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Chapter 6

Thermochemistry

TermDefinition
Thermochemistry The study of relationships between chemistry and energy
Energy The capacity to do work
Work The result of a force acting through a distance
Heat The flow of energy caused by a temperature difference. The transfer of thermal energy
Kinetic Energy The energy associated with the motion of an object
Thermal Energy The energy associated with the temperature of an object and is a form of kinetic energy
Potential Energy The energy associated with the position or composition of an object
Chemical Energy The energy associated with the relative position of electrons and nuclei in atoms and molecules, is also a form of potential energy
The Law Of Conservation of Energy(a.k.a First Law of Thermodynamics) Energy can be neither created nor destroyed
A System's Surrondings Are everything with which the system can exchange energy
Energy Transfer The surroundings gain the exact amount of energy lost by the system
joule(J) The SI unit for energy... 1J = 1Kg((m^2)/(s^2))
calorie(cal) The amount of energy required to raise the temperature of 1g of water by 1*C... 1 cal = 4.184 J(exact)
Calorie(Cal) Also called the Kilocalorie(kcal) Equivalent to 1000 cal
Kilowatt-hour(kWh) 1kWh = 3.60 * 10^6
The First Law of Thermodynamics The law of energy conservation, which states, "The total energy of the universe is constant"
Internal Energy(E) of a System The sum of the kinetic and potential energies of all of the particles that compose the system. Is a state funtion
State Function Its value depends only on the state of the system, not on how the system arrived at that state. Change in aptitude depends only on the difference between the initial and final values, not on the path taken
Calculating Internal Energy Change Δ”E = Efinal - Einitial.......... ΔE = Eproducts - Ereactants
Amount of Energy Lost by a System ΔEsystem = - ΔEsurroundings
Summarizing Energy Flow 1. If the reactants have higher E than products, ΔEsystem is negative and energy flows out of the system into surroundings.. 2. If the reactants have lower E than the products, ΔEsystem is positive and energy flows into system from surroundings
q(heat) if +q then system gains thermal energy... if -q then system loses thermal energy
w(work) if +q then work done on the system... if -q then work done by the system
ΔE(change in internal energy) if +ΔE then energy flows into the system... if -ΔE then energy flows out of the system
Internal Energy equation Energy entering the system through heat or work carries a positive sign, and energy leaving the system through heat or work carries a negative sign... ΔE = q + w
Energy Flow Out of System(System ---> Surroundings) ΔEsys < 0(negative)... ΔEsurr > 0(positive)
Energy Flow Into System(Surroundings ---> System) ΔEsys > 0(positive)... ΔEsurr < 0(negative)
Temperature A measure of the thermal energy within a sample of matter
Thermal Equilibrium The point at which there is no additional net transfer of heat between a system and its surroundings
Relationship Between Change in Temperature(ΔT) and Heat(q) Heat absorbed by a system and its corresponding temperature change are directly proportional... q { ΔT
Heat Capacity(C) The quantity of heat required to change the system's temperature by 1*C. Is an extensive property... C = (q/ΔT) = (J/*C)
Extensive Property It depends on the amount of matter being heated
Specific Heat Capacity or Specific Heat(Cs) The amount of heat required to raise the temperature of 1 gram of a substance by 1*C of 1K. Intensive property. Is dependent on type of substance... q = m x Cs x ΔT
Specific Heat Units J/(g*(*C)) or J/(g*K)
Molar Heat Capacity The amount of heat required to raise the temperature of 1 mole of a substance by 1*C or 1K. Intensive property. Is dependent on type of substance... q = n x Cs x ΔT
Intensive Property Does not depend on the amount of matter
Thermal Energy Transfer Between System and Surroundings If we define one substance as the system and the other as the surrounding, we can quantify the heat exchange as... qsyst = -qsurr
Combination of Equations mass(m)syst x specific heat(Cs)syst x change in temperature(”T)syst = -mass(m)surr x specifc heat(Cs)surr x change in temperature(ΔT)surr
Pressure-Volume Work Occurs when the force is caused by a volume change against an external pressure
Equation for Value of Pressure-Volume Work w = F x D... With w as work, F as force, and D as distance
Equation for Force P = F/A... or F = P x A... With force(F), pressure(P), and area(A)
Equation for Work Using Force w = F x D = P x A x D
