Chapter 4: Reaction in Aqueous Solution

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1 Chapter 4: Reaction in Aqueous Solution What is a Solution? Solute + substance dissolved typically smaller quantity Solvent! dissolving medium typically larger quantity Solution homogeneous mixture variable composition When water is used as the solvent, the solution is an aqueous solution Speciation in Solutions: Ionic Compounds Soluble Ionic Compounds - result in aqueous solutions with individually dispersed ions must consider stoichiometry as well: AgNO3 (s)! Ag+ (aq) + NO3 (aq) vs. Ba(NO3)2 (s)! Ba2+ (aq) + 2 NO3 (aq) Speciation in Solutions: Molecular Compounds Soluble Molecular Compounds - result in aqueous solutions with individually dispersed, neutral molecules.

2 Solution Composition: How much solute is present? qualitatively is a solution dilute or concentrated? quantitatively molar concentration, or molarity (M) mol solute molarity = volume solution units of molarity are mol/l Solution Calculations 1. Calculate the molarity of a solution. OR Calculate the molarity of ions in a solution. 2. Calculation to determine how to prepare a solution of known concentration. 3. Solution reaction stoichiometry calculations; including identifying limiting reactants. 4. Dilution calculations. Solution Calculations Calculate the molarity of a solution or of ions in a solution. ex g CoCl3 6H2O is dissolved in water resulting in ml of solution. Determine the molar concentration of this solution, and the molar concentration of chloride ions in this solution. [CoCl3] =???? [Cl ] =???? Solution Calculations Calculation to determine how to prepare a solution of known concentration. ex. Determine the mass (in g) of NaNO3 required to prepare ml of M NaNO3 (aq). ex. Determine the mass (in g ) of solute in 75.0 ml of 88.8 mm NiSO4 (aq).

3 Process of Preparing a Solution of Specific Concentration Solution Calculations Solution reaction stoichiometry calculations: ex. Determine the volume (in ml) of M Ba(OH)2 (aq) required to react completely with ml of 1.07 M HCl (aq): 2 HCl (aq) + Ba(OH)2 (aq)! BaCl2 (aq) + 2 H2O (l) ex ml of M AgNO3 (aq) and ml of M BaCl2 (aq) are combined and allowed to react accoding to the following equation: 2 AgNO3 (aq) + BaCl2 (aq)! 2 AgCl (s) + Ba(NO3)2 (aq) Determine the theoretical yield (in g) of AgCl (s). Dilution Calculations Dilution is the procedure of making a solution of lower concentration (i.e. more dilute) from a solution of higher concentration (i.e. more concentrated). mol mol solute M = = volume solution V so: mol solute = M! V How do each of the following change from Solution 1 (concentrated) to Solution 2 (dilute)?! Solution 1 concentrated Dilution Calculations Solution 2 dilute mol solute volume solution (V) molarity (M)

4 Dilution Calculations Aqueous Solutions and Electrolyte Classifications ex ml of 12.0 M HCl (aq) is transferred to a flask and an additional ml of water is added. Determine the molar concentration of the new, dilute HCl (aq). Classification of solutes and resulting solutions as: strong electrolytes weak electrolytes nonelectrolytes based on the ability of the solution to conduct electricity. ex. How would you prepare ml of M NH3 (aq) from a stock solution with [NH3] = 4.85 M? strong electrolytes are completely ionized in aqueous sol n solutions of strong electrolytes conduct electricity The extent to which a solution can conduct electricity is directly dependent on the number of charged particles (ions) present weak electrolytes are only partially ionized in aqueous sol n solutions of weak electrolytes conduct electricity only a little

5 Identification of Electrolytes compounds that are strong electrolytes: soluble ionic compounds strong acids and strong bases compounds that are weak electrolytes: weak acids and weak bases compounds that are nonelectrolytes: nonelectrolytes are not ionized in aqueous sol n solutions of nonelectrolytes do not conduct electricity molecular compounds (even if soluble) insoluble ionic compounds Precipitation Reactions: Formation of an Insoluble Solid Consider the reaction that occurs between aqueous solutions of NaCl & AgNO3: NaCl (aq) + AgNO3 (aq) AgCl (s) + NaNO3 (aq) Precipitation Reactions: Predicting Products If aqueous solutions of soluble ionic compounds MX and NZ are mixed, what new ionic compounds can form as products? in general: MX (aq) + NZ (aq)!??? +??? need to determine: chemical formulas of products physical states... Are these ionic compounds soluble or insoluble?

