothermic processes. Name of phase change States of matter involved? Exothermic or endothermic? Melting Solid to liquid Endothermic Freezing Boiling Condensing Sublimation* Deposition** *Sublimation: solid to gas phase change without passing through the liquid phase (Examples: dry ice, solid air fresheners, mothballs, Shrinking ice cubes) **Deposition: gas to solid phase change without passing through the liquid phase (Example: frost on a windshield water vapor in the air crystallizes on the cold glass) Vaporization, evaporation, and boiling: What s the difference? is the process by which a liquid changes to a gas. Evaporation and boiling are two types of vaporization. is vaporization only at the of the liquid, at temperatures below the boiling point. Rate of evaporation depends on temperature, and also on intermolecular forces. A use of evaporation in our bodies is perspiration. How does perspiration help your body cool? How does a fan or a cool breeze help you cool even more? 15

16 During boiling, vaporization occurs the liquid. The bubbles you see are bubbles of forming from the liquid (it s not ). The pressure inside the bubbles equals atmospheric pressure. The vapor then escapes into the atmosphere. The boiling point of a liquid at a pressure of 1 atmosphere (sea level) is called the. For water, that is. The boiling point of a liquid changes as the external pressure changes. If the external pressure above the liquid is higher than normal, the liquid boils at a temp. If the external pressure above the liquid is lower than normal, the liquid boils at a temp. Why don t foods cook the same at high altitudes? Example: making spaghetti Phase changes can be represented on a. The heating curve below is for water. Show where each state of matter exists, label the phase changes, provide values for the temperatures at each phase change and label the direction arrows as endo- or exothermic. Then write the formula used to show change in heat for each portion of the graph. Assume standard pressure (1 atm). H f = heat of fusion (heat per mass needed to melt a substance) = 334 J/g (for water) H v=heat of vaporization (heat per mass needed to vaporize a substance) = 2260 J/g (for water) Phase changes always occur at temperature. For example, the freezing/melting point of water is 0 C. If the temperature is exactly 0 C, there will be a mixture of liquid water and ice present. Because we have both states of matter (solid and liquid) present at the freezing/melting point, we sat the solid is in with the liquid. If you add heat at this point, you can melt all the ice and then heat the liquid water further if you want. If you take away heat (cool it), you can freeze the rest of the liquid and then cool the ice further if you want. 16

17 Specific Heat: Thermal Energy Calculations (use specific heat table in the back of the packet) Guided Practice: 1. How much thermal energy (J) is needed to raise the temperature of 50.0g H 2O from 14 C to 83 C? 2. How much thermal energy (J) must be added to 50.0 kg Al at -5 C to raise its temperature to 125 C? 3. A 500g block of metal absorbs 5016 Joules of thermal energy when its temperature changes from 20 C to 30 C. Calculate the specific heat of the metal. Homework: 1. A copper wire has a mass of 165 grams. An electric current runs through the wire for a short time and its temperature rises from 21 C to 39 C. What quantity of thermal energy has the copper absorbed? 2. How much thermal energy is absorbed by 250 g H 2O when it is heated from 10 C to 85 C? 3. A 38kg block of metal is heated from -26 C to 180 C. It absorbs 1,957,000 J of thermal energy during the heating. What is the specific heat of this metal? 4. A 200 g glass at room temperature, 20 C, is plunged into a hot dishwasher at 80 C. If the temperature of the glass reaches that of the dishwasher, how much thermal energy does the glass absorb? 5. Five kilograms of ice cubes are moved from the freezer of a refrigerator into a deep freeze. The refrigerator s freezing compartment is kept at -4 C. The deep freeze is kept at -17 C. How much thermal energy does the deep freeze s cooling system remove from the ice cubes? 17

18 ENTHALPY OF PHASE CHANGE WS Use the following information to solve the phase change problems: Specific heat of water = 4.18 J/g C Specific heat of ice = 2.03 J/ g C Specific heat of steam = 1.97 J/g C H fusion = 334 J/g H vaporization = 2259 J/g Guided Practice: 1. How much heat is required to melt 233 grams of ice into water, from -15 C to room temperature (25 C)? Ice: Melting: Water: Total: 2. How much heat is required to change 32.5 grams water into steam, from room temperature (25 C) to 115 C? Water: Vaporization: Steam: Total: 18

