Chem I and II All Notes
Chem I and II All Notes CHEM 1210
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Chemistry I and II Notes Naming Prefixe Number of s Atoms I. Naming molecular Mono 1 compounds (non- Di 2 metal and non- metal) a. Name the first Tri 3 element in the compound Tetra 4 b. Name the Penta 5 second element Hexa 6 ending in ‘ide’ Hepta 7 c. Use prefixes to Octa 8 indicate the number of atoms of each Do not use Nona 9 the prefix –mono Deca 10 for the first element II. Naming Ionic Compounds: Forms One Cation a. Name of metal b. Name of anion/polyatomic ion (Table 1) III.Naming Ionic Compounds: Forms Two or More Cations a. Name of metal b. Roman numeral indicating charge of metal c. Name of anion/polyatomic ion (Table 1) IV.Naming Acids: Contains 2 Elements a. Name Hydro b. Name non-metal ending in ‘ic’ c. Last word is acid V. Naming Acids: Contains 3 Elements a. Do not name hydrogen b. Name polyatomic ion If it ends in ‘ate’ then it goes to ‘ic’ (“ateic”) If it ends in ‘ite’ then it goes to ‘ous’ (“iteous”) c. Last word is acid Significant Figures I. All numbers from 1-9 are Prefi Symb Valu significant. II. Zeros between numbers x ol e are significant. III.Zeros at the Giga G 109 beginning of a number are not 6 significant. Mega M 10 IV.Zeros at the end of a Kilo k 103 number are only significant if there is Deci d 10-1 decimal in the number. -2 V. Multiplication and Centi c 10-3 Division Rule: The Milli m 10 result must contain Micro μ 10-6 the same number of sig figs as the number Nano n 10-9 with the least amount of sig figs. VI. Addition/Subtraction Rule: The result must contain the same number of decimal places as the number with least amount of decimal places. Atoms I. Atomic Theory: a. Atoms are the smallest part of an element that retains elemental properties. b. All atoms of an element are identical. (wrong due to isotopes) Isotopes: An atom with the name number of protons but a differing amount of neutrons. c. Atoms cannot be created or destroyed. -10 II. The units used for an atom are Å (anstrom)= 10 m and pm (picometer)= 10 m.-12 III.Average Atomic Mass: Σ(Isotope mass*Isotope abundance) IV.Periodic Table Trends: Figure 1.0 V. Avogadro’s Number: One mole of any substance= 6.022×10 23atoms or particles Combustion and Neutralization Reactions I. Combustion: Molecule + O CO2+ H O 2 2 II. Neutralization Reactions: Acids a. Acid + Base H O2+ Salt b. Acid + Sulfides Salt + H S2 c. Acid + Carbonates CO + H2O + S2lt III.Neutralization Reactions: Oxides a. Metal Oxide + H O Metal Hydroxide 2 b. Non-Metal Oxide + H O 2Acid c. Metal Oxide + Acid H O 2 Salt d. Non-metal Oxide + Base H O + 2alt Limiting Reactant I. To find the limiting reactant: a. Balance the chemical equation given. b. Convert the given mass of each reactant into moles. c. Use stoichiometry for each reactant to find the mass of product it produces. d. The reactant that produces a lesser amount of product is the limiting reagent. e. The reactant that produces a larger amount of product is the excess reagent. II. Percent Yield: (Actual yield)/(Theoretical yield) * 100 Quantum Numbers I. Principle Quantum Valu Lett Number (n): Main energy level/shell. e er 0 s II. Angular Momentum Q.N. (ℓ): Sublevel in each level. III.Magnetic 1 p Quantum Number (m ): ℓ Orbitals in each 2 d sublevel. a. m ℓ 2ℓ+1 and 3 f a range of -ℓ to +ℓ IV.Spin Quantum Number (m ):sDirection the electron spins. a. m san be +1/2 or -1/2 V. Pauli’s Exclusion Principle: No more than two electrons can occupy an orbital and they have opposite spins. VI. Electrons fill the orbitals Orbit Electro in this order: 1s 2s 2p al ns 3s 3p 4s 3d 4p 5s 4d 5p s 2 a. Group number tells us how many valence electrons there are. p 6 b. Period number tells d 10 us how many valence shells there are. f 14 VII. When orbitals are half filled they are stable. VIII. Electronegativity: How badly an atom wants an electron. IX. Ionization Energy: The energy required to remove an electron from an atom X. Electron Affinity: Energy released when adding an electron to an atom. Lewis Symbol I. Drawing a Lewis Symbol: a. Find the total number of valence