6 Solubility Rules Predict the Products: Are these precipitation reactions? ex: Pb(NO3)2 (aq) + KI (aq)!???? +???? ex: (NH4)2S (aq) + CuSO4 (aq)!???? +???? ex: K2CO3 (aq) + Ca(OH)2 (aq)!???? +???? a few notes and additions: rule 1: applies to all Group IA cations & NH4 +, so include Rb + and Cs + rule 2: applies to chlorates (ClO3 ) and perchlorates (ClO4 ) too rule 3: F salts tend to be less soluble than Cl, Br and I to rules 5, 6 & 7: applies to oxalates (C2O4 2 ) and chromates (CrO4 2 ) too ex: NH4NO3 (aq) + MgCl2 (aq)!???? +???? Writing Ionic Equations Balanced Chemical Equation focus: the big picture identity of reactants and products physical states stoichiometry Full (or Complete, or Total) Ionic Equation focus: complete speciation identity of all species present in solution Net Ionic Equation focus: actual chemical change do not include spectator ions in net ionic equation Writing Ionic Equations an example: Pb(NO3)2 (aq) + KI (aq)!???? +???? balanced chemical equation: Pb(NO3)2 (aq) + 2 KI (aq)! PbI2 (s) + 2 KNO3 (aq) full ionic equation: Pb 2+ (aq) + 2 NO3 (aq) + 2 K + (aq) + 2 I (aq)! PbI2 (s) + 2 K + (aq) + 2 NO3 (aq) net ionic equation: Pb 2+ (aq) + 2 I (aq)! PbI2 (s) spectator ions: K + and NO3 notes: all equations are balanced in terms of atoms & charge ions shown with charges polyatomic ions do not split up further all physical states included

7 Writing Ionic Equations What do you split up into ions when writing ionic equations? only strong electrolytes in (aq) soluble ionic cmpd s strong acid strong bases What do you not split up when writing ionic equations? anything that is (s), (l), (g) weak electrolytes and nonelectrolytes - even if in (aq) Acids and Bases Arrhenius Definitions: acids - compounds that produce an increase in [H+] when dissolved in water Acids and Bases bases - compounds that produce an increase in [OH ] when dissolved in water Brønsted-Lowry Definitions: acids - H+ donors bases - H+ acceptors Lewis Definitions: acids - electron pair acceptors bases - electron pair donors Acids and Bases Brønsted-Lowry acid-base behavior: acid ionization equation: HNO2 (aq) + H2O (l)! NO2 (aq) + H3O+ (aq) notes: HNO2 - acid; H2O - base + H3O (hydronium ion) and H+ used interchangeably Water is amphoteric; it can act as an acid or a base. base ionization equation: NH3 (aq) + H2O (l)! NH4+ (aq) + OH (aq) notes: H2O - acid; NH3 - base

8 Acids and Bases Identify the Brønsted-Lowry acid and base in the following reaction: H2SO4 (aq) + HPO4 2 (aq)! HSO4 (aq) + H2PO4 (aq) Strong vs. Weak Acids and Bases Acid and base strength is based on the extent of ionization that occurs when the substance is dissolved in water. Strong Acids: strong electrolytes - completely ionized in solution there are 6 strong acids - KNOW THEM! HCl, HBr, HI, HNO3, HClO4, H2SO4 (diprotic) Weak Acids: weak electrolytes - partially ionized (typically < 5%) in aqueous solution any acid that is not a strong acid is a weak acid some examples: HF, H2CO3, H3PO4, HNO2, HBrO4 Strong vs. Weak Acids and Bases Strong Bases: strong electrolytes - completely ionized in solution the strong bases are the hydroxides of the alkali metals & hydroxides of most alkaline earth metals: Weak Bases: LiOH, NaOH, KOH, RbOH, CsOH Ca(OH)2, Sr(OH)2, Ba(OH)2 weak electrolytes - partially ionized (typically < 5%) in aqueous solution weak bases tend to be organic compounds that contain nitrogen; ammonia and substituted amines some examples: NH3, (CH3)NH2, (CH3)3N C5H5N, N2H4, NH2OH

