Energy and Chemical Change Thermodynamics Energy and Chemical Change

Energy The ability to do work or produce heat. Two basic forms of energy: potential and kinetic energy. Potential energy is due to the composition or position of an object. Kinetic energy is the energy of motion.

Law of Conservation of Energy Energy can be converted from one form to another, but can neither be created nor destroyed. Chemical potential energy is the energy stored in a substance due to the arrangement of atoms and the strength of the chemical bonds.

Heat q Energy that flows from a warmer object to a cooler object. As a warmer object loses heat, the temperature decreases. SI unit of heat and energy is the Joule (J).

Calories are also often used as heat units. A calorie is defined as the amount of energy required to raise the temperature of one gram of pure water by one degree Celsius. Food calories (C) are the equivalent of 1000 calories. 1 cal = 4.184 J Convert 86.5 J to cal; convert 142 Cal to Joules.

Specific Heat Amount of energy required to raise one gram of a substance one degree Celsius. Water has an extremely high specific heat (4.18 J) which means that a large amount of energy must be used to raise the temperature of a quantity of water.

In order to calculate the heat energy that must be absorbed or released by a substance, you need to know the amount of the substance (mass in grams), the specific heat capacity (Cp) and the temperature change that the substance will undergo. q = m Cp T

What is the specific heat of iron if a 10.0 g sample changes from 50.4oC to 25.0oC with the release of 114 J of heat?

If the temperature of 34.4 g of ethanol increases from 25.0oC to 78.8oC, how much heat is absorbed by the ethanol? The specific heat of ethanol is 2.42 J/goC.

What temperature change will 100.0 mL of water undergo when it absorbs 1360 Joules of heat?

Thermochemistry: the study of heat changes that accompany chemical reactions and phase changes. System: the specific part of the universe under study. Surroundings: everything in the universe other than the system. Universe = System + Surroundings

The heat required for one mole of a substance to change phase is called the molar heat of: fusion when s l or vaporization when l  g Energy is released when l  s or g  l and Energy is absorbed when s  l or l  g

Calculate the heat required to melt 25.7 g of solid methanol at its melting point. The molar heat of fusion for methanol is 3.22 kJ/mol.

How much heat is evolved when 275 g of ammonia gas condenses to a liquid at its boiling point? The molar heat of vaporization for ammonia is 23.3 kJ/mol.

A thermochemical equation is a balanced chemical equation that includes the physical states of all reactants and products and the energy change. C6H12O6 (s) + 6O2(g)  6CO2(g) + 6H2O(l) H = -2808 kJ

Enthalpy (H) is the heat content of a system at constant pressure. The change in enthalpy for a reaction is called the Heat of reaction (Hrxn = Hfinal – Hinitial ) or Hrxn = Hproducts – Hreactants

The sign for the enthalpy change depends on whether the products or reactants have more energy.

Endothermic Reactions: Reactants + heat  Products Hrxn = + Hproducts > Hreactants

Exothermic Reactions: Reactants  Products + heat Hrxn = - Hproducts < Hreactants

If A  B and B  C what can you conclude?

Hess’s Law: In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or a series of steps. A  C = A  B B  C A  C

Given: 2CO(g) + O2 (g)  2CO2(g) H = ? Use these: C(s) + O2(g)  CO2(g) H = -393.5 kJ 2C(s) + O2(g)  2CO(g) H = +172.5 kJ

Calculate Hrxn for the reaction: HCl(g) + NH3(g)  NH4Cl(s) given the following: H2(g) + Cl2(g) 2HCl(g) H=-184kJ N2(g) + 3H2(g)  2NH3(g) -92 kJ N2(g) + 4H2(g) + Cl2(g)  2NH4Cl(s) -628 kJ

Calculate ΔH for the reaction: 2C(s) + H2(g)  C2H2(g) Given: ΔH= C2H2(g) + 5/2 O2(g)  2CO2(g) + H2O(l) -1299.6 C(s) + O2(g)  CO2(g) -393.5 H2(g) + ½ O2(g)  H2O(l) -285.8

Given the following reaction: Ca(s) + H2O(l) + CO2(g)  H2(g) + CaCO3(s) Use the following reactions: CaCO3(s) CaO(s) + CO2(g) ΔH = +177.8 kJ Ca(s) + 2 H2O(l)  Ca(OH)2(s) + H2(g) -414.3 kJ Ca(OH)2(s)  CaO(s) + H2O(l) +65.3 kJ

Hess’s Law Quiz Tomorrow Find ΔHorxn for C2H2(g) + 2 H2(g)  C2H6(g) Using the standard enthalpies of reaction below: 2 H2(g) + O2(g)  2 H2O(l) -572 kJ 2 C2H2(g) + 5 O2(g)  4 CO2(g) + 2 H2O(l) -2598 kJ 2 C2H6(g) + 7 O2(g)  4 CO2(g) + 6 H2O(l) -3122 kJ

Enthalpies of Formation The standard enthalpy of formation of a compound is called ΔHof ΔHof is the change in enthalpy for the reaction that forms 1 mole of the compound from its elements. All elements must be in their standard states at 298 K.

The enthalpy of formation for ethanol is -277.7 kJ. Write the equation for the formation of ethanol.

Standard enthalpies of formation can be used to determine the enthalpy of reaction using: Hrxn =  Hproducts –  Hreactants H2S(g) + 4F2(g)  2HF(g) + SF6(g)

Write the thermochemical equation for the formation of zinc nitrate. Determine ΔH for the following reaction: NH4NO3(s)  N2O(g) + 2H2O(l)

Laws of Thermodynamics 1st Law- the energy of the universe is constant. Keeps track of thermodynamics doesn’t correctly predict spontaneity. Entropy (S) is disorder or randomness 2nd Law: the entropy of the universe increases.

Entropy S Measure of the randomness or disorder of the particles in a system. Gas > Liquid > Solid Substances at higher temperatures have a higher entropy than substances at lower temperatures.

Determine whether the following results in an increase or decrease in the entropy: CO2(s)  CO2(l) CaO(s) + CO2(g)  CaCO3(s) Ag+(aq) + Cl-(aq)  AgCl(s)

For each of the following pairs, choose the substance with the higher entropy: Ar(l) or Ar(g) 1 mol As at 400K or 1 mol As at 298K 100g Na2SO4(s) at 30oC or 100g Na2SO4(l) at 30oC or 100g

Spontaneous A reaction that will occur without outside intervention. We can’t determine how fast. We need both thermodynamics and kinetics to describe a reaction completely. Thermodynamics compares initial and final states. Kinetics describes pathway between.

Spontaneous Reactions Most spontaneous reactions are exothermic. However, some endothermic reactions occur spontaneously due to an increase in entropy of the system. Two factors need to be considered, enthalpy and entropy in deciding spontaneity

Gibbs Free Energy (G) is the energy available to do work. G = H - T S G = - reaction is spontaneous G = + reaction is not spontaneous

H S G Charac-teristics - + Spon. Non- Spon @ Low T High T

The enthalpy of formation of hydrogen chloride is -92.31 kJ/mol. 1. Write the thermochemical equation for the formation of hydrogen chloride. 2. Determine the heat of reaction when hydrogen chloride reacts with ammonia to form ammonium chloride with an enthalpy of formation of -314.4 kJ/mol.