DATOATGENERALCHEMISTRYOUTLINES2-210826-081643.pdf

DAT / OAT GENERAL CHEMISTRY OUTLINES Table of Contents 1 - General Concepts & Laboratory Concepts Page 1 2 - Stoichiometry Page 5 3 - Atomic and Molecular Structure Page 7 4 - Periodic Properties...

DAT / OAT GENERAL CHEMISTRY OUTLINES Table of Contents 1 - General Concepts & Laboratory Concepts Page 1 2 - Stoichiometry Page 5 3 - Atomic and Molecular Structure Page 7 4 - Periodic Properties Page 12 5 - Gases Page 14 6 - Liquids and Solids Page 16 7 - Solutions Page 18 8 - Chemical Kinetics Page 20 9 - Chemical Equilibrium Page 22 10 – Acids & Bases Page 24 11 - Thermodynamics and Thermochemistry Page 28 12 - Oxidation-Reduction Reactions Page 31 13 - Nuclear Reactions Page 34 ChadsPrep.com 0  1 - General Concepts & Laboratory Concepts General Concepts Atoms, Molecules, and Ions Elements vs Compounds Ionic compounds 1) Name the metal (or polyatomic cation) Polyatomic Ions 2) State its oxidation state as a roman numeral in parenthesis SO42- sulfate (except Group I/II, Al, Zn, Cd, Ag) SO32- sulfite 3) Name the non-metal with an –ide ending (or name the polyatomic anion) NO3- nitrate NO2- nitrite PO43- phosphate CO32- carbonate HCO3- bicarbonate OH- hydroxide MnO4- permanganate CrO42- chromate Cr2O72- dichromate NH4+ ammonium Molecular compounds CN- cyanide 1) Give the numerical prefix of the first element (omit if there is only one) 2) Name the first element ClO- hypochlorite 3) Give the numerical prefix for the second element ClO2- chlorite 4) Name the second element with the –ide ending ClO3- chlorate ClO4- perchlorate (for Br and I also) Acids 1) Binary Acids (only hydrogen and one other element) Hydro-element-ic acid 2) Oxyacids HClO hypochlorous acid HNO2 nitrous acid HClO2 chlorous acid HNO3 nitric acid HClO3 chloric acid HClO4 perchloric acid ChadsPrep.com 1 Metric Units and Conversions Metric Prefixes 1cm3 = 1mL Terra 1012 Giga 109 Convert 1.3x108ns to seconds. Mega 106 kilo 103 centi 10-2 milli 10-3 icro 10-6 nano 10-9 pico 10-12 Convert 247,000 mg to kg. femto 10-15 Convert 2.0 miles to mm (1in = 2.54cm, 1mile = 5280ft). Density (d = m/v) What is the density (in g/cm3) of 3kg cube with 10cm edges? Atomic/Molecular Weights Bromine has two abundant isotopes; 50% of Br atoms have a mass of 79a.m.u, 50% have a mass of 81 a.m.u. What is the atomic mass (also called atomic weight) of bromine? ChadsPrep.com 2 Laboratory Concepts Significant Figures Zeroes in between significant figures are significant Zeroes right of the decimal, at the end are significant Precision and Accuracy Accuracy – how close to the true value Precision – how close measurements are to each other Weight Measurements -using an analytical balance -how to remove from the stock bottle -took a little too much? Volume Measurements Volumetric Pipets (0.006 to 0.05ml) Burets (0.02 to 0.20ml) Volumetric Flasks (0.02 to 0.5ml) Graduated Cylinder (0.1 to 0.2ml) Beakers/flasks -Measuring at the meniscus -How to determine the density of a solid ChadsPrep.com 3 Dilutions M1V1 = M2V2 (or C1V1 = C2V2) % Error Calculations 𝐴𝑐𝑡𝑢𝑎𝑙 − 𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 % 𝐸𝑟𝑟𝑜𝑟 = × 100 𝑇ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 Spectrophotometer Abs = cl Abs = Absorbance  = molar extinction coefficient (molar absorptivity) l = path length pH probes -2 and 3 point calibrations are common (use calibration buffers of known pH) -range of calibration -rinse with deionized water in between samples -dab dry with a kimwipe Safety How to heat a test tube How to dilute strong acids (A in W) Data Analysis