AP Chemistry Unit 1 Review Guide: IUPAC Naming, Stoichiometry, Solution Chemistry

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1 I. IUPAC Naming AP Chemistry Unit 1 Review Guide: IUPAC Naming, Stoichiometry, Solution Chemistry For Ionic Compounds: Formula to Name: 1. Identify the cation (positive ion) by name, then identify the anion by name. Then, put the two names together. 2. For metals that can have more than one oxidation state, it is important to identify the oxidation state in the name, using Roman numerals. Name to Formula: 1. Write the formula of the cation 2. Write the formula of the anion 3. Balance the formula electrostatically, using subscripts If you need more than one of a polyatomic ion, you must place the subscript outside of a set of parentheses that surround the polyatomic ion. For binary Acids: hydro the root of the nonmetal element ic acid. IE HCl = hydro chlor ic acid For Binary Molecular Compounds: 1. The naming system is for compounds composed of two nonmetallic elements. 2. The first element keeps its name a. The first element gets a prefix if it has a subscript in the formula 3. The second element gets the ide suffix (ending) a. The second element ALWAYS gets a prefix Prefixes 1 mono 6 hexa 2 di 7 hepta 3 tri 8 octa 4 tetra 9 nona 5 penta 10 - deca For Oxyacids: ate polyatomic ions (sulfate, nitrate, etc) make the ic acids (sulfuric acid, nitric acid) Acids with one more oxygen than the -ic acid is called the per- -ic acid. Acid with one less oxygen then the -ic acid is called the -ous acid. Acid has one less oxygen than the -ous acid, it is called the hypo- -ous acid. Examples: H 2 SO 5 = persulfuric acid HNO 4 = pernitric acid H 2 SO 4 = sulfuric acid HNO 3 = nitric acid H 2 SO 3 = sulfurous acid HNO 2 = nitrous acid H 2 SO 2 = hyposulfurous acid HNO = hyponitrous acid The KEY: All you really need to know are the ate ions. sulfate = SO 4 2-, nitrate = NO 3 -, chlorate = ClO 3 -, bromate = BrO 3 -, phosphate = PO 4 3-, carbonate = CO Compound Name Iron(III) nitride Calcium thiocyanate Sodium sulfite Copper(I) phosphate Silicon dioxide Cesium nitrite Formula PbCl 2 Na 2C 2O 4 KCN (NH 4) 2CO 3 Cu(NO 3) 2 Acid Name sulfurous acid Nitric acid Hypobromic acid Compound Name Carbon dioxide Carbon monoxide Diphosphorus pentoxide Dinitrogen monoxide Silicon dioxide Carbon tetrabromide Formula HF H 3PO 4 HClO Formula As 2O 5 PCl 3 CCl 4 H 2O SeF 6

2 II. Stoichiometry 1. Counting by weighing: atomic mass, formula mass A. Average atomic mass (mass# 1 x %abundance 1) + (mass# 2 x %abundance 2) + etc The Mole A. Avogadro s number x 1023 units = 1 mole B. An element's atomic mass expressed in grams contains 1 mole of atoms of that element grams of carbon is 1 mole of carbon grams of carbon-12 is 1 mole of carbon Molar Mass A. Molar Mass (Gram molecular weight) 1. The mass in grams of one mole of a compound 2. The sum of the masses of the component atoms in a compound a. Molar mass of ethane (C 2H 6): Mass of 2 moles of C = 2(12.01 g) Mass of 6 moles of H = 6(1.008 g) g 4. % Composition = (mass of part) x 100 (mass of whole) 5. Determining the Formula of a Compound A. Empirical formula (Smallest ratio of atoms in a compound) 1. Determine the percentage of each element in your compound 2. Treat % as grams, and convert grams of each element to moles of each element 3. Find the smallest whole number ratio of atoms 4. If the ratio is not a whole number, multiply each by an integer so that all elements are whole numbers B. Determining the molecular formula 1. Find the empirical formula mass 2. Divide the known molecular mass by the empirical formula mass, deriving a whole number, n 3. Multiply the empirical formula by n to derive the molecular formula 6. Chemical Equations A. Balanced chemical equations follow Law of Conservation of Mass, reactants on left, products on right. B. (s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous C. Coefficients are used to balance chemical equations 7. Stoichiometric Calculations: Amounts of Reactants and Products A. Balance the chemical equation B. Convert grams of reactant or product to moles C. Multiply by the molar ratio from balanced chemical equation D. Convert to grams of desired substance if necessary 8. Limiting Reagent A. The limiting reactant controls the amount of product that can form 1. Solving limiting reactant problems 2. Convert grams of reactants to moles 3. Use stoichiometric ratios to determine the limiting reactant 4. Solve as before, beginning the stoichiometric calculation with the grams of the limiting reactant

