Mr. Iadarola/Mr Harrington
Chemistry Regents Review
Review sessions:
Mon. 6/12 9:30-10:30 am Tues. 6/13 1-2 pm Wed. 6/14 10-11 am
Thurs. 11:15-12:15 pm Mon. 6/19 10-11 am
“Stuff” to know by 8:00 am, June 20, 2017:
1. Solids, liquids, gases: which takes the shape of a container or fills it “uniformly”?, which has the greatest attractions between particles if they are all at the same temp?, which has a regular geometric pattern?, the greatest disorder? (entropy)
2. Phase changes: Melting, freezing, vaporization, condensation, sublimation, fusion. Carbon dioxide and Iodine undergo sublimation. Which phase changes are exothermic? Endo? The opposite of sublimation is deposition
3. Calorimetry: q=cxmxDT if no phase change;
for phase changes q = heat of fusion (or vaporization) x mass. q = DHf x m
4. Kinetic theory of gases/ideal vs. real gases. (no attractions/negligible volume)
5. Hydrogen and Helium behave most ideally, why? Gases deviate from ideal behavior at high P and Low T, why?
6. a) Substances = elements and compounds (see p.17-19 in review book); mixtures vs. compounds; homogeneous vs. heterogeneous matter.
b) Mixtures are two or more substances that can be separated physically and have no definite composition. Compounds can’t be separated physically and have a definite composition (fixed mass ratio).
7. The smaller the mass of the gas, the faster it effuses (Graham’s Law).
8. Boyle’s, Charles’ , Gay-Lussac, Combined Gas Law. Know mathematically and which are inverse and direct. Temps must be in Kelvin!!!
9. STP = 101.3 kPa (or 1 atm) and 273 K (or 0 deg. C).
10. Vapor Pressure – the pressure of the vapor above a liquid when the liquid is in equilibrium with its vapor phase. Temp affects V.P. directly. In order for liquid to boil, its V.P. must equal the pressure on the surface.
If the surface pressure is 1.0 atm, that temperature is called the Normal B.P.
11. Vapor pressure is also a function of intermolecular attractions. Strong attractions mean lower vapor pressure.
12. The atomic number (number of protons) determines the element.
13. Mass number = protons + neutrons. Atomic masses on the table are weighted averages of naturally occurring isotopes based on their abundance in nature.
14. Isotopes are atoms of the same element (same protons), with different masses (number of neutrons).
15. All elements beyond Pb (#82) have only radioisotopes.
16. History of the atom: Dalton’s Model, Plum Pudding, Rutherford’s experiment, Bohr’s Model and Bright-line Spectra, Wave-Mechanical Model. Know the history and specific features of each. How did each one change our concept of the atom.
17. Ground state vs. Excited State
18. Bright-line Spectra result from e- absorbing energy, “jumping” to higher energy levels (excited state), then returning to a lower energy level and emitting the energy released in the form of photons of light.
19. Metals tend to lose e- and form + ions (cations). Nonmetals gain. (Why???)
20. Metals have properties of malleability, thermal/electric conductivity, ductility, luster due to metallic bonding (positive kernels immersed in a sea of mobile e-).
21. What is needed for a substance or solution to conduct electricity?
22. Ionic bonding = transfer of e-, usually metal to nonmetal, eneg. Diff. > 1.7 except metal hydrides (NaH, KH, etc…) which are ionic with eneg diff. < 1.7
23. Covalent bonds = sharing of e-, between two nonmetals, eneg diff = 0 means nonpolar bonding (equal sharing), 0.4-1.7 means polar covalent (unequal sharing).
24. Polyatomic ions – a group of covalently bonded atoms with a charge. (Chart E) They behave as one single entity. When bonded to other ions these compounds contain more than one type of bonding: ionic and covalent. Ex. NaNO3
25. Ionic solids – ionic bonds, no molecules, crystal lattices, generally salts, water soluble, high M.P., poor conductors in solid state (ions can’t move), but good conductors in molten (liquid) or water solutions (ions free to move).
26. Molecular solids – covalent bonds, low M.P., soft, poor conductors, nonpolar
27. Intermolecular Forces (IMFs) –
Hydrogen Bonding – if H is bonded to N, O, or F = strongest type of attraction!
