chapter6

media type="custom" key="3846625" width="46" height="55"Chapter 6 Thermochemistry Thermochemistry is the study of the energy evolved or absorbed in chemical reactions and any physical transformations, such as boiling and melting. Thermochemistry, generally, is concerned with the energy exchange accompanying transformations, such as mixing, chemical reactions, and including calculations of such quantities as the heat capacity, heat of formation, and enthalpy. toc

=Energy= Energy is the capacity to do work. All forms of energy are capable of doing work, but not all forms of work pertain to chemistry (take physics). Work is subsequently a directed energy change from a process.

__Types of energy:__ -Radiant energy: solar energy; comes from the sun. It is the primary energy source of Earth, and allows life to exist. -Thermal energy: the energy associated with random motion of atoms and molecules. -Chemical energy: stored within structural units of chemical substances. -Potential energy: energy available by virtue of an object's position.

Law of Conservation of Energy. The total quantity of energy in the universe is constant. This is the single most important law of our universe. media type="custom" key="3846613" width="42" height="50"

=Energy Changes in Chemical Reactions= Heat - the transfer of thermal energy between two bodies that are different temperature. Thermochemistry - the study of heat change in chemical reactions System - specific part if the universe that is of interest to us. Surroundings - the rest of the universe outside the system Open system - can exchange mass and energy with the surroundings Closed system - allows transfer of energy but not mass Isolated system - does not allow transfer of either energy or mass

Exothermic - process that gives of heat or transfer energy to the Surroundings Endothermic - heat has to be supplied to the system by the surroundings

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=Laws of Thermodynamics=

In thermodynamics, there are four laws that describe the specifics for the transport of heat and work in thermodynamic processes. These laws can be applied to systems about which one knows nothing other than the balance of energy and matter transfer.

Zeroth Law - defines thermal equilibrium as a state of a system where there is no temperature change over time. First Law of Thermodynamics - energy can be converted from one for to another, but cannot be created or destroyed Second Law of Thermodynamics - entropy is always increasing Third Law of Thermodynamics - impossible to cool an object to absolute zero media type="custom" key="3846613" width="43" height="51"
 * Sign Conventions for Work and Heat ||
 * Process || Sign ||
 * Work done by the system on the surroundings || - ||
 * Work done on the system by the surroundings || + ||
 * Heat absorbed by the system from the surroundings (endothermic process) || + ||
 * Heat absorbed by the surroundings from the system (exothermic process) || - ||

=Work and Heat= We know from physics that: w = Fd where w is work, F is the force acting, and d is distance. The pressure is defined as P=F/A (the pressure P equals the force F applied over an area A), and when substituted into the previously stated work tells us that the work done on by the gas on the surroundings is: w = - P * (delta V) where P is pressure and, delta V is the change in volume.

The internal energy E of a system can be defined as the sum of the kinetic and potential energies of all the particles in the system. the internal energy can be changed by both work and heat. That is, Delta E = q + w where E is internal energy, q is heat, and w is work (previously defined).

Sample Question: A gas is compressed from an initial volume of 5.75 L to a final volume of 1.23 L by an external pressure of 1.00 atm. During the compression the gas releases 128 J of heat.

Answer: -1.00 atm (1.23L - 5.75L) = 4.52L*atm 4.52L*atm x 101.325 J/(L*atm) = 457.989J 457.989J - 128J = 330J media type="custom" key="3846613" width="37" height="43"

=Enthalpy=

In thermodynamics, the **enthalpy** (denoted as //H//, or specific enthalpy denoted as //h//) is a thermodynamic property of a fluid. It can be used to calculate the heat transfer during process taking place in a closed thermodynamic system under constant pressure. Enthalpy //H// is an arbitrary concept but the enthalpy change //ΔH// is more useful because it is equal to the change in the internal energy of the system, plus the work that the system has done on its surroundings. The enthalpy of a system H, can be defined as H = E + PV where E is the internal energy of the system, P is the pressure of the system, and V is the volume of the system. change in enthalpy is very similar to the heat of reaction, especially in a constant pressure system.



