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AP Chemistry Notes Stephen Akiki Colchester High School Download at http://akiscode.com/apchem ♥♠♣♦ Special Thanks to Stephen Bosley (Boser) Contents 1 FOREWORD/DISCLAIMER 4 2 Solubility Rules 5 2.1 Soluble . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5 2.2 Insoluble . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5 2.3 Naming Rules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 5 3 Periodic Table of Elements 5 4 Poly Atomic Naming 6 5 Common Units, Constants and Charges 6 5.1 Fundamental Constants . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 6 5.2 Charge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 5.3 Radius . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 6 Atomic Theory 7 6.1 J.J. Thompson . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 6.2 Robert Millikan . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 6.3 Ernest Rutherford . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 6.4 Chadwick . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 6.5 John Dalton . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 7 7 Naming 8 7.1 Binary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8 7.1.1 Greek Prefixes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8 7.2 Ionic . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8 7.3 Acids . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8 7.3.1 Polyatomic . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8 7.3.2 Binary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 8 8 Cations 9 9 Reaction Type 9 9.1 Combination (Synthesis) . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9 9.2 Decomposition . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 9 9.2.1 Special Binary Salt Splits . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10 9.3 Combustion . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10 10 Blackbody Radiation 10 11 Bohr Model 11 11.1 Energy Level Formula . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11 11.1.1 Energy Change during Level Jumps . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11 1 12 Wavelength 11 12.1 De Broglie Formulas . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11 13 Quantum Values 12 13.1 Quantum Value Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12 13.2 Special cases . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 12 14 Periodicity 13 14.1 Electron Configuration . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13 14.2 Isoelectricity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13 15 Nuclear Chemistry 13 15.1 Isotopes . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 13 15.2 Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14 15.2.1 Alpha Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14 15.2.2 Beta Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14 15.2.3 Gamma Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14 15.2.4 Positron Radiation . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15 15.2.5 Electron Capture . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15 15.3 Nuclear Equations . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15 15.3.1 Radiation Table . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 15 15.4 Nuclear Stability . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16 15.4.1 Forces Invloved . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16 15.4.2 Belt of Stability . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16 15.4.3 Magic Numbers . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 16 15.4.4 Half-Life . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17 16 Ionization and Affinity 17 16.1 Ionization Energy . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17 16.2 Electron Afinity . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17 17 Reactions of Metals 17 18 Chemical Bonds 17 18.1 Intramolecular . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 17 18.1.1 Ionic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.1.2 Covalent Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.1.3 Metallic Bonding . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.2 Intermolecular . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.2.1 Ion-Dipole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.2.2 Dipole-Dipole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.2.3 Hydrogen Bond . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.2.4 London Dispersion/Van der Waals . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 18 18.2.5 Intermolecular Flowchart . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 18.3 Rule of Octet . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 19 Lewis Structures 19 19.1 Structures for Atoms . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 19.2 Structures for Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 19.3 Structure for Ions of Molecules . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 19 19.4 Lewis Structures for Molecular Structures (Covalent) . . . . . . . . . . . . . . . . . . . . . . . 20 19.5 Resonance Structures . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21 19.5.1 Formal Charge . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 21 20 Lattice Energies of Ionic Solids 21 21 Bond Lengths of Covalent Bonds 22 22 Electronegativity 22 22.1 Dipole . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 22 22.1.1 Dipole Moment . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 22 2 2 Solubility Rules 2.1 Soluble • Nitrates NO−13 - All nitrates are soluble • Chlorates ClO−13 - All chlorates are soluble • Alkali metal Cations and Ammonium cation compounds NH+14 are all soluble • Chlorides, Bromides, and Iodides are all soluble EXCEPT Ag+1, Pb+2, and Hg+2 • Acetates - All are soluble except Ag+ • Sulfates - All are soluble except Ba+2, Pb+2, Hg+2, Ca+2, Ag+1, and Sr+2 2.2 Insoluble • Carbonates CO−23 - all carbonates are insoluble except alkali metals and ammonium compounds • Chromates CrO−24 - all chromates are insoluble except alkali metals, ammonium, Ca+2, and Sr+2 • Hydroxides OH−1 - all hydroxides are insoluble except alkali metals, ammonium, Ba+2, Sr+2, and Ca+2 although the last two (Sr+2 and Ca+2) are only slightly soluble so a precipitate can form. • Phosphates PO−34 all are insoluble except alkali metals and ammonium • Sulfites SO−23 all are insoluble except alkali metals and ammonium • Sulfides S−2 all are insoluble except Alkali metals, alkali earth metals and ammonium 2.3 Naming Rules • All strong acids and bases are soluble and should be written as the ions when completing net ionic reactions . Sulfuric acid (H2SO4) should be written as H+ +HSO−14 • The strong acids are: HCL, HBR, HI, HNO3, HClO4, and H2SO4 • Strong bases are any alkali metal hydroxides (LiOH, NaOH, etc) and Ca(OH)2, Sr(OH)2, Ba(OH)2 • All acids and bases should be left in their molecular form: . Acetic acid → HC2H3O2 3 Periodic Table of Elements 5 4 Poly Atomic Naming • Zinc Zn+2 • Silver Ag+1 • Ammonium NH+14 • Hydroxide OH−1 • Cyanide CN−1 • Nitrate NO−13 • Acetate C2H3O−12 • Chlorate ClO−13 • Bromate BrO−13 • Iodate IO−13 • Manganate MnO−13 • Sulfate SO−24 • Bisulfate (Hydrogen Sulfate) HSO−14 • Carbonate CO−23 • Bicarbonate (Hydrogen Carbonate) HCO−13 • Selenate SeO−24 • Biselenate (Hydrogen Selenate)HSeO−14 • Oxalate C2O−24 • Phosphate PO−34 • Hydrogen Phosphate HPO−24 • Dihydrogen Phosphate H2PO−14 • Chromate CrO−24 Per Ate Ate Ite Hypo Ite Per Ic Ic Ous Hypo Ous +1 Oxygen Most Common Ion -1 Oxygen -2 Oxygen 5 Common Units, Constants and Charges 5.1 Fundamental Constants • Avogadros Number (N) . 