9 general form: Acid - Base Neutralization Reactions: Acid + Base! Salt* + Water * salt will have cation related to the base + anion related to the acid example: Write the balanced chemical equation for the neutralization reaction that will occur between aqueous solutions of nitrous acid and calcium hydroxide. Acid - Base Neutralization Reactions: Writing Ionic Equations Strong Acid + Strong Base: balanced chemical equation: HNO3 (aq) + KOH (aq)! KNO3 (aq) + H2O (l) full ionic equation: H + (aq) + NO3 (aq) + K + (aq) + OH (aq)! K + (aq) + NO3 (aq) + H2O (l) net ionic equation: H + (aq) + OH (aq)! H2O (l) spectator ions: K + and NO3 Acid - Base Neutralization Reactions: Writing Ionic Equations Weak Acid + Strong Base: balanced chemical equation: HC2H3O2 (aq) + CsOH (aq)! CsC2H3O2 (aq) + H2O (l) full ionic equation: HC2H3O2 (aq) + Cs + (aq) + OH (aq)! Cs + (aq) + C2H3O2 (aq) + H2O (l) net ionic equation: HC2H3O2 (aq) + OH (aq)! C2H3O2 (aq) + H2O (l) spectator ion: only Cs + Acid - Base Neutralization Reactions: Writing Ionic Equations Strong Acid + Weak Base: balanced chemical equation: HCl (aq) + NH3 (aq)! NH4Cl (aq) full ionic equation: H + (aq) + Cl (aq) + NH3 (aq)! NH4 + (aq) + Cl (aq) net ionic equation: H + (aq) + NH3 (aq)! NH4 + (aq) spectator ions: only Cl

10 Other Examples of Acid-Base Chemistry: Acidic and Basic Oxides in general - oxides of metals are basic basic oxides or base anhydrides metal oxides form basic solutions when dissolved in water: CaO (s) + H2O (l)! Ca(OH)2 (aq) metal oxides react with acids: CaO (s) + 2 HCl (aq)! CaCl2 (aq) + H2O (l) Other Examples of Acid-Base Chemistry: Acidic and Basic Oxides in general - oxides of nonmetals are acidic acidic oxides or acid anhydrides nonmetal oxides form acidic solutions when dissolved in water: SO2 (g) + H2O (l)! H2SO3 (aq) nonmetal oxides react with bases: SO2 (g) + 2 NaOH (aq)! Na2SO3 (aq) + H2O (l) Other Examples of Acid-Base Chemistry: Neutralization Reactions that Produce Gases general form of reaction: acid + base! salt + water + gas carbonates & hydrogen carbonates: gas formed is CO2 sulfates & hydrogen sulfates: gas formed is SO3 sulfites & hydrogen sulfites: gas formed is SO2 sulfides: gas formed is H2S Other Examples of Acid-Base Chemistry: Neutralization Reactions that Produce Gases some examples: Na2CO3 (s) + 2 HCl (aq)! 2 NaCl (aq) + H2O (l) + CO2 (g) full ionic equation: Na2CO3 (s) + 2 H + (aq) + 2 Cl (aq)! 2 Na + (aq) + 2 Cl (aq) + H2O (l) + CO2 (g) net ionic equation: Na2CO3 (s) + 2 H + (aq)! 2 Na + (aq) + H2O (l) + CO2 (g) spectator ion: Cl (aq) MgSO4 (s) + 2 HCl (aq)! MgCl2 (aq) + H2O (l) + SO3 (g) Ba(HSO3)2 (aq) + 2 HBr (aq)! BaBr2 (aq) + 2 H2O (l) + 2 SO2 (g) K2S (s) + 2 HNO3 (aq)! 2 KNO3 (aq) + H2S (g)

11 The ph Scale logarithmic scale of [H + ] in solution ph = "log[h + ]; [H + ] = 10 ph Relationship Between [H + ] and ph as [H + ] changes by a factor of 10, the ph of the solution changes by 1 unit higher [H + ] corresponds to lower ph higher [H + ] corresponds to more acidic solution recall: ph Calculations: Relative Acidity and Basicity of Solutions in any aqueous solution at 25 C: [H + ][OH ] = 1 x ph = log[h + ]; [H + ] = 10 ph higher [H + ] " more acidic solution " lower ph higher [OH ] " lower [H + ] " more basic solution " higher ph ph Calculations: Relative Acidity and Basicity of Solutions ex: Calculate the ph of M HNO3 (aq). ex: Will the ph of M HNO2 (aq) be less than, greater than, or equal to the ph of M HNO3 (aq)? Why? ex: Calculate the [H + ] in a solution with ph = ex: Calculate the ph of 0.20 M Ba(OH)2 (aq) and 0.20 M NaOH (aq). Should they be the same? Why or why not?

12 Oxidation-Reduction Reactions aka Redox Reactions oxidation: loss of electrons reduction: gain of electrons Consider the reaction of calcium metal with oxygen: 2 Ca (s) + O2 (g)! 2 CaO (s) What happens to each element during the course of this reaction? Oxidation and Reduction Half Reactions for the reaction of calcium with oxygen: oxidation 1 2 reaction: Ca! Ca e note: in oxidation 1 2 rxn e s are products reduction 1 2 reaction: O2 + 4 e! 2 O 2 note: in reduciton 1 2 rxn e s are reactants net redox reaction: 2 Ca + O2! 2 Ca O 2 note: add together the half-reactions; multiply as necessary to have same number of electrons in each half reaction before adding Oxidation Numbers (Nox) An oxidation number indicates the amount of electropositive or electronegative character of an atom - particularly as part of a polyatomic species. Determining Oxidation Numbers: ex. Determine Nox of sulfur in SO3. Nox S + 3(Nox O) = 0 Nox S + 3( 2) = 0 Nox S = +6 ex. Determine Nox of sulfur in SO3 2. Nox S + 3(Nox O) = 2 Nox S + 3( 2) = 2 Nox S = +4 ex. Determine Nox N, H, P, and O in (NH4)3PO4. consider NH4 + and PO4 3 separately: Nox N + 4(Nox H) = +1 Nox P + 4(Nox O) = 3 Nox N + 4(+1) = +1 Nox P + 4( 2) = 3 Nox N = 3 Nox P = +5

13 Recognizing Oxidation and Reduction Compare oxidation numbers of elements in reactants and products: ex: if Nox increases - the element is oxidized if Nox decreases - the element is reduced 2 Al (s) + Cr2O3 (s)! Al2O3 (s) + 2 Cr (s) Recognizing Oxidation and Reduction Identify the oxidizing and reducing agents: oxidizing agent reactant that facilitates oxidation by taking on (i.e. gaining) electrons lost during oxidation process oxidizing agent is the reactant that contains the element that is reduced reducing agent reactant that facilitates reduction by providing (i.e. losing) electrons that are gained during reduction process reducing agent is the reactant that contains the element that is oxidized Recognizing Oxidation and Reduction Consider the following reaction: 2 Ca3(PO4)2 + 6 SiO C! P4 + 6 CaSiO CO Identify the following: element oxidized element reduced oxidizing agent reducing agent Balancing Redox Equations: the Half-Reaction Method in acidic solution 1. Identify the element oxidized and the element reduced. 2. Write the skeletal oxidation & reduction half-reactions. 3. Balance elements other than H and O. 4. Balance O s by adding H2O. 5. Balance H s by adding H Balance charge by adding e s. *at this point the individual half reactions are balanced* 7. Prepare to add the half-reactions together; multiply as necessary so that the # of e s in oxidation half reaction equals # of e s in the reduction half reaction. 8. Add half reactions together; clean up.

