Notes for General Chemistry

General Chemistry 1: Atomic Structure A Z X A = Mass number = protons + neutrons Z = Atomic number = # of protons Note: Atomic Weight = weighted average Scientist Contributions Rutherford Model: 1911. Electrons surround a nucleus. Bohr Model: 1913. Described orbits in more detail. Farther orbits = !Energy Photon emitted when n¯, absorbed when n! Heisenberg Uncertainty: It is impossible to know the momentum and position simultaneously. AHED Mnemonic Absorb light Higher potential Excited Distant from nucleus Hund’s Rule: e- only double up in orbitals if all orbitals first have 1 e-. Diamagnetic vs. Paramagnetic Pauli Exclusion Principle: Paired e- must be ! "$, % ". # # Diamagnetic: All electrons are paired REPELLED by an external magnetic field !¯ Constants Light Energy Avogadro’s Number: @ABCC$ D $ EB#F ' E$G:5 Planck’s (h): @A@C@$ D $ EBHFI J¥s Speed of Light (c) JAB$ D$ EBK G 9 &' ($) & ' *$+ l Paramagnetic: 1 or more unpaired electrons ! PULLED into an external magnetic field Follow Hund’s rule to build the atom’s electron configuration. If 1 or more orbitals have just a single electron, the atom is paramagnetic. If there are no unpaired electrons, then the atom is diamagnetic. Examples: He = 1s2 = diamagnetic and will repel magnetic fields. C = 1s22s22p2 = paramagnetic and will be attracted to magnetic fields. Quantum Numbers Quantum Number Possible Values Name What it Labels n Principal e- energy level or shell number 1, 2, 3, … l Azimuthal 3D shape of orbital 0, 1, 2, …, n-1 ml Magnetic Orbital sub-type ms Spin Electron spin Integers –l ® +l " " ! $, % # # Notes Except for d- and f-orbitals, the shell # matches the row of the periodic table. Atomic Orbitals on the Periodic Table 0 = s orbital 1 = p orbital 2 = d orbital 3 = f orbital 4 = g orbital Maximum e- in terms of n = 2n2 Maximum e- in subshell = 4l + 2 Free Radical: An atom or molecule with an unpaired electron. The Aufbau Principle 3D shapes of s, p, d, and f orbitals General Chemistry 2: The Periodic Table Alkali Metals Noble Gases Alkaline Earth Metals Halogens Non-metals Post Transition Metals Transition Metals Rare Earth Metal Rows Unchanged 8A Zeff IE Pull between nucleus & valence e- Lose e1st Ionization energies 0 Gain eDHrxn < 0 when gaining ebut EA is reported as positive value Of the Noble Gases, only Kr and Xe have an EN EN Atomic Size Kr Xe Common Electronegativities H Force the atom exerts on an e- in a bond C N O F Exact 2.20 2.55 3.04 3.44 3.98 » 2.0 2.5 3.0 3.5 EA Noble Gases have no affinity for e-. It would take energy to force an e- on them 4.0 Only trend this direction Cations < Neutral < Anions General Chemistry 3: Bonding and Chemical Interactions Bond Type According to DEN 0 Covalent Bonds Covalent Bond: Formed via the sharing of electrons between two elements of similar EN. 0.5 Polar covalent 1.7 Ionic Ionic Bonds Bond Order: Refers to whether a covalent bond is a single, double, or triple bond. As bond order increases bond strength#, bond energy#, bond length¯. Ionic Bond: Formed via the transfer of one or more electrons from an element with a relatively low IE to an element with a relatively high electron affinity DEN > 1.7. Nonpolar Bonds: DEN < 0.5. Cation: POSITIVE + Polar Bonds: DEN is between 0.5 and 1.7. Anion: NEGATIVE − Coordinate A single atom provides both bonding electrons. Covalent Bonds: Most often found in Lewis acid-base chemistry. Intermolecular Forces Strength Nonpolar covalent Crystalline Lattices: Large, organized arrays of ions. Sigma and Pi Bonds Formal Charge Hydrogen O-H, N-H, F-H 1s Formal'Charge = valence'e0 − dots − sticks Dipole-Dipole 1s 1p Dots: Sticks: London Dispersion 1s 2p Nonbonding ePair of bonding electrons Note: Van de Waals Forces is a general term that includes Dipole-Dipole forces and London Dispersion forces. H-Bond acceptor Valence Shell Electron Pair Repulsion Theory (VSEPR) H-Bond donor Electronic Geometry: Molecular Shape : Hybridization sp sp2 sp3 sp3d sp3d2 e- Groups Around Central Atom 2 3 4 5 6 Bonded and lone pairs treated the same. Lone pairs take up less space than a bond to another atom. Bonded Lone Pairs Pairs 2 1 0 1 3 2 1 4 3 2 1 5 4 3 2 6 5 4 0 1 2 0 1 2 3 0 1 2 3 0 1 2 Electronic Molecular Geometry Shape Linear Linear Linear Trig Planar Trigonal Planar Bent Linear Tetrahedral Trig Pyramidal Tetrahedral Bent Linear Bond Angle 180¡ 120¡ 109.5¡ Trigonal Bipyramidal 90¡ Trigonal Bipyramidal Seesaw T-Shaped Linear Octahedral 120¡ Octahedral Square Pyramidal 90¡ Square Planar & General Chemistry 4: Compounds and Stoichiometry Equivalents & Normality Naming Ions Equivalent Mass of an acid that yields 1 mole of H+ or Mass: mass of a base that reacts with 1 mole of H+. For elements (usually metals) that can form more than one positive ion, the charge is indicated by a Roman numeral in GEW = !"#$%&!$'' parentheses following the name of the !"#&() &"%&*+ element Equivalents = Molarity = Iron(II) Fe3+ Cu+ Cu2+ Iron(III) Copper(I) Copper(II) Older method: –ous and –ic to the atoms with lesser and greater charge, respectively Fe2+ Ferrous Fe3+ Cu+ Cu2+ Ferric Cuprous Cupric Monatomic anions drop the ending of the name and add –ide HFO2S2N3P3- Hydride Fluoride Oxide Sulfide Nitride Phosphide Oxyanions = polyatomic anions that contain oxygen. MORE Oxygen = –ate LESS Oxygen = –ite NO3- Nitrate NO2SO42SO32- Nitrite Sulfate Sulfite In extended series of oxyanions, prefixes are also used. MORE Oxygen = Hyper- (per-) LESS Oxygen = Hypo- ClO- Hypochlorite ClO2ClO3ClO4- Chlorite Chlorate Perchlorate Polyatomic anions that gain H+ to for anions of lower charge add the word Hydrogen or dihydrogen to the front. HCO3- Hydrogen carbonate or bicarbonate HSO4- Hydrogen sulfate or bisulfate H2PO4- Dihydrogen phosphate !$''&",&-"!."/01 234 Normality = Fe2+ 35 6 For acids, the # of equivalents (n) is the # of H + available from a formula unit. 0"%!$#789 !"#&() &"%&*+ Compound Formulas Empirical: Simplest whole-number ratio of atoms. Molecular: Multiple of empirical formula to show exact # of atoms of each element. Types of Reactions Combination: Two or more reactants forming one product 2H2 (g) + O2 (g) ® 2H2O (g) Decomposition: Single reactant breaks down 2HgO (s) ® 2Hg (l) + O2 (g) Combustion: Involves a fuel, usually a hydrocarbon, and O2 (g) Commonly forms CO2 and H2O CH4 (g) + 2O2 (g) ® CO2 (g) + H2O (g) Single-Displacement: An atom/ion in a compound is replaced by another atom/ion Cu (s) + AgNO3 (aq) ® Ag (s) + CuNO3 (aq) Double-Displacement: Elements from two compounds swap places (metathesis) CaCl2 (aq) + 2AgNO3 (aq) ® Ca(NO3)2 (aq) + 2AgCl (s) Neutralization: A type of double-replacement reaction Acid + base ® salt + H2O HCl (aq) + NaOH (aq) ® NaCl (aq) + H2O (l) Acid Names -ic: Have one MORE oxygen than -ous. -ous: Has one FEWER oxygen than -ic. General Chemistry 5: Chemical Kinetics m Order Rate Law Integrated Rate Law 0 zeroth order Units of Rate Constant Zeroth Order Reaction First Order Reaction Second Order Reaction Reaction Order and Michaelis-Menten Curve: At low substrate concentrations, the reaction is approximately FIRST-ORDER. At very high substrate concentration, the reaction approximates ZERO-ORDER since the reaction ceases to depend on substrate concentration. Types of Reactions Combination: Two or more reactants forming one product. 2H2 (g) + O2 (g) ® 2H2O (g) Decomposition: Single reactant breaks down. 2HgO (s) ® 2Hg (l) + O2 (g) Combustion: Involves a fuel, usually a hydrocarbon, and O2 (g). Commonly forms CO2 and H2O. CH4 (g) + 2O2 (g) ® CO2 (g) + H2 O (g) Single-Displacement: An atom or ion in a compound is replaced by another atom or ion. Cu (s) + AgNO3 (aq) ® Ag (s) + CuNO3 (aq) Double-Displacement: Elements from two compounds swap places. (metathesis) CaCl2 (aq) + 2AgNO 3 (aq) ® Ca(NO 3)2 (aq) + 2AgCl (s) Neutralization: A type of double-replacement reaction. Acid + base ® salt + H2O HCl (aq) + NaOH (aq) ® NaCl (aq) + H2O (l) Hydrolysis: Using water to break the bonds in a molecule. Equations ( Arrhenius: ! " # $ %&' )* Definition of Rate: For aA + bB ® cC + dD Rate = + D,-. = + D / D0 ,1. 2D 0 = = D,3. 4D0 = D,5. 6 D0 ? Reaction Mechanisms Overall Reaction: A 2 + 2B ® 2AB Step 1: Step 2: A 2 + B ® A2 B A2B + B ® 2AB slow fast A2B is an intermediate Slow step is the rate determining step Arrhenius Equation ( Arrhenius: ! " # $ % &' )* Gibbs Free Energy LM " NO + NO;PQR +LM " NS:7TUVWX YLM " NVZ:7TUVWX k = rate constant A = frequency factor Ea = activation energy G R = gas constant = BCDEF; HIJ;K T = temp in K Trends: #A Þ #k #T Þ #k (Exponent gets closer to 0. Exponent becomes less negative) General Chemistry 6: Equilibrium Kinetic (Ea) and Thermodynamic Equilibrium Constant aA + bB ⇌ cC + dD Equilibrium Constant (Keq): !"# $ Reaction Quotient (Qc): 01 $ %&'( )%*'+ %,'- )%.'/ %&'( )%*'+ %,'- )%.'/ G) Control Kinetic Products: HIGHER in free energy than thermodynamic products and can form at lower temperatures. “Fast” products because they can form more quickly under such conditions. Thermodynamic Products: LOWER in free energy than kinetic products, more stable. Slower but more spontaneous (more negative DG) Exclude pure solids and liquids Reaction Quotient DG = 0, equilibrium Le Châtelier’s Principle If a stress is applied to a system, the system shifts to relieve that applied stress. Example: Bicarbonate Buffer CO2 (g) + H2O (l) ⇌ H2CO3 (aq) ⇌ H+ (aq) + HCO3- (aq) Q > Keq DG > 0, reaction ¬ ¯pH Þ &respiration to blow off CO2 &pH Þ ¯respiration, trapping CO2 General Chemistry 7: Thermochemistry Systems and Processes Temperature (T) and Heat (q) Isolated System: Exchange neither matter nor energy with the environment. Temperature (T): Scaled measure of average kinetic energy of a substance. Closed System: Can exchange energy but not matter with the environment. Celsius vs 0¡C = 32¡F Fahrenheit: 25¡C = 75¡F % Open system: Can exchange BOTH energy and matter with the environment. Isothermal Process: Constant temperature. Adiabatic Process: Exchange no heat with the environment. Freezing Point H 2O Room Temp ! " #$ '$( ) *+ & 37¡C = 98.6¡F Body Temp Heat (q): The transfer of energy that results from differences of temperature. Hot transfers to cold. Isobaric Process: Constant pressure. Isovolumetric: Constant volume. (isochoric) Enthalpy (H) Enthalpy (H): A measure of the potential energy of a system found in intermolecular attractions and chemical bonds. States and State Functions State Functions: Describe the physical properties of an equilibrium state. Are pathway independent. Pressure, density, temp, volume, enthalpy, internal energy, Gibbs free energy, and entropy. Standard Conditions: 298 K, 1 atm, 1 M Note that in gas law calculations, Standard Temperature and Pressure (STP) is 0¡C, 1 atm. Fusion: Solid ® liquid Freezing: Liquid ® solid Vaporization: Liquid ® gas Sublimation: Solid ® gas Phase Changes: Solid ® Liquid ® Gas: ENDOTHERMIC since gases have more heat energy than liquids and liquids have more heat energy than solids. Gas ® Liquid ® Solid: EXOTHERMIC since these reactions release heat. Hess’s Law: Enthalpy changes are additive. ¡ $from heat of formations D,-./ ¡ Entropy (S) Deposition: Gas ® solid Triple Point: Point in phase diagram where all 3 phases exist. Supercritical Fluid: Density of gas = density of liquid, no distinction between those two phases. Entropy (S): A measure of the degree to which energy has been spread throughout a system or between a system and its surroundings. @ 0? " $ ABC D Standard entropy of reaction ¡ ¡ Note: Entropy is maximized at equilibrium. Gibbs Free Energy (G) Gibbs Free Energy (G): Derived from enthalpy and entropy. DH DS Outcome + + Spontaneous at HIGH temps + - Non-spontaneous at all temps - + Spontaneous at all temps - - Spontaneous at LOW temps Note: Temperature dependent when DH and DS have same sign. From equilibrium constant Keq 0LMNO From reaction quotient Q P ) $Q$R$ST$#X( 0LMNO " 0LMNO Y 0LMNO " Q$R$ST$# ( ZB[ D G < 0 : Spontaneous D G = 0 : Equilibrium D G > 0 : Non-spontaneous General Chemistry 8: The Gas Phase Ideal Gases Real Gases Ideal Gas: Theoretical gas whose molecules occupy negligible space and whose collisions are perfectly elastic. Gases behave ideally under reasonably !temperatures and ¯pressures. STP: 273 K (0¡C), 1 atm 1 mol Gas: At STP 1 mol of gas = 22.4 L Units: !"#$% & '()"%%*+ & '()"$,-- & !)!./"01# & !2.'"345 6"7 & 8"9": Density of Gas: r = Avogadro’s Principle: DH CH GH B Gay-Lussac’s Law: At Extremely !P, ¯ V, Real gases will occupy more volume than or ¯ T: predicted by the ideal gas law because the particles occupy physical space. gT h Van der Waals de f " T j Ji k glK & "g;Y Equation of State: i C = = DE FG = DI CI GI Kinetic Molecular Theory V2 = V1JDHK"J GI K L C D I = k or G LH CH (n is constant) H = LI CI (T and P are constant) Boyle’s Law: PV = k or P1V1 = P2V2 Charles’s Law: At Moderately ! P, ¯ V, Real gases will occupy less volume than or ¯ T: predicted by the ideal gas law because the particles have intermolecular attractions. a corrects for attractive forces b corrects for volume of the particles themselves Ideal Gas Law Combined Gas Law: Real gases deviate from ideal behavior at ¯temperature & !pressure C G D G = k or = k or CH GH C = GI I DH DI = GH GI (n and T are constant) (n and P are constant) (n and V are constant) Avg Kinetic QR & " S "U"VT & " W "Q "Y T X T !T = molecules move FASTER !molar mass = molecules move SLOWER Root-Mean/;Y Square Speed: ]^>_ & " ` a Diffusion: The spreading out of particles from [high] ® [low] Effusion: The mvmt of gas from one compartment to another through a small opening under pressure Graham’s Law: bH Other Gas Laws bI (total pressure from partial pressures) Dalton’s Law: PA = XAPT EI EH ¯molar mass = diffuse/effuse FASTER !molar mass = diffuse/effuse SLOWER Dalton’s Law: PT = PA + PB + PC + … (partial pressure from total pressure) =c (X = mol fraction) Henry’s Law: MNO = kH x PA or MPOH MPOI = DH DI = kH Diatomic Gases Exist as diatomic molecules, never a stand-alone atom. Includes H2, N2, O2, F2, Cl2, Br2, and I2 Mnemonic: “Have No Fear Of Ice Cold Beer” The 7 Diatomic Gases = General Chemistry 9: Solutions Terminology Solubility Rules Solution: Homogenous mixture. Solvent particles surround solute particles via electrostatic interactions. Solvation or The process of dissolving a solute in solvent. Most Dissolution: dissolutions are endothermic, although dissolution of gas into liquid is exothermic. Solubility: Maximum amount of solute that can be dissolved in a solvent at a given temp. Soluble Na+, K+, NH4+ NO3Cl-, Br-, I- Except with Pb2+, Hg22+, Ag+ SO42- Except with Ca2+, Sr2+, Ba2+, Pb2+, Hg22+, Ag+ Molar Solubility: Molarity of the solute at saturation. Complex Ions: Cation bonded to at least one ligand which is the epair donor. It is held together with coordinate covalent bonds. Formation of complex ions #solubility. Solubility in Water: Polar molecules (with +/- charge) are attracted to water molecules and are hydrophilic. Nonpolar molecules are repelled by water and are hydrophobic. Polar = Hydrophilic Nonpolar = Hydrophobic Insoluble S2- Except with Na+, K+, NH4+, Mg2+, Ca2+, Sr2+, Ba2+ O2- Except with Na+, K+, Sr2+, Ba2+ OH- Except with Na+, K+, Ca2+, Sr2+, Ba2+ CrO42- Except with Na+, K+, Mg2+, NH4+ PO43- & CO32- Except with Na+, K+, NH4+ Concentration % by mass: -.""/"01234 -.""/"0123506 -014"/"01234 Molality: @- $ Normality: D $ Colligative Properties 1534>" 0? "0123506 -014"/"01234 AB/0?/"01C463 E/0?/4F25C.1463" 1534>"/0?/"0123506 Can also just be a lowercase m For acids, the # of equivalents (n) is the # of H + available from a formula unit. Dilutions: M1 V1 = M2 V2 Solutions Equilibria Saturated solutions are in equilibrium at that particular temperature. Solubility Product Equilibrium expression for something that dissolves. Constant: For substance AaBb, !"# $ %&'( %)'* %&'( %)'* Ion Product: +, $ IP < Ksp unsaturated IP = Ksp saturated at equilibrium IP > Ksp supersaturated, precipitate Formation or Kf. The equilibrium constant for complex formation. Stability Constant: Usually much greater that Ksp. Common Ion ¯solubility of a compound in a solution that already Effect: contains one of the ions in the compound. The presence of that ion shifts the dissolution reaction to the left, decreasing its dissociation. Chelation: When a central cation is bonded to the same ligand in multiple places. Chelation therapy sequesters toxic metals. Colligative Properties: Physical properties of solutions that depend on the concentration of dissolved particles but not on their chemical identity. The presence of other solutes ¯evaporation rate of solvent, thus ¯Pvap. Boiling Point Elevation: GHI $ J/!I /@J $ KLMKNOPKLM/QORPLS !I $ TLKUKMV/WLKMP/XYWSYZZKLM/RLMZPOMP @- $ [LUOU/RLMRYMPSOPKLM Freezing Point Depression: GH? $ J/!? /@!? $ QSYYNKMV/WLKMP/XYWSYZZKLM/RLMZPOMP Osmolarity: The number of individual particles in solution. Example: NaCl dissociates completely in water, so 0"-01 8/\/]O^U $ _/ 1534> Osmotic Pressure: “Sucking” pressure generated by solutions in which water is drawn into solution. p $ J/=/`/H J $ aOMb P/cLQQ/QORPLS = $ [LUOS/RLMRYMPSOPKLM/LQ/ZLUdPY e $ VOZ/RLMZPOMP H $ PY[WYSOPdSY General Chemistry 10: Acids and Bases Definitions Polyvalence & Normality Arrhenius Acid: Produces H+ (same definition as Brønsted acid) Arrhenius Base: Produces OH- Equivalent: 1 mole of the species of interest. Normality: Concentration of equivalents in solution. Brønsted-Lowry Acid: Donates H+ (same definition as Arrhenius acid) Brønsted-Lowry Base: Accepts H+ Polyvalent: Can donate or accept multiple equivalents. Example: 1 mol H3PO4 yields 3 mol H+. So, 2 M H3PO4 = 6 N. Lewis Acid: Accepts e- pair Lewis Base: Donates e- pair Note: All Arrhenius acids/bases are Brønsted-Lowry acids/bases, and all Brønsted-Lowry acid/bases are Lewis acids/bases; however, the converse of these statements is not necessarily true. Amphoteric Species: Species that can behave as an acid or a base. Amphiprotic = amphoteric species that specifically can behave as a BrønstedLowry acid/base. Polyprotic Acid: An acid with multiple ionizable H atoms. Titrations Half-Equivalence Point: The midpoint of the buffering region, in which half the (midpoint) titrant has been protonated or deprotonated. 3.L5 # 3L& 5 and -. # -M) +and a buffer is formed. Equivalence Point: The point at which equivalent amounts of acid and base have reacted. N' +O' # NP +OP pH at Equivalence Point: Strong acid + strong base, pH = 7 Weak acid + strong base, pH > 7 Weak base + strong acid, pH < 7 Weak acid + weak base, pH > or < 7 depending on the relative strength of the acid and base Properties Indicators: Weak acids or bases that display different colors in the protonated and deprotonated forms. The indicator’s pKa should be close the pH of the equivalence point. &'( at 298 K Water Dissociation Constant: !" # $% !" # !) *+!, pH and pOH: -. # + /012+3.4 5 -8. # + /012+38.& 5 -. 9 -8. # $: Tests: Litmus: Acid = red; Base = blue; Neutral = purple Phenolphthalein: pH < 8.2 = colorless; pH > 8.2 = purple Methyl Orange: pH < 3.1 = red; pH > 4.4 = yellow Bromophenol Blue: pH < 6 = yellow; pH > 8 = blue 3 .45 # $%&67 Endpoint: When indicator reaches full color. Polyvalent Acid/Base Multiple buffering regions and equivalence points. Titrations: Strong Acids/Bases: Dissociate completely Weak Acids/Bases: Do not completely dissociate Acid Dissociation Constant: !) # 37FG H 53IJ 5 37I5 Base Dissociation Constant: ! # 3 KH53G7J5 , 3KG75 -!) # + /012+;!) > Titration Setup -!, # + /012+;!, > Burette -!) 9 -!, # -!" # $: Conjugate Acid/Base Pairs: Strong acids & bases / weak conjugate Weak acids & bases / weak conjugate Neutralization Reactions: Form salts and (sometimes) H2O Titrant (strong acid in this example) Analyte / Titrand (weak base in this example) Conical flask Buffers Buffer: Weak acid + conjugate salt Weak base + conjugate salt Buffering Capacity: The ability of a buffer to resist changes in pH. Maximum buffering capacity is within 1 pH point of the pKa. J 3I 5 Henderson-Hasselbalch -. # -M) 9 012 37I5 Equation: 3 KH5 -8. # -M, 9 012 37G75 When [A-] = [HA] at the half equivalence point, log(1) = 0, so pH = pKa Titration Curve When titrating a weak base with a strong acid Midpoint -8. # -M, Equivalence Point N' +O' # NP +OP General Chemistry 11: Oxidation-Reduction Reactions Definitions Oxidation: Loss of eReduction: Gain of eWith Respect to Oxidation is GAIN of oxygen Oxygen Transfer: Reduction is LOSS of oxygen Oxidizing Agent: Facilitates the oxidation of another compound. Is itself reduced Reducing Agent: Facilitates the reduction of another compound. Is itself oxidized Balancing via Half-Reaction Method ¥ ¥ ¥ ¥ ¥ ¥ Separate the two half-reactions Balance the atoms of each half-reaction. Start with all elements besides H and O. In acidic solution, balance H and O using water and H+. In basic solution, balance H and O using water and OHBalance the charges of each half-reaction by adding e- as necessary Multiply the half-reactions as necessary to obtain the same number of e- in both half-reactions Add the half-reactions, canceling out terms on both sides Confirm that the mass and charge are balanced Oxidation # Rules ¥ ¥ ¥ ¥ ¥ ¥ ¥ Any free element or diatomic species = 0 Monatomic ion = the charge of the ion When in compounds, group 1A metals = +1; group 2A metals = +2 When in compounds, group 7A elements = -1, unless combined with an element of greater EN H = +1 unless it is paired with a less EN element, then = -1 O = -2 except in peroxides, when it = -1, or in compounds with more EN elements The sum of all oxidation numbers in a compound must = overall charge Net Ionic Equations Complete Ionic Equation: Accounts for all of the ions present in a reaction. Split all aqueous compounds into their relevant ions. Keep solid salts intact. Net Ionic Equation: Ignores spectator ions Disproportionation Reactions: A type of REDOX reaction in which one element is both (dismutation) oxidized and reduced, forming at least two molecules containing the element with different oxidation states REDOX Titrations: Similar in methodology to acid-base titrations, however, these titrations follow transfer of charge Potentiometric Titration: A form of REDOX titration in which a voltmeter measures the electromotive force of a solution. No indicator is used, and the equivalence point is determined by a sharp change in voltage General Chemistry 12: Electrochemistry Galvanic Cell Electrochemical Cells Anode: Always the site of oxidation. It attracts anions. Electrolytic Cell Emf & Thermodynamics Electromotive force and change in free energy always have OPPOSITE signs. Cathode: Always the site of reduction. It attracts cations. Red Cat = Reduction at the Cathode Type of Cell '¡()** DG ¡ Galvanic + - Electrolytic - + Concentration 0 0 e- Flow Anode ® Cathode Current Flow: Cathode ® Anode ¡ Galvanic Cells: House spontaneous reactions. -DG, +Emf, +!%#&& (Voltaic) Anode = NEG, Cathode = POS ¡ Electrolytic Cells: House non-spont reactions. +DG, -Emf, -!%#&& Anode = POS, Cathode = NEG Concentration Specialized form of galvanic cell in which both electrodes are Cells: made of the same material. It is the concentration gradient between the two solutions that causes mvmt of charge. Rechargeable Can experience charging (electrolytic) and discharging Batteries: (galvanic) states. Lead-Acid: Discharging: Pb anode, PbO2 cathode in a concentrated sulfuric acid solution. Low energy density. ¡ ¡ ¡ !%#&& + !"#$,%-./0$# Faraday constant (F): 96,485 C E C3D + F Ni-Cd: Discharging: Cd anode, NiO(OH) cathode in a concentrated KOH solution. Higher energy density than lead-acid batteries. NiMH: More common than Ni-Cd because they have higher energy density. Cell Potentials Reduction Potential: Quantifies the tendency for a species to gain eand be reduced. More positive Ered = greater tendency to be reduced. ¡ Standard Reduction Potential: !"#$ . Calculated by comparison to the standard hydrogen electrode (SHE). ¡ Standard Electromotive Force: !%#&& . The difference in standard reduction potential between the two half-cells. ¡ Galvanic Cells: +!%#&& ¡ Electrolytic Cells: -!%#&& Nernst Equation Describes the relationship between the concentration of species in a solution under nonstandard conditions and the emf. ¡ When Keq > 1, then +!%#&& ¡ When Keq < 1, then -!%#&& ¡ =0 When Keq = 1, then !%#&& GH ¡ !%#&& + !%#&& I