3 III. Solution Chemistry 1. Water, the common solvent A. Polar, dissolves polar solutes & ionic compounds B. Positive ions are attracted to negative oxygen end of water C. negative ions are attracted to positive hydrogen ends of water D. Like Dissolves Like 1. ionic, polar solutes dissolve in polar solvents 2. nonpolar solutes dissolve in nonpolar solvents 2. Strong vs. Weak electrolytes A. Ionic compounds that dissolve completely are considered strong electrolytes 1. Solubility charts( Group 1 ions, ammonium, nitrates) 2. Chlorides, bromides, iodides (exceptions: silver, mercury, & lead) 3. Fluorides except Group 2, Pb2+ & Fe2+ 4. Sulfates except Ca, Sr, Ba, Hg, Pb2+ & Ag 5. Strong acids (HCl, H 2SO 4, HNO 3) B. Weak electrolytes do not dissolve in water completely 1. Most Carbonates, Phosphates, Chromates, except Group 1 ions, ammonium, nitrate 2. Most Hydroxides except Group 1, Ammonium, Ba, & Sr 3. Most Sulfides except Group 1, Group 2, and Ammonium 4. All oxides except Group 1, Ammonium 3. Concentration A. Molarity = moles of solute Liters solution Types of Solution Reactions 1. Precipitation reactions A. Ionic compounds dissolve in water and the ions separate and move independently B. When two solutions are mixed, an insoluble solid forms C. precipitate can be predicted using solubility rules Ex: NaCl(aq) + AgNO 3(aq) AgCl(s) + NaNO 3(aq) AgCl is insoluble, therefore forms a precipitate E. If there is no insoluble product, the reaction does not occur Ex. NaCl(aq) + KNO 3(aq) NaNO 3(aq) + KCl(aq) F. Molecular Equations contain all the molecules, formula units, and their states (s,l,g,aq) G. Full Ionic Equations display each cation and anion as a separate species E. Net Ionic Equations only contain the ions that result in a chemical reaction. Ex: NaCl(aq) + AgNO 3(aq) AgCl(s) + NaNO 3(aq) Molecular Equation Na + (aq) Cl - (aq) + Ag + (aq) + NO 3 -(aq) AgCl(s) + Na + (aq) NO 3 -(aq) Full Ionic Equation Cl - (aq) + Ag + (aq) AgCl(s) Full Ionic Equation 2. Acid-Base reactions A. A soluble hydroxide and a soluble acid react to form water and a salt 3. Red-Ox Reactions 1. Red-ox reactions involve the transfer of electrons. To predict the products of a redox reaction, look at the reagents given to see if there is both an oxidizing agent and a reducing agent. Rules for Assigning Oxidation Numbers 1) the oxidation number of the atom or molecule of a free element is zero, Ex: H2 = 0 2) the oxidation number of a monatomic ion equals its charge 3) In compounds, oxygen has an oxidation number of -2, except in peroxides, where it is -1 4) In compounds containing hydrogen, hydrogen has an oxidation number of +1 5) In compounds, fluorine is ALWAYS assigned an oxidation number of -1 6) The sum of the oxidation states for an electrically neutral compound must be zero