Dipole-Dipole (between polar molecules); Van der Waals or Dispersion Forces (between nonpolar molecules – tend to be very weak, get stronger with size); Molecule-Ion attractions (between an ion and water such as aqueous salt solutions).
28. Intermolecular Forces determine B.P., M.P., vapor pressure, and are related to solubility. The stronger the forces, the higher the B.P., M.P. and lower the Vapor Pressure.
29. Exothermic – energy is released, products lower in E than reactants, downhill P.E. diagram.
30. Endothermic – energy is absorbed, products higher in E than reactants, uphill P.E. diagram.
31. Periodic trends across a period are determined by the increasing number of protons in the nucleus (nuclear charge) while the number of occupied energy levels remains the same: smaller radii, higher Ionization E., more nonmetallic character or tendency to gain electrons).
32. Trends down a group are determined by the increasing number of energy levels resulting in more shielding and a weaker attraction between nucleus and outermost electron (larger radii, lower I.E., more metallic character or tendency to lose e-).
33. Transition Metals – form colored ions or compounds, form bonds by losing e- from more than one energy level, have multiple oxidation #s and form more than one binary compound with halogens (CuCl and CuCl2 for example).
34. Dot Structures only involve valence e-; Ions need brackets and charges.
35. Metals have ionic radii smaller than their atoms because they lose e- to form ions. Nonmetals gain electrons to form ions, so their ionic radii > atomic radii.
36. Avogadro’s Law – equal volumes of gases at same temp. and pressure have equal numbers of molecules.
37. Types of radioactive emissions (alpha, beta, gamma – know symbols and characteristics). They can be separated be passing them through an electric field.
38. Writing formulas (write the ions, cross the charges, reduce). Stock system naming uses the oxidation # of the first element in roman numerals (Iron (II) chloride = FeCl2). Remember for acids, ate-ic, ite-ous.
39. BrINClHOF are all diatomic elements. Allotropes are different forms of the same element with different chemical and physical properties (diamond and graphite…ozone O3 and oxygen O2)
40. Know the metalloids (properties between those of metals and nonmetals)
41. Periodic Law (properties of elements repeat periodically as a function of atomic number); elements in the same group have similar properties (Same # valence e-)
42. Noble gases are monatomic gases (1 atom = 1 molecule); Halogens are diatomic; alkali and alkaline earth metals are most reactive metals. Halogens are most reactive nonmetals. Reactive elements are prepared by electrolysis of fused salts.
43. Noble gases are largely inert, but the larger ones have been made to form compounds with Fluorine, under extreme conditions.
44. 1 mole = 1 gram atomic mass or gram formula mass or molecular mass.
45. Ionic compounds are electrolytes. If they dissolve in water they dissociate. The only molecular compounds that are electrolytes are acids (and organic bases). Acids and bases that are molecular IONIZE in water (react with water). Organic acids and amines are weak electrolytes (pretty much the only organic compounds that are electrolytes).
46. Solution = homogeneous mixture.
47. F.P. depression, B.P. elevation. The amount of change in B.P. or F.P. is proportional to the concentration of dissolved particles. Ionic compounds and electrolytes dissociate or ionize to form more particles than nonelectrolytes of similar concentration (1 M NaCl has a higher B.P. than 1M sugar solution).
48. a) Know your P.E. diagrams. Where is the activated complex, the forward and reverse activation energy, change in enthalpy (heat of reaction), exo vs. endo., catalyzed vs. uncatalyzed.
b) Collision Theory!!! What leads to effective collisions? What does a catalyst do? How do concentration, temp., pressure affect the frequency of effective collisions?
49. Equilibrium – forward and reverse processes occur at same rate, all concentrations are constant (not equal).
50. LeChatelier’s Principle – equilibrium systems will shift to alleviate a stress put on the system. Increasing or decreasing concentrations, pressure or temp. are typical stresses. Catalysts don’t shift the equilibrium. Increasing Pressure favors the side with fewer moles of gas. Increasing Temp. favors the endothermic rxn.
51. Arrhenius acids – form H+ ion; Arrhenius bases – form OH- ion.
H+ + OH- à water is the net ionic equation for a neutralization.
