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AP Chemistry Note Outline Chapter 4: Reactions Strong and Weak Electrolytes Types of Reactions Introduction Precipitation Reactions Formation of a Gas Acid-Base Reactions Oxides Hydrolysis Oxidation-Reduction Reactions Reaction Prediction--Overall Tuesday, September 24 - TEST Goal: To be able to predict the products of a chemical reaction and answer a simple question about it. Strong and Weak Electrolytes (4.2) Definitions: o Strong Electrolyte o Weak Electrolyte o Nonelectrolyte Types of Reactions Introduction Page 1 of 13 Reactions Look for the following: Look for the following: 1. Two Ionic Compounds This will be a double replacement reaction and either a ppt is formed or a gas. You MUST know your solubility rules. 1. Key words Acidified or basic solution 2. Oxides: know the oxide rules* 3. Hydrolysis of a salt. Water and a salt produces an acid/base combo. 4. Acid/Base Reactionswater may be formed. 5. Decomposition Reactions 6. Words like Equilmolar are most likely acid/base reactions 7. Words like Protic or Diprotic lead you to acids *Oxide Rules: 1. nonmetal oxide and water → acid 2. metal oxide and water → base Page 2 of 13 2. Know basic oxidizing and reducing agents like MnO4 Cr2O72 H2O2 SO323. Combustion Reactions: Reactions that involve air are combustion and the reactant is oxygen (O2). 4. Single Replacement Reactions: A pure metal (NOT an ion) present is a single replacement reaction 5. Any element like a halogen (Cl2, Br2) is S.R. 6. Something is electrolyzed. 3. nonmetal oxide + metal oxide → salt Page 3 of 13 Types of Reactions: I. Reactions with no changes in Oxidation Numbers - Metathesis A. Precipitation OR the formation of a ____________ _________________________ (4.4-6) 1. Definition - the reaction of two ionic compounds to form an insoluble solid 2. Writing Equations a. Molecular Equation Cd(NO3)2(aq) + Na2S(aq) b. Ionic Equations: All strong electrolytes are shown ionized Cd2+(aq) + 2 NO3-(aq) + 2 Na+(aq) + S2-(aq) c. Net Ionic Equations: Cd2+(aq) + S2-(aq) 3. When a Weak Electrolyte is a Product Molecular Equation NaC2H3O2(aq) + HCl(aq) Ionic Equation Na+(aq) + C2H3O2-(aq) + H+(aq) + Cl-(aq) Net Ionic Equation C2H3O2-(aq) + H+(aq) 4. What to do when a nonsoluble reactant is used: Molecular Equation Mg(OH)2(s) + 2 HCl(aq) Ionic Equation Net Ionic Equation Page 4 of 13 YOU MUST LEARN THE SOLUBILITY RULES: THEY MUST BE MEMORIZED!!! 5. Examples: a. Solutions of strontium nitrate and sodium sulfate are mixed. b. A solution of copper (II) chloride is added to a solution of sodium sulfide. c. Solutions of sodium iodide and lead (II) nitrate are mixed d. A solution of copper (II) sulfate is added to a solution of barium hydroxide. e. Solutions of tri-potassium phosphate and zinc nitrate are mixed. B. Formation of a Gas 1. There are four types of gases that can be formed. a. b. c. d. 2. Writing equations: a. Molecular: 2 HCl(aq) + Na2S(aq) b. Ionic c. Net Ionic Page 5 of 13 3. Examples: a. Solid calcium carbonate is placed in hydrochloric acid. b. Solutions of sodium hydroxide and ammonium nitrate are mixed. c. Solid sodium sulfite is added to hydrobromic acid. C. Acid – Base Reactions (Neutralization) (4.8) 1. Acid and a Base forms _____________ and a _______________. a. Hydrochloric acid reacts with sodium hydroxide 2. Anhydrides: a. Acid anhydride SO3(g) + H2O N2O5(g) + H2O w / H 2O b. Base anhydride Na2O (s) + H2O CaO (s) + H2O d. Anydrides with an acid or base produce ___________ and a ___________. SO2 (g) + NaOH (aq) BaO (s) + HCl (aq) c. Metallic oxides + nonmetallic oxides Salt CaO(s) + SO2(g) Page 6 of 13 D. Hydrolysis – salts reacting with _______________________. w / H2O Example : KF(s) + H2O E. Some Decomposition Reactions A. Base metal oxide + water Ca(OH)2 CaO + H2O B. Salt containing oxygen metal oxide + nonmetal oxide CaCO3 CaO + CO2 C. Acid containing oxygen water + nonmetal oxide H2CO3 H2O + CO2 F. Non-Redox Rxn Practice 1. Hydrogen sulfide is bubbled through a solution of silver nitrate. 2. Equal volumes of dilute equimolar solutions of sodium carbonate and hydrochloric acid are mixed. 3. Dilute acetic acid solution is added to solid magnesium carbonate. 4. Sulfur trioxide gas is added to excess water. 