Directly Proportional vs Inversely Proportional ChadsPrep.com 4  2 - Stoichiometry Empirical vs Molecular Formulas What is the empirical formula of P4O10? What is the empirical formula of C3H8O2? Percent Composition What is the percentage of carbon in a 300g sample of CaCO3? What is the %O? A certain compound is 80% carbon and 20% hydrogen by mass and has a molecular weight of 30a.m.u. a) What is its empirical formula? b) What is its molecular formula? The MOLE 1 mole = 6.02x1023 (It’s just a big number!) -Molecular Weight vs Molar Mass You have 300g CaCO3. 1) How many moles of CaCO3 do you have? 2) How many oxygen atoms are in the sample? 3) How many grams of carbon are in the sample? ChadsPrep.com 5 Balancing Reactions __C5H12 + __O2 → __CO2 + __H2O Limiting Reactants and Theoretical Yield Calculations N2(g) + 3H2(g) → 2NH3(g) If 112g of N2 reacts completely with 30.0g H2, 1) What is the limiting reactant? 2) What is the theoretical yield? 3) How much of the reagent in excess is left over? 4) If 112g of N2 reacts with 30.0g of H2 to produce 68g NH3, what is the percent yield? ChadsPrep.com 6  3 - Atomic and Molecular Structure Atomic and Electronic Structure Atomic Structure - nucleus composed of protons and neutrons surrounded by an electron cloud Dalton proposed his atomic theory Thompson discovered electrons Millikan determined the charge of an electron (oil drop experiment) Rutherford discovered the nucleus (scattering of alpha particles in thin gold foil) Isotope symbols (atomic number, mass number, isotopes) Atomic Orbitals Quantum Numbers # NAME WHAT? RANGE n principal Shell [1…] l azimuthal Subshell (type of orbital) [0…(n-1)] ml magnetic Specific orbital (orientation in space) [-l…+l] ms spin Up or down +1/2 or -1/2 ChadsPrep.com 7 Ground State Electron Configurations -Standard and Noble Gas Configurations Aufbau Principle – electrons fill lowest energy orbitals first Hund’s Rule – degenerate orbitals each get an electron -Exceptions (Cr, Mo and Cu, Ag, Au) before pairing Pauli Exclusion Principle – no 2 electrons in an atom have the same 4 quantum numbers -Ions (transition metals too) -Number of valence electrons -Ground state vs. excited state Paramagnetic – attraction to a magnetic field; occurs due to the presence of unpaired electrons Diamagnetic – very slight repulsion to a magnetic field; occurs when all electrons are paired ChadsPrep.com 8 Electromagnetic Radiation  = c  = wavelength = frequency c = 3.0x108m/s (speed of light) hc Ephoton = h = h = 6.626x10-34J.s (Planck’s constant)  Bohr Model of the Atom − 2.18x10 −18 J En = ( ) (Electron energies in a hydrogen atom) n2 Absorption and Emission Line Spectra E = Ephoton Lyman (nf = 1), Balmer (nf = 2), Paschen (nf = 3), Bracket (nf = 4) series Wave Behavior of Matter (de Broglie Relation) h = m = mass v = velocity  = wavelength h = 6.626 × 10−34 J ∙ s mv Heisenberg Uncertainty Principle ∆x∆p ≥ constant The position and momentum of a particle cannot simultaneously be known with infinite precision. Photoelectric Effect K.E.electron = Ephoton -   = work function i.e. the minimum energy to ionize an electron ChadsPrep.com 9 Molecular Structure and Geometry Bonding Ionic Bond – transfer of electrons between elements with significant differences in electronegativity -usually a metal with a non-metal (but don’t forget about polyatomic ions) Covalent Bond – sharing of electrons between elements with similar differences in electronegativity -between two non-metals -polar covalent bonds—partial ionic character IONIC MOLECULAR NETWORK COVALENT METALLIC High m.p and