Thermochemistry! AP Chapter 5. Temperature vs. Heat Temperature is the average kinetic energy of the...
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Transcript of Thermochemistry! AP Chapter 5. Temperature vs. Heat Temperature is the average kinetic energy of the...
![Page 1: Thermochemistry! AP Chapter 5. Temperature vs. Heat Temperature is the average kinetic energy of the particles in a substance. Heat is the energy that.](https://reader035.fdocuments.net/reader035/viewer/2022081603/56649f285503460f94c413d0/html5/thumbnails/1.jpg)
Thermochemistry!
AP Chapter 5
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Temperature vs. Heat
• Temperature is the average kinetic energy of the particles in a substance.
• Heat is the energy that is transferred from one object to another.
• Heat always flows from the hotter object to the colder object.
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Energy!
• Energy is the ability to do work.
• Kinetic Energy - the energy of motion
• Potential Energy – the energy that an object has as a result of its composition or its position with respect to another object.
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Units of Energy
• 1 Joule = 1 kg m2/s2 (1 kJ is 1000J)• Used to calculate the energies associated with chemical
reactions.
• Calorie – Amount of energy required to raise the temperature of 1 gram of substance 1 °C. (This is specific heat!)1 calorie will raise the temperature of 1 g of H2O
from 14.5 °C to 15.5 °C.
1 calorie is equal to 4.184 Joules (exactly!)
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Systems and Surroundings
• System – the portion used in a study.– It can be an open system or a closed system.
• Open system – matter and energy can interact with the surroundings.
• Closed system – the matter cannot interact with the surroundings.
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First Law of Thermodynamics
•Energy Is Conserved!
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Internal Energy
• Internal Energy is the sum of all the kinetic and potential energies of all its components.
ΔE = Efinal - Einitial
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ΔE
• A positive value for ΔE is when Efinal > Einitial
• If energy has been absorbed from its surroundings, it is endothermic.
• If energy is given off to the surroundings, it is exothermic.
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Initial state refers to the reactants, while final state refers to the products.
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Endothermic reaction Exothermic reaction
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A system composed of H2 (g) and O2 (g) has greater internal energy than a system composed of H2O (l).
Gases have greater kinetic energy and must lose some of that energy to change states back to the liquid state.
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Internal energy is a function of state.
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a) If a battery is shorted out and loses energy to the environment only as heat, no work is done.
b) If a battery is discharged and loses energy as work (to make the fan run) it also loses heat energy.
c) The value of ∆E is the same.
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Enthalpy
• The change in enthalpy for a reaction (∆H) is the overall measure of energy that is absorbed to break bonds and the energy that is released when new bonds form.
• A reaction is said to be spontaneous if it occurs without being driven by an outside force. (driving forces are enthalpy & entropy)
• ∆H = ΣH(products) - ΣH(reactants)
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In an endothermic system where it absorbs heat, ∆H will be positive (∆H > 0).
In an exothermic system, where heat is given off, ∆H will be negative (∆H < 0).
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Enthalpy Diagrams
• Enthalpy is an extensive property – it depends on how much you have. If 1mol of CH4 and 2 mol O2 yield -890 kJ, then 2 mol CH4 and 4 mol O2 would yield double that.
• The enthalpy change for a reaction is equal in magnitude, but opposite sign, for a reverse reaction.
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Calorimetry
• This is a measure of the amount of energy that is needed or lost when a certain mass of a substance changes temperature.
• q = mC∆T• q is the amount of energy (J)• m is the mass of the substance (g)• C is the specific heat capacity of the substance• ∆T is the change in temperature
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Calorimeters
• Calorimeters are devices that measure the transfer of heat from one object to another.
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Heat of Formation (∆H°f)
• The heat change that occurs when one mole of a compound is formed from its elements at 1 atm pressure.
• Generally, the standard enthalpy of formation for any element in its most stable form is 0. (i.e. O2 gas would have a standard enthalpy of 0.)
• Remember Appendix C!
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Standard Enthalpy Changes
• The standard enthalpy change can be calculated from the standard enthalpies of formation of the reactants and products in the reaction (see Appendix C for values.)
• ∆H°rxn = Σ n∆H°f (products) - Σm ∆H°f (reactants)
• The n and m refer to the molar coefficients in the chemical equation.
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Also refer to Appendix C!
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Hess’s Law
• If you can break a chemical reaction into several steps, add up all of the ∆H’s for each step to get the overall ∆H for the reaction.
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Entropy• Entropy is a measure of randomness or
disorder of a system. The greater the disorder, the greater the entropy.
• In terms of entropy, gases>liquids>solids.• When pure substance dissolves in a liquid, its entropy
increases.• When gas molecules escape a solvent, entropy increases.• Entropy increases with molecular complexity.• Reactions that increase the number of moles of particles
often increase the entropy of the system.
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Predict!
• Na+ (aq) + Cl- (aq) → NaCl (s) ∆S is negative
• NH4Cl (s) → NH3 (g) + HCl (g) ∆S is positive