Chem7 - Thermodynamics

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mse263
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Chem7 - Thermodynamics
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2014-12-09 21:19:23
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  1. Enthalpy
  2. State Function
    • something that only depends on the initial & final state of the system, THAT’S. IT.
    • b/c that’s all it depends on, it’s “pathway independent”
  3. What 2 things should you remember that are NOT State Functions?
    • q: change in heat
    • W: change in work energy
    • these values DEPEND on the pathway taken
    • [almost everything else that’s talked about in thermochemistry - eg. ΔG, ΔH, ΔS - are state functions]
  4. Enthalpy (H)
    • like to think of it as “change in heat”, however ΔH & q aren’t exactly the same thing
    • they ARE equal if a rxn’s carried out at constant pressure (ΔH = qp)
  5. What does it mean for a rxn if ΔH is negative (ΔH < 0)?
    • it means the reaction is Exothermic (-)
    • ‘Exo’ means outside; this type of rxn takes heat FROM the system & releases it TO the surroundings
  6. What does it mean for a rxn if ΔH is positive (ΔH > 0)?
    • it means the reaction is Endothermic (+)
    • ‘Endo’ means inside; this type of rxn takes/pulls/absorbs heat from the surroundings INTO the system
  7. What are the 3 ways to calculate ΔH?
    1. Hess’ Law

    2. Bond Energies (Bond Dissociation Energies, or Bond Enthalpies): least accurate b/c avg BEs are used

    • 3. ΔH°f (Formation Enthalpies, or Heats of Formation)
    • • all 3 can be used to calculate ΔH, just depends on what you’re provided with
  8. Hess’ Law
    to calculate ΔH using Hess’ Law, a problem will give:

    1. an overall reaction

    2. individual component reactions

    3. + their standard ΔH values (in K or kJ)

    4. BALANCE the individual component rxns so what’s left is the overall rxn

    • 5. change the sign/multiply the component rxns’ ΔH°rxn then ADD THEM UP
    • • the resulting ΔH°rxn corresponds to the overall rxn
  9. Bond Energies (way #2 of finding ΔH)
    • defined as the energy to BREAK a bond

    • to find ΔH, subtract bond energies of bonds formed FROM the bond energies of the bonds broken

    • DO NOT. MAKE THE FORMED BOND ENERGIES. NEGATIVE. You already accounted for that by SUBTRACTING them from the reactant bond energies of bonds broken

    • ΔH°rxn = BEreactants – BEproducts

    • ΔH°rxn = BEbroken – BEformed

    • for bond enthalpies, reactants come 1st in the equation b/c reactant bonds are what’s being BROKEN (the definition of BE)
  10. Breaking v. Making a Bond
    • breaking a bond releases energy, ΔH = -, exothermic
    • making a bond COSTS energy, ΔH = +, endothermic
  11. ΔH°f (Formation Enthalpies, or Heats of Formation)
    • ΔH°rxn = ΣnΔH°f, products – ΣnΔH°f, reactants

    • for formation enthalpies, products come 1st in the equation b/c products are being FORMED

    • in a ΔH°f problem, e/a product & reactant ΔH°f should be given so you can plug into eqn & solve for ΔH°rxn
  12. What is the ΔH°f for any element in its standard state [eg. Cl2(g)]?
    ΔH°f = ZERO
  13. Diatomics
    • elements that when combined w/ one other atom of themselves form diatomic gasses w/ a standard state = 0 (ΔH°f = 0)
    • Never Have Fear Of Ice Cold Beer
    • Nitrogen, Hydrogen, Fluorine, Oxygen, Iodine, Chlorine, Bromine
  14. What are the only 2 elements on the periodic table that are LIQUID in their standard state?
    • 1. Mercury (Hg)
    • 2. Bromine (Br) [can also be a gas in its standard state]
  15. A Formation Reaction:
    • forms only ONE mole of a single product
    • has a product formed from ONLY elements (NO compounds)
    • has all elements in their standard states
    • eg. 1/2H2(g) + 1/2Cl2(g) → HCl(g)
  16. Standard Conditions (°)
    • Pressure: 1 atm
    • Concentrationaq: 1 mole
    • Concentrationg: 1 atm
    • Temperature: 298 K (~25° C, room temp.)
    • used to talk about reactions in the context of thermodynamics; different from STP (T: 273 K, P: 1 atm) which is used for gasses mostly
  17. Calorimetry
  18. Calorimetry
    the measurement of heat changes, often associated with chemical reactions
  19. What’s the difference between Calorie & calorie?
    • calorie: the amount of HEAT it takes to raise 1g of H2O 1° Celcius; also the same as 4.184 J (1 cal = 4.184 J)