electrons. b. Draw a skeleton structure of the molecule. The least electronegative element goes in the center. c. Add electrons to form octets around each atom. d. If there are too many electrons then form double and triple bonds. Halogens can only form single bonds. II. Formal Charges: Charge on each individual atom. a. Formal Charge= (Valence e )-(Non-bonding e )-(½ Bonding e ) - b. We want Lewis structures in which the formal charges on atoms are closest to zero. c. If there are formal charges then we want the negative charges on the more electronegative atoms. III.Resonance Structures: Same compound but has a different chemical makeup. Bonds I. Valence Shell Electron Pair Repulsion Theory (VSEPR): Bonded Non-bonded Molecular Angle e- e- Geometry 2 0 Linear 180° 3 0 Trigonal Planar 120° 2 1 Bent <120° 4 0 Tetrahedral 109.5° 3 1 Trigonal Pyramidal <109.5° 2 2 Bent <109.5° 5 0 Trigonal 90°/120 Bipryamidal ° 4 1 See-Saw 90°/120 ° 3 2 T-Shaped <90° 2 3 Linear 180° 6 0 Octahedral 90° 5 1 Square Pyramid <90° 4 2 Square Planar 90° II. Hybridization: The overlapping of orbitals due to excited electrons. Electron Hybridizati Domain on 2 sp 3 sp2 4 sp3 3 5 sp3d2 6 sp d III. Polarity a. If the molecule is symmetrical it will not be polar. b. It will be polar if individual bonds differ by Bond Type Ionization more than 0.4. c. Lone pairs on the Energy central atom Non-Polar 0-0.4 usually mean it’s Covalent polar. IV.Multiple Bonds: Polar Covalent .5-1.9 a. Multiple bonds contain one σ bond Ionic 2-4 and 1-2 π bonds b. H-brid orbital’s form σ bonds only and may contain non-bonding e pairs c. Non-hybrid orbital’s form π bonds V. Bond Order: ½ (Bonding e - Antibonding e ) Bond Type Structure Conductivity Melting/Boiling Points Ionic Crystalline Low High Covalent Amorphous or Low Low Crystalline Network Covalent Amorphous or Low High Crystalline Metallic Crystalline High Low Gasses I. Ideal Gas Law: PV=nRT a. n is the moles of gas and R is 0.082. II. Gas Density: d=MP/RT a. M is molar mass and R is 0.082. III.Law of Effusion: (1 /2 )/(√M2/√M1) a. n is the rate of effusion and M is molar mass. b. They are inversely proportional. Intermolecular Forces I. Ion-Dipole (Strongest) II. Hydrogen Bonding (2 ndStrongest) a. A specific type of dipole-dipole bonding that is stronger than normal dipole-dipole interactions. b. Hydrogen bonrd with Fluorine, Oxygen, or Nitrogen (FON). III.Dipole-Dipole (3 Strongest) a. Dipole refers to a polar molecule. IV.London Dispersion (Weakest) a. As the atomic weight increases so does the dispersion force. b. As the number of atoms increases so does the dispersion force. c. As the surface area increases so does the dispersion force. V. Vapor Pressure: The weaker the intermolecular forces the higher the vapor pressure. VI. Boiling Point: The stronger the intermolecular forces the higher the boiling point. VII. Freezing Point: The stronger the intermolecular forces the higher the freezing point. Properties of a Liquid I. Viscosity: A resistance to flow determined by intermolecular forces. a. The stronger the intermolecular forces the greater the viscosity. II. Surface Tension: The force required to expand the surface area of a liquid. III.Capillary Action: The rise of liquids up a very narrow tube due to cohesion and adhesion. a. Cohesion: When similar substances stick together. (Stronger) b. Adhesion: When different substances stick together. (Weaker) Phase Change I. Phase Change Diagram: Figure 2.0 II. Heating Curve: Figure 2.1 III.Vapor Pressure: The pressure of gas in a closed system. IV.Phase Diagram: Figure 2.2 a. Triple Point: All three phases are at equilibrium. b. Supercritical Fluid: Substance above the critical temperature and pressure. Not in a distinct gas or liquid form. c. A gas may be liquefied only at temps at or below critical temperature. Crystalline Solids I. Simple Solid: One atom inside of the structure. a. Edge Length (a): a= 2r II. Body Centered Solid: Two atoms inside of the structure. a. Edge Length: a= 4r/√3 III.Face Centered: Four atoms inside of the structure. a. Edge Length: a= √8*r Colligative Properties I. Boiling Point Elevation: ΔT=K m,Badd