14 Balance the following redox reaction that occurs in acidic solution using the half-reaction method: I2 + NO3! IO3 + NO2 step 2: ox 1 2 rxn: I2! IO3 red 1 2 rxn: NO3! NO2 step 3: ox 1 2 rxn: I2! 2 IO3 red 1 2 rxn: NO3! NO2 step 4: ox 1 2 rxn: 6 H2O + I2! 2 IO3 red 1 2 rxn: NO3! NO2 + H2O step 5: ox 1 2 rxn: 6 H2O + I2! 2 IO H + red 1 2 rxn: 2 H + + NO3! NO2 + H2O Balance the following redox reaction that occurs in acidic solution using the half-reaction method (continued): I2 + NO3! IO3 + NO2 step 6: ox 1 2 rxn: 6 H2O + I2! 2 IO H e red 1 2 rxn: 1 e + 2 H + + NO3! NO2 + H2O step 7: ox 1 2 rxn: 6 H2O + I2! 2 IO H e red 1 2 rxn: 10!(1 e + 2 H + + NO3! NO2 + H2O) step 8: 6 H2O + I H NO3! 2 IO H NO H2O net rxn: I2 + 8 H NO3! 2 IO NO2 + 4 H2O Some Categories of Redox Reactions combination (or formation) reactions: elements compounds ex. 2 Al (s) + 3 Br2 (l)! 2 AlBr3 (s) decomposition reactions: compounds elements ex. 2 NaH (s)! 2 Na (s) + H2 (g) Some Categories of Redox Reactions combustion reactions: ex. CH4 (g) + 2 O2 (g)! CO2 (g) + 2 H2O (g) halogen displacement reactions: F2 > Cl2 > Br2 > I2 most least reactive reactive ex. F2 (g) + 2 NaCl (s)! 2 NaF (s) + Cl2 (g) Br2 (l) + 2 KI (s)! 2 KBr (s) + I2 (s) but Br2 (l) + 2 LiF (s)!# 2 LiBr (s) + F2 (g)

15 Some Categories of Redox Reactions displacement reactions with metals: reactions of metals with water: 2 M + 2 H2O! 2 MOH + H2 M + 2 H2O! M(OH)2 + H2 Volumetric Analysis and Titrations Titration - laboratory technique usually used to determine the composition of a reactant or a solution. reactions of metals with acid: unbalanced M + H +! M n+ + H2 metal displacement reactions: M + M X! MX + M more reactive metals are higher in the activity series any metal can displace or reduce any metal below it in the activity series add one reactant sol n to the other in a slow, controlled way stop when the stoichiometric point is reached at the stoichiometric point - no limiting reactant and no excess reactant; reactants completely consumed and converted to products Acid - Base Titration example: ml of HNO3 (aq) is titrated with Ba(OH)2 (aq) ml of M Ba(OH)2 is required to reach the stoichiometric point. Determine [HNO3]. Gravimetric Analysis lab technique used to analyze and determine composition of a compound or mixture Redox Titration example: 24.0 ml of M KMnO4 (aq) are needed to reach the stoichiometric point in the titration of a 1.00 g sample of impure H2C2O4. Determine the mass % H2C2O4 in the sample. 2 MnO H C2O4 2! 2 Mn CO2 + 8 H2O involves specific reaction of one species to form an insoluble solid that can be isolated, and its mass determined

16 ex: An excess of potassium chromate solution is added to 30.0 ml of a solution containing a soluble barium salt. All the Ba2+ is removed from solution as a yellow precipitate of BaCrO4 is formed. Determine [Ba2+] in the original solution if 2.58 g BaCrO4 is formed. Ba2+ (aq) + CrO42 (aq)! BaCrO4 (s) Some Additional Problems: ex. Saccharin, C7H5NO3S, is sometimes dispensed in tablet form. Ten of these saccharin-containing tablets have a total mass of g. These ten tablets were dissolved in water, and then oxidized so that all the sulfur present is converted to sulfate ion. Next, excess BaCl2 (aq) was added so that the sulfate ion was removed from solution as BaSO4 (s). The mass of BaSO4 collected was g. Ba2+ (aq) + SO42! BaSO4 (s) a. Determine the average mass of saccharin per tablet. b. Determine the average mass % saccharin in these tablets. Some Additional Problems: ex. A mixture contains only NaCl and Fe(NO3)3. A g sample of this mixture is dissolved in water, and then an excess of NaOH is added, resulting in the formation of Fe(OH)3 precipitate: Fe3+ (aq) + 3 OH (aq)! Fe(OH)3 (s) The precipitate is collected by filtration and dried; the mass of Fe(OH)3 collected is g. a. Determine the mass of iron in the original sample. b. Determine the mass of Fe(NO3)3 in the original sample. c. Determine the mass % Fe(NO3)3 in the original sample.

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