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Just a Few Things To Remember for the 2015 Regents Chemistry Exam —Florian König Hastings High School 2014/2015 1. Homogeneous mixtures are called solutions. example: NaCl (aq), air Solid solutions are often between metals alloys Liquid solutions are transparent. 2. Temperature reprents (NOT: is equal to) the average molecular kinetic energy. Heat is the amount of total molecular kinetic energy. Objects with the same temperature but different mass or heat capacity have different amounts of total energy. 3. Endothermic reaction: A + B + energy (Energy is absorbed.) Exothermic reaction: A + B (Energy is released.) C C + energy 4. # of Joules = mass of water (in gram) x T x 4.2 J / (g K) 5. Kelvin scale: no negative numbers, based on molecular motion K = oC + 273 0 K = -273 oC 6. Boyle’s Law: The volume of a gas at constant temperature is inversely proportional to its pressure. (pV = constant p1V1 = p2V2) 7. Charles’ Law: The volume of a gas at constant pressure is directly proportional to its Kelvin temperature. (V/T = constant V1/T1 = V2/T2) 8. Real gases differ from ideal gases when the molecules are close together, i.e. under conditions of high pressure and low temperatures (best approximations to ideal gas: H2, He) 9. vapor pressure: depends only on temperature of a substance, not on its amount (Ref. H) 10. Boiling occurs when the vapor pressure = atmospheric pressure “normal boiling point” = boiling point at 1 atm pressure (101.3 kPa) 11. energy for (solid liquid) = “heat of fusion” (change of EPOT) energy for (liquid gas) = “heat of vaporization” (change of EPOT) 12. Sublimation: solid gas (no liquid phase) examples: dry ice CO2, iodine I2 13. nucleons: protons and neutrons mass number = # (protons) + # (neutrons) atomic number = # (protons) 14. Mass number - atomic number = # (neutrons) 15. In a neutral atom the positive and negative charges balance each other. #(protons) = #(electrons) 16. Isotopes are different kinds of atoms of the same element. (same number of protons, different number of neutrons) e.g. 12C/14C, 235U/238U 17. Atomic mass unit (amu) = 1/12 of the mass of 12C 18. Atomic mass of an element = weighed average of isotope masses 19. Ground state: all lower energy levels are filled, e.g., 2-8-18-4 (Germanium Ge) (electron configuration on Periodic Table: ground state) Excited state: one or more lower energy levels are not filled, e.g., 2-7-18-2 20. Spectral lines are produced when electrons fall back from higher to lower energy levels (e.g., PEL 3 PEL 2) and emit energy 21. Period number = principal energy level (PEL) 22. Valence electrons: (max. 8) electrons of the outermost principal energy level (valence shell) 23. Ionization energy is the energy to remove the 1st, 2nd, 3rd, ... valence electron (Ref. S) 24. Electronegativity is the attraction of an atom to other electrons (Ref. S) High electronegativity = nonmetals (fluorine 4.0), low = metals 25. Radioactivity (Ref. N & O) alpha emission (alpha decay) beta emission (beta decay) 26. atomic number decreases by 2, mass number by 4 atomic number increases by 1, mass number constant The rate of radioactive decay (half-life!) is independent of any conditions. 27. Nonpolar covalent bonds: Polar covalent bonds: Ionic bonds: EN = 0 - 0.4 EN = 0.4 - 1.7 EN > 1.7 28. Nonpolar molecules: e.g. H2, Cl2, N2, ... (only nonpolar bonds) CO2, CBr4, CCl4 (polar bonds, but symmetrical) 29. Polar molecules (dipoles): e.g. HCl, H2O, NH3, CHCl3, CH3COOH 30. Network solids = covalent bonds, no distinct molecules in the solid phase, e.g., C (graphite, diamond), SiO2 (quartz), asbestos 31. Molecular solids = covalent compounds, distinct molecules in the solid phase, e.g., H2O, CO2, CH4, SO3, etc. 32. Metallic bonds = cations (kernels) in a “sea of valence electrons” (“chocolate chip cookie” model) 33. Van