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| 6571039385 | Key Oxidation Number Rules | F=-1 (when not a lone element/diatomic) O=-2, =-1 when peroxide, =+2 in OF2 H=+1, =-1 when bonded to metal 1A Metals=+1 2A Metals=+2 Al=+3 | 0 | |
| 6571043098 | Electrochem Mnemonics | LEO GER-Loss of Electrons is Oxidation & Gain of Electrons is Reduction OIL RIG-Oxidation Is Loss Reduction is Gain AN OX & A Red CAT-Anode/Oxidation & Reduction/Cathode FAT CAT- e- flow From Anode To CAThode cations also flow FATCAT but in salt bridge | 1 | |
| 6571044392 | Standard Potential | Eo= Reduction Potential of Cathode - Reduction Potential of Anode. AKA-Eo=Eoreduction+Eooxidation | 2 | |
| 6571051816 | Stoichiometry/Mole Conversions | Mole=22.4L @STP Mole=mass/molar mass Mole=6.02*1023 Find Limiting Reactant: 1.Convert to moles 2.Divide by Coefficient Molarity=Moles/Liters (of solution) Molality=Moles/Kg (of solvent) %Y=Act/Theo; Act is measured. Theo is Calc | 3 | |
| 6571053774 | Bond Angles/Molecular Geometry | VSEPR-e- pairs and bonds push on each other *2 Bond & 0 Pairs, Linear, 180, sp hybrid, Ex: CO2 *2B & 1P, Bent, 120, sp2 Hybrid, Ex:NO2 *2B & 2P, Bent, <109.5, sp3 Hybrid, Ex:SO2 *3B & 0P, Trigonal Planar, 120, sp2 Hybrid, Ex: SO3 *3B & 1P, Trigonal Pyr., <109.5, sp3 Hybrid, Ex:NH3 *4B & 0P, Tetrahedral, 109.5, sp3 Hybrid, Ex:SiH4 *5B & 0P, Tri. Bipyrim., 120/90, sp3d Hybrid, Ex:PF5 | 4 | |
| 6571057650 | Gas Laws | Ideal→ PV=nRT & MM=DRT/P (Dirt over Pee) Dalton's → PT=P1+P2+P3 + ... Combined→ P1V1/T1=P2V2/T2 Graham's→ Rate A/Rate B = Sqrt(MMB/MMA) R=.08206 in atm; R=8.314 for everything else | 5 | |
| 6672928199 | Solubility Rules | Top 4 for Ionic Compounds: 1. 1A metals and NH4+ compounds are soluble 2. Nitrates & Acetates are soluble 3. Cl-,Br-,I- are soluble except: Ag+, Pb+2, & Hg2+2 4. SO4-2 are soluble except: Ba+2, Sr+2, Pb+2, & Hg2+2 If a compound is soluble or SA or SB then breaks up when writing net Ionic equation. | 6 | |
| 6672930268 | Strong Acid/Bases | Acids: HCl, HBr, HI, HNO3, H2SO4, HClO4 Bases: LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 All weak acids/bases do not break up 100% SA & SB break up 100% | 7 | |
| 6672930269 | Periodic Trends II | e- Affinity=Similar to Ion. En. because more Zeff for smaller atoms. Exceptions: Low affinity for full S or P orbitals & low affinity for any p3 because e- repulsion Electroneg.=Similar to Ion. En. small radii & almost full orbital. Exception:Noble Gases=0 (full orbitals) Reactivity-Metals with lowest Ionization Energy; Nonmetals with highest electronegativity. | 8 | |
| 6672934130 | Periodic Trends | Atomic Radius ↑↓=Atoms get larger going down PT because energy levels are added
Atomic Radius→←=Atoms get smaller to the right of PT because of more Zeff (Effective Nuclear Charge)
Ionization Energy= opposite of radius because easier to remove electrons further from nucleus
Exceptions: B| 9 | | |
| 6672947908 | Gibbs Free Energy | ΔG=ΔH-TΔS H=Heat Energy S=Entropy/Disorder ΔH=+ Endothermic Reaction becomes colder includes melting, vaporization, and sublimation; ΔH=- Exothermic Reaction becomes warmer includes combustion, freezing, condensation, & deposition. ΔS=+ products>reactants, melt., vapor.,& sublim.; ΔS=- react>products, freezing, condensation, deposit. | 10 | |
| 6672950370 | Gibbs Free Energy II | ΔG ΔH TΔS + (themodynamic. unfavorable)+ - - (themodynamic. favorable) - + + at low temp; - at high + + - at low temp; + at high - - | 11 | |
| 6672952042 | Hess's Law & Heat of Formations | ΔGo=Gprod-Greact; ΔHo=Hprod-Hreact; ΔS=Sprod-Sreact Includes Coef. (*electrochem/Eo does not) G=0 & H=0 for any pure element including diatomics Hess's law- a reactions enthalpy (or free energy) is equal to the sum of all of the reactions that make it up. (this is where you multiply by coeff or switch the signs) | 12 | |
| 6672952043 | Bond Energy | Total Bonds broken- Total Bonds formed (if it helps you can consider this the one thing that is reactant minus products) Breaking bonds requires energy. Forming bonds releases energy. | 13 | |
