Aqueous Reactions and Solution Stoichiometry

4.1 General Properties of Aqueous Solutions Aqueous solutions: solutions in which water is the solvent Solvent: dissolving medium Solute: dissolved by solvent Electrolytic solutions: Aqueous ionic solns Conduct electrical current Electrolyte: solute that produces ions in solution

Electrolyte vs. Nonelectrolyte Electrolytes form ions in solution Conduct electrical current when dissolved NaCl, KNO3 Nonelectrolytes Do not form ions in solution Do not conduct electrical current sucrose C11H22O11

Conduction of Electrical Current by an Electrolytic Solution

Ionic Compounds in Water Dissociate into separate ions Because water is a polar solvent, ions are dispersed and surrounded by water molecules (fig 4.3a)

Electrical Current is a movement of charged particles Because ions can now move in solution, they can carry an electrical current (electrical current is a net movement of charges)

Molecular Cpds in Water Molecular = Covalent bonding Entire molecules are dispersed (fig 4.3b) Because molecules are uncharged, they cannot conduct an electrical current

Some molecular cpds form ions Exceptions: some molecular compounds react with water to form ions, such as acids HCl + H2O  H3O+ + Cl-

Strong vs. Weak Electrolytes Many ionic substances dissolve completely (≈100% dissociated) More ions in solution = greater conductivity These are strong electrolytes HCl + H2O → H3O+ + Cl- Substances that partially dissolve form fewer ions = lesser conductivity These are weak electrolytes CH3COOH + H2O ↔ H3O+ + CH3COO-

Identifying Strong and Weak Electrolytes

Strong Electrolytes Are… Strong acids Strong bases Soluble ionic salts

Strong Electrolytes Are… Strong acids, strong bases, soluble ionic salts

Relating relative numbers of ions to chemical formulas MgCl2  Mg 2+ + 2Cl- one mole  one mole + 2 moles = 3 moles How many cations & anions would be formed from the following? 6 NiSO4 6 Ca(NO3)2 6 Na3PO4 6Al2(SO4)3

4.2 Precipitation Reactions Precipitate: insoluble compound formed from a reaction of solutions. When KI dissolves  K+ (aq) + I- (aq) When Pb(NO3)2 dissolves Pb2+(aq) + 2NO3-(aq) When the solutions are added together: 2K+(aq) + 2I-(aq) + Pb2+(aq) + 2NO3-(aq)  2K+(aq) + 2NO3-(aq) + PbI2 (s)

Metathesis (Exchange) Reactions Metathesis comes from a Greek word that means “to transpose.” It appears the ions in the reactant compounds exchange, or transpose, ions. AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)

AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq) Molecular Equation The molecular equation lists the reactants and products in their molecular form. AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)

Ionic Equation In the ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions. This more accurately reflects the species that are found in the reaction mixture. Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)  AgCl (s) + K+ (aq) + NO3- (aq)

Net Ionic Equation To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right. Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)  AgCl (s) + K+(aq) + NO3-(aq)

Solubility of Ionic Compounds Solubility is measured as the mass of solute that will dissolve in a given volume of solvent For solids & liquids it is often measured as g solute/100 g H2O @ 25°C Examples CaCl2 74.5g per100g water MnSO4 85.27 KNO3 13.3 Ethanol infinite PbI2 0.044

Solubility Table

Effect of Temperature on Salt Solubility

The Dissolving Process: Solvation Dipoles of water molecules align with ions Ions removed from solid Ions surrounded by water molecules forming “hydration shell”

4.3 Acid-Base Reactions Acids Ionize in aqueous soln to form hydrogen ion (H+) HCl (aq) → H+ (aq) + Cl- (aq) Acids produce H+ (Arrhenius’ definition) Acids donate H+ in chem rxns (Bronsted-Lowry) Monoprotic acids produce one H+ per molecule Diprotic acids produce two H+ per molecule H2SO4(aq) → H + + HSO4- HSO4- (aq) ↔ H + + SO42- Polyprotic acids? Reduce pH pH inversely proportional to [H+] pH = -log [H+]

Bases Produce OH- ions in aqueous solutions. Arrhenius’ definition Accept H+ in aqueous solutions. Bronsted-Lowry definition NaOH (aq)  Na+ (aq) + OH- (aq) NH3 (aq) + H2O > NH4+ (aq) + OH- (aq) Increase pH by neutralizing H+

Strong & Weak Acids & Bases Strong = highly ionized, strong electrolyte Weak = low ionization, weak electrolyte