der Waals’ forces (vdW) = weak attraction between nonpolar molecules, e.g., He, CO2, hydrocarbons 34. Dipole-dipole interactions = attraction between polar molecules H bonds = dipole-dipole interactions between H2O/HF/NH3 molecules 35. Ionic solids = high m.p and b.p, hard, conductivity only when molten or dissolved Molecular solids = low m.p. and b.p., soft, no conductivity Metallic solids = high m.p. and b.p., varying hardness, high conductivity Empirical formula Molecular formula = = lowest atom-atom ratio multiple of empirical formula 36. e.g. CH (empirical formula) REMEMBER: Polarity is a gradually changing property. C2H2, C6H6 (molecular formula) 37. nonmetals: metals r(atom) < r(anion) r(atom) > r(cation) 38. Most active nonmetal (strongest oxidizing agent): fluorine F 2 Most active metal (strongest oxidizing agent): cesium Cs (or francium Fr) (Ref. J) 39. Chemical groups are the columns in the Periodic Table. Elements in groups have similar chemical properties. Group 1 (I A) Group 2 (II A) Group 17 (VII A) Group 18 (VIII A) alkali metals alkaline earth metals halogens noble gases 40. Transition metals = metals with e- in d-orbitals Transition metal compound are often colored (CuSO4 blue, K2CrO4 yellow) and have more than one oxidation state. 41. Monoatomic molecules: noble gases He, Ne, Ar, Kr, Xe, Rn Diatomic molecules: Br2, I2, N2, Cl2, H2, O2, F2 42. 1 mole = 6.02 x 1023 particles (atoms, ions, molecules, formula units) mass of 1 mole = atomic (formula) mass in gram mass of element 43 Percentage composition = __________________________________ (Ref. T) formula mass of compound 44. Molar volume = volume of 1 mole of any gas at STP = 22.4 liters o (STP = 0 C and 1 atm) 45. Avogadro’s Law: Equal volumes of gas at equal temperature and pressure contain equal numbers of molecules. 46. You should be able to solve mass-mass, mass-volume and volume-volume problems (or do the same thing with moles) 47. Balanced equation gives: 48. Molarity M = a) mole ratio b) volume ratio (for gases only) moles solute ___________________________ (Ref. T) liters solution (solute: smaller part of solution; solvent: bigger part) 49. Dissolved solute LOWERS freezing point and RAISES boiling point. Freezing and boiling point constants for water: (first page of Ref. tables) 50. Heat of reaction = H = Hproducts – Hreactants H < 0 exothermic (see Ref.I for examples) H > 0 endothermic 51. Catalyst lowers activation energy for both forward and reverse reaction speed up reaction do not change position of equilibrium 52. LeChatêlier’s Principle: A system at equilibrium reacts to an external stress in such a way that the stress is relieved. e.g. temperature increases temperature decreases pressure increases pressure decreases reactant is added etc. pp. 53. Two driving forces: 54. Entropy S increases with: 55. endothermic reaction favored exothermic reaction favored side with fewer gas molecules favored side with more gas molecules favored shift towards product side a) enthalpy H tends to be minimal b) entropy S tends to be maximal phase: solid < liquid < solution < gas temperature: S (20 oC) < S (70 oC) # of particles, amount Acids and bases (Ref. K & L) Strong acids: hydrochloric acid HCl, nitric acid HNO3, sulfuric acid H2SO4 Weak acids: carbonic acid H2CO3, acetic acidCH3COOH Strong bases: alkali hydroxides (LiOH, NaOH, KOH) Weak bases: ammonia NH3 56. Brønsted acid = proton donor Brønsted base = proton acceptor BAAD (Bases Accept, Acids Donate) 57. acid 58. Amphoteric substances can act as acids or bases. conjugate base + H+ 59. Titration: MA x VA = MB x VB 60. Acidic salts: NH4Cl, (NH4)2SO4 Basic salts: K2CO3, Na2SO3 Neutral salts: KNO3, Li2SO4 61. Strong acids: H2SO4, HNO3, HCl, HI 62. pH = -log[H+] pH [H+] [OH-] The solution is… <7 10-7 10-7 acidic > M < M =7 10-7 M 10-7 M neutral >7 < 10-7 M > 10-7 M basic 63. Oxidation