| 6672955417 | Specific Heat | c-specific heat-The heat energy need to raise 1 gram of a substance by one degree celsius. q=mcΔT q=heat energy; m-mass; T-Temperature If a reaction is causing the temperature change then ΔH=-q. | 14 | |
| 6672957574 | Overlap Equations | ΔG ΔH TΔS ΔGo Eo Keq Favors neg pos >1 Products pos neg <1 reactants ΔGo= -RTln(K) ΔGo= -nfEo | 15 | |
| 6672959633 | Rate Law Equations | Rate=k[A]Coef[B]coef The order of rate can be calculated from a table using: Rate 1/Rate 2 = ([A1]/[A2])n 0th order:[x] vs time is linear & slope=-k 1st order:ln[x] vs time is linear & slope=-k 2nd order:1/[x] vs time is linear & slope=+k | 16 | |
| 6672959667 | Units | G, H, and q→ joules or kJ S→ joules/kelvin→ J/K T→ Kelvin m→ grams c→ J/(g*oC) Eo→ volts Instead of specific heat molar heat capacity could be used then c→ J/(moles*oC) and moles (n) will replace mass (m) | 17 | |
| 6672962992 | Rate Mechanism | Slow step is the rate determining step. Only the reactants that are at or before the slow step are used in the rate law. Catalyst appear in the beginning and end of the rate mechanism. They lower activation energy. Intermediates appear in the middle and disappear before the end of the rate mechanism. | 18 | |
| 6672964842 | Rate Law Units | [x]=concentration→ M or moles/L Rate= M/s k→ rate constant→ 1/(s*M^(n-1)) or M^(1-n)/s n is the overall order. | 19 | |
| 6684379889 | London Dispersion Forces | Usually the weakest Intermolecular force. Depends on the number of electrons/size of molecule. Occurs in all molecules | 20 | |
| 6684383347 | Ion Induced Dipole | Occurs when an ion molecule is close to a nonpolar molecule and creates a dipole. More likely to occur to large molecules because they are more polarizable | 21 | |
| 6684385217 | Dipole Induced Dipole | Occurs when a polar molecule is close to a nonpolar molecule and creates a dipole. More likely to occur to large molecules because they are more polarizable. Usually weaker than ion induced. | 22 | |
| 6684387597 | Dipole Dipole | Typically stronger than LDF and Weaker than H-Bond. Involves two separate polar molecules. Polar molecules only. | 23 | |
| 6684389452 | Hydrogen Bond | The Strongest of all IMFs. Not a Bond. Occurs between the Hydrogen of one molecule and the F, O, or N of another molecule. | 24 | |
| 6684392632 | Ionic Bond | Forms Ionic Solids with very high MP and BP Repetitive crystal structure caused by repeated bonds Conducts Electricity when melted or dissolved in water Ionic forces increases with larger charge difference and decreases with radius Examples: MgS>NaCl & NaCl>KBr. | 25 | |
| 6684395592 | Covalent Network Solid | Very strong repetitive covalent bonds in a network/lattice. Very High melting point/Boiling point Examples:Diamond(C) and Quarts(SiO2) | 26 | |
| 6684398531 | Metallic Bond | Delocalized sea of electrons Great at conducting electric current. Nucleus is stationary. Two types of alloys that can vary metals properties-Substitutional and Interstitial. | 27 | |
| 6684401080 | Vapor Pressure | Depends on Intermolecular forces VP↓ as IMFs ↑ Depends on temperature VP ↑ as T↑ VP must be > atm pressure to cause a substance to boil. | 28 | |
| 6684451108 | Equilibrium Expression | Kc=[products]coef/[reactants]coef Kp=P(products)coef/P(reactants)coef Ksp=[products]coef → (sp-solubility product) Ka=[H+][A-]/[HA] Kb=[B+][OH-]/[BOH] or [BH+][OH-]/[B] K has no units in equilibrium (k does in kinetics.) Kp is the only one without brackets. | 29 | |
| 6684454118 | Le Chatelier's Principle | Equilibrium shifts to the side with less stress. *Increasing temperature will cause a shift away from side with heat (shift right for endo & shift left for exo) *Increasing pressure will shift away from side with more gases. Solids and liquids have no effect. *Inert gas has no effect at constant volume & same effect as lower pressure if system pressure is constant. *Add product/reactant to cause a shift to opposite side. | 30 | |
| 6684456226 | Q vs K | Q is for initial conditions. K is for equilibrium.