6.02214199 ∗ 1023mol−1 • Plancks Constant (h) . 6.62606876 ∗ 10−34J ∗ s • Speed of Light (c) . 2.99792458 ∗ 108m/s 6 5.2 Charge • e− charge = −1.602 ∗ 10−19 coulombs • p+ charge = 1.602 ∗ 10−19 coulombs • Atomic Mass Unit (amu) = 1.66054 ∗ 10−24 . p+ = 1.0073 amu . n◦ = 1.0087 amu . e− = 5.486 ∗ 10−4 amu 5.3 Radius Angstroms ( ◦ A) = 10−10 meters 6 Atomic Theory 6.1 J.J. Thompson • Discovered e− and chargemass ratio . Charge to Mass ratio: 1.76 ∗ 108 Coulombs/Gram (Charge of e−/mass) • Plum Pudding Model of atom 6.2 Robert Millikan • Found charge and mass of e− • Millikan Oil Drop: . Charge oil drops in a field and adjust field until drops levitate 6.3 Ernest Rutherford • Discovered 3 types of radiation (Decay Particles) . Alpha particles: He2+ size, very damaging, stoppable - α . Beta particles - e− size, damaging, hard to stop - β . Gamma particles - tiny, not so damaging, unstoppable - γ • Also discovered proton and new dense nucleus model . Rutherford worked with α particles most and discredited Thompsons model of the nucleus 6.4 Chadwick • Discovers neutron by shooting radiation at light elements and it watching it kick out a neutral particle 6.5 John Dalton • Four Postulates . Everything made of atoms . Atoms of one element differ from those of a different element . Atoms will combine in whole number ratios . Atoms can not be created or destroyed • Law of Constant Composition . In a compound, atom ratios are constant 7 • Metal Nonmetal ∆→ Metal + Nonmetal (diatomic in nature) . 2NaCl ∆→ 2Na+ Cl2 • Metal Chlorates ∆→ Metal Chlorides + O2 . Fe(ClO3)2 ∆→ FeCl3 +O2 9.2.1 Special Binary Salt Splits These binary salts split into different elements (NH4)2CO3 → NH3 +H2O + CO2 H2SO3 → H2O + SO2 H2CO3 → H2O + CO2 NH4OH → NH3 +H2O H2O2 → H2O +O2 9.3 Combustion Hydrocarbon+O2 → CO2 +H2O ....⇓.... CxHy → double x (multiply by 2) then add 2 • C1: meth • C2: eth • C3: pro • C4: bu • C5: pent • C6: hex • C7: hept • C8: oct • C9: non • C10: dec 10 Blackbody Radiation When an object is heated it will emmit radiant energy E = hν • E = Energy • h = Max Plancks constant (6.626 ∗ 10−34J ∗ s) • ν = frequency Photoelectric effect: Metal will give off e−s if light shines on it. Light shining on a clean sheet of metals will release e−s if ν is strong enough. 10 11 Bohr Model Neils Bohr: 1. Only orbits of certain radii, corresponding to certain definate energies are permitted for the electron in a hydrogen atom. 2. An electron in a permitted orbit has a specific energy and is in an allowed energy state. An electron in an allowed state will not radiate energy and therefore will not spiral into the nucleus. 3. Energy is emmitted or absorbed by the e− only as the e− changes from one allowed energy state to another. 