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Chemistry Exam 3 Study Guide

by: Simrat Kaur

Chemistry Exam 3 Study Guide CHEM 1332

Marketplace > University of Houston > CHEM 1332 > Chemistry Exam 3 Study Guide
Simrat Kaur

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This study guide covers the material that will be on exam 3.
Fundamentals of Chemistry 2
Simon Bott
Study Guide
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This 6 page Study Guide was uploaded by Simrat Kaur on Thursday April 21, 2016. The Study Guide belongs to CHEM 1332 at University of Houston taught by Simon Bott in Spring2015. Since its upload, it has received 164 views.

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Date Created: 04/21/16
Chapter 15- Solubility  K sp[A] aeq[B]beq  Q < K sp o Reaction goes to the right o Can dissolve more solid and add more ions o Unsaturated  Q = K sp o Reaction is at equilibrium o Cannot dissolve more solid and add more ions o Saturated  Q > K sp o Reaction goes to left o Precipitate in contact with saturated solution Chapter 16- Thermodynamics  First Law of Thermodynamics: ΔH o Heat of Reaction  Exothermic reactions:  Give off heat and warm up the surroundings  Ex) Combustion reactions  -ΔH  Heat of vaporization  Endothermic reactions:  Give off “cold” and cool down the surroundings, absorbs heat.  Ex) cold packs  +ΔH  Heat of fusion o Heat of formation  Standard State  Pure solid or liquid  1M solutions  Gases at 1 atm  Usually 25°C  ΔH° (x) kj/mol f  H 2 N2, O2, 2 , 2l → gas  Br2→ liquid  I2, 4 , 4s ,8S , C (graphite) → solid  Metals are solid except Hg (l)Cs and Ga], noble gases o Hess’s Law  State function  Change independent of path  Reversible change o A + B → C ΔH 1 o C → D ΔH 2 o A + B → D ΔH 1 ΔH 2 o A + B → C + D ΔH 1 o C + D → A + B -ΔH 1  Flip equation → change sign  Multiply equation, multiply ΔH  Add equations, add ΔH  ΔH reaction:  ∑{ΔH° fproducts)} - ∑{ΔH° (feactants)}  ΔH° felement) = 0  Second Law of Thermodynamics: ΔS to ΔG o Spontaneous reactions take place without continual input of energy o Entropy of Universe, Suniv  ΔSuniv> 0 → spontaneous  ΔSuniv< 0 → nonspontaneous  ΔSuniverse ΔSsystem ΔS surroundings  ΔSsurr ΔH rxn ΔG rxn = ΔH rxn+ TΔS rxn  ΔGrxn< 0 → spontaneous  ΔGrxn> 0 → nonspontaneous o ΔG (kj)  Free energy of reaction  + endergonic → nonspontaneous  - exergonic → spontaneous  State equation  ∑{ΔG° fproducts)} - ∑{ΔG° (feactants)}  ΔG°f(element) = 0 o Temperature of Spontaneity change  Set ΔG = 0  T = ΔH°/ΔS°  At equilibrium, ΔG = 0 o Phase change entropy  Boiling point  Before → nonspontaneous  After → Spontaneous  At BP → ΔG = 0  BP = ΔH° vapΔS° vap  Third Law of Thermodynamics: Entropy ΔS o S of a substance cannot be zero, except on a perfect crystal at 0 K. o Dispersal of energy over the system; randomness or disorder o Units:  J/mol k o Phase  S < L < G o Pure vs. Mixture  Pure < Mixture o Temperature  Increase T, increase S o Pressure  Increase P, decrease S o Molecule Size  Increase size, increase S o Positive if S increases o Negative if S decreases o ΔS = ∑{ΔS°(products)} - ∑{ΔS°(reactants)} o ΔS (element) ≠ 0  Equilibrium and ΔG o Reaction line  A + B C + D  ΔG  Magnitude: energy change to get to equilibrium  Sign: direction to get to equilibrium o ΔG ∝ ln(Q/K) o ΔG = RTln(Q/K) o ΔG = RTln(Q) - RTln(K) o At Standard state:  Q = 1  ΔG° = -RTln(K) Chapter 17- Redox  Oxidation and Reduction o Oxidation  Gain oxygen  Lose hydrogen  Lose electrons o Reduction  Lose oxygen  Gain hydrogen  Gain electrons  Ionic redox reactions o Identify what is being oxidized and reduced o Balance equation → must have equal numbers of electrons lost and gained  Non-ionic redox reactions 1. Assign oxidation numbers (fake charges) Metal cations → charge on cations o Alkali = +1 o Alkaline earth = +2 o Al, Ga = +3 Non-metals o F = -1 o O = -2 o H = +1 If you still cannot assign oxidation numbers, assign appropriate negative charge to unassigned atom closest to F. 2. Identify what is being oxidized and reduced 3. Balance electron transfer Balance redox atoms using appropriate coefficients on whole species  Reducing agent o Loses electrons and is oxidized in the reaction.  Oxidizing agent o Gains electrons and is reduced in the reaction.  Half-equations 1. Eliminate spectator ions 2. Identify what is being oxidized and reduced 3. Balance electrons  In acidic/basic solution 1. Assign oxidation numbers 2. Identify if being oxidized or reduced 3. Add electrons on appropriate side 4. Balance charges with H (if in acidic solution) or OH (if in basic solution) 5. Balance number of oxygen atoms 6. Check number of hydrogen atoms  Potential o Positive potential is good o +E° red→ better oxidizing agent o +E° →oxetter reducing agent  Cell potential (V) o E° = E° + E° cell red ox  Galvanic Cells 2+ 2+ o Ex) Zn +(s) (aq)→ Zn (aq)Cu (s) o Zn |(s) 2+(aq)| Cu2+(aq)Cu (s)  E and ΔG o Potential  ΔG° = -nFE°  F (Faraday) = 96500 J/mol e V -  E° = (RT/nF)lnk  E = E° - (RT/nF)lnQ  Electrolytes  Electrolytic Cell o Metal → cation + electron(s)  Oxidation o Nonmetal + electron(s) → anion  Reduction o Metal + nonmetal → ionic compound  Spontaneous redox equation o Cation + electron(s) → metal  Reduction o Anion → nonmetal + electron(s)  Oxidation o Ionic compound → metal + nonmetal  Nonspontaneous redox equation  Need continual input of energy  Molten ionic compounds o Cathode: cation → element o Anode: anion → element o Ex) NaCl (l)  Aqueous ionic compounds o Cathode:  Cation → element  2H 2 + 2e → H + 2OH (aq) o Anode:  Anion → element - +  2H 2 → O +24e + 4H (aq) o Ex) NaCl (aq)  Equations: o Charge (C) = current (amps) x time (sec) o 1 mol e = 96500 C = 1 Faraday


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