ΔT II. Freezing Point Depression: ΔT=K m, fubtract ΔT III.Vapor Pressure Lowering: P =XAP A A° ° a. XAis mole fraction of solvent, P iA the vapor pressure of pure solvent b. PA<P A° IV.The melting point of an impure compound is generally lower than that of a pure solid. V. Osmosis: The net movement of solvent is always toward the solution with the higher concentration of solute. VI. Osmotic Pressure: The pressure required to stop the osmosis of pure solvent into a solution. a. Osmotic Pressure Equation: II=MRT M is Molarity, R is 0.0821 VII. Colloids: Particles that are intermediate between heterogeneous mixtures and solutions. a. Tyndall Effect: Colloids ability to scatter visible light. VIII. Van hoff Factor (i): How many ions in a compound. Solutions I. Solution: A homogenous mixture that can be broken into two parts- the solute (less) and the solvent (more). II. Entropy: The measure of disorder in a system. III.Solvation: When solute molecules are completely surrounded by solvent molecules. a. Hydration: When the solvent molecules are water molecules. IV.Energy in a Solution: ΔH Solution +1H +ΔH2 3 a. ΔH :1solute-solute bonds (endo) b. ΔH :2solvent-solvent bonds (endo) c. ΔH :3solute-solvent bonds (exo) V. Saturated: Contains the max amount of solute for a given solvent. VI. Unsaturated: Does not contain the max amount of solute. VII. Super Saturated: Contains more than the max amount of solute. a. Crystallize: A solid that forms out of a solution. VIII. Concentrations of Solutions: a. Molarity (M): (Moles of Solute)/(Liters of Solution) b. Molality (m): (Moles of Solute)/(kg of Solvent) c. Mole Fraction: (Moles of Solute)/(Moles of Solution) d. Mass Percent: (Mass of Solute)/(Mass of Solution) * 100 IX. Dilution: M Concentrationlume ConcentrationDilutionlume Dilution Reaction Rates I. Rate of a reaction: Δ[A]/Δt a. Δ[A] is the change in concentration. II. Specific rate of a reaction: (-1/a)*(Δ[A]/Δt) a. a is the coefficient of the molecule. III.Rate Law: rate=k[A] [B]m n a. K is a rate constant, m and n are the orders of the molecules. IV.Integrated Rate Law First Order: ℓn [A] = -ktt+ ℓn[A] 0 a. Half Life: ℓn2= kt 1/2 V. Integrated Rate Law Second Order: (1/[A] )= kt + t1/[A] ) 0 a. Half Life: t 1/21/k[A] ) 0 VI. Integrated Rate Law Zero Order: [A] = -ktt+ [A] 0 a. Half Life: t 1/2[A] /20) VII. Integrated Rate Law Order Graphs: Figure 3.0 VIII. Factors that affect the rate of a reaction: a. Orientation: The molecules need to be facing the right way to interact with each other. b. Activation Energy (E ): Theaenergy required to initiate a reaction. c. Collide: The particles have to collide for a reaction to occur. More collisions occur at higher temperatures. IX. Energy Graph: Figure 3.1 a. The rate-determining step is the one that requires the most E . a b. Intermediate: A molecule that is produced and consumed by a multistep reaction. Not a part of the reactants or products in the final equation. c. Catalyst: Speeds up the reaction by providing another pathway for the reaction to occur. X. Fraction of Molecules That Will React Equation: f= e (-Ea/RT) (-Ea/RT) XI. Arrhenius Equation: k=Ae a. A is the frequency factor. XII. ℓn(k1/k2)= (E aR)(1/T 21/T 1 Equilibrium I. No solids or liquids are included in equilibrium expressions. m n II. Keq (Products) / (Reactants) a. K>1 there are more products than reactants (shift left). b. K<1 there are more reactants than products (shift right). c. K=1 there is an equal amount of products and reactants. III.Kp=K cRT) Δn a. K ps the Pressure Constant. b. K cs the Concentration Constant. c. R is 0.0821. d. Δn is the difference in the moles of gas (products-reactants). IV.Finding K for multistep reactions or manipulated reactions: a. K = 1/K Reversen Forward b. K c(K )cwhere n is the coefficient change. c. K eqK *K1 2 V. Use ICE charts to solve equilibrium expressions (Initial, Change, Equilibrium). VI. Q=(Products) / (Reactants) not at equilibrium. a. Q>K there are more products than reactants (shift left). b. Q<K there are more reactants than products (shift right). c. Q=K the reaction has reached equilibrium. VII. Le Chatiliers Principle: A system at equilibrium when stressed will attempt to regain equilibrium. a. When the concentration of reactants increases then more products will be made (shift right) and vice versa. b. Exothermic reactions have heat as a product therefore increased temperature will shift the reaction towards the reactants (left). Shifts right for a decrease in temperature. c. Endothermic reactions have heat as a reactant therefore increased temperature will shift the reaction towards the products (right). Shifts left for a decrease in temperature. d. An increase in partial pressure will shift the equilibrium to the side with the fewest moles of gas. An increase in overall pressure will have no effect on equilibrium. e. A decrease in volume will shift the reaction to the side with fewer moles of gas. f. The addition of an inert gas will have no effect on equilibrium. g. The addition of a catalyst will have no effect on equilibrium. Acids and Bases Strong Strong Bases Acids HCl Group 1 I. Arrhenius Definition: Hydroxides HBr Sr(OH) 2 HI Ca(OH) 2 HNO 3 Ba(OH) 2 H2SO 4 HClO 3 a. Acid: Increases the concentration of H ions in a solution. + Contains H . b. Base: Increases the concentration of OH ions in a solution. Contains OH . - II. Bronstead-Lowry Definition: a. Acid: Proton donor. b. Base: Proton acceptor. III.Lewis Definition: a. Acid: Electron pair acceptor. b. Base: Electron pair donor. IV.A strong acid/base will completely dissociate in a solution. V. Acid-Base Conjugate Pairs: The acid has more protons than its conjugate base, the base has less protons than its conjugate acid. VI. Auto ionization of water: K =[H O ][OH ] K =1×10 -14 + -w 3 w VII. pH=-ℓog[H O 3 pOH=-ℓog[OH ] a. pH + pOH=14 VIII. Dissociation constants: K (acia constant) K (base cbnstant) a. K a[H O3][A ]/[HA] - [HA] is the acid, [A ] is its conjugate base Ka>1 strong acid, K <1aweak acid - + b. K b[OH ][BH ]/[B] [B] is the base, [BH ] is the conjugate acid c. K wK *a b d. Simplification rule: If the difference between the initial concentration and K /K as 1b00 or greater assume x is negligible in the ICE chart. e. Percent Ionization: ([H ]/[HA] )*100 IX. Acid-Base Salts: How will a compound if put into a solution affect its pH? a. If the ion is a strong acid/base ion then it will have no affect on the pH. b. If it is a cation of a weak base then it will cause the solution to be acidic. c. If it is an anion of a weak acid then it will cause the solution to be basic. X. Periodic Trends in acidity: a. For binary acids, acidity increases down and to the right. b. For oxy acids the more oxygen’s there are the more acidic the molecule. If the number of oxygens are equal then the more electronegative molecule is more acidic. c. For organic acids the more resonance structures that can be made, the more acidic the molecule. XI. Acid/Base Titrations: To reach the equivalence point (neutralization) the moles of the base must be equal to the moles of the acid. a. The titration of a polyprotic acid contains more than one ionizable H atom and the reaction with OH in a series of steps. Buffered Solutions I. Common Ion Effect: Whenever a weak electrolyte and a strong electrolyte containing a common ion are together in a solution, the weak electrolyte ionizes less than if it were alone in the solution. II. Buffered Solutions: Solutions that contain a weak conjugate acid- base pair that resist change in pH. III.Buffer Capacity: A buffers ability to resist pH. a. The greater the concentration the better able the buffer is to resist change in pH. IV.Henderson-Hasselbalch Equation: pH=pK + ℓog[base]/aacid] a. Optimal pH of a buffer is when pH=pK a V. Buffers Effective pH range: pK + 1 a VI. pH Titration Curve: Figure 4.0 Solubility I. Solubility Product Constant (K ): IndiSptes how soluble the solid is in water. n m a. K SpCation] [Anion] II. Solubility is the quality of a substance that dissolves to form a saturated solution. III.Factors that affect solubility: a. Common-ion effect b. pH: The solubility of a compound containing a basic anion increases as the solution gets more acidic. c. Complex ion: A complex ion is the assembly of a