4 Practice Problems For Stoichiometry Average Atomic Mass: 1. What is average atomic mass of Lithium if 7.42% exists as 6 Li (6.015 g/mol) and 92.58% exists as 7 Li (7.016 g/mol)? 2. Magnesium has three naturally occuring isotopes % of Magnesium atoms exist as Magnesium-24 ( g/mol), 10.03% exist as Magnesium-25 ( g/mol) and 11.17% exist as Magnesium-26 ( g/mol). What is the average atomic mass of Magnesium? Empirical Formula: 3. Find the empirical formula for the following molecular composition: 40% carbon; 6.7% hydrogen; 53.3% oxygen 4. If the molar mass of the compound in problem 3 is 92g/ mol, what is the molecular formula? Stoichiometry: 5. Balance the following chemical equation: C 2H 4O 2 + O 2 CO 2 + H 2O How much carbon dioxide will be produced from the combustion of grams of C 2H 4O 2? 6. Balance the following chemical equation: Al(s) + Cl 2 AlCl 3 How much chlorine gas will be needed to react with grams of aluminum? 7. If you have 40.0 grams of each of the reactants in problem 5, which is the limiting reagent and how much, in grams, of water will be produced? 8. If you have 0.75 grams of each of the reactants in problem 6, which is the limiting reagent and how much, in grams, of aluminum chloride will be produced? Write Molecular, Full Ionic, & Net Ionic equations for the following reactions. Label as No Reaction if no precipitate is produced. 9. Aqueous solutions of potassium iodide and silver nitrate are mixed. 10. Aqueous solutions of ammonium phosphate and sodium sulfate are mixed

5 Practice Problems For Stoichiometry Average Atomic Mass: 1. What is average atomic mass of Lithium if 7.42% exists as 6 Li (6.015 g/mol) and 92.58% exists as 7 Li (7.016 g/mol)? 2. Magnesium has three naturally occurring isotopes % of Magnesium atoms exist as Magnesium-24 ( g/mol), 10.03% exist as Magnesium-25 ( g/mol) and 11.17% exist as Magnesium-26 ( g/mol). What is the average atomic mass of Magnesium? Empirical Formula: 3. Find the empirical formula for the following molecular composition: 40% carbon; 6.7% hydrogen; 53.3% oxygen 4. If the molar mass of the compound in problem 3 is 92g/ mol, what is the molecular formula? Stoichiometry: 5. Balance the following chemical equation: C 2H 4O 2 + O 2 CO 2 + H 2O How much carbon dioxide will be produced from the combustion of grams of C 2H 4O 2? 6. Balance the following chemical equation: Al(s) + Cl 2 AlCl 3 How much chlorine gas will be needed to react with grams of aluminum? 7. If you have 40.0 grams of each of the reactants in problem 5, which is the limiting reagent and how much, in grams, of water will be produced? 8. If you have 0.75 grams of each of the reactants in problem 6, which is the limiting reagent and how much, in grams, of aluminum chloride will be produced? Write Molecular, Full Ionic, & Net Ionic equations for the following reactions. Label as No Reaction if no precipitate is produced. 9. Aqueous solutions of potassium iodide and silver nitrate are mixed. 10. Aqueous solutions of ammonium phosphate and sodium sulfate are mixed

6 Practice Problems For Red-Ox Reactions Assign oxidation states for each of the elements in the compounds below: Ex: NO 2 N = 4 +, O = 2-1. O2 5. NO3-2. H2O2 6. AlBr3 3. H2SO4 7. CrO42-4. Cr2O72-8. C2H5OH Given the following chemical equations, determine if the reaction is an oxidation-reduction reaction, identify the element that is being oxidized, the one being reduced, and then balance the chemical equation. You will not be required to write the halfreactions, but they can help! Example: Zn(s) + HNO 3(aq) Zn(NO 3) 2(aq) + H 2(g) Oxidation States 0 1+, 5+, 2-2+, 5+, 2-0 Oxidized: Zinc Zn Zn e- Reduced: Hydrogen 2 H e- H 2 Zn(s) + 2 HNO 3(aq) Zn(NO 3) 2(aq) + H 2(g) 1. N 2H 4 + Cu(OH) 2 + H + N 2 + Cu + H 2O 2. KMnO 4 + C 2O 4 + H 2O MnO 2 + CO 2 + KOH 3. S 8(s) + Fe(s) Fe 2S 3(s)