b.p. Lower m.p and b.p. High m.p and b.p. Conduct electricity Brittle C(diamond) Conduct heat Lattice energy  q1q2/r C(graphite) Luster SiO2(quartz) Malleable Ductile Lattice Energy NaCl(s) → Na+(g) + Cl-(g) -higher for higher charges and smaller ions Molecular Structures (Lewis Dot Structures) Octet Rule Exceptions 1) Third row and lower can exceed the octet rule 2) H only 2, Be only 4, B and Al only 6 (usually) 3) Molecules with odd numbers of electrons -Formal Charge [Normal Valence – (“dots and lines”)] DRAWING LEWIS STRUCTURES 1 Set up skeleton with single bonds (central atom is the atom which can make the most bonds). 2 Fill the octets of the outside atoms with lone pairs of electrons. 3 Any remaining valence electrons go on central atom. 4 Once all electrons are used, form multiple bonds to central atom if not “full.” CCl4 NF3 HCN CO2 SF4 NO3- ChadsPrep.com 10 Molecular Geometry VSEPR Theory (Valence Shell Electron Pair Repulsion) Electron Electron Domain Bond Angles Hybridization of Domains Geometry Central Atom 2 Linear 180 sp 3 Trigonal Planar 120 sp2 4 Tetrahedral 109.5 sp3 5 Trigonal Bipyramidal 90, 120, 180 sp3d 6 Octahedral 90, 180 sp3d2 TABLE OF MOLECULAR GEOMETRIES Electron No lone pairs 1 lone pair 2 lone pairs 3 lone pairs Domains 2 3 4 5 6 Polarity Possibly Nonpolar - Linear, Trigonal Planar, Tetrahedral, Trigonal Bipyramidal, Octahedral, Square Planar Sigma and Pi Bonds Sigma Bonds – result from the end-to-end overlap of atomic orbitals (including hybrid orbitals) -All single bonds are sigma bonds. -The first bond of a double or triple bond is also a sigma bond. Pi Bonds – result from the sideways overlap of ‘p’ orbitals -The second bond in a double bond or the second and third bond in a triple bond ChadsPrep.com 11  4 - Periodic Properties The Periodic Table and Descriptive Chemistry PROPERTIES OF VARIOUS GROUPS FROM THE PERIOD TABLE Alkali Metals highly reactive with water due to low ionization energies; form ionic compounds Alkaline Earth Metals lower ionization energies; reactive with water (increasing reactivity down the group) Halogens have high (very negative) electron affinities Noble Gases Chemically Inert (mostly) Transition Metals Often form brightly colored compounds Metal oxides are basic (examples: Na2O, BaO) Non-metal oxides are acidic (examples: NO2, SO3, CO2) Periodic Trends Atomic Size (Radius) Effective Nuclear Charge Zeff = Z – S Z = # protons S = # shielding electrons Bond Length Ionic Radii Isoelectronic Series S2- > Cl- > Ar > K+ > Ca2+ ChadsPrep.com 12 Ionization Energy Ionization Energy - Energy required to remove an electron 2nd, 3rd, 4th, etc. Ionization Energies Electron Affinity Electron Affinity - Energy change for gaining an electron Electronegativity ChadsPrep.com 13  5 - Gases Characteristics of Gases Temperature C vs K Pressure P = F/A 1atm = 760 torr = 760mmHg = 101,325Pa -molecular basis of pressure Volume Liters; 1cm3 = 1mL -Mercury Manometer and Barometer ChadsPrep.com 14 Kinetic Molecular Theory (Ideal Gas Assumptions) 1) The volume of gas molecules is insignificant. -most accurate at low pressures 2) No intermolecular attractions between molecules; i.e. all collisions are elastic -most accurate at high temperatures and for gases with lower intermolecular forces 3) Avg. Kinetic energy  T 3𝑅𝑇 K.E.avg = 3⁄2 RT (per mole of gas) 𝑣𝑟𝑚𝑠 = √ 𝑀 Gas Laws Ideal Gas Law PV = nRT 1 Boyle’s Law P V Charle’s Law VT Avogadro’s Law Vn P1V1 P2V2 Combined gas law = n1T1 n2T2 1 mole gas = 22.4L @ STP Dalton’s Law of Partial Pressures PTot = PA + PB + … PA = APTot A = mol fraction A Graham’s