metal ion and the Lewis bases bonded to it. The solubility of metal salts increases in the presence of suitable Lewis bases provided the metal forms a complex ion with the base. Formation constant (K f: [Complex ion]/[Cation][Base] d. Amphoterism: Metal oxides and hydroxides that are soluble in strong acids/bases because they themselves behave as an acid or a base. e. Solute-Solvent interactions: like dissolves like f. Temperature: For solids and liquids an increase in temperature means increased solubility. For gasses an increase in temperature means decreased solubility. g. Pressure: As pressure increases so does gasses solubility. IV.Q=[Cation] [Anion] m a. Q>K Spprecipitation occurs, reducing ion concentrations until Q=K Sp b. Q=K Spis a saturated solution. c. Q<K Spsolid dissolves increasing ion concentrations until Q=K Sp V. Selective Precipitation: The separation of ions in an aqueous solution by using a reagent that forms a precipitate with one or more of the ions. VI. Qualitative Analysis: Determines only the presence or absence of a particular metal ion. VII. Quantitative Analysis: Determines how much a given substance is present. Thermodynamics I. First Law of Variab Situation Sig le n q Heat is gained by the + system Heat is lost by the - system w Work is done on the + system Work is done by the - system Thermodynamics: Energy is conserved. II. Energy-Work Equation: ΔE=q+w a. ΔE is the change in energy. b. q is heat. c. w is work. III.Exothermic: Heat is lost by the system to the surroundings. (-ΔH) IV.Endothermic: Heat is gained by the system from the surroundings. (+ΔH) V. Enthalpy (H): The amount of heat content used or released in a system at a constant pressure. a. Enthalpy is a state function (value doesn't depend on path taken). b. Enthalpy Formula: H=E+PV. E is internal energy, P is pressure, V is volume. c. The change in enthalpy is equal to the heat. ΔH=q d. ΔH=H ProductsReactants e. Hess’s Law: If a reaction involves several steps, ΔH for the overall reaction is equal to the sum of the enthalpy changes for each step. ° VI. Standard Enthalpy of Formation (ΔH ): The fhange in enthalpy associated with the formation of one mole of a compound from its elements in their standard states. a. ΔH rxnΔH (froducts)-ΔH (Rfactants) VII. Heat Capacity: The amount of heat required to increase the temperature of a substance by 1°C or 1K. VIII. Specific Heat (C ):SThe amount of heat required to raise the temperature of one gram of a substance by 1K or 1°C. a. CS=q/(m*ΔT) C S J/(g*°C) IX. Spontaneous Processes are processes that proceed without any intervention. X. Irreversible Processes require work to put the system and surroundings back to their original state. a. Spontaneous processes are irreversible. XI. Reversible Processes: The system changes in such a way that the system and surroundings can be put back in their original state by exactly reversing the process. Entropy I. Entropy (S) is the measure of disorder in a system. a. S=q/T S=J/mol*K b. Entropy is a state function. ΔS=S -S FinalInitial c. As a result of all spontaneous processes the entropy of the universe increases. II. Second Law of Thermodynamics: Any irreversible process increases total entropy, whereas reversible processes result in no overall change in entropy. a. Reversible processes: ΔS UnivS SysΔS Surr0 b. Irreversible processes: ΔS UnivS SysΔS Surr0 III.Microstate: A particular microscopic arrangement of the atoms that corresponds with the physical stat of the system. a. Vibrate: Gas, Liquids, Solids b. Rotation: Gas, Liquids c. Translation: Gas d. Entropy increases with the number of microstates of the system. e. Entropy is a measure of how many microstates are associated with a particular macroscopic state. IV.Decrease in entropy is –ΔS, increase in entropy is ΔS. a. Entropy increases based on physicals states. Gas>Liquid>Solid b. If the physical states are the same then entropy increases as the number of atoms increases. c. If the physical state and atoms are the same then entropy increases as mass increases. d. Entropy increases with increasing temperature because the increased motion leads to a