Chemistry Equation Sheet

DAT General Chemistry Equation Sheet Chapter 0: General and Lab Concepts Review Dilutions Percent Error Absorbance (Spectrophotometer) 𝑀1 𝑉1 = 𝑀2 𝑉2 or 𝐶1 𝑉1 = 𝐶2 𝑉2 (𝐴 − 𝑇) × 100 𝑇 𝐴𝑏𝑠 = 𝜀𝑐𝑙 Absorption/ Emission Line Spectra Kinetic Energy of an electron (Photoelectric Effect) 𝐸𝑝ℎ𝑜𝑡𝑜𝑛 = ℎ𝑓 = Average Kinetic Energy 𝐾𝐸𝑎𝑣𝑔 = 3⁄2 𝑅𝑇 𝜀 = 𝑚𝑜𝑙𝑎𝑟 𝑒𝑥𝑡𝑖𝑛𝑐𝑡𝑖𝑜𝑛 𝑐𝑜𝑒𝑓𝑓𝑖𝑐𝑖𝑒𝑛𝑡 (𝑚𝑜𝑙𝑎𝑟 𝑎𝑏𝑠𝑜𝑟𝑝𝑡𝑖𝑣𝑖𝑡𝑦) 𝑐 = 𝑠𝑎𝑚𝑝𝑙𝑒 ′ 𝑠 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 𝑙 = 𝑝𝑎𝑡ℎ 𝑙𝑒𝑛𝑔𝑡ℎ Root-MeanSquare Speed (𝒗) Ideal Gas Law 3𝑅𝑇 𝑣= √ 𝑀𝑚 𝑅 = 8.314 𝑃𝑉 = 𝑛𝑅𝑇 𝑛 = # 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝐿 ⋅ 𝑎𝑡𝑚 𝑅 = 0.0821 𝑚𝑜𝑙 ⋅ 𝐾 𝑓 = 𝑝ℎ𝑜𝑡𝑜𝑛′ 𝑠 𝑓𝑟𝑒𝑞𝑢𝑒𝑛𝑐𝑦 𝑐 = 𝑠𝑝𝑒𝑒𝑑 𝑜𝑓 𝑙𝑖𝑔ℎ𝑡 (3.0 × 108 𝑚⁄𝑠) 𝜆 = 𝑝ℎ𝑜𝑡𝑜𝑛′ 𝑠 𝑤𝑎𝑣𝑒𝑙𝑒𝑛𝑔𝑡ℎ 𝛥𝐸 = 𝐸𝑝ℎ𝑜𝑡𝑜𝑛 𝐾𝐸𝑒− = 𝐸𝑝ℎ𝑜𝑡𝑜𝑛 − 𝜙 𝜙 = 𝑤𝑜𝑟𝑘 𝑓𝑢𝑛𝑐𝑡𝑖𝑜𝑛 (𝑚𝑖𝑛𝑖𝑚𝑢𝑚 𝑒𝑛𝑒𝑟𝑔𝑦 𝑛𝑒𝑒𝑑𝑒𝑑 𝑡𝑜 𝑖𝑜𝑛𝑖𝑧𝑒 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛) Molality Henry’s Law Freezing Point Depression Boiling Point Elevation Vapor Pressure Depression (Raoult’s Law) Osmotic Pressure (𝝅) 𝑃= Boyle’s Law 𝐴 = 𝑎𝑟𝑒𝑎 𝑅 = 8.314 𝐽 𝑚𝑜𝑙 ∙ 𝐾 𝐽 𝑚𝑜𝑙 ∙ 𝐾 𝑀𝑚 = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 1 𝑃 𝑉∝𝑇 𝑉∝𝑛 𝑃1 𝑉1 𝑃2 𝑉2 = 𝑛1 𝑇1 𝑛2 𝑇2 P=1 atm T=273 K 𝑉∝ Charles’ Law Avogadro’s Law Combined Gas Law Standard Temp. & Pressure (STP) Standard Conditions *1 mol of gas = 22.4 L at STP All aqueous species @ 1M All gaseous species @1 atm T=298 K 𝑃(𝑀𝑀) 𝑚 = 𝑅𝑇 𝑣 Density 𝑀𝑀 = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝐿 ⋅ 𝑎𝑡𝑚 𝑅 = 0.0821 𝑚𝑜𝑙 ⋅ 𝐾 𝑚 = 𝑚𝑎𝑠𝑠 𝑣 = 𝑣𝑜𝑙𝑢𝑚𝑒 Chapter 7: Chemical Solutions Molarity 𝐹 = 𝑓𝑜𝑟𝑐𝑒 𝑇 = 𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝐴 = 𝑎𝑐𝑡𝑢𝑎𝑙 ℎ = 𝑃𝑙𝑎𝑛𝑐𝑘 ′𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 (6.63 × 10−34𝐽 ∙ 𝑠) ℎ𝑐 𝜆 𝐹 𝐴 Pressure Chapter 2: Atomic and Electronic Structure Energy of a photon Chapter 5: Gases 𝑀 𝑜𝑟 𝐶 = 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 𝑉 = 𝑣𝑜𝑙𝑢𝑚𝑒 𝑚𝑜𝑙𝑒𝑠𝑠𝑜𝑙𝑢𝑡𝑒 𝑀= 𝐿𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑚𝑜𝑙𝑒𝑠𝑠𝑜𝑙𝑢𝑡𝑒 𝑚= 𝑘𝑔𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑃𝐴 = 𝑘𝐻 [𝐴] 𝛥𝑇𝐹 = −𝑖𝐾𝐹 𝑚 𝛥𝑇𝐵 = 𝑖𝐾𝐵 𝑚 𝑃𝑠𝑜𝑙𝑛 = 𝜒𝑠𝑜𝑙𝑣 𝑃𝑠𝑜𝑙𝑣 0 𝜋 = 𝑖𝑀𝑅𝑇 𝑃𝐴 = 𝑝𝑎𝑟𝑡𝑖𝑎𝑙 𝑝𝑟𝑒𝑠𝑠𝑢𝑟𝑒 𝑜𝑓 𝑔𝑎𝑠 𝐴 𝑘𝐻 = 𝐻𝑒𝑛𝑟𝑦 ′ 𝑠 𝐿𝑎𝑤 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 (𝑣𝑎𝑟𝑖𝑒𝑠 𝑝𝑒𝑟 𝑝𝑟𝑜𝑏𝑙𝑒𝑚) [A] = conc. of gas A 𝑖 = 𝑣𝑎𝑛′ 𝑡 𝐻𝑜𝑓𝑓 𝑓𝑎𝑐𝑡𝑜𝑟 𝐾𝐹 = 𝐹. 𝑃. 𝑑𝑒𝑝𝑟𝑒𝑠𝑠𝑖𝑜𝑛 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑚 = 𝑚𝑜𝑙𝑎𝑙𝑖𝑡𝑦 𝑖 = 𝑣𝑎𝑛′ 𝑡 𝐻𝑜𝑓𝑓 𝑓𝑎𝑐𝑡𝑜𝑟 𝐾𝐵 = 𝐵. 𝑃. 𝑒𝑙𝑒𝑣𝑎𝑡𝑖𝑜𝑛 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑚 = 𝑚𝑜𝑙𝑎𝑙𝑖𝑡𝑦 𝑃𝑠𝑜𝑙𝑛 = 𝑉𝑃 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝜒𝑠𝑜𝑙𝑣 = 𝑚𝑜𝑙 𝑓𝑟𝑎𝑐𝑡 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑃𝑠𝑜𝑙𝑣 0 = 𝑉𝑃 𝑜𝑓 𝑝𝑢𝑟𝑒 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 𝑀 = 𝑚𝑜𝑙𝑎𝑟𝑖𝑡𝑦 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑖 = 𝑣𝑎𝑛′ 𝑡 𝐻𝑜𝑓𝑓 𝑓𝑎𝑐𝑡𝑜𝑟 𝐿 ⋅ 𝑎𝑡𝑚 𝑅 = 0.0821 𝑚𝑜𝑙 ⋅ 𝐾 𝑇 = 𝑡𝑒𝑚𝑝. 