HT2 Thermochemistry PDF

# Thermal Energy Unit ## HT2 | Thermochemistry **HT2.1 | Thermochemistry** is the study of energy changes that occur during **chemical reactions.** The energy stored in the chemical bonds of a substance is called **chemical potential energy (PE).** For example, when you by gasoline for your car, y...

# Thermal Energy Unit ## HT2 | Thermochemistry **HT2.1 | Thermochemistry** is the study of energy changes that occur during **chemical reactions.** The energy stored in the chemical bonds of a substance is called **chemical potential energy (PE).** For example, when you by gasoline for your car, you are actually buying the chemical potential energy it contains. The controlled explosions of the gasoline transform the potential energy into **heat (thermal energy)** and **work (kinetic energy).** **Energy** is often defined as the capacity for doing **work.** Energy transfers as either **heat** or **work**, or a combination of both. In physical processes (HT1), **Heat** is energy that transfers from one object to another because of a **temperature difference** between the objects. Heat flows from a warmer object to a cooler object until the **temperatures** are the same. **Kinetic Energy** ~ motion of particles (movement) **Heat** **Potential Energy** ~ spatial arrangement of particles (spacing) **Chemical reactions** generally involve either the **absorption** or the **release** of heat, not from temperature differentials, but from bonds breaking and reforming. ### The Law of Conservation of Energy: During any chemical or physical process, energy is neither **created** nor **destroyed.** This means that the energy in the universe **remains**. | System | Surroundings | |---|---| | Potential energy (Reactants) Δ(PE) Energy released to the surrounding as heat (Products) | The system is the reactants and the products. The surroundings are everything else. ΔPE means change in potential energy. This graph shows the system with a decrease in PE. We call this exothermic. | ### Conservation of Energy in a Chemical Reaction: - **Endothermic Reaction** - Reactant + Energy -> Product - In this example, the energy of the system (reactants and products) increases, while the energy of the surroundings decreases. - **Exothermic Reaction** Reactant -> Product + Energy - In this example, the energy of the system (reactants and products) decreases, while the energy of the surroundings increases. In every case, however, the total energy **DOES NOT CHANGE!** **HT2.2 | Heat of Reaction:** **Heat of reaction** (ΔH, in J or kJ) is the amount of heat released (neg) or absorbed (pos) during a chemical reaction. We call this term **enthalpy**. ### Endothermic vs. Exothermic Processes **Endothermic** - Reactions in which energy is **absorbed** as the reaction proceeds. Heat flows **into** the system from the surroundings. - Kinetic energy (and temperature) of the surroundings decreases and is stored in the system as potential energy. **Exothermic** - In an **exothermic** process, energy is **released** as the reaction proceeds. Heat flows **out of** the system into the surroundings. - Potential energy is converted into heat energy (which is released to the surroundings; kinetic energy (and temperature) of the surroundings increases! Energy + Reactants -> Products - ΔH is **positive** for an endothermic reaction and **negative** for an exothermic reaction. ### Potential Energy Diagrams: A potential energy diagram visually represents the energy change in the reaction system. - If the arrow goes **downward**, the reaction is **exothermic**. Potential energy stored in the bonds of the reactants is released as heat to the surroundings. ΔH is negative. - If the arrow goes **upward**, the reaction is **endothermic**. Heat form the surroundings is absorbed and stored as potential energy in the products. ΔH is positive. **CH_4(g) + 2O_2(g) -> CO_2(g) + 2H₂O₍₁₎** ΔH = - 890.3 kJ mol^-1 **H_2(g) + 1/2I_2(g) -> HI_(s)** ΔH = + 26.5 kJ mol^-1 In a chemical equation, the enthalpy change for the reaction can be written as either a **reactant** or a **product.