Module 1 General Chemistry PDF

MODULE 1│CHEM 1 GENERAL CHEMISTRY I. MATTER Chemical compounds always contain the exact proportion of...

MODULE 1│CHEM 1 GENERAL CHEMISTRY I. MATTER Chemical compounds always contain the exact proportion of element in fixed ratio (by mass) Mass + Volume Ex. H2O →2H + O, C6H12O6 = CH2O STATES SOLID LIQUID GAS Indefinite 3. Law of Multiple Proportion – John Dalton Shape Definite *non- *assumes Indefinite When 2 elements form more than 1 compounds, it can be compressible container shape *compressible expressed in a fixed whole number (by mass) Volume Definite Ex. CO → 28g/mole, CO2 → 44g/mole Molecular Vibration Gliding *ex. Constant C = 12g/mole motion *2 stones water falls random O = 16g/mole Plasma/ Ionized Gas – 4th state; most abundant state of matter. 4. Law of combining weights Has p+ and e- (thus, greatly affected by magnetic field) Proportions by weight when chemical reaction takes place Ex. ionized Ne light, Aurora, Stars, Sun can be expressed in small integral unit IFA Strength most ↑ or strongest: S > L > G > P Ex. MgO → 40g/mole (100%) Enthalpy (heat/ reaction energy): P > G > L > S Mg = 24g/mole (60%); O = 16g/mole (40%) A. PHASE CHANGE III. ATOMIC STRUCTURE 1. Democritus – Atomos Melting (Solid to Liquid) “Indivisible” *aka: Fusion, Liquefaction, Thawing 2. John Dalton – Billiard ball Matter is made up of atoms Freezing (Liquid to Solid) Postulates: Elements are composed of indivisible, indestructible atoms Evaporation (Liquid to Gas) Atoms alike for a given element (isotopes) Atoms of different elements differ in size, mass & other properties (isobars) Condensation (Gas to Liquid) Compound are formed form 2 or more atoms at different elements Sublimation Atoms combined in simple numerical ratios to form (Solid to Gas) *moth/naphthalene balls compounds Deposition 4. J.J. Thompson – Plum Pudding/ Raisin bread (Gas to Solid) *dry ice/ cardice e- in (+) framework Recombination 5. Ernest Rutherford (discoverer of proton) – Nuclear (Gold foil/ α- (Plasma to Gas) *aka: Deionization scattering experiment) atom is mostly empty; (+) particles in nucleus Ionization (Gas to Plasma) 6. Neil Bohr – Planetary mostly used B. MATTER CLASSIFICATION 7. Erwin Schrodinger – Quantum/ Mechanical/ e- cloud 1. Pure substance Modern atomic Model; estimates the probability of finding an Element – simplest form of substance. e- in certain position (i.e. at e-cloud/ orbital) Compound – 2 or more chemical united (separated via chemical means) Atoms 2. Mixture – 2 or more substance wherein individual substance Proton – (+) ion identifies are retained (separated via physical means. Alcohol + Atomic number (basis of electronic configuration) Water via distillation) Ernest Rutherford Homogeneous – 1 phase; solution *clear colored Heterogeneous – 2 phases; suspension, colloid *ex. milk Electrons – (-) ion p+ in uncharged state C. CLASSIFICATION BASED ON DEPENDENT TO THE AMOUNT negligible weight 1,836x lighter that p+ OF MATTER PRESENT J.J. Thompson Cathode ray tube: e- m/2 ratio 1. Extrinsic Property “Dependent” R.A. Millikan Length, mass/weight, volume, pressure, entropy, enthalpy, Oil drop experiment: measure accurate charge and electrical resistance