4. Flawed theory because it only works for hydrogen 11.1 Energy Level Formula En = (−2.18 ∗ 10−18J)( 1n2 ) • E1: −2.18 ∗ 10−18J • E2: −5.45 ∗ 10−19J • E3: −2.42 ∗ 10−19J • E4: −1.36 ∗ 10−19J • E5: −8.72 ∗ 10−20J • E6: −6.056 ∗ 10−20J • E∞: 0 11.1.1 Energy Change during Level Jumps ∆E = EF − E0 • n = 3→ 2 | −3.03 ∗ 10−19J • n = 4→ 2 | −4.09 ∗ 10−19J • n = 5→ 2 | −4.578 ∗ 10−19J • n = 6→ 2 | −4.844 ∗ 10−19J 12 Wavelength 12.1 De Broglie Formulas λ = hmv or λ = hp • λ = Wavelength • h = Plancks Constant (6.626 ∗ 10−34J ∗ s) • m = Mass of particle in Kg • v = Velocity of particle (meterssecond ) • p = Momentum Example m = 9.11 ∗ 10−28g v = 5.97 ∗ 106m/s λ = 6.626∗10 −34J∗s (9.11∗10−31Kg)(5.97∗106m/s) = 1.22 ∗ 10 −10m 11 13 Quantum Values 1. Principle Quantum number - (n) n = 1 (lowest) n = ∞ (at 8 or 9) Follows Bohrs En = (−2.18 ∗ 10−18J)( 1n2 ) 2. Azimuthal Quantum number - (l) l = n - 1 if... • l = 0 → S shape • l = 1 → P shape • l = 2 → D shape • l = 3 → F shape Example n = 3 l = 2 ⇓ 3d 3. Magnetic Quantum number (orbital) - (ml) -l and l including zero m0 = 0 m1 = −1, 0, 1 m2 = −2,−1, 0, 1, 2 4. Spin magnetic quantum number - (ms) + 12 or - 1 2 13.1 Quantum Value Table n Possible l values Subshell ml values # of orbitals in subshell total # of orbitals in shell e− in shell 1 0 1s 0 1 1 2 2 0 2s 0 1 4 8 1 2p -1,0,1 3 - - 3 0 3s 0 1 9 18 1 3p -1,0,1 3 - - 2 3d -2,-1,0,1,2 5 - - 4 0 4s 0 1 16 32 1 4p -1,0,1 3 - - 2 4d -2,-1,0,1,2 5 - - 3 4f -3,-2,-1,0,1,2,3 7 - - 13.2 Special cases • Chromium has 6 half-filled orbitals • Copper has one half-filled orbital and 5 filled orbitals 12 15.2.4 Positron Radiation When a positively charged nucleus emits its p+ leaving only the n◦. 15.2.5 Electron Capture When an electron in orbit falls into the nucleus (positively charged) and makes it neutrally charged. 15.3 Nuclear Equations 15.3.1 Radiation Table Neutron: 10n Proton: 11p + Electron: 0−1e − Positron: 01e − Alpha Particle: 42He or 4 2α Beta Particle: 0−1e − or 0−1β Example Alpha 238 92 U →23490 Th+42 He Beta 131 53 I →13154 Xe+ 0−1e− 1 0n→11 p+ 0−1 e− Positron 11 6 C →115 B+01e− 1 1p→10 n+01e− Electron Capture 81 37Rb+ 0 −1e − →8136 Kr 1 1p+ 0 −1e − →10 n Positron-Electron Collision (Gamma) 0 1e+ 0 −1e − →00 γ 15 15.4 Nuclear Stability Understanding why are some nuclides are radioactive while others are not. 15.4.1 Forces Invloved • Electrostatic . Try to rip apart the nucleus because of like charges • Strong Nuclear . Try to pull together the nucleus because subatomic particles naturally stick together • The Glue . Neutrons act as the glue and more of it is required when the electrostatic force gets really strong 15.4.2 Belt of Stability • Area A . More neutrons than protons - Beta decay → creates protons • Area B . Less neutrons than protons - Positron emission (Smaller B) or Electron Capture (Larger B) • Area C . Every element above 83 p+ is radioactive and no glue can hold it together - Alpha decay 15.4.3 Magic Numbers The Magic Numbers tend to be stable if you have either a proton or neutron in those numbers. If you have both, they are very stable. (p+) 2 8 20 28 50 82 - (n◦) 2 8 20 28 50 82 126 • If (p+) and (n◦) even → likely stable • If either is odd → could go either way • If (p+) and (n◦) odd → likely unstable 16 15.4.4 Half-Life The time it takes 12 the amount of a substance to decay. Example 5g of nuclide 1 2 life of 15 years How much of the original nuclide remains after 45 years? 5 ⇓ (15 years) 2.5 ⇓ (30 years) 1.25 ⇓ (45 years) 0.625g 16 Ionization and Affinity 16.1 Ionization Energy The energy needed to remove an e− (how easy it is to lose an e−). Needs energy (+). 16.2 Electron Afinity How much a gaseous atom will be attracted to a free e− (how easy it is to gain an e−). Releases energy (-). 17 Reactions of Metals Metal Oxides = Basic • Metal + Water → Metal Hydroxide + H2 • Metal + O2(Li or any non-Alkali metal) → Metal Oxide • K + O2(Any other Alkali metal) → Metal Peroxide (O−12 ) . K + O2 → KO2 • Metal Oxide + H2O → Metal Hydroxide . Na2O +H2O → NaOH • Metal Oxide + Acid → Salt + H2O . Na2O +HCL→ NaCl +H2O Nonmetal Oxides = Acidic • Nonmetal Oxide + H2O → Acid . CO2 +H2O → H2CO3 . SO2 +H2O → H2SO3 . P4O10 +H2O → H3PO4 • Nonmetal Oxide + Base → Salt + H2O . CO2 +NaOH → Na2CO3 +H2O 18 Chemical Bonds When 2 or more atoms are strongly attached (attracted) to each other. 