** A chemical equation that includes the enthalpy change is called a **thermochemical equation.** ### Below is a thermochemical equation the describes an exothermic reaction of calcium oxide and water. In that reaction, enthalpy is a product because heat was released (created) **CaO_(s) + H₂O₍₁₎ -> Ca(OH)_2(s) + 65.2kJ** In the example above, the **heat of reaction** (the enthalpy change for the equation), which is denoted as ΔH, is **-65.2 kJ/mol**. Note that ΔH is **negative** since the reaction is exothermic. Other reactions absorb heat from surroundings. For example, baking soda (sodium bicarbonate) decomposes when it's heated. The carbon dioxide released in the reaction causes the muffins to rise while baking. This process is endothermic. **2NAHCO_3(s) + 86kJ -> Na₂CO_3(s) + H₂O_(l) + CO_2(g)** ΔΗ for the above equation is **+86KJ/mol** ### Practice: Indicate whether each reaction is **endothermic** or **exothermic**. Rewrite the equation as a " ΔΗ = equation. Draw a potential energy diagram for each! 1. **Br₂₍₁₎ + Cl_2(g) + 29.4 kJ -> 2 BrCl_(g)** Endothermic * **Br₂₍₁₎ + Cl_2(g) -> 2 BrCl_(g) ΔH=+29.4kJ/mol** 2. **NH_3(g) + HCl_(g) -> NH₄Cl_(s) + 176 kJ** Exothermic * **NH_3(g) + HCl_(g) -> NH₄Cl_(s) ΔH=-176kJ/mol** 3. **N₂O_4(g) + 58.0 kJ -> 2 NO_2(g)** Endothermic * **N₂O_4(g) -> 2NO_2(g) ΔH= + 58.0kJ/mol** 4. **CS₂₍₁₎ + 3 Cl_2(g) -> CCl₄₍₁₎ + S₂Cl₂₍₁₎ + 112 kJ** Exothermic * **CS₂₍₁₎ + 3 Cl_2(g) -> CCl₄₍₁₎ + S₂Cl₂₍₁₎ ΔH = -112kJ/mol** ### Heat of Reaction Calculations: Chemistry problems involving enthalpy changes are similar to stoichiometry problems. - The amount of heat released or absorbed depends on the number of moles of the reactants involved (the limiting reactant). - For example, the decomposition of 2 mol of sodium bicarbonate requires 85kJ of heat; therefore, 4 mol of sodium bicarbonate requires 88 kJ of heat. - We show these calculations with our times-bar conversions! **Example 1:** Calculate the amount of heat (in kJ) required to decompose 2.24 mol NaHCO_3(s). **2NaHCO_3(s) + 85 kJ -> Na₂CO_3(s) + H₂O_(l) + CO_2(g)** **2.24 mol NaHCO₃ x ^{85 KJ}/_{2 mol NaHCO₃} = 95 KJ** As per usual, it is more common to start with mass in **g** and then convert to **mol** to use the ratio in the equation. **Example 2:** The reaction below is highly exothermic. Calculate the amount of heat (in kJ) released when 36.0 g of Al reacts with excess Fe₂O3_(s). **2 Al_(s) + Fe₂O_3(s) -> 2 Fe_(s) + Al₂O_3(s) ΔΗ = -848 kJ/mol** **36.0g Al x ^{1mol Al}/_{26.989 Al} x ^{-848 KJ}/_{2 mol Al} = -566KJ** **Example 3:** Given the thermochemical equation: **2 NH_3(g) + 2 N₂O_(g) -> 4 N_2(g) + 3 H₂O_(l) + 1010 kJ** a) What quantity of heat (in kJ) is liberated by the reaction of 50.0 g of N₂O_(g) with excess NH_3(g)? **50.0g N₂O x ^{1mol N₂O}/_{44.029 N₂O} x ^{-1010 KJ}/_{2 mol N₂O} = -574 KJ** b) What quantity of heat (in kJ) is liberated by the reaction that produces 50.0 g of N_2(g)? **50.0g N₂x ^{1mol N₂}/_{28.02 g N₂} x ^{-1010 KJ}/_{4 mol N₂} = -451 KJ** ## HT2.2| Thermochemistry Assignment: 1. Hydrogen gas and oxygen gas release 482.6kJ of heat when they combine to form