Chapter 5 Thermochemistry PDF

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Chapter 5 Thermochemistry 1. Energy and Energy Changes 1. Forms of Energy 1. Energy: the capacity to do work or to transfer heat, all forms are either kinetic or potential 2. Kinetic energy: energy that results from motion 2

1. Equation: Ek = 

3. Potential energy: the energy possessed by an object by virtue of its position 1. Chemical energy: energy stored within the structural units 2. Electrostatic energy: potential energy that results from the interaction of charged particles 1. Equation: Eel ∝  1. Oppositely charged particles attract each other, whereas particles of like charges repel eachother. The magnitude of the resulting electrostatic potential energy is proportional of the 2 charges and divided by the distance between them. 2. If the charges are opposite (one positive and one negative), the result is a negative Eel value which indicates attraction 3. If the charges are the same (both positive or both negative), the result is a positive Eel value which indicates repulsion 4. Law of conservation of energy: Kinetic and potential energy are interconvertible (one can be converted from one to the other. Energy can assume many different forms that are interconvertible, the total amount of energy in the universe is constant. When one form of energy disappears, the same amount of energy must appear in another form or forms. 2. Energy changes in chemical reactions 1. System: the specific part of the universe that is of interest to us 2. Surroundings: part of the universe not included in the system 3. Heat: the transfer of energy between two bodies that are at different temperatures 1. Heat absorbed or heat released 4. Thermochemistry: the study of heat, the transfer of thermal energy in chemical reactions 1. Exothermic process: a process that gives off heat

1. From the system to the surroundings 2. Endothermic process: a process that absorbs heat 1. Heat has to be supplied to the system by the surroundings for the reaction to occur 3. Units of energy 1. Joule (J): SI unit of energy Equation(s): 2 2 2 2 1J = Ek =  = (2kg)(1m/s) = (1kg)(m /s ) = 1 N・m 

2

1J = Ek = 

2

1J = (2kg)(1m/s)

2 2 1J = (1kg)(m /s )

1J = 1 N・m  2

Where 1 N・m  = (kg)(m/s)

1. Kilojoules; because joules are very small, we often express it in kJ where 1. 1kJ= 1000 J 2. calorie (cal) 1. 1 cal = 4.184J 2. 1Cal= 1000cal 3. 1Cal= 4184 J 1. Introduction to thermodynamics 1. Thermodynamics: the scientific study of the interconversion of heat and other kinds of energy. 1. There are three types of systems

2. 1. Open system: a system that can exchange mass and energy with its surroundings

2. Closed system: a system that can exchange energy (but not mass) with the surroundings 3. Isolated system: a system that does not exchange neither mass nor energy with its surroundings 2. State and state functions 1. State of a system: the values of all relevant macroscopic properties, such as composition, energy, temperature, pressure and volume 2. State functions: properties that are determined by the state of the system, independent of how the state was achieved. 3. The first law of thermodynamics: based on the law of conservation of energy, states that energy can be converted from one form to another but cannot be created or destroyed 1. Energy is a state function, so we can therefore demonstrate the first law by measuring the change in the energy of a system between its initial state and final state in a process. The change in internal energy, ∆U , is given by: ∆U= Uf  - Ui 1. where ∆ means final minus initial 2. f can also be products where i can also be reactants ∆U= U( products) - U( reactants) 1. Thermal energy released by the system is absorbed by the surroundings. The transfer of energy from the system to the surroundings does not change the total energy of the universe. Therefore, the sum of the energy changes is zero: ∆Usystem   + ∆Usurroundings =0 = -∆Usurroundings ∆Usystem   1. Work and heat 1. Energy is defined as the capacity to do work or transfer heat. When a system absorbs or releases heat, its internal energy also changes. The overall change is given: ∆U= q  +w 1. Q is heat (released or absorbed by the system) 2. W is work done (on the system or by the system) 3. Neither q nor w is a state function. There sum, ∆U, d  oes not depend on the path between initial and final states because U is a state function 1. ∆U u  sually (unless otherwise specified) will refer to ∆U system  2. The sign conventions for q and w are as follows: q is + for an endothermic process and for an exothermic process. w is + for work done on a system by the surroundings and for work done by the system on the surroundings. 3. Process

Sign

Heat absorbed by the system (endothermic)

q is positi ve

Heat released by the system (exothermic)

q is nega tive

Work done on a system by the surroundings (volume decrease)

w is positi ve

Work done by a system by the surroundings (volume increase)

w is nega tive

1. 1. Enthalpy 1. Reactions carried out at constant volume or at constant pressure 1. The amount of work done by a pressure-volume, or PV work is given by: w = -P∆V 1. P is the external, opposing pressure and ∆V is the change in volume of the container as the result 1. When a chemical reaction is carried out at a constant volume, then no PV work is done because ∆V=0 ∆U= q  - P∆V 1. And because P∆V= 0 at constant volume; 1. qv= ∆U 2. Constant volume-pressure are often inconvenient and sometimes impossible to achieve. Most reactions occur in an open container under conditions of constant pressure (usually at whatever that atmospheric pressure is. In general, for a constant pressure process we write: ∆U = q + w ∆U = qp - P∆V

qp = ∆U + P∆V 1. Enthalpy and enthalpy changes 1. Enthalpy (H) 1. Equation: H = U + PV 1. U is the internal energy of the system and P and V are the pressure and volume of the system 2. U, P, V are all state functions 1. For any process, the change in enthalpy is given by: 1. Equation: ∆H = ∆U + ∆(PV) qp = ∆H 1. q is not a state function 1. If a reaction occurs under constant- volume conditions, the heat change (qp) is equal to ∆U 2. If a reaction occurs under constant- pressure conditions, the heat change (qp) is equal to 3.

1.

2. 1.

∆H Enthalpy of a reaction is equal to the difference between the enthalpies of the products minus the enthalpy of the reactants ∆H = H(products) - H(reactants) The enthalpy can be positive or negative depending on the process 1. For endothermic process (heat is absorbed) ∆H is positive 2. For exothermic process (heat is released) ∆H is negative Enthalpy changes two common processes, the first involving a physical change, and the second involving a chemical change Thermochemical equations: chemical equations that show the enthalpy changes as well as mass relationships 1. Guidelines for interpreting, writing, and manipulating thermodynamic equations: 1. When writing thermochemical equations, specify the physical states of all reactant and products 2. If we multiply both sides of a thermochemical equation by a factor 'n', then ∆H must also by multiplied by that same factor. 3. When chemical equations are reserved, the roles of the reactant and products are changed. The magnitude for ∆H stays the same, but its sign is inverted. Therefore an exothermic process becomes and endothermic process and an endothermic process becomes and exothermic process...