18.1 Intramolecular These forces act inside an atom or molecule: 17 19.4 Lewis Structures for Molecular Structures (Covalent) 1. Add valence e−s from all the atoms. 2. Write the symbols for the atoms. If there are more than 2 atoms, identify the central atom. Connect them with a single line which represents 2 shared e−s. Subtract the number of e−s from total found in step 1. . Central atom will be closest to Si, P or Metaloid staircase. 3. Complete octets around the atoms bonded to the central atom (Hydrogen does not get more than 2). 4. Place the remaining pairs around the central atom even if doing so gives more than an octet to the central atom. 5. If there are not enough pairs to complete an octet in the central atom, then you ned to try using double or triple bonds. CH4 CH2Cl2 HNO3 CO2 HCN 20 19.5 Resonance Structures Benzene 19.5.1 Formal Charge Valence e−s of an atom - (total unbonded e−s + 12 total bonded e −s) Molecular structures that tend to be the common one have a formal charge is closest to zero and any negative charge is on the most electronegative element. 20 Lattice Energies of Ionic Solids Coulombs Law E = KQ1Q2 d • Q1/Q2 = ion charges • d = Distance between ions of the final crystalized lattice form. a The greater the charge, the higher the energy. a The closer the ions, the higher the energy. Example Which has a greater lattice energy? +1 Na −1 Cl vs +2 Mg −1 Cl2 +2 Mg −1 Cl2 has greater charges thus a higher lattice energy. +1 Li −1 Cl vs +1 Na −1 Cl +1 Li is smaller than +1 Na so +1 Li wil be closer to −1 Cl than +1 Na will so +1 Li −1 Cl will have a higher lattice energy. 21 21 Bond Lengths of Covalent Bonds • Single - Longest • Double - Medium • Triple - Shortest Length Single CO−44 1.42 ◦ A Double CO2 1.24 ◦ A Triple CO 1.13 ◦ A 22 Electronegativity Difference in electronegativity determines the character of the bond. • Large difference → Ionic Bond . Biggest difference is 3.3 • Medium difference → Polar Covalent . HF - 1.8 • Small/No difference → Non-Polar Covalent . H2 - 0 22.1 Dipole → H − F Arrow points towards more electronegative atom. 22.1.1 Dipole Moment Numeric value that represents how strong the dipole is Example Which has the greater dipole moment? OR Which has greater electronegative difference? HI or HF Answer: HF 22 Branch Structure Naming Table Clg }-ethylpentane CrHig isoheptane 2,2—dlimethylpentane CH, CH, C(CH,),CH,CH, 3.3~<climethylpenrane (CH,),CHC(CH,), 2.2,3-trimethy bute CH,(CH,);CH, heptane CH, CHICH,)CH(CH,), limethylpentane (CH), CHCH, CHICH,), 24-dimetbploentare SICH JCC, 2-racthyhesaac 3 methythemae 25.4 Alkenes a Spotted by seeing a double bond • CH2 → Methene • C2H4 → Ethene • C3H6 → Propene • C4H8 → Butene • C5H10 → Pentene • C6H12 → Hexene • C7H14 → Heptene • C8H16 → Octene • C9H18 → Nonene • C10H20 → Decene 25.4.1 Alkene Naming Naming Alkenes is similar to naming Alkanes save for the naming of the root chain. To name the root chain you must give side where the double bond is the lowest number and name all branches after using this number scheme. You should end up with something like 2 Pentene 2 Pentene 25.5 Alkynes a Spotted by seeing a triple bond • CH → Methyne • C2H2 → Ethyne • C3H4 → Propyne • C4H6 → Butyne • C5H8 → Pentyne • C6H10 → Hexyne • C7H12 → Heptyne • C8H14 → Octyne • C9H16 → Nonyne • C10H18 → Decyne 26 25.5.1 Alkyne Naming Naming Alkynes is similar to naming Alkenes. Identify the root chain as you would using Alkenes except now you identify the triple bond instead of the double bond. 2 Hexyne 26 Functional Groups  When discussing functional groups, the letter R is used to signify any hydrocarbon or hydrocarbon chain. 