    • Calorie: seen on food labels

    • 1 Calorie = 1000 calories
  20. q = CΔT
    • q: amount of heat released or absorbed during a reaction (in Joules)
    • C: heat capacity of the calorimeter
    • ΔT: change in temperature
    • *can turn a temperature change into a heat change if you know the heat capacity of the calorimeter
  21. q = mcΔT
    • q: heat of a rxn (in Joules)
    • m: mass (in GRAMS)
    • c: specific heat or specific heat capacity of a substance (in J/g°C or K)
    • eg. cH2O(l) = 4.184 J/g°C
    • ΔT: the CHANGE will be the same whether K or °C is used
  22. In a graph of q (heat, x-axis) v. T (temperature, y-axis), does the temperature change DURING a PHASE change (from solid to liq, liq to gas, etc.)?
    • NO - looks like a flat line w/ a slope of 0
    • the temperature stays constant b/c the energy is used to break intermolecular forces (for a covalently bonded compounds) & ionic bonds (if the compound is ionic)
  23. Why can’t the equation q = mcΔT be used to determine the heat associated with a phase change in/during a reaction?
    •because there is NO change in temperature (ΔT = 0) during a phase change; heat is used to break IMFs, not change the temperature of a reaction

    • instead use nΔH°f or mΔH°f (given, & depending on units of ΔH°f whether moles or mass should be used) to find the heat associated w/ a phase change
  24. ΔH°f
    • change in heat associated w/ turning a solid into a liquid (fusion = melting)
    • ΔH°vap = change in heat associated w/ turning a liquid into a gas
    • ΔH given will depend on phase change; change sign of value if going in the reverse direction
  25. Entropy
  26. Entropy (S)
    directly correlated to randomness/disorder; when Entropy increases, so do randomness/disorder
  27. Rank Phases based on their Entropy:
    3. Solid: least entropy; molecules are stationary (perhaps vibrating slightly) in crystal structures

    2. Liquid: intermediate entropy

    1. GAS: most entropy & SIGNIFICANTLY more than most corresponding liquids or solids
  28. What do you want to focus on when looking to see if the change in Entropy of a rxn is +, -, or close to 0?
    want to focus on the GAS
  29. When given the reaction 2C(s) + O2(g) → 2CO(g), what is ΔS?
    1. ignore the solid (C(s)) - gases have way more energy in comparison