Law of Effusion r1 M2 = r2 M1 Real Gases Van der Waals equation:  an 2  an 2  P + 2 (V − nb) = nRT + corrects for intermolecular attractive forces V2  V  −6nb- Liquids and corrects for Solids (because molecules do have volume) repulsions ChadsPrep.com 15  6 – Liquids & Solids Intermolecular Forces DECREASING STRENGTH Hydrogen Bonding -- F-H, O-H, or N-H bond required in a pure substance F, O, or N to hydrogen bond with H2O Ion-Dipole Forces – Intermolecular force between ions and polar substances (ex. hydration of NaCl) Dipole-dipole -- Intermolecular force between polar molecules Ion-Induced Dipole Forces – Intermolecular force between ions and nonpolar substances London Dispersion Forces (Van der Waals) -A temporary or transient dipole -All molecules have London Dispersion forces. -The greater the weight and surface area, the greater the London dispersion forces Higher IM Forces Lead to: Higher Boiling Pt Higher Heat of Vaporization Higher Viscosity Higher Surface Tension Lower Vapor Pressure General Rule for solubility: “Like dissolves like.” Structures of Solids Cubic Unit Cells 1) Simple Cubic – atoms at the corners of the unit cell 1 atom per unit cell 2) Body-Centered Cubic – atoms at the corners and body center of the unit cell 2 atoms per unit cell 3) Face-Centered Cubic - atoms at the corners and face centers of the unit cell 4 atoms per unit cell Close Packing 1) Hexagonal Close Packing (ABAB…) 2) Cubic Close Packing (ABCABC…) ChadsPrep.com 16 Phase Changes Vapor Pressure and Boiling Pt. Phase Diagrams Be able to identify the following: -lines of equilibrium Liquid Critical -triple point Solid Pressure Point -critical point -normal m.p. and b.p. 1 atm Gas Triple Point Temperature -phase diagrams for CO2 and H2O ChadsPrep.com 17  7 - Solutions Vocabulary Solvent/solute – solute is dissolved in the solvent Saturated – maximum amount of a solute is dissolved in a solvent Unsaturated – less than the maximum amount of a solute is dissolved in a solvent Supersaturated – more than the maximum amount of a solute is dissolved in a solvent Strong Electrolytes – dissociate into ions completely in H2O -include strong acids, strong bases and soluble ionic compounds Weak Electrolytes -- partially dissociate into ions in H2O -include weak acids and weak bases Nonelectrolytes – don’t dissociate into ions to any significant extent in H2O Concentration 𝑚𝑜𝑙𝑒𝑠𝑠𝑜𝑙𝑢𝑡𝑒 𝑚𝑜𝑙𝑒𝑠𝑠𝑜𝑙𝑢𝑡𝑒 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦 (𝑀) = 𝑚𝑜𝑙𝑎𝑙𝑖𝑡𝑦 (𝑚) = 𝐿𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑘𝑔𝑠𝑜𝑙𝑣𝑒𝑛𝑡 Solubility of Ionic Compounds in Water 1) All Group I metal, NH4+, NO3- ClO4-, and C2H3O2- (acetate) salts are soluble. 2) Most Ag+, Pb2+, Hg22+ and S2- salts are insoluble. AgNO3 + KCl → AgCl + KNO3 What is the net ionic equation? Solubility of Solids and Gases in Liquids 1) Solids are typically more soluble at higher temperatures. 2) Gases are less soluble at higher temperatures. 3) Gases are more soluble at higher pressures. Henry’s Law PA = kH[A] ChadsPrep.com 18 Colligative Properties van’t Hoff factor (i) – number of ions a compound dissociates into per formula unit 1) Freezing Pt. Depression TF = −iK F m 2) Boiling Pt. Elevation TB = iK B m 3) Vapor Pressure Depression (Raoult’s Law) PA =  A PA* (*Denotes pure) 4) Osmotic Pressure  = iMRT ChadsPrep.com 19  8 - Chemical Kinetics Thermodynamics – spontaneity, heat change, equilibrium, etc. but not how fast (can’t tell us about the rate) Kinetics – study of the rates of reactions Rate Expressions 2NH3(g) → N2(g) + 3H2(g) − [ NH 3 ] [ H 2 ] [ N 2 ] Rate = = = 2t 3t t Rate