water. Is the reaction endothermic or exothermic? Is the change in enthalpy (ΔH) positive or negative? Write out and balance the reaction, and draw a potential energy diagram. * **2 H_2(g) + O_2(g) → 2 H₂O_(g) ΔΗ = -482.6 kJ/mol** 2. a. Write the above as a thermochemical equation 3. b. How much heat is released if we begin with 2.0087 g of O₂? (-30.29 kJ) 4. c. How much heat is released if we begin with 1.5021 g of H₂? (-179 kJ) 5. Hydrazine, N₂H₄, is used in rocket fuels. The thermochemical equation is as follows: **N₂H_4(l) + O_2(g) -> N_2(g) + H₂O_(l) ΔΗ = -622.4 kJ/mol** How much heat is liberated by the combustion of 1.000 g of N₂H₄ (l) ? (-19.41 kJ) 6. Glucose, C₆H₁₂O₆, is converted into ethanol, C₂H₅OH (l), in the fermentation of juice to make wine. The equation is as follows: **C₆H₁₂O₆ → 2 C₂H₅OH(l) + 2 CO_2(g) ΔΗ = -67.0 kJ/mol** a. Draw a potential energy diagram for the reaction. b. How much heat is released when a litre of wine containing 95.0g of C₂H₅OH(l) is produced? (-69.1 kJ) 7. Given the heat value ΔH = 42.7 kJ for the following equation: **2 NaN₃ -> 2 Na_(s) + 3 N_2(g)** a) Write the heat value into the equation to make a thermochemical equation b) Draw a potential energy diagram for the system c) What is the value of the ΔH for the use of 1.50 kg of N_2(g)? (762 kJ) d) How many grams of NaN₃ (s) would be decomposed by 125 kJ of heat? (381 g) 8. Given the thermochemical equation: **2 NH_3(g) + 3 N₂O_(g) -> 4 N_2(g) + 3 H₂O_(l) + 622.4 kJ** a. What is the heat of reaction? (ΔΗ = -622.4 kJ) b. How much heat is liberated by the reaction of 50.0 g of N₂O_(g)? (-236 kJ) c. What quantity of heat is released from the reaction that produces 50.0 g of N_2(g)? (-278 kJ) ## HT2.3|Measuring Enthalpy Change in a Thermochemical Equation: Using a similar principle to heat exchange in HT1, we can determine the enthalpy change in a system by examining the energy changes in the surroundings. We know that energy change in the system will be equal and opposite to energy change in the surroundings. **Q_surroundings = mcΔT** **Q_system = -Q_surroundings** **Q_system = nΔH** (where n = number of moles) ### Q_surroundings: To determine the energy change in the surroundings, chemists use substances that will change temperature. They can therefore calculate the energy change in the surroundings using the specific heat formula. **Q = mcΔT** - Q = energy or heat change (J or kJ) - m = mass of substance in surroundings (g or kg) - c = specific heat capacity in (J/g°C or J/kg°C) - ΔT = temperature change in surroundings (°C) <h3>Q_system:</h3> To determine the energy change in the system, chemists use the following formula: **Q = nΔH** - Q = energy or heat change (J or kJ) - n = number of moles of the limiting reactant (mol) - ΔH = molar enthalpy of the limiting reactant (kJ/mol) **Where:** - n = number of moles of the thing that reacts - ΔH = molar enthalpy for the reaction - m = mass of the surroundings (usually water) - c = specific heat capacity of the surroundings (usually water) - ΔT = temperature change of the surroundings We can rearrange this formula to solve for molar enthalpy of reaction: **Q_system = -Q_surroundings** **nΔH = -mcΔT** ### Example 1: Calculate the molar enthalpy of the reaction if, when 0.25 moles of a substance is reacted, the heat released produces a 4.5°C temperature increase in 325 = 325g mL_ of water in a coffee cup calorimeter. **System:** - ΔΗ= ^-mcΔT/_n - ΔΗ=? - n=0.25mol - ΔΗ = -(325g) (4.19J/g℃) (+4.5℃)/_0.25mol - ΔΗ=-25000J/mol or △H=-24,511.5J/mol = -25KJ/mol **Surroundings:** - m= 3259 - C=4.19 J/g°C - ΔT=+4.5°C ### Example 