26.1 Alcohol • Root Name: -ol • Identification: R-OH Ethanol 26.2 Aldehyde • Root Name: -al • Identification: R-CHO Ethanal 27 27.2 Anions [Al(OH)4]−1 a The charge of a anion is determined by the individual charges of the elements. . Al+3 + 4(OH)−1 . 3 - 4 . -1 27.3 Coordination Number Generally (especially with cations) the coordination number is twice the charge of the transition metal. Example [Cr(H2O)6]+3 Cr+3 → 3 ∗ 2 = 6 27.4 Naming 27.4.1 Cations • Give the prefix associated with the coordination number • Give appropriate name for ligand • Name the transition metal • Give roman numeral of transition metal Example [ Chromium︷︸︸︷ Cr (H2O)︸ ︷︷ ︸ Aqua Hexa︷︸︸︷ 6 ]+3 Hexa Aqua Chromium (III) 27.4.2 Anions • Give prefix associated with the coordination number • Give appropriate Ligand name • Name transition metal with -ate ending • Give roman numeral Example [Al(OH)4]−1 Tetra Hydroxo Aluminate No roman numeral because Al is always +3 28 Acidic and Basic Redox 28.1 Acidic • Find oxidation number • Write 12 reaction with e −s • Add H2O, then H+ and balance accordingly • Balance for e−s and everything else • Add together both balanced 12 reactions and cancel out where possibly to simplify 30 28.2 Basic • Find oxidation number • Write 12 reaction with e −s • Add H2O, then H+ and balance accordingly • Add OH amounts to both sides equal to the number of H+ • Cancel out the H+ with the OH to form H2O • Move all H2O to one side • Balance for e−s and everything else • Add together both balanced 12 reactions and cancel out where possibly to simplify Example +7 Mn O−4 + +3 C2 −2 O4 −2 → +1 Mn O−22 + +4 C O −2 3 (4OH + C2O4 → 2CO3 + 2e− + 2H2O) ∗ 3 (2H2O +MnO4 + 3e− →MnO2 + 4OH) ∗ 2 12OH + 3C2O4 → 6CO3 + 6e− + 6H2O 4H2O + 2MnO4 + 6e− → 2MnO2 + 8OH 4OH + 3C2O4 + 2MnO4 → 2MnO2 + 6CO−23 29 Thermodynamics The study of energy and its transformations Units of Energy: • Joules and Calories . 1 cal = 4.184 J The two main driving forces of thermodynamics is Enthalpy and Entropy: 29.1 Enthalpy Enthalpy stands for the Heat of the reaction and is denoted by ∆H. If: • ∆H<0 . Reaction is exothermic • ∆H>0 . Reaction is endothermic There are 4 ways to find ∆H. 29.1.1 Stoichiometry Problems Example How much heat is released when 3.2 grams of Hydrogen is reacted with excess Oxygen? 2H2 +O2 → 2H2O ∆H◦ = −572 KJ 3.2 g H2 1 ∗ 1 mole H2 2.02 g H2 ∗ −572 KJ2 mole H2 = −453.069 KJ Ratio = Energy ReleasedCoefficient of Hydrogen in formula 31 29.1.2 Calorimetry Find the ∆H by running a reaction and heating or cooling a substance. q = m ∗ c ∗∆T • q = Heat released or absorbed • m = Mass of what is being heated (grams) • c = Specific heat. Unique to every substance ( Jg∗C ) . Specific heat of water is 4.184 • ∆T = Change in temperature Example Burn 0.1 grams of CH4 and it heats 100 grams H2O from 20◦ C to 33.29◦ C. q = 100 ∗ 4.184 ∗ 13.29 = 5560 J = 5.560 KJ 0.1 grams CH4 1 ∗ 1 mole CH4 16 g CH4 = 0.00625 moles CH4 5.560 0.00625 = 889.6 KJ Mole 29.1.3 Hess Law Multiple reactions can be added together then ∆Hs can be added together. Example Si+ 2H2 → SiH4 ∆H = +34 KJMole Si+O2 → SiO2 ∆H = −911 KJMole H2 + 12O2 → H2O ∆H = −242 KJ Mole Find ∆H for: SiH4 + 2O2 → SiO2 + 2H2O SiH4 →Si+2H2 ∆H = −34 KJMole Si+O2 → SiO2 ∆H = −911 KJMole 2H2 + 2O2 → H2O ∆H = −484 KJMole SiH4 + 2O2 → SiO2 + 2H2O ∆H = −1429 KJMole 29.1.4 Standard Heat of Formation Standard heat (enthalpy) of formation (∆H◦f ) 2 is the energy involved in forming one mole of a chemical from its elements under standard conditions. a Elemental substances (O2, H2, etc.) always have a ∆H of zero. Example Find the ∆H for: 2H2O2 → 2H2O +O2 ∆HfH2O2 = −187 ∆HfH2O = −285 2∗(−187) 2H2O2︸ ︷︷ ︸ −374 → 2∗(−285) 2H2O +O2︸ ︷︷ ︸ −570 ∆H = ∑ product− ∑ reactant ∆H = −570− (−374) = −196 KJMole 2This symbol may be shortened to ∆H or ∆Hf in this subsection. 