Laws Overall Reaction: A + 2B → C rate = k[A]m[B]n -‘m’ and ‘n’ are the reaction orders and are often integers (0, 1, and 2 are most common) -orders and rate constant must be determined experimentally (except for elementary reactions) [A]0 (M) [B]0 (M) Initial Rate (M/s) 0.20 0.10 0.011 0.20 0.30 0.033 0.40 0.10 0.044 What is the rate law for this reaction? What is the overall reaction order? What is the rate constant? Reaction Coordinate Diagrams IDENTIFY 1) H or G 2) Ea (G‡) 3) Transition states (activated complex) 4) Intermediates 5) Slow step (RDS) ChadsPrep.com 20 Catalyst – 1) speeds up a reaction 2) by lowering the activation energy 3) by providing an alternate mechanism (pathway) for the reaction to occur -A catalyst is NOT consumed in the reaction. -A catalyst does NOT shift the equilibrium (you just reach equilibrium faster) Collision Theory (Requirements for a Reaction) 1) Collision INCREASE IN TEMPERATURE 2) Proper orientation of reactant molecules 1) Increases the collision frequency 3) Sufficient energy for a reaction to occur 2) Increases the percentage of high energy collisions. k = Ae-Ea/RT (Arrhenius Equation) Mechanisms -Elementary Reactions -Unimolecular vs. Bimolecular vs. Termolecular X + X → X2 (slow) X + X → X2 (fast or rapid equilibrium) X2 + Y → Z (fast) X2 + Y → Z (slow) What is the overall reaction? What is the intermediate? What are the overall rate laws respectively? ChadsPrep.com 21  9 - Chemical Equilibrium Definition of Dynamic Equilibrium: FORWARD RATE = REVERSE RATE Equilibrium Constant (Law of Mass Action) [𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠] 𝑃𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠 𝑘𝑓 Keq Meaning 𝐾𝑐 = 𝐾𝑝 = 𝐾𝑒𝑞 = [𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠] 𝑃𝑟𝑒𝑎𝑐𝑡𝑎𝑛𝑡𝑠 𝑘𝑟 K >> 1 (>103) Products favored at eq. K Ssolid Saq > Ssolid Srxn > 0 Increase in number of moles of gas S = nSproducts - nSreactants Srxn > 0 Sublimation, vaporization, fusion Srxn > 0 s → aq ChadsPrep.com 28 Enthalpy (H) Endothermic (H > 0) vs Exothermic (H < 0) Hess’s Law 3NO(g) → N2O(g) + NO2(g) Hº = ? 2NO(g) + O2(g) → 2NO2(g) Hº = -113kJ 2N2(g) + O2(g) → 2N2O(g) Hº = 163kJ N2(g) + O2(g) → 2NO(g) Hº = 181kJ Enthalpies of Formation Ho = nHof,products - nHof,reactants N2H4(g) + O2(g) → N2(g) + 2H2O(g) Hº = ? Hfº (kJ/mol) N2H4(g) 95.4kJ H2O(g) -242kJ Formation Reactions Forms one mole of a single substance from its component elements in their standard states. STANDARD STATES Gases He, Ne, Ar, Kr, Xe, Rn, H2, N2, O2, F2, Cl2 Liquids Br2, Hg Solids Everything Else Special C(graphite), S8 Bond Dissociation Energy Bond breaking is endothermic Bond making is exothermic H = Dreactants - Dproducts = Dbroken - Dformed N2H4(g) + O2(g) → N2(g) + 2H2O(g) H = ? Bond Enthalpies (kJ/mol) N-H 391 O-H 463 O=O 495 N-N 163 N≡N 941 ChadsPrep.com 29 Gibbs Free Energy (G) -A measure of spontaneity MEANING OF G VALUES G < 0 Spontaneous G > 0 Nonspontaneous G = 0 At Equilibrium G = G + RTlnQ (G under nonstandard conditions) Standard Conditions (º) G = -RTlnKeq All aqueous species 1M G and Keq All gaseous species 1atm G < 0 Keq > 1 T = 298K (not really) G > 0 Keq < 1 STP G = 0 Keq = 1 T = 273K P = 1atm Go = Ho - TSo Ho So -TSo - + - Spontaneous at all temperatures + - + Nonspontaneous at all temperatures - - + Spontaneous at low temperatures + + - Spontaneous at high temperatures ChadsPrep.com 30  12 - Oxidation-Reduction Reactions Vocabulary Oxidation – loss of electrons LEO says GER Reduction – gain of electrons OIL RIG Oxidizing Agent (Oxidant) – species that is reduced Reducing Agent (Reductant) – species that is oxidized Anode – the site of oxidation An Ox Cathode –the site of reduction Red Cat Oxidation States 1) Elements in their elemental form are in the zero oxidation state. 