mass of e- 2. Intrinsic Property “Independent” Neutrons – no charge Density/ SpGr (water = 1g/ml or cc), viscosity (resistance to Atomic mass (Nucleon) = p+ + n0 flow), velocity (m/sec), temperature, color James Chadwick II. FUNDAMENTAL CHEMISTRY LAWS # p+ = Atomic # = 11 # n0 = Atomic mass – p+ = 23-11 = 12 1. Law of Conservation of Mass/ Matter – Antoine Lavoiser # e- = p+ in uncharged stated: 11-1 = 10 Mass/ Matter is always constant (neither created nor destroyed) P #p = 15 #n = 16 2. Law of Definite/ Constant Proportions – Joseph Proust #e = 18 (Proust’s law) Module 1 – General Chemistry Page 1 of 8 RJAV 2022 Find: Atomic no. = 15 Atomic mass = 15 + 16 = 31 Charge = 15 – 18 = -3 Eugene Gold Stein – discovered anode rays Electrochemistry – particle separation based on e- Ex: Capillary electrophoresis – separation of compounds based on electrophoretic mobility 4A 6A Electrode: Anode Cathode Ionic bonding Charge: + electrode - electrode Undergoes: Oxidation Reduction RED CAT ELECT IN REDuction happens in CAThode where ELECTrons get IN 1A VILEORA Valence Increase, Loses e-, undergoes Oxidation, Reducing Valence shell electron pair repulsion (VSEPR) theory agent Predicts the geometry of the molecule as well as any bonded and unbonded electron pair VDGEROA Valence Decrease, Gains e-, undergoes Reduction, Oxidizing Linear (180˚) - CO2 agent (KMnO4-, Na2Cr2O7) Alkynes (Sp) Isotopes Trigonal planar (120˚) - BF3 same p+/atomic number/ element Alkenes (Sp2) differ in atomic mass Non-isotopes: 19F, 127I, 31P, etc. Main isotopes: +1: 1H, 12C, 14N, 32S, 35Cl ; +2: 16O, 79Br Isobars same atomic mass differ in elements Isomers Tetrahedral/bent (109.5˚) - CCl4 , H2O same molecular formula Alkanes (Sp3) differ in structure IV. CHEMICAL BONDS * 2 bonded pair, 2 unbonded pair Molecule – aggregate of 2 or more atoms in definite arrangement held together by chemical bonds Ions – with net (+) or (-) charge Empirical formula – simplest whole number ratio (might be same Trigonal bipyramid - PF5 with MF). Ex: CH2O vs. C6H12O6 A. FORCES OF ATTRACTION Intermolecular FA/ Van der Waals/ Electrostatic Between molecule; weak and short-lived Created by “molecule’s polarizability”; exerted when 2 uncharged atoms (n0) approach very closely Octahedral - SF5 H-bonding Keesom Debye London orientation Induction Dispersion (D-D) (D-ID) (ID-ID) Strongest IFA > > Weakest IFA H + S, O, N, X Water – Water – Aromatics (electronegative Water Benzene (Benzene – atoms) Benzene) Dipole (D) – Polar Dashed line – away Induced Dipole (ID) – Nonpolar Wedged line – toward Trigonal and Octahedral – are exemption to the octet rule Intramolecular FA Within molecule Covalent Ionic Valence bond theory Sharing of e- Transfer of e- States that bonds are formed by sharing of electron from overlapping atomic orbitals (covalent) Nonmetal + Nonmetal Metal + Nonmetal (Glycosidic & Peptide bond) (NaCl) Glycosidic – ether bond S─O─S Peptide bond – amide bond AA─peptide─AA bond Covalent Bonding Lone pair Pair of valence electrons that are not shared with another atom in covalent bond s = spherical (sigma bond – stronger bond formed; headways overlap) p = dumbbell (pi bond – weaker; sideways overlap) Module 1 – General Chemistry Page 