Thermochemistry The study of energy during physical changes and chemical changes.

Objectives Be able to define and correctly use energy-related terminology. Identify and understand endothermic and exothermic processes.

Energy energy: the ability to do work (push things) chemical bonds store chemical potential energy. Some substances store much energy (example: fuels, explosives) but others contain little energy (H2O, CO2).

Heat Terms thermal energy: total KE of particles in a substance (depends on #g and #oC) heat: thermal energy that flows from warmer to cooler areas. Heat transfers by: (1) conduction—colliding particles (2) radiation—electromagnetic waves (3) convection—currents enthalpy: a measure of energy stored in a “system” at constant pressure

Energy Transfer exothermic process: system releases energy and the surroundings typically get warmer Example: coffee cools down, air gets warmer. surroundings (includes thermometer) energy system “enthalpy” decreases -DH

Energy Transfer endothermic process: system gains energy from the surroundings (the surroundings may get colder) Example: the surroundings get cooler as ice melts. surroundings energy system “enthalpy” increases +DH

Objectives Be able to define and correctly use the common units of thermal energy. Be able to make calculations related to thermal energy and temperature changes. Be able to define and understand the concept of specific heat.

Calorimetry calorimetry: calculating or measuring heat transfer during various processes common energy units: calories (cal) or joules (J) It takes about 4000 J to heat 1 liter of water by 1oC. (4 J for 1 mL) Q = m · DT · Cp “Q” represents energy or enthalpy (in joules) +Q = enthalpy increase = endothermic -Q = enthalpy decrease = exothermic

Specific Heat specific heat: the amount of energy required to raise 1.00 g of a substance by 1.00oC. molecules = higher Cp (change T slowly) metals = lower Cp (change T quickly) Substance Cp (J/goC) water 4.184 rubber 2.0 plastic 1.7 concrete 0.9 steel 0.5 mercury 0.14

Energy and Temperature Changes temperature changes involve changes in kinetic energy Q = m · DT · Cp Example: What is the change in enthalpy when a cup of water (227 g) of water cools from boiling to room temperature (97oC to 22oC)? Example: A wedding ring absorbs 16.4 J of energy when it is placed on a finger (the temperature rises from 21oC to 38oC). If the mass of the ring is 4.80 g, what is the “specific heat” of the metal?

Objectives Understand the concept of latent heat and how it corresponds to potential energy. Be able to make latent heat calculations.

Endothermic Processes Latent Heat Changes in state involve changes in potential energy. This stored energy is often called latent heat. Temperature (KE) is constant during a phase change. Endothermic Processes Exothermic Processes VAPOR Latent Heat of Vaporization +DH -DH Latent Heat of Condensation WARM LIQUID COOL LIQUID Latent Heat of Fusion +DH -DH Latent Heat of Solidification SOLID

Latent Heat Values Substance ammonia, NH3 5.65 23.4 ethanol, CH3CH2OH DH (kJ/mol) (+ melt, - solid) + vapor, - cond ammonia, NH3 5.65 23.4 ethanol, CH3CH2OH 4.60 43.5 hydrogen, H2 0.12 0.90 methanol, CH3OH 3.16 35.3 oxygen, O2 0.44 6.82 water, H2O 6.01 40.7 Size of value depends on the strength of intermolecular bonds!

Latent Heat Calculations Temperature remains constant during a phase change, so don’t use Q = m · DT · Cp use: Q = (m/M) · DH Example: How much energy is needed to boil 19.75 g of ethanol (CH3CH2OH)? Example: How much water (at 0oC) is freezing if 2.5 kJ of energy is released?

Latent Heat and Flathead Cherries! Sometimes water is sprayed on cherries to protect them from freezing. Why? As the water freezes, it releases latent heat which warms/protects the fruit.

Objectives Be able to draw a heating curve or cooling curve for a substance. Be able to correctly label the regions where DKE and DPE are occurring on a heating or cooling curve.

Heating and Cooling Curves Imagine heating an ice cube. What energy changes take place as it is continually heated? changes of state Q = m/M · DH D PE = molecules pulled apart when boiling D KE = molecules speed up D PE = molecules pulled apart when melting TEMPERATURE D KE = molecules speed up heating or cooling Q = m · DT · Cp D KE = molecules speed up TIME

Objectives Understand the concept of a standard heat of formation. Be able to calculate the heat of reaction using Hess’s Law and determine if a reaction is endothermic or exothermic.

Standard Heat of Formation standard heat of formation (DHf0): change in enthalpy that accompanies the formation of one mole of a compound from its elements at 25oC and 101.3kPa. DHf0 for any uncombined element in its normal state = 0 kJ/mol

Hess’s Law heat of reaction (DH0 ): energy absorbed or released in a chemical reaction Use Hess’s Law to calculate the heat of reaction: DH0 = SDHf0 [products] - SDHf0 [reactants] Calculate DH0, state whether endothermic or exothermic, and write the “thermochemical equation” (A) CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) (B) 2CO2(g) → 2CO(g) + O2(g)

Objective Be able to draw a “potential energy diagram” for a chemical reaction using a thermochemical equation. Be able to correctly determine the temperature changes that occur in the surroundings for both endothermic and exothermic reactions.

Kerr Dam (slide in progress) Energy is needed in the refining process for copper production; very ENDOthermic Cu2+ (aq) + 2e- + energy → Cu (s) Kerr Dam was built to produce electricity for copper production in Anaconda. Plus, the copper could be used to make wiring for homes in the region.

Thermochemical Equations and PE Diagrams endothermic reaction (+DH0) 2CO2(g) + energy → 2CO(g) + O2(g) 2CO2(g) + 566 kJ → 2CO(g) + O2(g) PE diagram 2CO(g) + O2(g) + 566 kJ +DH (enthalpy increases) 2CO2(g) endothermic—products have more energy, so they are less stable

PE Diagrams exothermic reaction CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + energy CH4(g) + 2O2(g) → CO2(g) + 2H2O(g) + 890 kJ CH4(g) + 2O2(g) 890 kJ - DH (enthalpy decreases) CO2(g) + 2H2O(g) exothermic—products have less energy stored in their bonds and are more stable

PE and KE a thermometer is a part of the “surroundings” exothermic reactions: PE → KE, so temperature of surroundings increase endothermic reactions: temperature of surroundings may decrease if KE → PE