2) Group 1 metals are +1 and Group 2 metals are +2 in compounds. 3) Hydrogen is +1 except when bonded to metals (when it’s –1). 4) Transition elements must be determined from anion’s charge (except Al=+3, Zn=+2, Cd=+2, Ag=+1). 5) The most electronegative elements get their typical oxidation state. 6) The last element not assigned balances the charge of the compound/ion. Balancing Redox Rxns Cr3+ + Zn  Cr + Zn2+ MnO4-(aq) + Br-(aq)  MnO2(aq) + Br2(l) ChadsPrep.com 31 Electrochemical Cells 1) The anode is ALWAYS the site of oxidation. (an ox) 2) The cathode is ALWAYS the site of reduction. (red cat) 3) Electrons ALWAYS flow from anode to cathode through the wire. 4) Anions flow to the anode and cations to the cathode through the salt bridge. 5) In galvanic/voltaic cells (spontaneous), the cathode is + and the anode is –. In electrolytic cells (nonspontaneous), the cathode is – and the anode is +. 6) The cathode gains mass, while the anode loses mass (for metal/metal salt galvanic cells). GALVANIC (VOLTAIC) CELLS ELECTROLYTIC CELLS Spontaneous (G < 0, Ecell > 0) Nonspontaneous (G > 0, Ecell < 0) Produce Electricity Consume Electricity Function as Batteries Produce Elements Standard Cell Ptoentials ( = red + ox or  = cat - an) REDUCTION POTENTIALS The Standard Hydrogen Electrode (SHE) K+ + 1e- → K -2.93V 2H+ + 2e- → H2 0.00V Na + 1e → Na + - -2.71V Determine standard cell potentials of the following: Al + 3e → Al 3+ - -1.66V Fe2+ + Cu → Cu2+ + Fe Zn + 2e → Zn 2+ - -0.76V Cr + 3e → Cr 3+ - -0.74V Fe + 2e → Fe 2+ - -0.44V Ni + 2e → Ni 2+ - -0.25V 2H + 2e → H2 + - 0.00V 2Cr + 6H+ → 3H2 + 2Cr3+ Cu2+ + 2e- → Cu +0.34V Ag + 1e → Ag + - +0.80V I2 + 2e → 2I - - +0.54V ChadsPrep.com 32 Non-standard cell potentials 0.0592 E = E − log Q (Nernst Equation) n Qualitative Effects according to Le Chatelier’s Principle 1) Shift to the right – increases potential 2) Shift to the left – decreases potential Reduction Potentials What is the strongest oxidizing agent? What is the strongest reducing agent? Which pairs will react spontaneously? Electrolytic Cells K+ + 1e- → K -2.93V 2I-  I2 + 2e- -0.54V 2H2O + 2e  H2 + 2OH - - -0.83V 2H2O  O2 + 4H+ 4e- -1.23V What are the products of electrolysis of KI(l)? What are the products of electrolysis of KI(aq)? Quantitative Calculations (Faraday’s Law) (I)(t s )(MWproduct ) = mass of product (n)(F) (I)(t s ) = moles of product (n)(F) ChadsPrep.com 33  13 - Nuclear Reactions The Nucleus Nuclear Structure – protons and neutrons Atomic Number – number of protons or charge Atomic Weight vs Mass Number Isotopes – same element but different mass number Nuclear Reactions Parent → daughter(s) + energy mass is always lost (E = mc2) 233 92U → 42α + ? Radioactive Decay ROUTE OF DECAY RESULT LIKELY FOR? -decay Reduces Mass # Large nuclei (Z>83) - decay ( emission) n→p N/Z is too high + decay (Positron emission) p→n N/Z is too low Electron Capture p→n N/Z is too low -decay Kinetics of Radioactive Decay Always 1st order (N = N0e-kt or lnN = lnN0 – kt) Half-life (fractions, percents, masses, activities) Nuclear Binding Energy mass defect (m) – a nucleus always weighs less than its constituent nucleons E = mc2 m must be in kg 56 Fe has the highest nuclear binding energy per nucleon Fission and Fusion Fission – the splitting of larger nuclei (Z>56) into smaller nuclei (nuclear reactors, fission bombs) Fusion – the combination of smaller nuclei (Z