2 of 8 RJAV 2022 Molecular orbital theory B. Multivalent (with variable charges) States that bonds are formed from interaction of atomic orbitals +1, +2 = Hg, Cu from molecular orbitals +1, +3 = Au +2, +3 = Fe, Co, Ni +3, +5 = Bi, Sb F. POLYATOMIC IONS A. O-containing polyatomic anions (Oxyanions) Oxyanions Salt Oxyacid (Aq) ClO- Hypochlorite Hypochlorous acid (HClO) ClO2- Chlorite Chlorous acid (HClO2) ClO3- Chlorate Chloric acid (HClO3) ClO4- Perchlorate Perchloric acid (HClO4) NO2- Nitrite Nitrous acid (HNO2) NO3- Nitrate Nitric acid (HNO3) SO32- Sulfite Sulfurous acid (H2SO3) SO42- Sulfate Sulfuric acid (H2SO4) PO43- Phosphate Phosphoric acid (H3PO4) Bonding – lower energy (stable) –ate: common form Antibonding – higher energy (unstable) –ite: -1 O to –ate form hypo…ite -1 O to –ite form B. REACTION TYPES per…ate +1 O to –ate form Synthesis/ Combination/ Direct Union B. H-containing polyatomic anions A + B → AB Monohydrogen/ bi: with 1H+ ions Dihydrogen: with 2H+ ions Decomposition/ Analysis Anion Salt AB → A + B HCO3- Bicarbonate (Hydrogen carbonate) i.e. HSO3- Bisulfite Complete combustion: HSO4- Bisulfate CH4 O2 CO2 + H2O HPO4-2 Biphosphate Incomplete combustion: H2PO4-1 Dihydrogen phosphate CH4O2 CO + C(5) + H2O G. MOLE RELATIONSHIPS Single Displacement AB + X → AX + B Avogadro’s # 1 mole = 6.022 x 1023 atoms/ molecules Double Displacement/ Metathesis/ Exchange 𝑔 AB + CD → AC + BD 𝑀𝑊 = 𝑚𝑜𝑙 i.e. Neutralization: Ex. Calculate the no. of NaOH atoms using Avogadro’s no. (Mass = NaOH + HCl → NaCl + H2O 20g, MW = 40/mol) Precipitation: AgNO3 + NaCl → AgCl2 + NaNO3 20𝑔 6.022𝑥1023 𝑎𝑡𝑜𝑚𝑠 𝑥 40𝑔/𝑚𝑜𝑙 1 𝑚𝑜𝑙 AgCl2 – white curdy ppt. Ans = 3.011 x 1023 atoms C. REACTIVITY SERIES Molarity/ Formality (M) 3N HCl or 3K HCl Metals 𝑛 𝑠𝑜𝑙𝑢𝑡𝑒 Li > K > Ba > Ca > Na > Mg > Al > Mn > Zn > Cr > Fe > 𝑀= Cd > Co > Ni > Sn > Pb > H2 > Cu > Ag > Hg > Pt > Au 𝐿 𝑠𝑜𝑙𝑛 Nonmetals (bases on electronegativity) Molality (m) F > Cl > Br > I 𝑛 𝑠𝑜𝑙𝑢𝑡𝑒 Examples: 𝑚= Co + MgCl2 → NR 𝑘𝑔 𝑠𝑜𝑙𝑣𝑒𝑛𝑡 Zn + CuSO4 → ZnSO4 + Cu NaBr + Cl2 → NaCl + Br Normality (N) 𝐺𝐸𝑊 𝑁= 𝑜𝑟 𝑀 𝑥 𝑓 D. NOMENCLATURE OF INORGANIC COMPOUNDS 𝐿 Factor (f) 1. Covalent compounds Acid (H): HCl = 1; H2SO4 = 2 CO: Carbon monoxide Base (OH): Al(OH)3 = 3 SiO2: Silicon dioxide Salt (M): Al2O3 = Al: 3x2 = 6; Na2SO4 = 2 N2O: Dinitrogen monoxide CCl4: Carbon tetrachloride V. ELECTRONIC CONFIGURATION 2. Ionic compounds Ex: Pb(NO3)4 Aufbau Principle Classical: Plumbic nitrate Atoms may be built by progressive filling of energy of main Stock: Lead(IV) nitrate energy sub level (i.e., levels of lower energy levels are occupied first) E. MONOATOMIC IONS s=2, p=6, d=10, f=14 A. Monovalent +1 = Group 1 (H, Li, Na, K ‫ ׀‬Ag) +2 = Group 2 (Be, Mg, Ca, Sr, Ba ‫ ׀‬Zn, Cd) -2 = Group 6A (Oxide, Sulfide) -1 = Group 7A (Fluoride, Chloride, Bromide, Iodide) Module 1 – General Chemistry Page 3 of 8 RJAV 2022 Ex. Calcium = Atomic # 20; Atomic mass 40 Quantum theories 1s2 2s2 2p6 3s6 4s2 [Argon] 4s2 1. Pauli’s exclusion theory No 2 e- will have same set of quantum number (“exclusive” Shortcut – Noble gas: He = 2 2. Heisenberg’s uncertainty theory Ne = 10 Impossible to predict/ accurately determine the particle’s Ar = 18 velocity (position & momentum) Kr = 36 Xe = 