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Thermochemistry

Thermochemistry. Nature of Energy. Kinetic Energy E k = ½mv 2 m=mass (kg) v=velocity (m/s) Potential Energy E p = mgh g = gravity (9.8 m/s 2 ) h = height (m) Chemical Energy. Units of Energy. Joule (J) – 1 kg-m 2 /s 2 Small amount of energy so we use kJ

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Thermochemistry

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  1. Thermochemistry

  2. Nature of Energy • Kinetic Energy • Ek = ½mv2 • m=mass (kg) • v=velocity (m/s) • Potential Energy • Ep = mgh • g = gravity (9.8 m/s2) • h = height (m) • Chemical Energy

  3. Units of Energy • Joule (J) – 1 kg-m2/s2 • Small amount of energy so we use kJ • Calorie (cal) – amount of energy needed to raise the temperature of 1 gram of water by 1 degree • Normally uses Cal = 1000 cal = 1kcal • 1 cal = 4.184 J • How many calories are there in 15 joules?

  4. Transfer of Energy • To understand energy transfer we must define two quantities: • System • Closed vs Open • Surroundings • Energy transferred in two ways • Work (w) – causes motion • W = F x d • Heat (q) – causes temperature change • Energy is the ability to do work or transfer heat.

  5. First Law of Thermodynamics • 1st Law = Energy is conserved. • Internal energy of a system = sum of all kinetic and potential energy of the components of the system • Internal Energy = E • We measure ΔE = Efinal – Einitial • ΔE has two parts • Magnitude = number, unit • Direction = + or - • For a chemical reaction: • What is the initial state? • What is the final state?

  6. Heat and Work • ΔE = q + w • q > 0 when heat is added to the system • w > 0 when work is done on the system • Endothermic • q > 0 • Exothermic • q < 0

  7. Heat, Work and Internal Energy

  8. Energy Practice • If hydrogen and oxygen gases are trapped in a cylinder and ignited water is produced. As the reaction occurs, the system loses 1150J of heat to the surroundings. The reaction also causes a piston to rise as the hot gases expand. The expanding gas does 480J of work on the surroundings as it pushes against the atmosphere. What is the change in internal energy?

  9. Energy Practice • Calculate the change in the internal energy of the system for a process in which the system absorbs 140J of heat from the surroundings and does 85J of work on the surroundings. • Calculate ΔE and determine where the process is exo- or endothermic for the following cases. • A system releases 133kJ of heat to the surroundings and does 39 kJ of work on the surroundings • The system absorbs 77.5kJ of heat while doing 63.5kJ of work on the surroundings

  10. State Functions • State function = a function whose value only depends on its present condition not on the history of the sample

  11. State Functions • ΔE = q + w

  12. Enthalpy • Usually the only kind of work done by a chemical reaction is mechanical work • Reactions normally done at constant P • Zn(s) + 2H+(aq) → Zn2 +(aq) + H2(g) • W = -PΔV

  13. Constant Volume vs. Pressure • ΔE = q + w • W = -PΔV • qv= ΔE • qp = ΔE + PΔV

  14. Enthalpy • H = enthalpy • Enthalpy measures the heat flow in chemical changes occurring at constant pressure • H = E + PV • ΔH = ΔE + PΔV • ΔH = qp • When is ΔH positive and when is it negative?

  15. Exo- and Endothermic

  16. Enthalpy Practice • Indicate the sign of the enthalpy change, ΔH, in each of the following processes carried out under atmospheric pressure and indicate whether the process is exo- or endothermic. a) an Ice cube melts b) 1g of butane (C4H10) of combusted in sufficient oxygen to give complete combustion to CO2 and H2O c) a Bowling is dropped from a height of 8 ft into a bucket sand

  17. Enthalpy Practice • Suppose we confine 1g of butane and sufficient oxygen to completely combust it in a cylinder. The cylinder is perfectly insulating, so no hear can escape to the surrounding. A spark initiates combustion of the butane, which forms carbon dioxide and water vapor. If we used this apparatus to measure enthalpy change in the reaction, would the piston rise, fall, or stay the same? • Hint write a balanced reaction

  18. Enthalpies of Reaction • ΔH = Hfinal – Hinitial • Enthalpy of products – enthalpy of reactants • Heat of Reaction (ΔHrxn) = Enthalpy change of the reaction • 2H2(g) + O2(g) → 2H2O(g) ΔH = -483.6 kJ • Thermochemical equation • Enthalpy Diagram

  19. Enthalpy Guidelines • Enthalpy is an extensive property • CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH = -890kJ • 1 mol of CH4 w/ 2 mol of O2 releases 890kJ • 2 mol of CH4 w/ 4 mol of O2 releases 1780kJ • How much heat is released when 4.50g of methane is burned in a constant pressure system?

  20. Concept Practice • Hydrogen peroxide can decompose to water and oxygen by the following reaction: • 2H2O2(l) → 2H2O(l) + O2(g) ΔH = -196kJ Calculate the value of q when 5.00g of H2O2(l) decomposes at constant pressure.

  21. Enthalpy Guidelines • The enthalpy change for a reaction is equal in magnitude but opposite in sign to ΔH for the reverse reaction. • The enthalpy change for a reaction depends on the state of the reactants and products.

  22. Concept Practice • Consider the following reaction: 2Mg(s) + O2(g) → 2MgO(s) ΔH = -1204kJ a) Is the reaction exo- or endothermic? b) Calculate the amount of heat transferred when 2.4g of Mg(s) reacts at constant pressure. c) How many grams of MgO are produced during an enthalpy change of -96.0kJ? d) How many kJ of heat are absorbed when 7.50g of MgO(s) are decomposed into Mg(s) and O2(g) at constant pressure?

  23. Homework Set 1 • 12, 17, 20, 21, 22, 25, 29, 34, 35, 41

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