    2. notice this rxn has 1 mole of gas turning into 2 moles of gas

    • * this rxn would definitely have a +ΔS (ΔS>0)
    • there would be an increase in disorder
  30. What happens if a reaction has no gaseous components?
    • then add up the moles of all reactants & the moles of all products
    • whichever there are MORE MOLES of, that’s the side w/ more Entropy
    • the more molecules there are, the more Entropy
  31. Vapor Deposition
    • when a gas turns into a solid (g → s)
    • represents a DECREASE in Entropy (1 mole of g → 0 moles of g, OR just know that solids have less entropy than gasses)
  32. If you dissolve something that is fairly soluble in water there will be a(n) _________ in entropy.
    • Entropy will INCREASE when a compound that’s fairly water soluble is dissolved in water.
    • Entropy/order will DECREASE if the compound is NOT soluble in water (eg. oil droplets, anything non-polar, something that induces water to form a solvent cage, etc.)
  33. Formula for a Reaction’s Change in Entropy
    • products – reactants
    • (same as the formula using heats of formation ΔH°f where products go 1st b/c they’re what’re fORMED)
    • ΔS°rxn = ΣnΔS°products – ΣnΔS°reactants
    • subtract the entropy the reaction started with from the entropy the reaction ended with
  34. Gibbs Free Energy
  35. Gibbs Free Energy
    • thermodynamic property that tells you whether a reaction is SPONTANEOUS or not
    • when ΔG < 0 (when ΔG is negative) the rxn is spontaneous
    • when ΔG > 0 (when ΔG is positive) the rxn is nonspontaneous OR spontaneous in the reverse direction
    • when ΔG = 0 the rxn is at equilibrium
  36. Spontaneous Definition
    • if a reaction is spontaneous, it occurs without any “help” from the surroundings
    • has nothing to do with reaction speed
  37. How to Calculate ΔG
    ΔG°rxn = ΣnΔG°f, products – ΣnΔG°f, reactants
  38. ΔG = ΔG° + RT*lnQ
    • ΔG: Gibbs free energy
    • ΔG°: ΔG standard
    • R: 8.314 J/mol*K
    • T: temperature
    • RT*lnQ: takes into account any concentrations that aren’t 1 molar (i.e. standard conditions like in ΔG°)
    • *when the rxn reaches equilibrium, it’s the NONstandard value of ΔG (NOT ΔG°) that’s equal to 0
  39. The above Formula at Equilibrium equals:
    • ΔG° = –RT*lnK
    • b/c ΔG = 0 & at equilibrium Q = K
    • when ΔG° < 0, K > 1 → products favored
    • when ΔG° > 0, K < 1 → reactants favored
  40. Natural Log Logistics
    • when you take the natural log (ln) of a number greater than 1 (eg. K > 1) → positive #
    • when you take the natural log (ln) of a number less than 1 (eg. K < 1) → negative #
    • that’s the basis of how the sign of ΔG° is determined if the value of the equilibrium constant K is known
  41. ΔG = ΔH – TΔS
    • “Get Higher Test Scores”

    • ΔG & ΔH (both in Joules) are typically much bigger than ΔS (Joules/Kelvin), therefore they’re often given in kJ while ΔS is given in Joules

    • 1000 J = 1 kJ
  42. ΔG = ΔH – TΔS at Equilibrium
    • ΔG = 0 → 0 = ΔH – TΔS → ΔS = ΔH/T

    • using this relationship is often how the temperature of a PHASE CHANGE is calculated
  43. ΔH
    • refers to the transfer of enthalpy & at a constant pressure, we can think of that as heat

    • the universe PREFERS exothermic reactions b/c they lower the energy of the system; lower energy = more stable

    • universe prefers ΔH < 0
  44. ΔS
    • the universe prefers disorder

    • if given a choice it would prefer for a system to have a ΔS > 0 (positive)
  45. –ΔH, +ΔS
    • this reaction will be spontaneous at ALL temperatures (universe is given exactly what it wants)

    • –ΔH – T(+ΔS) = –ΔG

    • adding 2 negative numbers → ΔG will always be negative aka the rxn is always spontaneous
  46. +ΔH, –ΔS
    • this reaction will be NONspontaneous at all temperatures (universe is given the exact opposite of what it wants)

    • + ΔH + TΔS = +ΔG

    • adding 2 positive numbers → ΔG will always be positive aka the rxn is always nonspontaneous