Neutralization Reactions & Salts Reaction of an acid with a base Acid + base  water + salt HCl (aq) + NaOH (aq)  H2O (l) + NaCl (aq) Salt is a category of substances: An ionic cpd whose cation came from a base and whose anion came from an acid Total ionic equation: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  H2O (l) + Na+(aq) + Cl-(aq) Net ionic equation: H+(aq) + OH-(aq)  H2O (l)

Ionic Equations HCl (aq) + NaOH (aq)  H2O (l) + NaCl (aq) Write all aqueous substances as dissociated ions Total ionic equation: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  H2O (l) + Na+(aq) + Cl-(aq) Cancel spectator ions (ions common to both sides) Net ionic equation for neutralization reaction H+(aq) + OH-(aq)  H2O (l)

Ionic Equations Write the molecular, total ionic and net ionic equations for the reaction of magnesium hydroxide and hydrochloric acid. Mg(OH)2 + 2HCl  2H2O + MgCl2 Total ionic? Net ionic?

Acid-Base Rxns with Gas Formation Two ions, in particular, react with acids to produce a gas Sulfide, S2- and bicarbonate HCO3- These reactions could be used as a test for acids 2HCl(aq) + Na2S(aq)  H2S(g) + 2NaCl(aq) Write the total and net ionic equations.

Acid-Base Rxns with Gas Formation HCl(aq) + NaHCO3(aq)  NaCl(aq) + H2O(l) + CO2(g) Write the total ionic equation Write the net ionic equation

4.4 Oxidation-Reduction Reactions: Oxidation is a loss of electrons Redox reactions involve the transfer of e-s A substance is oxidized when it loses e-s Ca(s) + 2HCl(aq)  CaCl2 (aq) + H2(g) In this reaction, Ca is oxidized… Ca(s) + 2H+(aq)  Ca2+(aq)+ H2(g) Ca0  Ca2+ + 2e-

Reduction is a gain of electrons A substance is reduced when it gains e-s Ca(s) + 2HCl(aq)  CaCl2 (aq) + H2(g) In this reaction, H+ is reduced… Ca(s) + 2H+(aq)  Ca2+(aq)+ H2(g) 2H1+ (aq) + 2e-  H20(g)

Redox reactions are coupled Oxidation & reduction always occur together One cannot occur without the other Ca(s) + 2H+(aq)  Ca2+(aq)+ H2(g) Ca0  Ca2+ + 2e- 2H+(aq) + 2e-  H2(g)

Oxidation States Assigning oxidation numbers is a way of keeping track of electrons in a cpd Some rules… Oxidation number of elements = 0 e.g. Mg0 Oxidation number of monatomic ions = charge of the ion e.g. Al3+ Nonmetals usually have negative ox #’s examples: O2-, F 1- Exception: H1+ Sum of ox #’s in a neutral cpd = 0 e.g. MgSO4 Sum of ox #’s for a polyatomic ion = charge of the ion e.g. SO42-

Determining Oxidation States Using the above rules, determine the oxidation state of the following Mn in MnO2 Mn in MnO41- Cl in HCl Cl in HClO Cl in HClO4 Cl in ClO41- Cl in ClO31- Fe in FeBr2 Fe in FeSO4

Oxidation Numbers & Redox Rxns The oxidation number is reduced in reduction 2H1+ (aq) + 2e-  H20(g) The oxidation number is increased in oxidation Ca0  Ca2+ + 2e-

Oxidation Numbers & Redox Rxns Identify the atoms that are oxidized and reduced in the following reactions. Identify the oxidizing and reducing agents. Ni(s) + Cl2(g)  NiCl2(s) 3Fe(NO3)2(aq) + 2Al(s)  3Fe(s) + 2Al(NO3)3 (aq) Cl2 (aq) + 2NaI (aq)  I2 (aq) + 2NaCl(aq) PbS(s) + 4H2O2 (aq)  PbSO4 (s) + 4H2O(l)

Metathesis Reactions Displacement reactions Some ion is displaced by another AX + BY  AY + BX AgNO3 + KCl  AgCl + KNO3 For the reaction to lead to a net change, ions have to be removed from solution Driving forces for metathesis reactions: Formation of precipitate, gas, weak or nonelectrolyte

Oxidation of Metals by Acids Single displacement rxn A + BX  AX + B Zn + 2HBr  ZnBr2 + H2 Write the ionic eqn

Oxidation of Metals by Acids Write the half-reactions (oxidation, reduction) Zn0  Zn2+ + 2e- 2H1+ + 2e-  H20 Write the redox reaction by adding the half reactions Zn0 + 2H1+ + 2e-  Zn2+ + 2e- + H20 Zn0 + 2H1+ +  Zn2+ + H20