54 3. Hund’s rule Rn = 86 Orbitals are filled up singly before pairing up Most stable arrangement of e- in subshells is the one with greatest no. of parallel spins. VI. GAS LAWS Boyle's/Mariotte P₁𝑽₁ = 𝑷₂𝑽₂ 𝑜𝑟 𝑷 ∝ 1𝑣 Temperature (in K) Charles 𝑻₁ 𝑻₂ = 𝑜𝑟 𝑽 ∝ 𝑻 𝑽₁ 𝑽₂ Pressure (in atm) Gay-Lussac's 𝑷₁ 𝑷₂ = 𝑜𝑟 𝑷 ∝ 𝑻 𝑻₁ 𝑻₂ Volume (in L) Combined 𝑷₁𝑽₁ 𝑷₂𝑽₂ A. QUANTUM NUMBERS “fingerprints” = 𝑻₁ 𝑻₂ Principal Quantum Number (n = 1 to 7) Ideal main energy level; size of orbital (electron cloud), distance of e- 𝑷𝑽 = 𝒏𝑹𝑻 from nucleus 𝐿.𝑎𝑡𝑚 Ex. O2 = 1s2. 2s2. 2p4 (n=2) R = 0.08205 𝑚𝑜𝑙.𝐾 At STP Azimuthal/ Angular Momentum (ℓ = 0 to 3) T = 273.15 K Angular momentum & shape of orbital; subshell P = 1 atm ℓ = 0 ─ s : sharp (spherical shape) V = 22.4 L ℓ = 1 ─ p : principle (dumbbell shape) ℓ = 2 ─ d : diffuse (clover leaf) Avogadro’s ℓ = 3 ─ f : fundamental Equal volumes of different gases have same no. of moles at Ex. O2 = ℓ = 1 STP 𝑽₁ 𝑽₂ 𝑽 = 𝑜𝑟 𝑽 ∝ 𝒏 𝑜𝑟 = 𝒌 𝒏₁ 𝒏₂ 𝒏 Magnetic Quantum Number (mℓ = -ℓ, 0, +ℓ) k = 6.022 X 1023 Orientation of orbital in space Ex. O2 = mℓ = -1, 0, +1 Dalton’s Law of Partial Pressures n = 1, ℓ = 0 (s) : mℓ = 0 Total pressure in a mixture (non-interacting gases) is equal to n = 2, ℓ = 0,1 (s,p) : mℓ = -1,0,+1 [3 degenarate orbitals] = same energy levels the sum of the partial pressures of each gas. 𝑃𝑡 = 𝑃1 + 𝑃2 + 𝑃3 n = 3, ℓ = 0,1,2 : mℓ = -2,-1,0,+1,+2 [5 DO] n = 4, ℓ = 0,1,2,3 : mℓ = -3,-2,-1, 0, +1,+2,+3 [7 DO] Graham’s Rate of effusion (diffusion) and speed gas are inversely Magnetic Spin (ms = + ½ , - ½ ) proportional to the square root of their density providing the Magnetic moment/ Rotation temperature and pressure are same for 2 gases Spin Diffusion – rate at which 2 gases mix ↑ = Incomplete; clockwise + ½ Effusion – rate at which gas escapes through a pinhole ↑↓ = Complete; counterclockwise = - ½ vacuum. Rate 𝑑𝑖𝑓𝑓𝑢𝑠𝑖𝑜𝑛 ∝ 1√𝑑𝑒𝑛𝑠𝑖𝑡𝑦 Ex. Oxygen = ms= + ½ Fick’s 1st Law Diffusion rate (flux) of liquid or gas is directly proportional to the concentration gradient (ftom high concentration to low n ℓ mℓ ms concentration) 1 0=s 0 Henry’s Law of Gas Solubility 𝑷𝒓𝒆𝒔𝒔𝒖𝒓𝒆 ∝ 𝑺𝒐𝒍𝒖𝒃𝒊𝒍𝒊𝒕𝒚 2 0,1 = p -1,0,+1 Decrease temperature, Increase Pressure (i.e., sealed container), more CO2 is dissolved in water. 3 0,1,2 = d -2,-1,0,+1,+2 Real/Van der Waals 𝒂𝒏𝟐 4 0,1,2,3 = f -3,-2,-1, 0, +1,+2,+3 (𝑷 + ) (𝑽 − 𝒏𝒃) = 𝒏𝑹𝒕 𝒗𝟐 an2 = internal pressure per mole nb = incompressibility Raoult’s 𝑷𝒔𝒐𝒍𝒖𝒕𝒊𝒐𝒏 = 𝑿𝒔𝒐𝒍𝒗𝒆𝒏𝒕 𝑷𝒔𝒐𝒍𝒗𝒆𝒏𝒕 Magnetism types: X = mole fraction o Diagmagnetism – no unpaired e- Temperature o Paramagnetism – at least 1 unpaired e- Module 1 – General Chemistry Page 4 of 8 RJAV 2022  A. TEMPERATURE Entropy (∆S) = measure of system’s thermal energy per unit temperature; degree of disorderliness or randomness °C = (°F – 32) / 1.8 ∆S = (+) → spontaneous; increase (irreversible) – real case °F = (°C x 1.8) + 32 ∆S = (-) → non spontaneous; constant (reversible) – ideal case (in a K = °C + 273.15 steady state/ equilibrium) Absolute temperature ∆H → does not predict spontaneity 0K = absolute zero (lowest possible temperature) 3RD LAW: VII. SOLUTION If an object reaches absolute zero temperature (0 K = -273.15 = -459.67 °) Solute + Solvent Entropy of perfect, solid, crystalline