# Chem7 - Thermodynamics

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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?
• 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)
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
 Author: mse263 ID: 291185 Card Set: Chem7 - Thermodynamics Updated: 2014-12-10 02:19:23 Tags: GeneralChemistry Folders: General Chemistry Description: CS Show Answers: