Chapter 4 Types of Chemical Reactions
Classifying Chemical Reactions by What Atoms Do
Classification of Reactions + Synthesis reaction + Decomposition reaction + + Single displacement reaction + + Double displacement reaction
Classification of Reactions 2 H2 (g) + O2 (g) ---------> 2 H2O (g) Synthesis reaction CaCO3 (s) ---------> CaO (s) + CO2 (g) Decomposition reaction 2 Al (s) + Fe2O3 (s) ---------> Al2O3 (s) + 2 Fe (l) Single displacement reaction Ba(NO3)2 (aq) + Na2SO4 (aq) ---------> BaSO4 (s) + 2 NaNO3 (aq) Double displacement reaction
Chemical Reactions Classified by Reaction Type
Precipitation Reactions
Precipitation Reactions Precipitation reactions are reactions in which a solid forms when we mix two solutions. 1) reactions between aqueous solutions of ionic compounds 2) produce an ionic compound that is insoluble in water 3) The insoluble product is called a precipitate.
Precipitation Reactions 2 KI(aq) + Pb(NO3)2(aq) PbI2(s) + 2 KNO3(aq)
No Precipitate Formation = No Reaction KI(aq) + NaCl(aq) KCl(aq) + NaI(aq) No precipitate forms, therefore, no reaction. KI(aq) KCl(aq) + NaI(aq) NaCl(aq)
Process for Predicting the Products of a Precipitation Reaction 1. Determine which ions are present in each aqueous reactant. 2. Determine formulas of possible products. 3. Determine solubility of each potential product in water. 4. If neither product will precipitate, write no reaction after the arrow. 5. If any of the possible products are insoluble, write their formulas as the products of the reaction using (s) after the formula to indicate solid. Write any soluble products with (aq) after the formula to indicate aqueous. 6. Balance the equation. Remember to only change coefficients, not subscripts
Solubility Rules Compounds Containing the Following Ions Are Mostly Soluble Li +, Na +, K +, NH4 + NO3 -, C2H3O2 -, ClO4 - Exceptions None None Cl -, Br -, I - Ag +, Hg2 2+, Cu +, Pb 2+ SO4 2 - Ca 2+, Sr 2+, Ba 2+,Pb 2+ Compounds Containing the Following Ions Are Mostly Insoluble Exceptions OH - Group I cations, Ca 2+, Sr 2+, Ba 2+ S 2 - Group I cations, Ca 2+, Sr 2+, Ba 2+, NH4 + CO3 2 -, PO 4 3 - Group I cations, NH4 +
Practice Predict the products and balance the equation K2CO3(aq) + NiCl2(aq) K2CO3(aq) + NiCl2(aq) KCl (?) + NiCO3(?) K2CO3(aq) + NiCl2(aq) 2 KCl (?) + NiCO3(?) K2CO3(aq) + NiCl2(aq) 2 KCl (aq) + NiCO3(s)
Practice Predict the products and balance the equation KCl(aq) + AgNO3(aq) KCl(aq) + AgNO3(aq) KNO3(?) + AgCl(?) KCl(aq) + AgNO3(aq) KNO3(aq) + AgCl(s)
Practice Predict the products and balance the equation Na2S(aq) + CaCl2(aq) Na2S(aq) + CaCl2(aq) NaCl(?) + CaS(?) Na2S(aq) + CaCl2(aq) 2 NaCl(?) + CaS(?) Na2S(aq) + CaCl2(aq) 2 NaCl(aq) + CaS(aq) No Reaction!!!!!
Practice Predict the products and balance the equation (NH4)2SO4(aq) + Pb(C2H3O2)2(aq) (NH4)2SO4(aq) + Pb(C2H3O2)2(aq) NH4C2H3O2(?) + PbSO4(?) (NH4)2SO4(aq) + Pb(C2H3O2)2(aq) 2 NH4C2H3O2(?) + PbSO4(?) (NH4)2SO4(aq) + Pb(C2H3O2)2(aq) 2 NH4C2H3O2(aq) + PbSO4(s)
Ionic Equations Equations that describe the chemicals put into the water and the product molecules are called molecular equations. 2 KOH(aq) + Mg(NO3)2(aq) 2 KNO3(aq) + Mg(OH)2(s) Equations that describe the material s structure when dissolved are called complete ionic equations. Aqueous strong electrolytes are written as ions. Insoluble substances, weak electrolytes, and nonelectrolytes are written as molecules. 2K + (aq) + 2OH (aq) + Mg 2+ (aq) + 2NO3 (aq) 2K + (aq) + 2NO3 (aq) + Mg(OH)2(s)
Ionic Equations Equations that describe the chemicals put into the water and the product molecules are called molecular equations. 2 KOH(aq) + Mg(NO3)2(aq) 2 KNO3(aq) + Mg(OH)2(s) Equations that describe the material s structure when dissolved are called complete ionic equations. Aqueous strong electrolytes are written as ions. Insoluble substances, weak electrolytes, and nonelectrolytes are written as molecules. 2K + (aq) + 2OH (aq) + Mg 2+ (aq) + 2NO3 (aq) 2K + (aq) + 2NO3 (aq) + Mg(OH)2(s)
Ionic Equations Ions that are both reactants and products are called spectator ions. 2 K + (aq) + 2 OH (aq) + Mg 2+ (aq) + 2 NO3 (aq) 2 K + (aq) + 2 NO3 (aq) + Mg(OH)2(s) An ionic equation in which the spectator ions are removed is called a net ionic equation. 2 OH (aq) + Mg 2+ (aq) Mg(OH)2(s)
Write the ionic and net ionic equation K2SO4(aq) + 2 AgNO3(aq) 2 KNO3(aq) + Ag2SO4(s) 2K + (aq) + SO4 2- (aq) + 2Ag + (aq) + 2NO3 - (aq) 2K + (aq) + 2NO3 - (aq) + Ag2SO4(s) 2 Ag + (aq) + SO4 2 (aq) Ag2SO4(s)
Write the ionic and net ionic equation Na2CO3(aq) + 2 HCl(aq) 2 NaCl(aq) + CO2(g) + H2O(l) 2Na + (aq) + CO3 2- (aq) + 2H + (aq) + 2Cl - (aq) 2Na + (aq) + 2Cl - (aq) + CO2(g) + H2O(l) CO3 2 (aq) + 2 H + (aq) CO2(g) + H2O(l)
Chapter 4 Acid/Base Reactions
Acids and Bases in Solution Acids ionize in water to form H + ions. (More precisely, the H+ from the acid molecule is donated to a water molecule to form hydronium ion, H3O + ) Bases dissociate in water to form OH- ions. (Bases, such as NH3, that do not contain OH- ions, produce OH- by pulling H + off water molecules.) In the reaction of an acid with a base, the H + from the acid combines with the OH- from the base to make water. The cation from the base combines with the anion from the acid to make a salt. acid + base salt + water
Molecular Models of Selected Acids
Acid-Base Reactions Also called neutralization reactions because the acid and base neutralize each other s properties 2 HNO3(aq) + Ca(OH)2(aq) Ca(NO3)2(aq) + 2 H2O(l) Note that the cation from the base combines with the anion from the acid to make the water soluble salt. The net ionic equation for an acid-base reaction is H + (aq) + OH-(aq) H2O(l) (as long as the salt that forms is soluble in water)