Equation for Work with pistons The distance through which the force acts is the change in the height of the piston as it moves during the expansion(Δh)... w = P x A x Δh
Equation for Work with pistons and A x Δh = ΔV w = -PΔV
Equation for ΔErxn with Constant-Volume Calorimetry If reaction is carried out at constant volume... ΔErxn = qv + w(w = 0 at constant volume)... ΔErxn = qv(heat at constant volume)
Heat at Constant Volume(qv) The heat evolved (or given off)
Calorimetry The experimental procedure used to measure the heat evolved in a chemical reaction. We measure the thermal energy exchanged between the reaction(defined as the system) and the surroundings by observing the change in temperature of the surroundings
Bomb Calorimeter A piece of equipment designed to measure ΔE for combustion reactions. Measures changes in internal energy for combustion reactions
Bomb Calorimeter equation The temperature change(ΔT) is related to the heat absorbed by the entire calorimeter assembly(qcal) by the equation... qcal = Ccal x ΔT... with Ccal as the heat capacity of the entire calorimeter assembly
Energy Transfer in Bomb Calorimetry qcal = -qrxn
Enthalpy(H) The sum of a systems internal energy and the product of its pressure and volume, Is a state function H = E + PV
Change in Enthalpy for any Process Occurring Under Constant Pressure ΔH = ΔE + PΔ”V
Change in Enthaly(ΔH) Using Heat at Constant Pressure(qp) ΔH = qp
Endothermic Reaction Absorbs heat from surroundings; feels cool to the touch; Δ”H > 0(positive)
Exothermic Reaction Releases heat to its surroundings; feels warm to the touch; ΔH < 0(negative)
1. Summarizing Enthalpy 1. The value of ΔH for a chemical reaction is the amount of heat absorbed or evolved in the reaction under conditions of constant pressure
2. Summarizing Enthalpy 2. An endothermic reaction has a positive ΔH and absorbs heat from the surroundings. An endothermic reaction feels cold to the touch
3. Summarizing Enthalpy 3. An exothermic reaction has a negative ΔH and gives off heat to the surroundings. An exothermic reaction feels warm to the touch.
Enthalpy of Reaction or Heat of Reaction(ΔHrxn) The enthalpy change for a chemical reaction. Is an extensive property, so depends on amount
Constant Pressure Measures ΔHrxn
Coffee-cup Calorimeter For many aqueous reactions, we can measure ΔHrxn fairly simply using a coffee-cup calorimeter
Constant Value Calorimetry(Forgot To Put Higher Up) Measures ΔErxn
Equation for Heat Absorbed by or Lost from the Solution(acting as surroundings) qsol'n = msol'n x Cssol'n x ΔT
Energy Transfer In Constant-Pressure Caolrimetry qrxn = -qsol'n
Summarizing Calorimetry 1. Bomb calorimetry occurs at constant volume and measure ΔE for a reaction.. 2. Coffee-cup calorimetry at constant pressure and measures ΔH for a reaction
1st Relationships Involving ΔHrxn 1. If a chemical equation is multiplied by some factor, then ΔHrxn is also multiplied by the same factor... (A + 2B --> C ΔH1).. (2A + 4B --> 2C ΔH2 = 2 x ΔH1)
2nd Relationship Involving ΔHrxn 2. If a chemical equation is reversed, then ΔHrxn changes sign... (A + 2B --> C ΔH1).. (C --> A + 2B ΔH2 = -ΔH1)
3rd Relationship Involving Δ”Hrxn(Hess's Law) 3. If a chemical equation can be expressed as the sum of a series of steps, then ΔHrxn for the overall equation is the sum of the heats of reaction for each step... (A + 2B --> C Δ”H1) + (C --> 2D ΔH2) = (A + 2B --> 2D Δ”H3 = ΔH1 + ΔH2
Hess's Law The change in enthalpy for a stepwise process is the sum of the enthalpy changes of the steps
1. Standard State(NOT STP) Standard State for Gas, Solid, Liquid: 1 atm @ 25*C.. Stardard State for Solution: 1 atm, 25*C, 1 M
2. Standard Enthalpy Change(ΔH*) The change in enthalpy for a process when all reactants and products are in their standard states. The degree sign indicates standard states
3. Standard Enthalpy(Heat) of Formation(ΔHf*) For a Pure Compound: The change in enthalpy when one mole of the compound forms from its constituent elements in their standard states.. For a Pure Element in Its Standard State: ΔHf* = 0
Formation of Methane(CH4) C(s,graphite) + 2H2(g) --> CH4(g) ΔHf* = -74.6 kJ/mol
Formation of Water(H2O) H2(g) + (1/2)O2(g) --> H2O(g) ΔHf* = -285.8kJ/mol
Formation of Carbon Dioxide(CO2) C(s, graphite) + O2 --> CO2(g) ΔH*f = -393.5 kJ/mol
Calculating the Standard Enthalpy Change for a Reaction To calculate ΔH*rxn, subtract enthalpies of formation of reactants times their stoichiometric coefficients from enthalpies of formation of products times their stoichiometric coefficients. ΔH*rxn = Σnp * ΔH*f(products) - Σnf * ΔH*f(reactants)
Created by: TimChemistry1