52. Alternate Theory acids (Bronsted)– proton donors; bases – proton acceptors.
53. Titrations – MaVa x #H+ = MbVb x # OH-
54. Indicators, table M.
55. pH is based on conc. of H+ ion, it is an exponential scale. More hydronium ion (H3O+)means lower pH. 10X increase in H+ ion means pH decreases by 1 pH unit.
56. Hydrolysis of salts – know the parent acid and base to determine the type of solution formed when the salt is dissolved in water.
57. Ox. #s – know the rules and exceptions.
58. LEO-GER, or OIL RIG.
59. AnOx RedCat
60. Voltaic cells (batteries) produce electricity from spontaneous redox reactions. A salt bridge allows ions to migrate and maintains electrical neutrality of the half cells.
61. Electrolytic cells require electrical energy (electricity) be applied to the cell to make a nonspontaneous reaction occur. Typically used for electroplating.
62. The object to be plated is placed at the cathode (reducing on the object surface).
63. The two driving forces in nature are a decrease in energy (or enthalpy, DH= -) and an increase entropy (order to disorder, DS =+).
64. Equations must be balanced by mass and charge!!! Redox involves both the gain and loss of electrons simultaneously.
65. The higher the element is on table J, the more active it is. A more active element will replace a less active element in a single replacement rxn. Cu, Ag, Au will not react with H+ (acid) because they are less active than H2 according to the table. The more active the metal (nonmetal) the more likely it is to oxidize (reduce).
66. The strongest oxidizing agent is the element most easily reduced (Fluorine).
67. Mass, Charge and Energy are conserved in all chemical and physical changes.
68. Carbon forms 4 bonds. If single bonds only (saturated), the bonds are directed to the corners of a tetrahedron.
69. Unsaturated hydrocarbons have double or triple bonds (alkenes, alkynes).
70. Know examples of molecules that are nonpolar by symmetry (CH4, CO2). Hydrocarbons are generally nonpolar by symmetry. Molecules that have asymmetrical distributions of charge (electrons) are dipoles (polar).
71. Alkanes undergo substitution (form many products), alkenes and alkynes undergo addition (the multiple bond breaks and only one product forms).
72. Relying solely on the reference table for naming and identifying functional groups is a recipe for disaster!!!
73. Esterification = acid + alcohol à water + ester
74. Saponification = Fat + NaOH à soap + glycerol; Fat = glycerol + 3 Fatty acids à fat + 3 water
75. Benzene and Toluene (recognize their structure and that they are aromatic compounds).
76. Recognize polymers and the type (addition or condensation polymers). Any reaction that is a condensation reaction involves the formation of water.
77. Fission – splitting heavy nuclei into lighter ones, releasing Energy
Fusion – combining light nuclei into heavier ones, releasing Energy
78. In any nuclear process by which energy is released the mass of the products is less than the mass of the reactants (mass is converted to Energy by E=mc2).
79. Fissionable material – U-235, U-233, Pu-239.
80. Fusion requires high T and P to overcome repulsions between nuclei. Fission requires fissionable material and slow neutrons. Know the parts of a Fission reactor.
81. Uses of radioisotopes: tracers, medical, carbon dating, food irradiation, etc…
82. Sig figs!!! How many sig figs does 0.004010 have?
83. Water is polar (a dipole)!!! Anything polar dissolves in it. Like dissolves like.
84. Br and Hg are the only two liquid elements at STP. Know which elements are gases at STP and use ref. table BP and MP data to determine an element’s phase.
85. The larger the number of Carbons in a compound the greater the number of isomers, the stronger the intermolecular forces (dispersion), and the higher the B.P./M.P. and lower the vapor pressure.
86. Typical dipoles – H2O, HF, NH3, HCl
87. Elements can’t be decomposed by chemical means.
88. Compounds are made of two or more elements combined in a definite ratio.
Mixtures are made up of two or more substances physically combined (not bonded to one another) and the composition can vary. The components of a mixture retain their properties.
89. Nuclear changes have only to do with the nucleus; chemical changes have to do with rearranging electrons.
90. Nuclear reactions (fission and fusion) release much more energy per gram than chemical reactions.
91. Be familiar with the reference tables in order to efficiently find information.
92. Study, bring a #2 pencil, black pen, and a non-programmable calculator.