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Chapter 6

Chapter 6.docx

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Department
Chemistry
Course Code
CHEM 1F92
Professor
Lydia W.L.Chen

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Chapter 6 – Thermochemistry 6.1 – The Nature of Energy and Types of Energy • Energy is the capacity to do useful work.All forms of energy are capable of doing work. o Ex: The energy in a tidal wave can be harnessed to do useful work. • Work is total energy change that results from a process. • Kinetic Energy is the energy possessed by a moving object. • Radiant energy (or solar energy) heats the atmosphere and Earth’s surface. It used by plants and vegetation during the process of photosynthesis. It also influences global climate patterns. • Thermal Energy is the energy associated with the motion of atoms and molecules. o The faster the motion of atoms and molecules in a substance, the warmer it is and the greater its thermal energy. o The greater the number of particles in a substance, the greater its thermal energy (but not necessarily the greater its temperature). o Temperature is the average K of all the particles in a substance. • Chemical energy is stored in the bonds within a chemical substance. During a chemical reaction, chemical energy is stored, released, and converted into other forms of energy. • Potential energy is the energy stored in an object due to its position. • Chemical energy is a form of potential energy because it depends on the positions and arrangements of atoms within a substance. Law on Conservation of Energy: The total energy of the universe remains constant. If it disappears in one part of the universe, it will appear in another form, somewhere else in the universe. 6.2 – Energy Changes In Chemical Reactions • Heat is the transfer of thermal energy between two bodies at different temperatures. Heat is transferred from a hot object to a cold object. • Thermochemistry is the study of heat change in chemical reactions. • The system is the part of the universe that is being studied or observed. In a chemical reaction, the system includes the reactants (and the container in which they are held). • The surroundings are the rest of the universe, outside the universe. • There are 3 types of systems o An open system is one in which both mass and energy (in the form of heat) can be transferred with its surroundings. o Aclosed system only allows the transfer of energy (in the form of heat), not mass. o An isolated system does not allow the transfer of mass or energy. • An exothermic reaction is a reaction in which thermal energy is transferred from the system to the surroundings. 2H 2 (g)O 2 (g)2H O 2 E(l)gy • An endothermic reaction is a reaction in which thermal energy is transferred from the surroundings to the system. Energy + 2HgO (s) 2Hg + (l) 2 (g) • In an exothermic and endothermic reaction. The difference in energy between the products and the reactants is equal to the amount of energy transferred to or from the surroundings. 6.3 –Introduction to Thermodynamics • Thermodynamics is the study of the conversion between heat and other forms of energy. • The state of a system is the values of all the macroscopic properties relevant to the system. o They include composition, energy, temperature, pressure, volume. • Astate function is a macroscopic property that depends only on the current state of a system, not the way that the state was achieved (independent of path). o The volume of a system only depends on the system’s current volume, not the changes that led to this volume. • Changes in a state function only depend on the initial and final conditions. o ∆V = V -2V 1 • Some state functions are volume, temperature, pressure and energy. The First Law of Thermodynamics The First Law of Thermodynamics: Energy cannot be created or destroyed. It can only be converted from one form to another form. • Since determining the total internal energy of a system (E) is difficult, the change in the internal energy of a system is often measured (∆E). o Total internal energy is the sum of the kinetic energy of all the atoms/molecules in a system and the potential energy in the bonds between molecules and within molecules. o Hence, it is very difficult to measure. However, the ∆E is easier to measure. ∆E = E products reactants • If ∆E is (-), some of the internal energy (chemical energy in bonds) of the reactants is converted to E th and released to the surroundings. o Therefore, the energy lost by the system is gained by the surroundings in the form of E .th ∆E system ∆E surroundings ∆E system -∆E surroundings • The energy lost by a system in one form can appear in the surroundings in a different form. o The energy lost as E th burning oil in a power plant (system) can be converted to electrical energy, heat, and light in homes (surroundings). • The first law of thermodynamics can be represented by the equation below. ∆E = q + w The change in internal energy of a system is the sum of the heat exchange (q) between the system and the surroundings as well and the work (w) done on or by the surroundings. q and w are in lowercase because they are not state functions. State functions are represented by uppercase letters. Work and Heat Work • Imagine a gas in a cylinder fitted with a weightless, frictionless movable piston at a certain temperature, pressure, and volume. • As the gas expands, the gas pushes the piston upward against a constant opposing external atmospheric pressure P. The work done by the gas on the surroundings is: w = -P∆V The work done by a gas on its surrounding is the (-) of the product of the opposing external atmospheric pressure and the change in volume of the gas. w = -F∆d The work done by a gas on its surroundings is the (-) of the product of the opposing external force and the change in distance of the gas in the container. • For gas expansion (work done by system), w < 0. For gas compression, (work done on system), w > 0. • P must be a positive value. If P = 0 (only occurs in a vacuum), no work will be done when the gas expands. 1 L ∙ atm = 101.325 J • Work is not a state function because the same gas held in a cylinder a piston can do different amounts of work when pushing the piston upwards if the opposing pressure changes. o Work depends on the initial state, final state, and how the process is carried out. Heat • Heat (q) is also not a state function. This is because the temperature of a sample does not only increase by the addition of heat. Work can be done on the sample to increase its temperature. If this happens, the change in temperature is not only caused by heat. o Hence, the amount of heat absorbed/released depends on the process that is carried out (does the process involve any work???).  ∆E = q + w 6.4 – Enthalpy of Chemical Reactions • The total internal energy of a system changes if volume is constant or if pressure is constant: If Volume is Constant (∆V = 0): ∆E = q + w The subscript v represents that this is a constant volume process.Although heat (q)
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