Enthalpy is a common state function, since internal energy pressure and volume are all state functions. It is the heat at constant pressure. Delta H = qmedia type="custom" key="3948141" width="12" height="16" media type="custom" key="3846613" width="42" height="44"

=Calorimetry=

A **calorimeter** is a device that is used to determine the heat associated with a chemical reaction. **Calorimetry** is the science of measuring heat, and is based on observing changes in temperature due to heat absorption and dispersion. The **heat capacity** C of substance is defined as: If the heat capacity is given //per gram// of a substance, it is called the **specific heat capacity**. The units are The measurement of heat changes is performed using calorimetry, usually an enclosed chamber within which the change to be examined occurs. The temperature of the chamber is monitored using a thermometer. The energy released by the reaction is equal to the energy absorbed by the solution. This is equal, qualitatively, to the specific heat capacity x mass of solution x increase in temperature. More simply: E = //s// x //m// x (Delta T)
 * Specific Heats of Some Common Substanaces ||
 * Substance || Specific Heat (J/gC) ||
 * Al || .9 ||
 * Au || .129 ||
 * C (graphite) || .72 ||
 * C (diamond) || .502 ||
 * Cu || .385 ||
 * Fe || .444 ||
 * Hg || .139 ||
 * H2O || 4.184 ||
 * C2H5OH (ethanol) || 2.46 ||

Sample Question: .35 grams of water are placed in a calorimeter. The water temperature drops from 34.79 C down to 30.01 C when .81 grams of a substance is placed into the calorimeter with it. The substance has a specific heat of .63 J/g-C. What was the initial temperature of the substance? Answer: 0.35 x 4.18 ( 30.01 - 34.79) + 0.81 x 0.63 ( 30.01 - T)=0 - 6.99 + 15.3 - 0.510 T = 0 8.31 = 0.510 T T = 16.3 °C media type="custom" key="3846613" width="43" height="46"

=Standard Enthalpy=

Enthalpy is a state function, as such the change in enthalpy is going from some initial state to some final state, is independent of the pathway. This means that 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 one step or a series of steps. This principle is Hess's Law.

The total enthalpy of a system cannot be measured directly; the //enthalpy change// of a system can be however. Enthalpy change is defined by the following equation: Δ//H// = //H final// − //H initial// where Δ//H// is the //enthalpy// change, //H// final is the final enthalpy of the system, measured in joules. In a chemical reaction, //H final// is the enthalpy of the products. //H initial// is the initial enthalpy of the system, measured in joules. In a chemical reaction, //H initial// is the enthalpy of the reactants.

For an endothermic reaction, the system's change in enthalpy is equal to the energy //absorbed// in the reaction, including the energy //lost by// the system and //gained// from compression from its surroundings. A relatively easy way to determine whether or not a reaction is exothermic or endothermic is to determine the sign of Δ//H//. If Δ//H// is positive, the reaction is endothermic, heat is absorbed by the system due to the products of the reaction having a greater enthalpy than the reactants. On the other hand if Δ//H// is negative, the reaction is exothermic, that is the overall decrease in enthalpy is achieved by the generation of heat.

Hess's Law The process works by treating each chemical equation as a mathematical equation, adding and subtracting multiples of each to cancel each variable (chemical). The yield arrow acts as an equals sign, and different states of elements count as different variables. media type="custom" key="3846613" width="36" height="42"

=Laboratory=

Coffee Cup Calorimetry For this lab you will need:
 * two or more styrofoam cups
 * a lid with holes for a thermometer and stirrer
 * a thermometer
 * a stirrer
 * hydrochloric acid
 * magnesium strips (approximately .1g)

First, nestle your styrofoam cups into each other. The more cups, the more isolated the system will be from it's surroundings. Fill your calorimeter with 100 ml of .1M hydrochloric acid. Set your thermometer into the acid, and allow it to become stable. Drop a magnesium strip into the acid and seal the cup. Watch the temperature rise, while stirring, and wait for it to cap. Record temperature change. Try to calculate the kcal/mole for the reaction. It should be around .5. Clean out your cup using standard laboratory cleanup procedures.

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=References= All images and information was drawn from myself or the following sources: Chang, Chemistry, 8th Edition Zumdahl, 4th Edition