numbers (see separate handout with rules and common oxidation states) 64. -4 -3 -2 ———————— <———————— 65. Hydrides: Compounds with negatively polarized hydrogen (ox. state = -1) (LiH, NaH, KH) 66. Oxidizing agents = oxidize something else, get reduced Reducing agents = reduce something else, get oxidized 67. Reduction strengths (Ref. J) Strong oxidizing agents are on top right (F2, Cl2), strong reducing agents on top left (Li, Rb, K, …) 68. Electrochemical cell: Electrolytical cell: 69. In BOTH types of cell the same types of reaction occur at the same electrode: ANode — OXidation 70. -1 0 +1 +2 +3 +4 oxidation ———————————> reduction ——————————— chemical energy electrical energy electrical energy chemical energy REDuction — CAThode Objects being plated in an electrolytical cell are always wired as the (negative) cathode. 71. The farther away two metals are from each other in Ref. J, the more energy is released when they react with each other. 72. Organic compounds = covalent, molecular compounds based on carbon (Ref. P, Q, R) usually nonpolar, low m.p./b.p., low reactivity 73. Isomers = same molecular formula, but different structure 74. Saturated = only single bonds (alkanes) Unsaturated = multiple bonds (alkenes or alkynes) 75. Alkanes Alkenes Alkynes Benzenes CnH2n+2 CnH2n CnH2n-2 CnH2n-6 (cyclic, alternating single/double bonds) 76. Ethyne = acetylene = C2H2 H—CC—H 77. Benzene series: 78. Alcohols R — OH do not ionize (OH is covalently bound) benzene C6H6 toluene C6H5—CH3 Type of alcohol # of OH groups Example monohydroxy ~ 1 CH3OH methanol CH3CH2OH ethanol dihydroxy ~ 2 C2H4(OH)2 ethylene glycol trihydroxy ~ 3 C3H5(OH)3 glycerol Type of alcohol Structural characteristic | R — C — OH | Primary secondary R1 | R2 — C — OH | tertiary R1 | R2 — C — OH | R3 Example ethanol CH3CH2OH propanol CH3CH2CH2OH 2-propanol (CH3)2CH — OH 2-methyl-2-propanol (CH3)3C — OH 79. Organic acids (carboxylic acids) = R — COOH 80. Acid + CH3CH2COOH + 81. Aldehydes Ketones Ethers Amines 82. Addition Substitution Polymerization Cracking Saponification 83. Nuclear Chemistry (Ref. N & O): a) All isotopes with an atomic number bigger than 83 are unstable. b) Neutrons (no charge!) are not accelerated in a particle accelerator. c) Mass defect = sum of neutron and proton masses is less than nucleus mass Alcohol CH3CH2OH Ester + water CH3CH2COOCH2CH3 + H2O R—CHO R1—CO—R2 R1—O—R2 R—NH2 to multiple bonds (e.g. hydrogenation of unsaturated fats) at single bonds (saturated compounds) small building blocks large molecules large molecules small building blocks hydrolysis of fats with lye fatty acids (soaps) + glycerol 84. Nuclear reaction: a) fission = splitting of nuclei (naturally or after neutron bombardment) b) fusion = two nuclei are combined 4 c) types of radiation He nuclei (mass = 4, charge = +2) electrons (mass = 0, charge = -1) high-energy radiation (no mass or charge) d) nuclear reactors U-235 or U-238 as fuel moderator slow down neutrons control rods absorb neutrons 85. Radioisotopes 86. Significant figures: see separate handout 87. Do as you oughta, acid to water! 88. Percentage error: Ref. T 89. Applications of chemical principles Tc-99 to detect brain tumors I-131 to treat thyroid gland Haber Process N2 + 3 H2 2 NH3 (Ammonia is used for fertilizers and explosives.) 90. Refining of metal ores: reaction of metal oxide with carbon, e.g., Fe3O4 + 2 C reaction of metal sulfide with oxygen, e.g., CuS + O2 3 Fe + 2 CO2 Cu + SO2 91. Fractional distillation = separates petroleum into different parts with different boiling points 92. READ THE QUESTIONS CAREFULLY! 93. Read the Reference Tables. Sing the Reference Tables. Be the Reference Tables. 94. Do not study the night before or the morning of the exam. Get a good night’s sleep. Have breakfast. 95. BRING A PEN, A NO. 2 PENCIL, A NON-GRAPHING CALCULATOR, SOME WATER AND A SNACK!