If Q| 31 | | |
| 6684461651 | pH Equations | pH=-log[H+] pOH=-log[OH-] & pKa=-log[Ka] [H+]*[OH-]=10-14=Kw pH+pOH=14 Buffers: pH=pKa + log([A-]/[HA]) [H+]=Ka*[HA]/[A-] | 32 | |
| 6685481443 | Acid Base Reactions | Strong Acid + Strong Base → Neutral Salt (pH=7) HCl + NaOH → NaCl + H2O Weak Acid + Strong Base → Weak Base (pH>7) HF + NaOH → NaF + H2O Strong Acid + Weak Base → Weak Acid (pH<7) HCl + NH3 → NH4Cl Weak Acid + Weak Base → Depends on Ka Value HF + NH3 → NH4F | 33 | |
| 6685483182 | Rice Table | Reaction A +2B → 3C + D Initial 5 5 5 5 Change -x -2x +3x +x Equilibrium 5-x 5-2x 5+3x 5+x Used for solving equilibrium expressions. | 34 | |
| 6685483183 | Conjugate & Buffers | NH3 → base NH4+ → conjugate acid (gained H+) HCN→ acid CN- → Conjugate base (lost H+) Ka*Kb=10-14 A buffer can be formed from an acid base conjugate pair. **It can also be formed by a strong acid/base reacting with excess base/acid. | 35 | |
| 6685487950 | Nuclear | Nuclear notation: AZX where A=atomic mass; Z=atomic number, X=element symbol When balancing nuclear reactions make sure the numbers on either side of the equation are equal. Fusion:combining nuclei; Fission-splitting nuclei Alpha: 42He or 42 α least penetrating Beta:0-1e or 0-1β Positron:0+1e or 0+1β Gamma:00γ most penetrating power | 36 | |
| 6685487985 | Factors Affecting Rate | For a reaction to be successful a collision must occur with enough energy and the correct orientation. Concentration-the higher the concentration the more frequent collision Temperature-high T→ faster molecules & more energy Catalyst-lowers act. energy needed by alt. pathway Surface Area-high surface area/small particles means more collisions | 37 | |
| 6685491020 | Bond Type & Strength | σ sigma (bond directly) π-reach around with p orbitals
Singl-σ only; double 1 σ & 1 π, triple 1σ 2π
In order of strength: single< double| 38 | | |
| 6685534444 | Other Equations | Kp=Kc(RT)Δn Δn=prod of gas-react of gas %Yield=(act/theo)*100; Rate= Δ[x]/t Calculate energy needed for a phase change: q=m*ΔHF ΔHF-enthalpy of fusion q=m*ΔHv ΔHv-enthalpy of vaporization Know how to solve Enthalpy of fusion/vaporization for either mass or grams. | 39 | |
| 6685538095 | Ions Vs Isotopes | Ions: Change in number of electrons: can affect periodic trends like atomic size but no effect on mass. Relevant to PES. Not to mass spectroscopy. Isotopes: Change in atomic mass which depends on number of protons and neutrons NOT electrons. Neutrons have no effect on periodic trends relevant to mass spectroscopy not PES. | 40 | |
| 6685542260 | Boiling & Freezing Point | ΔTB=iKBm ΔTF=iKFm i is the dissociation factor (how many pieces does the ion break into.Ex: K2O→ 2K+ + O-2 3 pieces) K is a constant and m is molality (moles/kg). When a solute is dissolved in water freezing point goes down (depresses) and boiling point goes up (elevates) **Exception a volatile liquid with a low boiling point will make the boiling point drop.** | 41 |