19 Homework: 3. How much heat is needed to melt 1.43 grams of ice into water from C to 84.3 C? Ice: Melting: Water: Total: 4. How much heat is needed to convert grams water into steam, from 32.5 C to 112 C? 5. How much heat do you need to add to 3.22 grams H2O to raise the temperature from -23 C to 152 C? 6. How much heat is needed to raise the temperature of 199 grams H2O from C to 154 C? 19

20 Thermochemistry Test Review Energy (in general) 1) There are two types of energy: (energy due to motion) & (energy due to position). Thermal energy (heat) is kinetic/potential (circle one). Chemical energy (energy stored in bonds) is kinetic/potential (circle one). 2) Define the Law of Conservation of Mass: 3) A toaster is powered with 1500 J of electric energy. When on, it converts 1000 J to thermal energy, 300 J to light energy, and the remaining portion to sound energy. How much sound energy is produced? Enthalpy 4) In your body, blood is at a higher temperature than any other body tissue. So when your hands are cold, how does your body warm them up? because heat always travels from to objects. 5) What is the heat change for the above reaction, at constant pressure if you begin with g of K2O? Is this endothermic or exothermic? K2O (s) + H2O (l) 2KOH (aq) H = 215 kj 6) What is the heat of reaction for the following: C 3H 8 (g) C 3H 8 (l) kj ΔH = Energy Diagrams 7) The energy diagram below represents the equation: 2NO(g) + O 2(g) 2NO 2(g). a) Label the reactants, products, activation energy b) Is the reaction endothermic or exothermic? c) Explain the energy transfer (Which form of energy existed first, and to which kind of energy did it change?) d) Would ΔH be positive or negative? 20

21 8) Breaking bonds requires/releases (chose one) energy and would therefore be an exothermic/endothermic (choose one) process. 9) Use the diagram below to answer the following questions: a) What is value of the change in enthalpy? What letter represents this change? b) Is this process endothermic or exothermic? c) Which letter represents the heat required to break bonds? What is this called? d) Which letter represents the activation energy of the reverse reaction? e) Which letter represents the enthalpy of the reactants? Which letter represents the enthalpy of the products? Heat of Formation For questions 10-12: (a) Calculate the standard enthalpy of the reaction for the following reactions using the standard enthalpies of formation chart in the back of the packet or the other information given in the problem, (b) classify each as either endothermic or exothermic, and (c) determine which energy diagram best describes the reaction: Energy Diagram 1: Energy Diagram 2: Energy Diagram 3: 10) 4NH3(g) + 5O2(g) 6H2O(g) + 4NO(g) a) b) 21

22 c) Diagram # 11) SiO2(g) + 3C (s) kJ SiC (s) + 2CO (g) a) b) c) Diagram # 12) Magnesium reacts with hydrochloric acid in a single replacement reaction. Balanced Equation: a) b) c) Diagram # Specific Heat 13) Define specific heat capacity: 14) The that the specific heat capacity, the more resistant the object is to a change in temperature (i.e. it requires more energy to change the temperature). 15) Determine the initial temperature of an 84.3 gram sample of water after 14,500 J of heat is applied. The final temperature of the water sample is 100 ºC. The specific heat of water is 4.18 J/gºC. 22

23 16) A sample of iron is placed in a hot water bath with an initial temperature of ºC causing 3.50 x 10 3 J of heat to be transferred. After several minutes the temperature of the water was recorded at 22.0 ºC. Calculate the mass of the copper sample. Metal Copper Aluminum Specific Heat J/g x C J/g x C 17) Using the specific heat data from above, will aluminum or copper reach the higher temperature assuming they gain the same amount of heat? Phase Changes 18) How much heat is required to change 32.5 grams ice into steam, from -10 C to 115 C? Specific heat of water = 4.18 J/g C Specific heat of ice = 2.03 J/ g C Specific heat of steam = 1.97 J/g C H fusion = 334 J/g H vaporization = 2259 J/g Calorimetry 19) A solution s temperature increases as the frequency of collisions between reactants increases. Temperature is a measure of the molecules energy. 20) According to the law of conservation of energy, heat lost by the reaction in a calorimeter must the heat gained/lost (chose one) by the water. 21) Students conduct an experiment where a reaction occurs in a calorimeter. Calculate the heat released in Joules to the nearest whole number. The specific heat capacity of water is J/g⁰C. Mass Initial Final (g) Temperature (ºC) Temperature (ºC)

24 Table of Specific Heats Substance Specific Heat (C) (J/gK) H 2O (l) Al (s) Cu (s) H 2O (s) 2.05 Glass (s) 0.84 Table of Heats of Formation 24

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