32 30.6 Rate Laws A+B → C +D rate = k[A]m[B]n • k = Constant • m = Order of A • n = Order of B a Order of 0 → No effect a Order of 1 → Linear - Double the concentration and you double the rate a Order of 2 → Squared - Double the concentration and you quadruple the rate Example: Trial [A] [B] Rate 1 0.1 M 0.1 M 0.04 M/s 2 0.2 M 0.1 M 0.08 M/s 3 0.1 M 0.2 M 0.04 M/s Solve for m: trial 2 trial 1 = ( [] [] )m = rate rate = ( 0.2 0.1 )m = 0.08 0.04 2m = 2 m = 1 Solve for n: ( 0.2 0.1 )n = 0.04 0.04 1n = 1 n = 0 rate = k[A]1[B]0 Solve for k: 0.04 = k[0.1]1[0.1]0 k = 0.4 30.6.1 Order Table Comments Zero Order First Order Second Order Rate Law rate = k rate = k[A]1 rate = k[A]2 Integrated Rate law [A]− [A]0 = −kt ln[A]− ln[A]0 = −kt 1[A] − 1 [A]0 = kt [A] = −kt+ [A]0 ln[A] = −kt+ ln[A]0 1[A] = kt+ 1 [A]0 Graph [A] vs Time ln[A] vs time 1[A] vs time K = Slope Slope = −k Slope = −k Slope = k Half-Life (t 1 2 ) t 1 2 = [A]02k t 12 = 0.693 k t 12 = 1 k[A]0 35 Example: 2N2O5 → 4NO2 +O2 [N2O5] Time (s) 0.1 0 0.0707 50 0.05 100 0.025 200 0.0125 300 0.00625 400 1. What is the order of the reaction? [A] 6= straight 1 [A] 6= straight ln[A] = straight Order of 1 2. What is the k constant value? ln(0.0707)−ln(0.1) 50−0 = −0.347 50 = 0.00693 k = 0.00693 3. What is the concentration of N2O5 at t = 150? ln[A] = −(0.00693)(150) + ln(0.1) ln[A] = −3.34 [A] = 0.0354 M 4. What is the rate at 150 seconds? rate = k[A] rate = 0.00693 ∗ [0.0354] rate = 2.45 ∗ 10−4 M/s 5. What is the half life? t 1 2 = 0.693k t 1 2 = 0.6930.00693 t 1 2 = 100 s 31 Reaction Mechanisms Many/most reactions do not take place in one step. If a reaction were to react in one step, then you could use the balanced reaction to determine the rate law. For example, assume the following occured in one step. MgCl2 + 2Hbr → 2HCl +MgBr2 rate = k[MgCl2]1[HBr]2 In reality though, things are not always as easy. Through experimentation we figure out that the rate law for: NO2 + CO → NO + CO2 is rate = k[NO2]2 Because the rate law does not link up with the equation, it is not a single step reaction. 36 31.1 Elementary Steps • Unimolecular - 1 reactant • Bimolecular - 2 reactants • Terrmolecular - 3 reactants 32 Equilibrium The state where the concentration or partial pressures (if it is a gas) of all the reactants and products remain constant with time. For equilibrium to occur, the forward reaction rate must equal the reverse rate. In other words, the amounts do not have to be equal, but the rates must be. 32.1 Types of Equilibrium • Static → No movement • Dynamic → Movement such as a sealed container of water 32.2 Equilibrium Constant Expressions aA+ bB ⇀↽ cC + dD Kc = [C]c[D]d [A]a[B]b Kp = (PCc)(PDd) (PAa)(PBb) • Kc = Concentration constant • Kp = Partial Pressure constant 32.2.1 Converting Constants To convert between the two constants Kc and Kp use the formula: Kp = Kc(RT )∆n • ∆n = ∑ Product Coefficients− ∑ Reactant Coefficients 33 Gas Laws 33.1 Gas Units and Conversions 1 Atm = 760 Torr (mmHg) = 101.3 kPa = 14.7 PSI 33.2 Ideal Gas Law Pv = nRT • P = Pressure (Atm) • v = Volume (L) • n = Number of moles • R = 0.0821 (constant) • T = Temperature (Kelvin) Example 3 grams of HCl at 