𝑖𝑛 𝐾𝑒𝑙𝑣𝑖𝑛 Dalton’s Law of Partial Pressures 𝑃𝑡𝑜𝑡𝑎𝑙 = 𝑃𝐴 + 𝑃𝐵 + ⋯ Dalton’s Law of Partial Pressures Graham’s Law of Effusion 𝑃𝐴 = 𝜒𝐴 𝑃𝑡𝑜𝑡𝑎𝑙 Real Gas Equation 𝑟1 𝑀𝑚2 =√ 𝑟2 𝑀𝑚1 𝑎𝑛2 )(𝑉 𝑉2 − 𝑛𝑏) = 𝑛𝑅𝑇 (𝑃 + 𝜒𝐴 = 𝑚𝑜𝑙 𝑓𝑟𝑎𝑐𝑡𝑖𝑜𝑛 𝑜𝑓 𝑔𝑎𝑠 𝐴 𝑟 = 𝑟𝑎𝑡𝑒 𝑜𝑓 𝑒𝑓𝑓𝑢𝑠𝑖𝑜𝑛 𝑀 = 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠 𝑎 & 𝑏 = 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡𝑠 𝑠𝑝𝑒𝑐𝑖𝑓𝑖𝑐 𝑡𝑜 𝑒𝑎𝑐ℎ 𝑔𝑎𝑠 𝑎𝑛2 𝑐𝑜𝑟𝑟𝑒𝑐𝑡𝑠 𝑓𝑜𝑟 𝑉2 𝐼𝑀𝐹𝑠 −𝑛𝑏 𝑐𝑜𝑟𝑟𝑒𝑐𝑡𝑠 𝑓𝑜𝑟 𝑣𝑜𝑙𝑢𝑚𝑒 Chapter 8: Chemical Kinetics General Rate Law A+B→C+D 𝑟𝑎𝑡𝑒 = 𝑘[𝐴]𝑚 [𝐵]𝑛 Rate Constant Units 0 𝑜𝑟𝑑𝑒𝑟: 𝑘 = 𝑀1 ∙ 𝑠 −1 1𝑠𝑡 𝑜𝑟𝑑𝑒𝑟: 𝑘 = 𝑠 −1 2𝑛𝑑 𝑜𝑟𝑑𝑒𝑟: 𝑘 = 𝑀 −1 ∙ 𝑠 −1 3𝑟𝑑 𝑜𝑟𝑑𝑒𝑟: 𝑘 = 𝑀 −2 ∙ 𝑠 −1 Arrhenius 𝑘 = 𝐴𝑒 −𝐸𝑎 ⁄𝑅𝑇 Equation 𝑘 = 𝑟𝑎𝑡𝑒 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑚 & 𝑛 = 𝑑𝑒𝑡𝑒𝑟𝑚𝑖𝑛𝑒𝑑 e𝑥𝑝𝑒𝑟𝑖𝑚𝑒𝑛𝑡𝑎𝑙𝑙𝑦 𝑘 = 𝑟𝑎𝑡𝑒 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑀 = 𝑚𝑜𝑙𝑎𝑟𝑖𝑡𝑦 s= 𝑠𝑒𝑐𝑜𝑛𝑑𝑠 𝑘 = 𝑟𝑎𝑡𝑒 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝐴 = 𝑢𝑛𝑖𝑞𝑢𝑒 𝑡𝑜 𝑒𝑎𝑐ℎ 𝑟𝑥𝑛 𝐸𝑎 = 𝑎𝑐𝑡. 𝑒𝑛𝑒𝑟𝑔𝑦 𝐽 𝑅 = 8.314 𝑚𝑜𝑙 ∙ 𝐾 𝑇 = 𝑡𝑒𝑚𝑝. 𝑖𝑛 𝐾𝑒𝑙𝑣𝑖𝑛 Bootcamp.com DAT General Chemistry Equation Sheet Chapter 9: Chemical Equilibria Equilibrium Constant Expressions 𝐾𝑐 = Solubility Product Constant (𝑲𝒔𝒑 ) 𝑘𝑓𝑜𝑟𝑤𝑎𝑟𝑑 𝑘𝑟𝑒𝑣𝑒𝑟𝑠𝑒 𝑃𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 𝐾𝑃 = 𝑄= 𝑘 = 𝑟𝑎𝑡𝑒 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑃 = 𝑝𝑟𝑒𝑠𝑠𝑢𝑟𝑒 [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠] 𝐾𝑒𝑞 = Reaction Quotient (Q) Chapter 11: Thermodynamics & Thermochemistry [𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠] 𝑃𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 [𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠] [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠] 𝐾𝑠𝑝 = 𝑄 > 𝐾 = 𝑠ℎ𝑖𝑓𝑡 𝑙𝑒𝑓𝑡 𝑄 < 𝐾 = 𝑠ℎ𝑖𝑓𝑡 𝑟𝑖𝑔ℎ𝑡 𝑄 = 𝐾 = 𝑒𝑞𝑢𝑖𝑙𝑖𝑏𝑟𝑖𝑢𝑚 [𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠] [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠] Chapter 10: Acid-Base Equilbria & Titrations Ionization Constant of Water pH & pOH 𝐾𝑤 = [𝐻3 𝑂+ ][ 𝑂𝐻− ] = 1 × 10−14 @ 25 °C 𝑝𝐻 = −log [𝐻+ ] 𝑝𝑂𝐻 = −log [𝑂𝐻− ] 𝑝𝐻 + 𝑝𝑂𝐻 = 14 [𝐻+ ] = 10−𝑝𝐻 [𝑂𝐻− ] = 10−𝑝𝑂𝐻 [𝐻+ ][𝑂𝐻− ] = 1 ∗ 10−14 [𝐻+ ] = 𝑐𝑜𝑛𝑐. 𝑜𝑓 𝑝𝑟𝑜𝑡𝑜𝑛𝑠 [𝑂𝐻− ] = 𝑐𝑜𝑛𝑐. 𝑜𝑓 ℎ𝑦𝑑𝑟𝑜𝑥𝑖𝑑𝑒 Weak Acids 𝐻𝐴 + 𝐻2 𝑂 ⇌ 𝐻3 𝑂+ + 𝐴− [𝐻3 𝑂+][𝐴−] 𝐾𝑎 = [𝐻𝐴] [𝐻+] = √𝐾𝑎 [𝐻𝐴] 𝐾𝑎 = 𝑎𝑐𝑖𝑑 𝑑𝑖𝑠𝑠𝑜𝑐𝑖𝑎𝑡𝑖𝑜𝑛 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 Weak Bases 𝐴− + 𝐻2 𝑂 ⇌ 𝐻𝐴 + 𝑂𝐻− [𝑂𝐻− ][𝐻𝐴] 𝐾𝑏 = [𝐴− ] − [𝑂𝐻 ] = √𝐾𝑏 [𝐴− ] 𝑝𝐾𝑎 = −log [𝐾𝑎 ] 𝑝𝐾𝑏 = −log [𝐾𝑏 ] 𝑝𝐾𝑎 + 𝑝𝐾𝑏 = 14 𝐾𝑏 = 𝑏𝑎𝑠𝑒 𝑑𝑖𝑠𝑠𝑜𝑐𝑖𝑎𝑡𝑖𝑜𝑛 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 [H+] & [OH-] pKa & pKb 𝐾𝑤 = 𝐾𝑎 × 𝐾𝑏 = 1 ∗ 10−14 Neutralization Reaction Buffers 𝑛𝐴 𝑀𝐴 𝑉𝐴 = 𝑛𝐵 𝑀𝐵 𝑉𝐵 𝑝𝐻 = 𝑝𝐾𝑎 + log [𝐴− ] [𝐻𝐴] Chapter 13: Nuclear Reactions Kinetics 𝑁 = 𝑁0 𝑒 −𝑘𝑡 (always 1st ln 𝑁 = ln 𝑁 − 𝑘𝑡 0 order) 0.693 𝑡1/2 = 𝑘 Nuclear Binding Energy 𝐸 = 𝛥𝑚𝑐 2 Larger 𝐾𝑎 = smaller p𝐾𝑎 =stronger acid Larger 𝐾𝑏 = smaller p𝐾𝑏 =stronger base 𝑛𝐴 = # 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝐻 + 𝑛𝐵 = # 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠 𝑂𝐻− [𝐴− ] = 𝑐𝑜𝑛𝑐. 