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Unit 1: Thermochemistry

Unit 1: Thermochemistry. Introduction The First Law of Thermodynamics Enthalpy Enthalpy of Reaction Calorimetry Hess’s Law Enthalpy of Formation. Introduction. Most daily activities involve processes that either use or produce energy: Activities that produce energy Metabolism of food

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Unit 1: Thermochemistry

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  1. Unit 1: Thermochemistry • Introduction • The First Law of Thermodynamics • Enthalpy • Enthalpy of Reaction • Calorimetry • Hess’s Law • Enthalpy of Formation

  2. Introduction • Most daily activities involve processes that either use or produce energy: • Activities that produce energy • Metabolism of food • Burning fossil fuels • Activities that use energy: • Photosynthesis • Pushing a bike up a hill • Baking bread

  3. Introduction • Thermodynamics • The study of energy and its transformations • Thermochemistry: • A branch of thermodynamics • The study of the energy (heat) absorbed or released during chemical reactions

  4. Introduction • Objects can have two types of energy: • Kinetic energy • Energy of motion • Thermal energy • The type of kinetic energy a substance possesses because of its temperature • Potential energy • Energy of position • “stored” energy resulting from the attractions and repulsions an object experiences relative to other objects

  5. Introduction • Attractive and repulsive forces include: • Gravity • Electrostatic forces between charged particles • e- has potential energy due to its position near the positively charged nucleus • Most important attractive/replusive forces in chemistry

  6. Introduction • Attractive and repulsive forces within a substance lead to a type of potential energy called chemical energy • The potential energy stored in substances resulting from the arrangements of the atoms in the substance

  7. Introduction • Units of Energy • SI unit = joule (J) • 1 J = the kinetic energy of a 2 kg mass moving at a speed of 1 m/s • A very small quantity • Kilojoule (kJ) • 1 kJ = 1000 J

  8. Introduction • Units of Energy (cont) • Calorie (cal) • Originally defined as the amount of energy needed to raise the temperature of 1g of water from 14.5oCto 15.5oC. • 1 cal = 4.184 J (exactly) • Kilocalorie (kcal) • 1 kcal = 1000 cal

  9. Introduction • On the exam, you must be able to convert from one set of energy units to the other using dimensional analysis. • You must know the conversion factors given on the previous slides!!!

  10. Introduction Example: Convert 725 cal to kJ.

  11. Introduction Example:A particular furnace produces 9.0 x 104 BTU/hr of heat. If 1.00 BTU = 251.9958 cal, use dimensional analysis to calculate the number of kJ of heat delivered by the furnace after running for 2.50 hours.

  12. Introduction • When using thermodynamics to study energy changes, we generally focus on a limited, well-defined part of the universe. • System: • The portion of the universe singled out for study • Surroundings: • Everything else

  13. Introduction The system is usually the chemicals in the flask/reactor. The system The flask and everything else belong to the surroundings.

  14. Introduction • Open system: • A system that can exchange both matter and energy with the surroundings • Closed system: • A system that can exchangeenergywith the surroundings butnot matter A cylinder with a piston is one example of a closed system.

  15. Introduction • In a closed system energy can be gained from or lost to the surroundings as: • Work • Heat • Work: • Energy used to cause an object to move against a force • Lifting an object • Hitting a baseball

  16. Introduction • Heat: • The energy used to cause the temperature of an object to increase • The energy transferred from a hotter object to a cooler one • Energy: • The capacity to do work or to transfer heat

  17. work Introduction • The potential energy of a system can be converted into kinetic energy and vice versa. • Energy can be transferred back and forth between the system and the surroundings as work and/or heat. Potential energy Kinetic energy

  18. The First Law of Thermodynamics • Although energy can be converted from one form to another and can be transferred between the system and the surroundings: Energy cannot be created or destroyed. (First Law of Thermodynamics) • Any energy lost by the system must be gained by the surroundings and vice versa.

  19. The First Law of Thermodynamics • The First Law of Thermodynamics can be used to analyze changes in the Internal Energy (E) of a system. • The sum of all kinetic and potential energy of all components of a system • For molecules in a chemical system, the internal energy would include: • the motion and interactions of the molecules • the motion and interactions of the nuclei and electrons found in the molecules

  20. The First Law of Thermodynamics • Internal Energy: • Extensive property • depends on mass of system • Influenced by temperature and pressure • Has a fixed value for a given set of conditions • State function

  21. The First Law of Thermodynamics • The internal energy of a system is a state function. • A property of the system that is determined by specifying its condition or its state in terms of T, P, location, etc • Depends only on its present condition • Does not depend on how the system got to that state/condition

  22. The First Law of Thermodynamics • The internal energy of a system can change when: • heat is gained from or lost to the surroundings • work is done on or by the system. • The change in the internal energy D E = Efinal - Einitial DE = change in internal energy Efinal = final energy of system Einitial = initial energy of system

  23. The First Law of Thermodynamics • If Efinal > Einitial, • DE >0 (positive) • the system has gainedenergy from the surroundings. • endergonic

  24. The First Law of Thermodynamics • The decomposition of water is endergonic (DE > 0): 2 H2O (l) 2 H2 (g) + O2 (g) H2 (g), O2(g) final Energy must be gained from the surroundings. E H2O (l) initial

  25. The First Law of Thermodynamics • If Efinal < Einitial, • DE < 0 (negative) • the system has lostenergy to the surroundings. • exergonic

  26. H2 (g), O2(g) H2O (l) The First Law of Thermodynamics • The synthesis of water is exergonic (DE < 0) 2 H2 (g) + O2 (g) 2 H2O (l) Energy is lost to the surroundings in this reaction. initial E final

  27. The First Law of Thermodynamics • The internal energy of a system can change when energy is exchanged between the system and the surroundings • Heat • Work • The change in internal energy that occurs can be found: D E = q + w Where q = heat w = work

  28. The First Law of Thermodynamics • By convention: • q = positive • Heat added to the system • w = positive • Work done on the system by the surroundings • q = negative • Heat lost by the system • w = negative • Work done by the system on the surroundings

  29. The First Law of Thermodynamics Example: Calculate the change in internal energy of the system for a process in which the system absorbs 240. J of heat from the surroundings and does 85 J of work on the surroundings.

  30. The First Law of Thermodynamics

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