2: Calculate the molar enthalpy for the solidification of gallium metal (Ga) if 10.0 g of gallium causes 50.0 mL of water to change temperature from 24.0°C to 27.8°C when it solidifies. **Ga** - ΔΗ = ^-mcΔT/_n - △H=? - n= 10.0g Gax ^{1mol Ga}/_{69.729 Ga} = 0.14343... mol Ga - ΔΗ=-(50.0g)(4.19J/gc) (3.8℃)/_{0.14343... mol} - ∆H = -5550J/mol) (or - 5600J/mol if following addition/subtraction rules) **Water:** - m=50.0g - C=4.19J/gc - AT= 27.8°C-24.0°C = 3.8℃ ### Example 3: The molar enthalpy of methane, CH4, is Hr = -803 KJ/mol. What is the minimum mass of methane that must be burned to warm 4.00 L of water from 22.4°C to 87.6°C, assuming no heat losses? * ΔΗ=-803KJ = -803000J/mol * CHY * mass:? * n=? **water:** - m=4.00kg - C=4190J/kgc - AT=87.6℃-22.4℃ = +65.2°C **n = ^-mcΔT/_ΔH = (4.00kg) (4190J/kg°c) (+65.2°C)/_{-803 000J/mol} =1.3608...mol x ^{16.059 CHy}/_{1 mol CHy} = 21.8g CHy** ## HT2.3|Calorimetry Worksheet 1. 9.0 grams of charcoal, C, were completely consumed in a bomb calorimeter. If we assume that the 2.0 L of water absorbed all the heat released by the charcoal, and if the temperature of the water increased from 20.25 °C to 56.04 °C, what is the molar enthalpy of carbon? (ΔΗ = -4.0x102 kJ/mol) 2. Find the temperature increase expected for 1.00 L of water when it absorbs all of the energy from the combustion of 1.00 g of acetylene, C2H2 (g). The molar enthalpy of combustion for acetylene is -1,290 KJ/mol. (AT = +11.8°C) 3. In a chemistry experiment, 10.0 g of urea (NH2COHN2) is added to 150 mL of water in a simple coffee cup calorimeter. A temperature decrease of 3.7°C is noticed. Calculate the molar enthalpy of urea. (ΔΗ = 17 kJ/mol) 4. If 0.315 moles of hexane (C6H14) is combusted in a bomb calorimeter containing 5.65 liters of water, calculate the molar heat of combustion of hexane if the water temperature rises 55.4 °C? (ΔΗ = -4160 kJ/mol) 5. CS2, a very flammable liquid, has a molar enthalpy of -1028 kJ/mol. What do you expect aluminum's final temperature to be if 1.0 kg of Al is initially at 20.0 °C, and it absorbs all the heat from the following sample of CS2? (Tf = 180°C) - mass of CS2 before burning: 22.6 g - mass of CS2 after burning: 11.6 g ## HT | Unit Exam Review Be sure to review your notes as well as your assignments. Notes are fair game for multiple choice. ### HT1 | Heat of Physical Change 1. How much energy is required to heat 400.0 g of brass (c = 3.8x102 J/kg°C) from 15°C to 82°C? 2. If 63 kJ of heat are removed from 2.5 kg of iron (c = 4.5x102 J/kg°C) at an initial temperature of 91°C, what is its final temperature? 3. 0.200 kg of an unknown metal at 93°C is placed in 0.200 kg of water (c = 4.19x103 J/kg°C) at 10.0°C. The final temperature of both the metal and water is 17.0°C. What is the heat capacity of the metal? 4. A furnace in a home raises the temperature of a home by 2.5°C after adding 750. kJ of heat. What is the mass of the air in the home? 5. How much heat energy is required to change 250 g of steam at 108.0 degrees Celsius to ice at -4.0 degrees Celsius? ### HT2 | Heat of Chemical Change 6. The molar enthalpy of combustion of methane, CH4, is Hr = -803 KJ/mol. What is the minimum mass of methane that must be burned to warm 6.00 L of water from 22.4°C to 90.0°C, assuming no heat losses? 