substance is zero at Colligative properties (See Physical Pharmacy) absolute 0 temperature Dependent on the amount of solute present in the solution Gibb’s free energy (∆G) Vapor pressure lowering Thermodynamic state function that combines enthalpy and Boiling point elevation (Ebullition) entropy Freezing point depression ∆G = ∆H ‒ T∆S Osmotic pressure (Π) ∆G < 0 (-) → spontaneous ∆G > 0 (+) → non spontaneous Vapor Pressure Lowering ∆G = 0 → equilibrium (no more work to be done) Raoult's Law : vapor pressure of a solution is dependent IX. CHEMICAL KINETICS on the amount of nonvolatile solute added to solution Boiling Point Elevation ∆Tb = 𝑖𝐾𝑏𝑚 Study of reaction rates and reaction mechanism Freezing Point Depression ∆Tf = 𝑖𝐾𝑓𝑚 Reaction Rate (M/s) Change in concentration of a reactant or product concentration Osmotic Pressure (Π) Π= 𝑖𝑛𝑅𝑇 with time - pressure needed to stop osmosis 𝑉 aA + bB → cC + dD *small letters: coefficient that balance the chemical reaction 1 ∆[A] 1 ∆[B] 1 ∆[C] 1 ∆[D] VIII. THERMODYNAMICS 𝑅𝑎𝑡𝑒 = − =− = = 𝑎 ∆t 𝑏 ∆t 𝑐 ∆t 𝑑 ∆t Study of energy conversion/ transformation in the universe Rate Law Expresses relationship of the rate of reaction to the rate A. PARTS OF THE UNIVERSE constant (K) and concentration of reactants raised to some power A. System aA + bB → cC + dD Rate = K [A]x [B]y (xy=order of reaction) Open System - allows exchange of energy and matter Where x & y = order of reaction (0th, 1st, or 2nd) Closed System - allows exchange of energy but not matter\ Isolated System "Adiabatic Walls" - does not allow exchange A. REACTION RATE THEORIES of both energy and matter 1. Collision Theory B. Surrounding – everything outside the system rate of chemical reaction is proportional to the number of collisions per time B. PATH DEPENDENCE Requirements for effective collision: Proper orientation State Function Activation energy (Ea) – minimum amount of energy required to Independent (depends only on initial & final states of system) initiate chemical reaction Ex. Enthalpy (H), Internal energy (U), Gibb’s Free Energy (G), Entropy (S) 2. Transition Theory (Formation of Intermediate Complex) - rate depends on Ea Non-State Function required to form intermediate state (where new bonds are Dependent formed and old bonds are broken) Work and Heat B. FACTORS AFFECTING REACTION RATE (Directly Zeroth Law proportional) If two systems are in thermal equilibrium respectively with a third system, they must be in thermal equilibrium with each Nature of Reactants other = ↑ reactivity ↑ reaction rate (faster) a=c b=c Concentration of Reactants (except Zero order) a=b = ↑ concentration ↑ reaction rate C. LAWS OF THERMODYNAMICS Catalyst (Enzyme – Michaelis Menten Kinetics) = ↑ reaction rate 1ST LAW: Law of conservation of Energy Enzymes – speeds up the chemical reaction by lowering Ea Energy is neither created nor destroyed but can be transformed from one form to another Surface Area Enthalpy (H) = U, P, V = ↑ SA ↓ particle size ↑ reaction rate ↓ reaction time Hess' Law: ∆H is independent of reaction/steps that occurred (only the initial and final steps is the basis) Temperature q = Heat = ↑ Temp ↑ KE ↑ mobility of molecules ↑ collision ↑ reaction rate; ∆H = (+) → heat is