Information Required for Memorization Your students are required to memorize the following information for Chem 10. This information must not be supplied on Cheat Sheets for your Semester Exams or Final Exam. SI Prefixes Prefixes greater than the Base Unit Symbol Scientific Notation Tera T 10 12 Giga G 10 9 Mega M 10 6 kilo k 10 3 hecto h 10 2 deka da 10 1 Prefixes less than the Base Unit Symbol Scientific Notation deci d 10 centi c 10 milli m 10-3 micro µ 10-6 nano n 10-9 pico p 10 2 femto f 10 5 Element Symbols and Names Group 1A: H, Li, Na, K, Rb, Cs Group 2A: Be, Mg, Ca, Sr, Ba Group 3A: B, Al, Ga Group 4A: C, Si, Ge, Sn, Pb Group 5A: N, P, As, Sb Group 6A: O, S, Se, Te Group 7A: F, Cl, Br, I Group 8A: He, Ne, Ar, Kr, Xe, Rn Group 3B: Sc Group 4B: Ti Group 5B: V Group 6B: Cr, W Group 7B: Mn Group 8B: Fe, Co, Ni Group 1B: Cu, Ag, Au Group 2B: Zn, Cd, Hg
Monatomic Ions Non-metal Anions Non-metals will form anions with only one possible negative charge. The names of these anions are based on the element root name, but the ending is changed to ide. The following Periodic Table shows the charges for nonmetal anions commonly found in ionic compounds. 1A 2A 3A 4A 5A 6A 7A 8A H B C -4 N -3 O F Si P -3 S Cl As Se Br Te I At Metal Cations Most (but not all) main group metals will form cations with only one possible charge. Most (but not all) transition metals will form cations with more than one possible charge. Metal cation names are the same as the original element names. When metals form more than one cation, then the cation charge must be indicated in the name as Roman Numerals in brackets. The following Periodic Table shows the charges for metal cations commonly found in ionic compounds: 1A 2A 3A 4A 5A H +1 (acids) Li +1 Be Na +1 Mg Transition Metals (B) Al +3 K +1 Ca Ti Ti +4 Cr Cr +3 Cr +6 Mn Mn +3 Mn +4 Fe Co Ni Cu +1 Fe +3 Co +3 Ni +3 Cu Zn Ga +3 Rb +1 Sr Ag +1 Cd In +1 Sn In +3 Sn +4 Cs +1 Ba Au +1 Au +3 Hg Pb Bi +3 Pb +4 Bi +5
Polyatomic Ions OH Hydroxide O 2 CN Cyanide CO 3 SCN Thiocyanate SO 3 HCO 3 Hydrogen Carbonate (Bicarbonate) SO 4 HSO 3 Hydrogen Sulfite (Bisulfite) S 2 O 3 HSO 4 Hydrogen Sulfate (Bisulfate) C 2 O 4 C 2 H 3 O 2 Acetate CrO 4 NO 2 Nitrite Cr 2 O 7 NO 3 Nitrate MnO 4 Permanganate -3 PO 3 ClO Hypochlorite -3 PO 4 ClO 2 Chlorite ClO 3 Chlorate +1 NH 4 ClO 4 Perchlorate Hg 2 Peroxide Carbonate Sulfite Sulfate Thiosulfate Oxalate Chromate Dichromate Phosphite Phosphate Ammonium Mercury (I)
Rules for Assigning Oxidation Numbers This is a prioritized list. If two rules contradict each other, follow the rule that appears higher on the list. 1. The atoms in pure elements are assigned an oxidation number of zero. 2. Monatomic ions are assigned an oxidation number equal to their charge. 3. For atoms in covalent molecules and polyatomic ions: a. The sum of all the oxidation numbers of the atoms in a covalent molecule must equal zero. The sum of all the oxidation numbers of the atoms in a polyatomic ion must equal the charge on the ion. b. Fluorine is assigned an oxidation number of 1. c. Oxygen is assigned an oxidation number of 2 (an exception to this is when oxygen occurs as the peroxide ion, O 2, where it is assigned an oxidation number of 1). d. Hydrogen is assigned an oxidation number of +1. e. For all other elements: the element with the greater electronegativity is typically assigned a negative oxidation number equal to its charge as an anion in ionic compounds. The Half Reaction Method for Balancing Redox Reactions in Acidic Solution This method assumes the reaction occurs in aqueous, acidic solution where H 2 O and H + are plentiful. 1. Write two half-reactions, one for oxidation and one for reduction. 2. Balance each half-reaction as follows: a. Balance all elements other than oxygen and hydrogen. b. Balance oxygen by adding the appropriate number of water molecules (H 2 O). c. Balance hydrogen by adding the appropriate number of hydrogen ions (H + ). d. Balance the charge by adding the appropriate number of electrons (e - ). 3. Multiply each half-reaction by a whole number so that the number of electrons lost in the oxidation halfreaction equals the number of electrons gained in the reduction half-reaction. 4. Add the two half reactions together, keeping all the reactants together on the left of the yield arrow and all the products together on the right of the yield arrow. The electrons will cancel out so they are not shown in the final equation. 5. Cancel any species that appear on both sides of the equation.
VSEPR Theory # of effective electron groups Electronic Geometry Molecular Shape 2 bond angles =180 Linear bond angles =180 Linear 3 bond angles =120 Trigonal Planar bond angles =120 Trigonal Planar bond angles <120 Bent (or V-shaped) 4 bond angles =109.5 Tetrahedral bond angles =109.5 Tetrahedral bond angles <109.5 Trigonal Pyramidal bond angles <109.5 Bent (or V-shaped)