26◦ C in a 3 Liter container. What is the pressure? P (3) 3 = ( 3 grams36.5 g/mole )(0.081)(26+273) 3 P = 0.0664 Atm 37 35.2 pH and pOH pH and pOH are measures of the amount of ions in a solution that either cause the solution to be acidic or basic. pH Scale Basic ⇒ 0↔ 14⇐ Acidic Important Formulas pH = −log[H+] pOH = −log[OH−] pH + pOH = 14 [H+] = 1 ∗ 10−pH [OH−] = 1 ∗ 10−pOH Example What is the concentration of HCl with a pH of 3? [HCl] = 0.001 M 35.2.1 Changing Concentrations M1V1 = M2V2 (0.25 M)(5 mL) = M2(50 mL) M2 = 0.025 M 35.3 Strong Acids and Bases Strong acids and bases completely dissociate in water. 35.3.1 Strong Acids • HCl • H2SO4 • HBr • HI • HNO3 • HClO4 35.3.2 Strong Bases • Group 1 - Hydroxides . NaOH . KOH • Group 2 - Heavier Hydroxides . Ca(OH)2 . Sr(OH)2 . Ra(OH)2 35.4 Weak Acids and Bases Weak acids and bases do not completely dissociate in water. 40 35.4.1 Ka Constant HA ⇀↽ H+ +A− HA+H2O ⇀↽ H3O+ +A− Ka = [H+][A−] [HA] Example Benzoic acid dissociates as follows: HC7H5O2 ⇀↽ x H+ + x C7H6O − 2︸ ︷︷ ︸ x2 [HC7H5O2] = 0.4 M Ka = 6.3 ∗ 10−5 What is the pH? Ka = [H+][C7H5O − 2 ] [HC7H5O2] 6.3 ∗ 10−5 = x 2 0.4 35.4.2 Kb Constant The Kb constant is used when bases are involved in a reaction (as opposed to Ka which is used in reactions with acids). To convert between Kb and Ka use the following formula: Ka ∗Kb = Kw • Kw = 1 ∗ 10−14 Example F− +H2O ⇀↽ HF +OH− Ka = 7.2 ∗ 10−4 What is the Kb constant? Kb = 1∗10 −14 7.2∗10−4 = 1.39 ∗ 10 −11 Find the pH and pOH. Kb = [HF ][OH−] [F−] 1.39 ∗ 10−11 = x 2 ( 0.00220+13.3 ) x = 9.13 ∗ 10−7 pOH = 6.04 pH = 7.96 35.5 Common Ion Effect The effect of ionization of a weak electrolyte (acid/base) is decreased by adding a strong electrolyte that has an ion in common with the weak electrolyte. 35.6 Buffer Made of 2 components: 1. Weak acid 2. The salt of that acid 41 36 Equilibrium of Saturated, Soluable Salts Solubility is how well a solute dissolves in a solvent4. Example: CaCO3 (s) ⇀↽ Ca +2 (aq) + CO −2 3 (aq) Ksp = [Ca+2][CO−23 ] • Ksp is the solubility product . A large Ksp means the solution is very soluable (meaning lots of products) . A small Ksp means the solution is not very soluable. 1. Given Ksp, find the ion concentration. Ksp = [Ca+2][CO−23 ] = 4.5 ∗ 10−9 [Ca+2] = [CO−23 ] = √ 4.5 ∗ 10−9 = 6.7 ∗ 10−5 M 2. Given Ksp, find the solubility (g/L). 6.7 ∗ 10−5 M = 6.7∗10 −5 1 ∗ 100.1 1mole = 6.37 ∗ 10 −3g/L 3. Given solubility, find ion concentration. Solubility of Silver Chloride at 25◦C is 1.3 ∗ 10−7 g100 mL 1.3 ∗ 10−7 g100 mL → g L ∗ 10 10 = 1.3 ∗ 10 −6 g L 1.3∗10−6 L ∗ 1 mole 143.35 g = 9.11 ∗ 10 −9 m L 4. Given solubility, find Ksp Ksp = [Ag+][Cl−] = (9.11 ∗ 10−9)2 = 8.3 ∗ 10−17 37 Kinetic Molecular Theory 37.1 Postulates: • The volume of the individual particales of a gas can be assumed to be negligible. . So volume is determined by the space between molecules • The gas particles are in constant motion. The pressure exerted by a gas is due to collisions of the gas with the walls of the container. • Gas particles are not attracted to one another. • The average kinetic energy of a gas is directly proportional to the Kelvin temperature. Kenergy = 32 (0.0821)T OR Kenergy = 12 (Molar Mass)(V elocity) 2 A) CO at 760 torr and 0◦C B) N2 at 760 torr and 0◦C C) H2 at 760 torr and 0◦C Q. Which will have the highest kinetic energy? A. All will have the same kinetic energy Q. Which will have a higher velocity? A. H2 will because if all kinetic energies are constant according to the formula k = 1 2 mv2 the smallest mass will yield the highest velocity to keep k constant. 4Virtually every salt is soluable to some degree. 42