Common Acids Name Formula Uses Strength Perchloric HClO4 explosives, catalysts Strong Nitric HNO3 explosives, fertilizers, dyes, glues Strong Sulfuric H2SO4 explosives, fertilizers, dyes, glue, batteries Strong Hydrochloric HCl metal cleaning, food prep, ore refining, stomach acid Strong Phosphoric H3PO4 fertilizers, plastics, food preservation Moderate Chloric HClO3 explosives Moderate Acetic HC2H3O2 plastics, food preservation, vinegar Weak Hydrofluoric HF metal cleaning, glass etching Weak Carbonic H2CO3 soda water, blood buffer Weak Hypochlorous HClO sanitizer Weak Boric H3BO3 eye wash Weak
Common Bases Name Formula Common Name Uses Strength Sodium Hydroxide NaOH Lye, Caustic Soda soap, plastic production, petroleum refining Strong Potassium Hydroxide KOH Caustic Potash soap, cotton processing, electroplating Strong Calcium Hydroxide Ca(OH)2 Slaked Lime cement Strong Sodium Bicarbonate NaHCO3 Baking Soda food preparation, antacids Weak Magnesium Hydroxide Mg(OH)2 Milk of Magnesia antacids Weak Ammonium Hydroxide NH4OH Ammonia Water fertilizers, detergents, explosives Weak
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) HCl(aq) NaOH(aq) NaCl(aq) + H2O(l)
Write the molecular, ionic, and net-ionic equation for the acid-base reaction HNO3(aq) + Ca(OH)2(aq) HNO3(aq) + Ca(OH)2(aq) Ca(NO3)2(aq) + H2O(l) 2HNO3(aq) + Ca(OH)2(aq) Ca(NO3)2(aq) + 2H2O(l) 2H + (aq) + 2NO3 - (aq) + Ca 2+ (aq) + 2OH - (aq) Ca 2+ (aq) + 2NO3 - (aq) + 2H2O(l) 2H + (aq) + 2OH - (aq) 2H2O(l)
Write the molecular, ionic, and net-ionic equation for the acid-base reaction HCl(aq) + Ba(OH)2(aq) HCl(aq) + Ba(OH)2(aq) BaCl2(aq) + H2O(l) 2HCl(aq) + Ba(OH)2(aq) BaCl2(aq) + 2H2O(l) 2H + (aq) + 2Cl - (aq) + Ba 2+ (aq) + 2OH - (aq) Ba 2+ (aq) + 2Cl - (aq) + 2H2O(l) 2H + (aq) + 2OH - (aq) 2H2O(l)
Write the molecular, ionic, and net-ionic equation for the acid-base reaction H2SO4(aq) + Sr(OH)2(aq) H2SO4(aq) + Sr(OH)2(aq) SrSO4(s) + 2 H2O(l) 2H + (aq) + SO4 2- (aq) + Sr 2+ (aq) + 2OH - (aq) SrSO4 (s) + 2H2O(l) 2H + (aq) + SO4 2- (aq) + Sr 2+ (aq) + 2OH-(aq) SrSO4 (s) + 2H2O(l)
Chapter 4 Acid/Base Titrations
Titration A solution s concentration is determined by reacting it with another solution and using stoichiometry this process is called titration. In the titration, the unknown solution is added to a known amount of another reactant until the reaction is just completed. At this point, called the endpoint, the reactants are in their stoichiometric ratio. The unknown solution is added slowly from an instrument called a burette.
Acid-Base Titrations The difficulty is determining when there has been just enough of one solution (the titrant) added to complete the reaction. In acid-base titrations, because both the reactant and product solutions are colorless, a chemical (indicator) is added that changes color when the solution undergoes large changes in acidity/alkalinity. At the endpoint of an acid-base titration, the number of moles of H + equals the number of moles of OH-. This is the equivalence point.
Titration The titrant is the base solution in the burette. As the titrant is added to the flask, the H + reacts with the OH to form water. But there is still excess acid present so the color does not change. At the titration s endpoint, just enough base has been added to neutralize all the acid. At this point the indicator changes color.
Titration
The titration of 10.00 ml of HCl solution of unknown concentration requires 12.54 ml of 0.100 M NaOH solution to reach the end point. What is the concentration of the unknown HCl solution? HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) ml of NaOH soln L of NaOH soln mol of NaOH mol of HCl L L mol ml mol mol ml of HCl soln L of HCl soln Molarity = mol of HCl L of HCl soln
The titration of 10.00 ml of HCl solution of unknown concentration requires 12.54 ml of 0.100 M NaOH solution to reach the end point. What is the concentration of the unknown HCl solution? HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) 0.001 L NaOH soln 12.54 ml NaOH solution x x 1.000 ml NaOH soln 0.100 mol NaOH 1.000 L NaOH soln x 1.00 mol HCl = 1.25 x 10-3 mol HCl in the sample 1.00 mol NaOH 10.00 ml HCl solution x 0.001 L HCl soln = 0.0100 L HCl soln 1.000 ml HCl soln 1.25 x 10 Molarity of HCl solution = -3 mol HCl = 0.0100 L HCl soln 0.125 M
What is the concentration of NaOH solution that requires 27.5 ml to titrate 50.0 ml of 0.1015 M H2SO4? H2SO4 (aq) + 2 NaOH (aq) Na2SO4 (aq) + 2 H2O (l) ml of H2SO4 soln L of H2SO4 soln mol of H2SO4 mol of NaOH L L mol m mol mol mol of NaOH ml of NaOH soln L of NaOH soln Molarity = L of NaOH soln
What is the concentration of NaOH solution that requires 27.50 ml to titrate 50.00 ml of 0.1015 M H2SO4? H2SO4 (aq) + 2 NaOH (aq) Na2SO4 (aq) + 2 H2O (l) 50.00 ml H2SO4 soln x 0.001 L H2SO4 soln x 0.1015 mol H2SO4 1.000 ml H2SO4 soln 1.000 L H2SO4 soln x 2.00 mol NaOH = 0.1015 mol NaOH in the sample 1.00 mol H2SO4 27.50 ml NaOH soln x 0.001 L NaOH soln = 0.02750 L NaOH soln 1.000 ml NaOH soln Molarity of NaOH soln = 0.1015 mol NaOH = 0.02750 L NaOH soln 0.3691 M
Chapter 4 Gas-Evolving Reactions
Gas-Evolving Reactions Some reactions form a gas directly from the ion exchange: K2S(aq) + H2SO4(aq) K2SO4(aq) + H2S(g) Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water. K2SO3(aq) + H2SO4(aq) K2SO4(aq) + H2SO3(aq) H2SO3 H2O(l) + SO2(g)
NaHCO3(aq) + HCl(aq) NaCl(aq) + CO2(g) + H2O(l) NaHCO3(aq) NaCl(aq) + CO2(g) + H2O(l) HCl(aq)
Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water. NaHCO3(aq) + HCl(aq) NaCl(aq) + H2CO3(aq) H2CO3 H2O(l) + CO2(g)
Compounds that Undergo Gas-Evolving Reactions Reactant Reactant Exchange Product Gas Formed Example Metal sulfide Metal hydrogensulfide Acid H2S H2S K2S (g) + HCl (aq) H2S (g) + KCl (aq) Metal carbonate Metal hydrogencarbonate Acid H2CO3 CO2 K2CO3 (aq) + HCl (aq) CO2 (g) + H2O (l) + KCl (aq) Metal sulfite Metal hydrogensulfite Acid H2SO3 SO2 K2SO3 (aq) + HCl (aq) SO2 (g) + H2O (l) + KCl (aq) Ammonium salt Base NH4OH NH3 KOH (aq) + NH4Cl (aq) NH3 (g) + H2O (l) + KCl (aq)
Practice Predict the products and balance the equations H2SO4(aq) + CaS(aq) CaSO4(aq) + H2S(aq) Na2CO3(aq) + 2 HNO3(aq) 2 NaNO3(aq) + H2CO3 2 NaNO3(aq) + H2O (l) + CO2(g) 2 HCl(aq) + Na2SO3(aq) 2 NaCl(aq) + H2SO3 2 NaCl (aq) + H2O (l) + SO2 (g)
Chapter 4 Redox Reactions
Oxidation/Reduction Basic Definitions e- X loses electrons Y gains electron X is oxidized Y is reduced X is the reducing agent Y is the oxidizing agent X increases its oxidation number Y decreases its oxidation number
Oxidation and Reduction - Symbolic Representation
Oxidation and Reduction at the Atomic Level
Redox Reactions Oxidation/reduction reactions involve transferring electrons from one atom to another. Also known as redox reactions Many involve the reaction of a substance with O2(g). 4 Fe(s) + 3 O2(g) 2 Fe2O3(s) Atoms in Elements-------> Ions in Compound
Combustion as Redox 2 H2(g) + O2(g) 2 H2O(g)
Redox without Combustion 2 Na(s) + Cl2(g) 2 NaCl(s)
Reactions of Metals with Nonmetals Consider the following reactions: 4 Na(s) + O2(g) 2 Na2O(s) 2 Na(s) + Cl2(g) 2 NaCl(s) The reactions involve a metal reacting with a nonmetal. In addition, both reactions involve the conversion of free elements into ions. Na2O = 2 Na + + O 2- NaCl = Na + + Cl -
Oxidation and Reduction To convert a free element into an ion, the atoms must gain or lose electrons (of course, if one atom loses electrons, another must accept them). Atoms that lose electrons are being oxidized, atoms that gain electrons are being reduced. 2 Na(s) + Cl2(g) 2 Na + Cl (s) Na Na + + 1 e oxidation Cl2 + 2 e 2 Cl reduction
Chapter 4 Redox Reactions Electron Bookkeeping
Electron Bookkeeping For reactions that are not metal + nonmetal, or do not involve O2, we need a method for determining how the electrons are transferred. Chemists assign a number to each element in a reaction called an oxidation state that allows them to determine the electron flow in the reaction. Even though they look like them, oxidation states are not ion charges! Oxidation states are imaginary charges assigned based on a set of rules. Ion charges are real, measurable charges.