5. Powdered magnesium oxide is added to a container of carbon dioxide gas. 6. Solid sodium acetate is added to water. Page 7 of 13 II. Oxidation-Reduction Reactions Oxidation Numbers Rules 1. Free elements have oxidation states of 0 2. Ions keep their charges 3. Oxygen in a compound is –2 unless as peroxide (then it is –1) 4. Fluorine is –1 5. Hydrogen is +1 unless as a hydride (then it is –1) 6. Sum of oxidation states equals charge of substance. If it is a compound than it is zero. Oxidation State Practice Definitions: Oxidation - Reduction - Oxidizing Agent - Reducing Agent - Examples: 2Pb + 3O2 2PbO + 2SO2 Cl2 + OH- ClO- + Cl- + H2O Page 8 of 13 Balancing Redox Reactions Rules—Acid solution 1. Write reactions as ½ reactions—One for the oxidation and one for the reduction 2. Balance all elements except H and O 3. Balance H with H+ 4. Balance O with H2O 5. Balance Charge with e6. Multiply reactions by factors such that the e- cancel 7. Add both ½ reactions Examples MnO4- +_ Fe2+ Fe3+ + Mn2+ SO42- + Cl- Cl2 + SO32- Page 9 of 13 (acid) Rules—Base solution 1. Write reactions as ½ reactions—One for the oxidation and one for the reduction 2. Balance all elements except H and O 3. Balance H with H+ 4. Balance O with H2O 5. Add OH- to both sides to cancel the H+ 6. Cancel out any extra water and OH7. Balance Charge with e8. Multiply reactions by factors such that the e- cancel Add both ½ reactions Al + MnO4- MnO2 + Al(OH)4- Cl2 + MnO4- MnO2 + ClO- Page 10 of 13 Reactions when Oxidation Numbers do change I. Rules for Oxidation Numbers The oxidation number of any free element (an element not combined chemically with a different element) is zero, regardless of how complex its molecules might be. The oxidation number for any simple, monoatomic ion is equal to the charge on the ion. The sum of all the oxidation numbers of the atoms in a molecule or ion must be equal the charge of the particle. In its compound, fluorine has an oxidation of -1 In its compounds, hydrogen has an oxidation number of +, except hen combined with metals when it is -1 In its compounds, oxygen has an oxidation number of -2, except in forming peroxides hen it is -1. All Group I metals in compounds have a charge of +1 All Group II metals in compounds have a charge of +2 All Halogens have a charge of -1 when combined with metals in binary compounds II. Balancing Redox Equations A. Oxidation States Method (we use the half-reaction method) 1. Assign oxidation numbers to atoms in equation 2. Identify the substance oxidized and determine the number of electrons lost 3. Balance the electrons gained 4. balance the non redox substances by inspection B. Half-Reaction (Acidic Solution) 1. Divide the skeleton equation into half-reactions 2. Balance atoms other than H & O 3. Balance oxygen by adding H2O to the side that needs O 4. Balance hydrogen by adding H+ to the side that needs H 5. Balance the charge by adding electrons 6. Make the number of electrons gained equal to the number lost and then add the two halfreactions 7. Cancel anything that is the same on both sides. C. Half-Reactions Method (Basic Solution) 1. 1-7 Same as B above 2. Add to both sides of the equation the same number of OH- as there are H+ 3. Combine H+ and OH- to form H2O 4. Cancel any H2O that you can Page 11 of 13 III. Types of Redox Reactions A. Simple Redox 1. Hydrogen Displacement Ca(s) + 2H2O(l) Ca(OH)2(s) + H2 2. Metal Displacement Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s) 3. Halogen Displacement Cl2(g) + KBr(aq) 2KCl(aq) + Br2(l) 4. Combustion CH4(g) + 2O2(g) CO2(g) + 2H2O(g) B. Disportionation This is where one substance both oxidizes and reduces Cl2(g) + 2OH-(aq) OCl-(aq) + Cl-(aq) + H2O(l) C. Reactions involving oxoanions such as Cr2O7214H+(aq) + Cr2O72- + 6 Fe2+ 2 Cr3+ + 7 H2O + 6 Fe3+ Prediction Practice with Redox 1. solid copper is added to a dilute nitric acid solution. 2. a solution of potassium permanganate is mixed with an alkaline solution of sodium sulfite. 3. ethanol is completely burned in air. 4. sodium metal is added to water. 5. hydrogen peroxide solution is added to a solution of iron(II) sulfate. Page 12 of 13 IV. Redox Reaction Prediction Important Oxidizers MnO4- (acid solution) MnO4- (basic solution) MnO2 (acid solution) Cr2O7-2 (acid) CrO4-2 HNO3, conc HNO3, dilute H2SO4, hot conc Metallic Ions Free Halogens (F2, Cl2, Br2, I2) HClO4 Na2O2 H2O2 Formed in reaction Mn+2 MnO2 Mn+2 Cr+3 Cr+3 NO2 NO SO2 Metallous Ions Halide ions (F-, Cl-, Br-, I-) ClOHO2 Important Reducers Halide Ions Free Metals Metalous Ions Nitrite Ions Sulfite Ions Free Halogens (dil, basic, sol) Free Halogens (conc, basic sol) C2O42- Formed in Reaction Halogens Metal Ions Metallic ions Nitrate Ions SO42Hypohalite ions Halate ions CO2 Redox reactions involve the transfer of electrons. The oxidation numbers of at least two elements must change. Single replacement, some combination and some decomposition reactions are redox reactions. 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. When a problem mentions an acidic or basic solution, it is probably is redox. Page 13 of 13