7. Find the temperature increase expected for 1.50 L of water when it absorbs all of the energy from the combustion of 3.00 g of acetylene, C2H2 (g). The molar enthalpy of combustion for acetylene is -1,290 KJ/mol. 8. Given the following reaction... **(CH₃)₂CO(l) + 4O₂(g) → 3CO₂(g) + 3H₂O(g)** **ΔΗ = -1760 kJ** a. Write the thermochemical equation for the above reaction. b. Is this reaction an exothermic or an endothermic reaction? c. Draw a potential diagram for this reaction. Make sure you label all the necessary components of an energy diagram. d. Calculate the energy change when 2.52 moles of liquid acetone is reacted with excess oxygen gas. e. Calculate the energy change when 78.5 g of acetone is reacted with excess oxygen gas. ## Thermal Energy Unit ### Calorimetry Lab **Purpose:** The purpose of this activity is to determine the amount of heat lost or gained in two separate reactions and to be able to write the thermochemical equation for each reaction. - Dissolving CaCl₂ in H₂O - Dissolving NH₄Cl in H₂O **Background Information:** Scientists measure the change in thermodynamic quantities in thermochemical equations using a device known as a calorimeter. One kind of calorimeter, known as a coffee cup calorimeter, is shown at right. Coffee cup calorimeters are usually used to measure changes that take place in solution. The insulation provided by the foam cups ensures that any heat absorbed or released by the system (reactants and products) goes only to the solvent in the cup. In other words, the surroundings are restricted to the solvent in the cup. The thermometer allows the AT of the surroundings to be measured. By the First Law of Thermodynamics, we know any heat lost by the system must be absorbed by the surroundings. Using AT, which can be measured, and the heat capacity of the solvent in the coffee cup, the heat lost by the system can be calculated. Since the pressure is constant, this is equal to the enthalpy change for the process (q = ∆Η). **Example:** When 1.00 g of NaOH(s) is dissolved in 100.0 mL of H2O (I) in a coffee cup calorimeter, the temperature of the water rises from 25.00 to 27.66 °C. a) Find AH for the dissolving process b) What is the enthalpy change per mole of NaOH(s)? c) Was this process endothermic or exothermic? d) Write the thermochemical equation expressing what happened in this reaction. e) What type type of reaction occurred? **General Lab Procedure:** 1. Review MSDS sheets for CaCl₂ and NH₄Cl. Put on proper safety equipment. 2. Measure 100.0 mL of deionized water and pour it into the calorimeter. 3. Record the initial temperature of the water. 4. Use a scale to weigh a small scoop of CaCl₂ (approximately 5.0 grams). 5. Cover the calorimeter if possible, leaving only holes for the thermometer and stir stick. 6. Add the chemical to the calorimeter. Stir with a stir stick until dissolved. 7. Record the final temperature. 8. Rinse out your calorimeter and clean your stir stick and thermometer. Then repeat this procedure with a fresh set up for a small approximately 5.0g of NH₄CI. (Make sure to dispose of chemical solutions in the proper waster beaker.) **Data:** | Trial | mass of chemical (g) | Moles of chemical (mol) | volume of H₂O (mL) | Mass of H₂O (g) | Ti (°C) | Tf (°C) | ΔT (°C) | |---|---|---|---|---|---|---|---| | CaCl₂ | | | | | | | | | NH₄CI | | | | | | | | **Calculations/Questions:** (Show all work on looseleaf) For each of the two reactions: a) Determine ΔH for dissolving, kJ (*note, do not divide -mcAT by moles for this) b) Determine the molar enthalpy of solution for the substance (kJ/mol) c) Write the thermochemical equation that describes what happened in the reaction (reactant as solid, same product is now aqueous) **Results Summary:** | Trial | ΔΗ for dissolving process (kJ) | Molar enthalpy ΔΗ (kJ/mol) | Endothermic or Exothermic? | Thermochemical Equation | |---|---|---|---|---| | CaCl₂ | | | | | | NH₄CI | | | | |