Rules for Assigning Oxidation States (in order of priority) 1. Free elements have an oxidation state = 0. In Na (s), Na = 0 ; In Cl2 (g), Cl2 = 0 2. Monatomic ions have an oxidation state equal to their charge. In NaCl, Na = +1 and Cl = 1 3. (a) The sum of the oxidation states of all the atoms in a compound is 0. Na = +1 and Cl = 1 in NaCl, (+1) + ( 1) = 0
Rules for Assigning Oxidation States (in order of priority) 3. (b) The sum of the oxidation states of all the atoms in a polyatomic ion equals the charge on the ion. In NO3, N = +5 and O = 2 [3 x (2-) + 1 x (5+) = -1] 4. (a) Group I metals have an oxidation state of +1 in all their compounds. (b) Group II metals have an oxidation state of +2 in all their compounds.
Rules for Assigning Oxidation States (in order of priority) 5. In their compounds, nonmetals have oxidation states according to the table below
Assign an oxidation state to each element in the following Br2 K + LiF CO2 SO4 2 Na2O2 Br = 0, (Rule 1) K = +1, (Rule 2) Li = +1, (Rule 4a) & F = 1, (Rule 5) O = 2, (Rule 5) & C = +4, (Rule 3a) O = 2, (Rule 5) & S = +6, (Rule 3b) Na = +1, (Rule 4a) & O = 1, (Rule 3a)
Determine the oxidation states of all the atoms in a propanoate polyatomic anion, C3H5O2 There are no free elements or free ions in propanoate, so the first rule that applies is Rule 3b (C3) + (H5) + (O2) = 1 Because all the atoms are nonmetals, the next rule we use is Rule 5, following the elements in order: H = +1 O = 2 (C3) + 5(+1) + 2( 2) = 1 (C3) = 2 C = ⅔
Oxidation and Reduction Another Definition Oxidation occurs when an atom s oxidation state increases during a reaction. Reduction occurs when an atom s oxidation state decreases during a reaction. -4 0 +4-2 CH4 + 2 O2 CO2 + 2 H2O oxidation reduction
Oxidation Reduction Oxidation and reduction must occur simultaneously. If an atom loses electrons another atom must take them. The reactant that reduces an element in another reactant is called the reducing agent. The reactant that oxidizes an element in another reactant is called the oxidizing agent. 2 Na(s) + Cl2(g) 2 Na + Cl (s) Na is oxidized Cl is reduced Na is the reducing agent Cl2 is the oxidizing agent
Assign oxidation states, determine the element oxidized and reduced, and determine the oxidizing agent and reducing agent in the following reactions: 0 Sn 4+ + Ca Sn 2+ + Ca 2+ Sn 4+ is being reduced; Sn 4+ is the oxidizing agent. Ca is being oxidized; Ca is the reducing agent. 0 0 S +4 2 F2 + S SF4 F is being reduced from F 0 to F-;F2 is the oxidizing agent. S is being oxidized from S 0 to S +4 ;S is the reducing agent. F -
Assign oxidation states, determine the element oxidized and reduced, and determine the oxidizing agent and reducing agent in the following reactions: 0 +7 +3 +4 Fe + MnO4 + 4 H + Fe 3+ + MnO2 + 2 H2O oxidation reduction Fe is the reducing agent. MnO4 is the oxidizing agent.