absorbed; COLD (endothermic) Arrhenius Equation (T, Ea, RR) ∆H = (-) → heat is released; HOT (exothermic) X. CHEMICAL EQUILIBRIUM 2ND LAW: Law of Entropy No way but UP aA + bB ⇌ cC + dD For an isolated system, Total entropy can never decrease over time Module 1 – General Chemistry Page 5 of 8 RJAV 2022  A. LAW OF MASS ACTION Acids & Bases Hard Soft Ionic radius Small Large reaction rate proportional to the product of the concentrate of Oxidation states High Low the reactants to the power of its coefficient in a balanced Polarizability Low High equation Electronegativity High Low 𝐶 𝑐 𝐷𝑑 Ex. Ions of alkali & Heavy metals: 𝐾𝑒𝑞 = 𝑎 𝑏 𝐴 𝐵 alkaline earth Ag+, Au+, Hg22+/ Keq = 1: No shift (in equilibrium) metals, H+, NH4, Hg2+, Cd2+ Keq > 1: Favors product formation (to the right/ forward reaction) Ti4+, Cr3+ Keq < 1: Favors reactant formation (to the left/ backward or reverse reaction) OH-, F-, Cl-, CO32-, H- (Hydride), I-, CH3COO- SCN- Le Chatelier’s Principle (# stress reliever) If an external stress is applied to a system at equilibrium, the A. ACID-BASE FORMULA system adjusts in such a way that stress is partially offset as the system reaches new equilibrium Acids & Bases General Ionic equilibria formula For Weak Acids & Bases (with External Stressors: Equilibrium Change constant) shift Equilibrium pH = -log[H+] pKa = -log[ka] constant (Kc) - 25°C pOH = -log[OH-] pKb = -log[kb] Concentration Yes No pH + pOH = 14 pKa + pKb = 14 Pressure & Volume Kw = [H+][OH-] = 1x10-14 Kw = Ka x Kb = 1x10-14 (For Gases only) *pKa is constant, while pH varies []–M ↑ P = shift to side with lesser gas Yes No Ionic equilibria moles Ex. HA (aq) + H2O (I) ↔ H2O H3O+ (aq) + A- (aq) ↓ V = shift to side [𝐻 +][𝐴−] with greater gas 𝐾𝑎 = [𝐻𝐴] moles Conjugate Base = -1H+ to subs in question Temperature Conjugate Acid = +1H+ to subs in question Yes Yes ↑ T, Endo: Right In writing Equilibrium constants: ↑ T, Exo: Left - Aq. & gaseous reacting spp. (constant: S & L) Catalyst No No - Unit should be in M XI. ACIDS AND BASES Common ion effect Acids Bases Addition of compound having an ion in common with the dissolved Taste Sour Bitter substance will result to: pH 7 Equilibrium shift (either to the left or right) + Litmus paper Red Blue Suppressed ionization of the dissolved substance (WA or WB) + Metals → H2 gas (corrode pH change - metals) + Carbonates/ → CO2 gas Ex. CH3COOH – CH3COONa Mixture - Bicarbonates (effervescence) CH3COONa (s) → Na+ (aq) + CH3COO- (aq) - ↑ pH + Fat → Soap: Slippery CH3COOH (aq) ↔ H+ (aq) + CH3COO- (aq) - suppressed (Saponification) ionization - Henderson-Hasselbalch/ Buffer pair equation Ex. Manuf NaOH – Hard soap KOH – Soft soap For buffer solutions (WA + CB or WB + CA) HAc + Ac- Theories Acid Base NH3 + NH4+ Arrhenius Liberates H+ Liberates OH- 𝑠𝑎𝑙𝑡 Bronsted-Lowry Donates p+ Accepts p+ Weak acids 𝑝𝐻 = 𝑝𝐾𝑎 + 𝑙𝑜𝑔 𝑎𝑐𝑖𝑑 Lewis e- pair acceptor e- pair donor 𝑏𝑎𝑠𝑒 Weak bases 𝑝𝐻 = 𝑝𝐾𝑏 + 𝑙𝑜𝑔 𝑠𝑎𝑙𝑡 Lewis Theory: *pH = pKa (@ half neutralization point) Lewis’s acid/ Electrophile e- loving Buffer Solution (+) iron or metal (e- poor spp.) has the ability to resist changes in pH upon addition of small amounts of either acid or base Lewis’s base/ Nucleophile Weak acid and its CB (salt of