    *another way to phrase this is that the rxn will be spontaneous in the REVERSE direction at all temperatures
  47. –ΔH, –ΔS
    • eg. freezing, deposition
    • so the reaction is only spontaneous at LOW temperatures
  48. +ΔH, +ΔS
    • this is what BOILING is
    • so the reaction is only spontaneous at HIGH temperatures
  49. When is this reaction spontaneous?
    2C(s) + O2(g) → 2CO(g)
    ΔH = -100 kJ
    • ΔH < 0 [negative]
    • 1 moles of gas → 2 moles of gas, so ΔS > 0
    • have –ΔH, +ΔS → therefore the rxn is spontaneous at ALL temperatures
  50. Laws of Thermodynamics
  51. Two objects at/under Thermal Equilibrium are at the same WHAT?
    Temperature
  52. 1st Law of Thermodynamics
    • Conservation of Energy: energy cannot be created or destroyed
    • but it can be converted from one form to another
  53. ΔE = q + w
    • ΔE: change in internal energy
    • q: the amount of heat transferred into or out of the system
    • w: work energy; the amount of work done ON the system by the surroundings or done BY the system ON the surroundings
    • this equation is another way of saying the 1st law of thermodynamics; if there’s an energy change in the system, it can be accounted for either by heat or by work
  54. Pressure/Volume Work (PV work)
    • w: change in work-Energy of the system (looking at this from the system’s perspective)
    • w = – PΔV
    • so ΔE = q – PΔV
  55. Expansion v. Compression of a System
    • Expansion: ΔV = positive
    • • Pressure is always positive (it can’t be negative)
    • • so w = –(+P)(+ΔV) = negative aka w < 0
    • • when a gas expands it pushes back/does work ON the surroundings, therefore the system will have LESS (–w) energy once expanded

    • Compression: ΔV = negative (volume decreases)
    • • so w = –(+P)(-ΔV) = positive aka w > 0
    • • when a gas IS compressed, the surroundings exert a force on the system, meaning the system will have MORE (+w) energy once compressed
  56. What is generally true of expanding v. compressed real, non-ideal gasses?
    • expanding gasses tend to cool down (some of their energy is used to expand, therefore heat is lost in the form of energy expenditure → temp decreases)
    • compressed gasses tend to heat up
    • *these 2 statements aren’t true for IDEAL gasses b/c ideal gasses have NO intermolecular attractive forces
  57. work-Energy aka PΔV Summary:
    • if the system does work on the surroundings, the system loses work-Energy → w < 0, negative

    • if the surroundings do work on the system, the system gains work-Energy → w > 0, positive
  58. Isochoric Process
    • a process in which the Volume remains CONSTANT (doesn’t change; ΔV = 0)
    • if heat is added to an ideal gas but there is no change in volume → both the temperature & subsequently the pressure (increase number of collisions) increase
    • also if ΔV = 0 then there’s no (PΔV) work done
  59. Isobaric
    • a process in which the Pressure remains CONSTANT (doesn’t change; ΔP = 0)
    • (iso: same, bar: used to measure pressure)
  60. Isothermal
    • a process in which there is no change in Temperature
    • the internal energy of a gas is directly proportional to its temperature; doubling the temperature doubles the average internal energy (ΔE)
    • in an isothermal rxn, ΔE = 0 b/c ΔT = 0
    • (can derive -w = q from ΔE = q – w)
  61. Adiabatic
    means q = 0, therefore ΔE = w (or – PΔV)
  62. What’s the difference between an Isothermal & Adiabatic reaction?
    • Isothermal: no temperature change
    • Adiabatic: no HEAT change when q = 0
    • for an Isothermal expansion to occur, heat needs to be added into the system b/c the gas would normally decrease in temperature during an expansion (q does not = 0)
    • an Adiabatic expansion would occur if a real gas expands & its temperature is allowed to decrease (like it logically would, ΔT does not = 0)
  63. 2nd Law of Thermodynamics
    • for a spontaneous process (including chemical reactions & phase changes), the entropy of the universe INCREASES
    • universe means system & surroundings TOTAL; entropy of BOTH increases
    • ΔSuniverse > 0
    • ΔSsystem + ΔSsurroundings > 0
    • [even if the individual ΔSsystem is negative, this law maintains that the ΔSuniverse is STILL positive]
  64. How would you know if a reaction was spontaneous?
    1. ΔG < 0 (negative)

    2. ΔSuniverse > 0 (positive - this is a new way*)

    3. ΔH – TΔS < 0 (b/c all that is = ΔG)
  65. In electrochemistry, what does it mean if a reaction’s Potential (aka Emf or Voltage) is positive?
    it means the reaction is SPONTANEOUS
  66. 3rd Law of Thermodynamics
    • “A perfect crystal at 0 Kelvin has 0 Entropy”
    • perfect means it has NO impurities & is at thermodynamic equilibrium

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