greater number of microstates. e. At melting and boiling points there is a large increase in entropy. V. Third Law of Thermodynamics: The entropy of a pure crystalline substance at absolute zero is zero. VI. Entropy Graph: Figure 5.0 VII. Standard Molar Entropies S° are molar entropies for substances in their standard state, 1 atm 298K. a. ΔS°=ΣnS°(Products)-ΣmS°(Reactants) m and n are the coefficients b. ΔSSurrΔH Sys ΔH is in kJ and needs to be converted to J Gibbs Free Energy I. Gibbs Free Energy Equation: ΔG=ΔH-TΔS, ΔG is in kJ a. –ΔG is a spontaneous reaction, ΔG is non-spontaneous, ΔG=0 is at equilibrium. ΔG ΔH ΔS Reaction Sign Sign Sign - - + Spontaneous + + - Non-spontaneous + or - - - Spontaneous at low temps + or - + + Spontaneous at high temps II. When ΔH and ΔS have the same sign then T=ΔH/ΔS. a. If ΔH and ΔS are positive then the reaction is spontaneous when T>ΔH/ΔS. b. If ΔH and ΔS are negative then the reaction is spontaneous when T<ΔH/ΔS. III.Free Energy for molecules not at standard states: ΔG=ΔG°-RTℓnQ a. R is 8.314, Q=[Products] /[Reactants] , ΔG° free energy at standard states. IV.At equilibrium ΔG=0 and Q=K: ΔG°=-RTℓnK a. K>1 Spontaneous, K<1 Non-spontaneous Redox Reactions I. Oil Rig: Oxidation is loss of electrons; Reduction is gain of electrons. II. Oxidation Number Rules: a. For atoms in elemental form the oxidation number is zero. b. Hydrogen is +1 in molecular compounds and -1 in ionic compounds. c. Oxygen is -2 except for peroxides where it is -1. d. Fluorine is always -1. e. Cl, Ba, I are -1 in binary compounds and have positive charges in other compounds. Balancing Electrochemical Equations I. Break up into half reactions. II. Balance the non O and H elements. III.Balance O with H2O IV.Balance H with H + V. Balance the electrical charge by adding electrons (each half reaction needs to have the same number of electrons) VI. Add equations together crossing out like terms Voltaic Cells I. Voltaic Cells: A device in which the transfer of electrons takes place through an external pathway. It is a spontaneous reaction. - II. Anode: The electrode at which oxidation occurs. e is a product. III.Cathode: The electrode at which reduction occurs. e is a reactant. IV.The salt bridge is used to provide the solutions with cations and anions to keep the solutions charge neutral. V. Voltaic Cell Diagram: Figure 6.0 VI. Oxidizing Agent: element that reduces; F is the best agent. VII. Reducing Agent: element that oxidizes; Li is the best agent. VIII. Electromotive force or Cell Potential: E° =E° cell cathodeE°anode a. The greater the E° cellhe better it is at reducing. Postive E° cells spontaneous so the E° cellust always be postive. Free Energy and Voltaic Cells I. ΔG°=-nFE° cell a. n is the number of electrons transferred, F is Faradays constant 96.485 kJ/V*mol II. Cell potential and Equilibrium Constant Equation: E° =(RT/nF) cell a. R and F should be in kJ. III.Nernst Equation: E cell° -cell592V/n) ℓogQ a. This is for cells at non-standard states. b. Concentration Cell: Transfer of electrons between the same elements so E° =0 bcell cellives us a voltage due to differing concentrations of ions. Batteries I. All batteries are voltaic cells therefore all batteries are spontaneous reactions. II. Primary cells are non-rechargable. a. Alkaline batteries, 1.5 V III.Secondary cells are rechargeable. a. Lead Acid, Lithium-Ion 2V each cell, 6 cells stacked, 12V total Electrolytic Cell I. Electrolytic Cell: A cell with a non-spontaneous reaction that must be powered by a battery. The battery has to have a voltage greater than that of the E° .cell II. Electrolytic Cell Diagram: Figure 7.0 III.Amperes (A): A measure of current. (colombs/sec) a. Faradays Constant can also be written: 96,485 colombs/mol of e - Nuclear Chemistry I. Element Notation: Figure 8.0 II. Stable nucli have a 1:1 nuetron-proton ratio. Atoms with 84 protons or greater are not stable and are radioactive (only undergo alpha emission). III.Radioactive Processes: Name Symb Emissi ol on 4 Alpha α 2He Beta β -1e Gamma γ 0γ 0 Positron - 1e Electron - -1e Capture
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