𝑜𝑓 𝑏𝑎𝑠𝑒 [𝐻𝐴] = 𝑐𝑜𝑛𝑐. 𝑜𝑓 𝑎𝑐𝑖𝑑 𝑁 = 𝑎𝑚𝑡 𝑜𝑓 𝑟𝑎𝑑𝑖𝑜𝑖𝑠𝑜𝑡𝑜𝑝𝑒 𝑎𝑓𝑡𝑒𝑟 𝑡𝑖𝑚𝑒 𝑡 𝑁0 = 𝑖𝑛𝑖𝑡𝑖𝑎𝑙 𝑎𝑚𝑡 𝑘 = 𝑟𝑎𝑡𝑒 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑡 = 𝑡𝑖𝑚𝑒 𝑡1/2 = ℎ𝑎𝑙𝑓 𝑙𝑖𝑓𝑒 *note that 𝑡1/2 is independent of concn for 1st order rxns 𝑚 = 𝑚𝑎𝑠𝑠 (𝑀𝑈𝑆𝑇 𝑏𝑒 𝑖𝑛 𝑘𝑔) 𝑐 = 𝑠𝑝𝑒𝑒𝑑 𝑜𝑓 𝑙𝑖𝑔ℎ𝑡 𝑚 (3.0 × 108 ) 𝑠 Enthalpy (H) (𝛥𝐻 > 0): 𝑒𝑛𝑑𝑜𝑡ℎ𝑒𝑟𝑚𝑖𝑐 (𝛥𝐻 < 0): 𝑒𝑥𝑜𝑡ℎ𝑒𝑟𝑚𝑖𝑐 Enthalpy of Formation 𝛥𝐻𝑓 = 𝛴𝑛𝛥𝐻°𝑓(𝑝𝑟𝑜𝑑𝑢𝑐𝑡) − 𝛴𝑛𝛥𝐻°𝑓(𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠) First Law of Thermodynamics PressureVolume Work Calorimetry Thermal Energy (q) Heat Curves & Thermal Energy (q) Entropy (S) Bond Dissociation Energy Gibb’s Free Energy (𝜟𝑮) Gibb’s Free Energy (𝜟𝑮) 𝑛 = 𝑐𝑜𝑒𝑓𝑓𝑖𝑐𝑖𝑒𝑛𝑡 𝑓𝑟𝑜𝑚 𝑏𝑎𝑙𝑎𝑛𝑐𝑒𝑑 𝑟𝑥𝑛 𝛥𝐸 = 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑖𝑛𝑡𝑒𝑟𝑛𝑎𝑙 𝛥𝐸 = 𝑞 + 𝑤 𝑒𝑛𝑒𝑟𝑔𝑦 𝑞 = ℎ𝑒𝑎𝑡 𝑤 = 𝑤𝑜𝑟𝑘 𝑃 = 𝑒𝑥𝑡𝑒𝑟𝑛𝑎𝑙 𝑝𝑟𝑒𝑠𝑠𝑢𝑟𝑒 𝑤 = −𝑃𝛥𝑉 𝛥𝑉 = 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑣𝑜𝑙𝑢𝑚𝑒 𝑞 = −𝐶𝑐𝑎𝑙𝑜𝑟𝑖𝑚𝑒𝑡𝑒𝑟 𝛥𝑇 𝐶𝑐𝑎𝑙𝑜𝑟𝑖𝑚𝑒𝑡𝑒𝑟 = 𝑠𝑝𝑒𝑐𝑖𝑓𝑖𝑐 ℎ𝑒𝑎𝑡 𝑜𝑓 𝑐𝑎𝑙𝑜𝑟𝑖𝑚𝑒𝑡𝑒𝑟 𝛥𝑇 = 𝑐ℎ𝑎𝑛𝑔𝑒 𝑖𝑛 𝑡𝑒𝑚𝑝. +𝑞: ℎ𝑒𝑎𝑡 𝑔𝑎𝑖𝑛𝑒𝑑 𝑏𝑦 𝑠𝑦𝑠𝑡𝑒𝑚 𝑞 = 𝑚𝐶𝛥𝑇 −𝑞: ℎ𝑒𝑎𝑡 𝑙𝑜𝑠𝑡 𝑓𝑟𝑜𝑚 𝑠𝑦𝑠𝑡𝑒𝑚 𝑞 = 𝑚𝛥𝐻𝑓𝑢𝑠𝑖𝑜𝑛 𝑚 = 𝑚𝑎𝑠𝑠 𝑞 = 𝑚𝛥𝐻𝑣𝑎𝑝𝑜𝑟𝑖𝑧𝑎𝑡𝑖𝑜𝑛 𝐶 = 𝑠𝑝𝑒𝑐𝑖𝑓𝑖𝑐 ℎ𝑒𝑎𝑡 𝛥𝑆 = Σ𝑛𝑆𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 − Σ𝑛𝑆𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 𝑆𝑔𝑎𝑠 > 𝑆𝑙𝑖𝑞𝑢𝑖𝑑 > 𝑆𝑠𝑜𝑙𝑖𝑑 𝑆𝑎𝑞 > 𝑆𝑠𝑜𝑙𝑖𝑑 𝛥𝐻 = 𝛴𝛥𝐻𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 − 𝛴𝛥𝐻𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 = 𝛴𝛥𝐻𝑏𝑟𝑜𝑘𝑒𝑛 − 𝛴𝛥𝐻𝑓𝑜𝑟𝑚𝑒𝑑 𝑚𝑎𝑘𝑖𝑛𝑔 𝑏𝑜𝑛𝑑𝑠 = 𝑒𝑥𝑜𝑡ℎ𝑒𝑟𝑚𝑖𝑐 (−𝛥𝐻) 𝑏𝑟𝑒𝑎𝑘𝑖𝑛𝑔 𝑏𝑜𝑛𝑑𝑠 = 𝑒𝑛𝑑𝑜𝑡ℎ𝑒𝑟𝑚𝑖𝑐 (+𝛥𝐻) 𝛥𝐺° = 𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 𝑐𝑜𝑛𝑑𝑖𝑡𝑖𝑜𝑛𝑠 𝛥𝐺° = 𝛥𝐻° − 𝑇𝛥𝑆° 𝛥𝐻° = 𝑒𝑛𝑡ℎ𝑎𝑙𝑝𝑦 𝑇 = 𝑡𝑒𝑚𝑝. 𝑖𝑛 𝐾𝑒𝑙𝑣𝑖𝑛 𝛥𝑆° = 𝑒𝑛𝑡𝑟𝑜𝑝𝑦 𝛥𝐺 = 𝑛𝑜𝑛𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 𝛥𝐺 = 𝛥𝐺° + 𝑅𝑇𝑙𝑛𝑄 𝑐𝑜𝑛𝑑𝑖𝑡𝑖𝑜𝑛𝑠 𝛥𝐺° = −𝑅𝑇𝑙𝑛𝐾𝑒𝑞 𝐽 𝑅=8.314 𝑚𝑜𝑙∙𝐾 𝑄 = 𝑟𝑒𝑎𝑐𝑡𝑖𝑜𝑛 𝑞𝑢𝑜𝑡𝑖𝑒𝑛𝑡 𝐾𝑒𝑞 = 𝑒𝑞𝑢𝑖𝑙𝑖𝑏𝑟𝑖𝑢𝑚 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 Chapter 12: Electrochemistry & Redox Reactions Standard Cell Potential Nernst Equation 𝐸° = 𝐸°𝑟𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛 + 𝐸°𝑜𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛 𝐸° = 𝐸°𝑐𝑎𝑡ℎ𝑜𝑑𝑒 + 𝐸°𝑎𝑛𝑜𝑑𝑒 Faraday’s 𝑚𝑎𝑠𝑠 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 = 𝐼 ∗ 𝑡𝑠 ∗ 𝑀𝑊𝑝𝑑𝑡 𝑛∗𝐹 Law 𝐸𝑐𝑒𝑙𝑙 = 𝐸° − 0.0592 log 𝑄 𝑛 𝑚𝑜𝑙𝑒𝑠 𝑜𝑓 𝑝𝑟𝑜𝑑𝑢𝑐𝑡 = 𝐼 ∗ 𝑡𝑠 𝑛∗𝐹 𝐸𝑐𝑒𝑙𝑙 = 𝑛𝑜𝑛𝑠𝑡𝑎𝑛𝑑𝑎𝑟𝑑 𝑐𝑒𝑙𝑙 𝑝𝑜𝑡𝑒𝑛𝑡𝑖𝑎𝑙 𝑛 = # 𝑜𝑓 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠 t𝑟𝑎𝑛𝑠𝑓𝑒𝑟𝑟𝑒𝑑 𝑄 = 𝑟𝑒𝑎𝑐𝑡𝑖𝑜𝑛 𝑞𝑢𝑜𝑡𝑖𝑒𝑛𝑡 𝑀𝑊 = 𝑚𝑜𝑙𝑒𝑐. 𝑤𝑒𝑖𝑔ℎ𝑡 𝐼 = 𝑐𝑢𝑟𝑟𝑒𝑛𝑡 (Amps) 𝑡𝑠 = 𝑡𝑖𝑚𝑒 (𝑠𝑒𝑐𝑜𝑛𝑑𝑠) 𝑛 = # 𝑜𝑓 𝑒𝑙𝑒𝑐𝑡𝑟𝑜𝑛𝑠 𝑡𝑟𝑎𝑛𝑠𝑓𝑒𝑟𝑟𝑒𝑑 𝐹 = 𝐹𝑎𝑟𝑎𝑑𝑎𝑦 ′ 𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 𝑐𝑜𝑢𝑙𝑜𝑚𝑏𝑠 (96485 ) 𝑚𝑜𝑙 𝑒 − Bootcamp.com

Chemistry Equation Sheet

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Chemistry Equation Sheet

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