WA) (-) ion or nonmetal (e- rich spp.) Weak base and its CA (salt of WB) Lewis’s acid (+) Lewis’s base (-) Van slyke Ni + CO Ni CO Buffer capacity Bmax (degree or magnitude of capability to resist Cl- + SnCl4 SnCl4 Cl- change in pH of the buffer) Pearson's Hard and Soft Acid and Base (HSAB) XII. SOLUBILITY PRODUCT CONSTANT (Ksp) Hard-Hard/ Soft-Soft ↑ Ksp = ↑ solubility Thermodynamically Stronger interaction Hard Acid + Hard Base → Ionic complexes Solubility (g/L) Soft Acid + Soft Base → Covalent complex Number of grams of solute dissolved in in 1L of saturated solution Hard-Soft/ Soft-Hard Thermodynamically Weaker interaction Molar solubility (mol/L) Number of moles of solute dissolved in 1L of saturated solution Module 1 – General Chemistry Page 6 of 8 RJAV 2022 Predicting formation of precipitate formation (Q ion product constant) – computed based on initial concentration: Q < Ksp → unsaturated Q = Ksp → saturated Q > Ksp → supersaturated ↓↓ Noyes Whitney equation Dissolution rate is directly proportional to the solute surface area, solute concentration at boundary layer, and diffusion coefficient XIII. ELECTROCHEMISTRY study of the production of electricity from energy released during spontaneous and nonspontaneous chemical reactions 1. Spontaneous Voltaic cells/ galvanic cells Johan Wolfgang Dobereiner REDOX reaction (Anode - Oxidation; Cathode – Reduction) "Law of Triads" Electrons migrate from Anode → Cathode John Newlands 2. Nonspontaneous "Law of Octaves" Electrolytic cells: Electric current is applied to remove e- and Periods transfer to another cell (Electroplating) Dmitri Mendeleev "Father of Modern Periodic Table" XIV. PERIODIC TABLE (Lothar Meyer) Atomic Mass/ Weight Antoine Lavoisier Henry Moseley first extensive list of elements (~ 33) "Created Modern Periodic Table" Metals vs Nonmetals property varies with increasing atomic number Metal Nonmetal Glenn Seaborg Oxides Basic Acidic Discovered transuranic elements. > Uranium Actinides below Good Reducing agent Oxidizing agents lanthanides (exhibit radioactivity; unstable proton-to-neutron Conductor ✓ x ratio) Malleable ✓ Brittle Ductile ✓ x Law of Octaves Metallic luster ✓ X (except I2) Every 8th element similar physicochemical property when State at RT Solid (except Hg) Solid, liquid, gas arranged according to increasing Atomic weight (Ex. H, F, Cl) Octet rule Hg – only liquid metal Elements (Atomic nos. 1-20) with < 8 electron “react” to Amphoteric – can act as acid or base achieve 8 electrons (stable) Malleable – ability to be pounced into thin sheets Ductile – ability to be drawn into wires Valence e- – electron found in outermost shell A. PERIODIC TABLE Group A Valence Valence/ Elements (together e- charge with other Elements: 118 valences, if any) Periods (Horizontal rows): 7 1A (Alkali M) 1 +1 H, Li, Na, K, Rb, Cs, Groups/ Family (Vertical columns): 18 Fr, NH4 Groups A: Representative elements (s & p block) 2A (Alkaline Earth M) 2 +2 Be, Mg, Ca, Sr, Ba, Ra Groups B: Transition elements (d block) 3A (Boron G.) 3 +3 B, Al, Ga, In, Tl Actinides & Lanthanides: Inner transition elements (f block) 4A (Carbon G.) 4 (+/- 4) C, Si │Ge, Sn, Pb (+2, +4) B. PERIODIC TRENDS 5A (Nitrogen G.) 5 -3 N, P │As, Sb, Bi (+3, +5) Ionization energy 6A (Oxygen G./ 6 -2 O, S, Se, Te, Po energy needed to remove outermost electron in neutral atom Chalcogens) →↑ 7A (Halogens) 7 -1 F, Cl, Br, I, At 8A/ 0 (Inert/ Noble/ 8 0 He, Ne, Ar, Kr, Xe, Stable Gases Rn Electron affinity energy given off when neutral atom gains extra electron Group B Valence Elements (making it more negative) →↑ 1B (Coinage M.) +1 Cu (+2), Ag, Au (+3) 2B (Volatile M.) +2 Zn, Cd, Hg & Hg2 Electronegativity 3B (Scandium Subgrp) - Sc, Y, Lanthanides (La-Lu), ability of an atom to attract electron pair to itself, forming Actinides (Ac-Ir) covalent bond. →↑ 4B (Titanium Subgrp) - Ti, Zr, Hf F: most electronegative (most reactive Oxidizing agent) 5B (Vanadium Subgrp) - V, Nb, Ta O2: 2nd most electronegative 6B (Chromium Subgrp) - Cr, Mo, W Note: Decrease – Top to Bottom; Increase – Left to Right 7B (Manganese - Mn, Tc, Re, Bh Subgrp) Atomic radius 8B (Iron Triad) - 1st Triad: Fe, Co, Ni (+2, +3) ½ difference between nucleus of 2 atoms ↓← 2nd Triad (Light): Rh, Ru, Pd 3rd Triad (Heavy): Os, Ir, Pt Metallic property New Elements Note: Increase – Top to Bottom; Decrease – Left to Right Nihonium 113Nh Atomic radius/ Metallic property Moscovium 115Mc Tennessine 117Ts Oganesson 118Og Module 1 – General Chemistry Page 7 of 8 RJAV 2022  XV. RADIOACTIVITY Spontaneous emission of particles/ ionizing radiation by unstable nuclei of heavier elements (p+-to-n0 ratio) (atomic # 92 and above: transuranic elements) Non-SI: Curie (Ci) │1 Ci = 3.7 x 1010 decay/ sec Discovered: Po & Ra SI: Becquerel (Bq) │1 Bq = 1 decay/ sec R.E.M. (roentgen equivalent in man) Unit of radiation damage Rad/ gray Unit of amount of exposure to radiation Radioactive Emissions Radioisotopes decay RANDOMLY Beta & Gamma can penetrate body tissue Rays/ Mass Velocity Penetrating Prevented Decay power by Alpha Heaviest Slowest Low Paper (4) (0.1 speed of light) Beta Light Fast Medium Al (1/2000) (0.9 speed of light) Gamma No mass Fastest High Pb & charge (speed of light) 0 A. MODES OF DECAY Alpha (α) decay is the emission of α particle from the nucleus. For example, polonium-210 undergoes α decay: 210 4 206 210 4 206 Po⟶ He + Pb or Po ⟶ α + Pb 84 2 82 84 2 82 Beta (β) decay is the emission of an electron from a nucleus. Iodine- 131 is an example of a nuclide that undergoes β decay: 131 0 131 131 0 131 I⟶ e + Xe or I ⟶ β + Xe 53 -1 54 53 -1 54 Gamma emission (γ emission) is observed when a nuclide is formed in an excited state and then decays to its ground state with the emission of a γ ray, a quantum of high-energy electromagnetic radiation. The presence of a nucleus in an excited state is often indicated by an asterisk (*). Cobalt-60 emits γ radiation and is used in many applications including cancer treatment: 60 0 60 Co* ⟶ γ + Xe 27 0 27 Positron emission (β+ decay) is the emission of a positron from the nucleus. Oxygen-15 is an example of a nuclide that undergoes positron emission: 15 0 15 15 4 15 O⟶ e + N or O ⟶ β + N 8 +1 7 8 +1 7 Electron capture occurs when one of the inner electrons in an atom is captured by the atom’s nucleus. For example, potassium-40 undergoes electron capture: 40 0 40 K + e ⟶ Xe 19 -1 18 Module 1 – General Chemistry Page 8 of 8 RJAV 2022