Oxidation of metals by acids or salts Oxidation of metals by acids & salts Mg(s) + 2HCl(aq)  MgCl2(aq) + H2(g) Fe(s) + Ni(NO3)2(aq)  Fe(NO3)2(aq) + Ni(s) How can you predict whether either of these reactions will occur? Use Activity Series (Table 4.5)

Activity Series A list of metals arranged according to the ease with which they are oxidized. If a metal is higher on the list, it is easier to oxidize If a metal is higher on the list, it will displace any ion lower on the list Elements at the top of the list are referred to as active metals

Activity Series A list of metals arranged according to the ease with which they are oxidized.

4.5 Concentrations of Solutions Molarity Moles solute / liters of solution What mass of NaCl is required to make 1.00L of 1.00M solution of NaCl? 58.45 g NaCl Calculate the molarity of a soln when 23.4g Na2SO4 is dissolved to a volume of solution of 125 mL 1.32M Na2SO4

Molar Concentration of Electrolytes Molar concentration of ions relates to the mole ratio of ions in the formula In 2.0M Na2SO4 .... [Na+] = 4.0M [SO42-] = 2.0M What is the [K+] in 0.025M potassium carbonate? K2CO3, therefore [K+] = 0.050M

Dilutions Mi x Vi = Mf x Vf Where Mi is initial concentration Vi is initial volume Mf is final concentration Vf is final volume C1V1 = C2V2

4.6 Solution Stoichiometry & Chemical Analysis Solution stoichiometry is used to determine unknown concentrations of solutions, quantities of reactants, products, etc. These are classical stoichiometry problems, except that molar quantities must be determined using molar concentrations and volumes

Using Molarities in Stoichiometric Calculations

Solution Stoichiometry Determine the number of moles of HNO3 in 25 mL of 6.0 M HNO3 M = moles solute / L of solution Moles = M x V Moles = 6.0 moles / L x .025 L Moles = 0.15 moles HNO3

Sample Problem Determine the moles of water formed when 25.0mL of 0.100 M HNO3 is completely neutralized by NaOH. Calculate moles of HNO3 (0.100 moles/L) x (0.025L) = 0.0025 mol Write balanced equation HNO3 + NaOH  NaNO3 + H2O Determine yield of water using stoich methods 2.5 E-3 mol HNO3 x (1 H2O /1 HNO3) = 2.5 E-3 mol H2O

Titrations Process of determining the unknown concentration of a solute by reacting it with one of known concentration (a standard solution) Equivalence point Stoichiometrically equivalent quantities are present End point Point at which an indicator changes color, approximating the equivalence point. Indicator A substance that tells when the equivalence point is achieved E.g. acid-base indicators, formation of a precipitate, etc.

Acid-Base Titration Titration is an analytical technique in which one can calculate the concentration of a solute in a solution.

Titration Example Determine the [Cl-] in a sample of water by titrating with 0.100 M AgNO3 Data: 22.5 mL of standard was required to titrate 45.0 mL water Write net ionic equation Ag+(aq) + Cl-(aq)  AgCl (s) Determine moles of Ag+(aq) used 0.0225L x 0.100 mol/L = 0.00225 mol Ag+ Determine moles of Cl-(aq) consumed 0.00225 mol Ag+ x (1 mole Cl- / 1mol Ag+) = 0.00225 mol Cl- Calculate [Cl-] 0.00225 mol Cl-/0.045L = 0.05 mol/L

Titration Problem 25.00 mL of 0.100 M HBr is titrated with 0.200 M NaOH. How many mL of the NaOH solution are required to reach the equivalence point? Write the net ionic equation. Determine moles of H+ ion present. Determine moles of OH- ion needed. Calculate volume of NaOH solution containing necessary amount of OH- ion.

Acid-Base Titration Indicator: phenolphtalein, colorless when acidic, pink when basic Equivalence point: [H3O+] = [OH-] Solution remains faintly pink 20.0 mL HCl is titrated A few drops of phenolpthalein added as indicator 0.100 M NaOH standard solution of is used to titrate the sample. 45.0 mL of NaOH solution was required to turn the acid solution pink What was the concentration of the acid?

Titration Example HCl + NaOH  NaCl + H2O Net eqn H+ + OH-  H2O Determine # moles OH- required to reach end point Use equation to determine # moles of H+ Determine [H+] by dividing moles H+ by volume of acid solution