Chapters 1 – 3 Summary Measurements Basics of atomic structure Chemical and molecular formulas The concept of “moles” Chemical equations and stoichiometry Calculations involving stoichiometry

Chapter 4 Chemical Reactions Ions in Aqueous Solutions Types of Chemical Reactions Working with Solutions Quantitative Analysis

Ionic Theory of Solutions Arrhenius Theory (1884): Aqueous solutions that are electrically conducting contain ions. (ie., The ions in aqueous solutions account for the electrical conductivity of the solution.)

Water is the dissolving medium, or solvent. Aqueous Solutions Water is the dissolving medium, or solvent.

A Solvent retains its phase (if different from the solute) is present in greater amount (if the same phase as the solute)

Some Properties of Water Water is “bent” or V-shaped. The O-H bonds are covalent (e-‘s are shared). Water is a polar molecule. Hydration occurs when salts, strong acids and strong bases dissolve in water.

Overall charge on a molecule is “zero”. Nuclear charge on hydrogen attracts negative charges Electrons on oxygen attract positive charges F = k (q1 * q2)/r12 2

Electrical conductivity of pure water The light does not glow, thus no current goes through the solution.

A Solute dissolves in water (or other “solvent”) changes phase (if different from the solvent) is present in lesser amount (if the same phase as the solvent)

Ionic solid is solvated by water molecules

Solvated ions Solvated ions (ie., solutes) are called electrolytes. Ionic solids and molecular compounds may both become solvated in water. Electrolyte = a compound that dissolves in water to give an electrically conducting solution.

Electrolytes Strong - conduct current efficiently; highly ionized NaCl, HNO3, HCl, NaOH Weak - conduct a small current; medium ionization vinegar (acetic acid), tap water Non-electrolyte - no current flows; no ions formed in solution pure water, sugar solution, CH3OH

Electrical Conductivity of a solution of NaCl Thus, NaCl is highly ionized in water.

Electrical conductivity of ammonia NH4OH is weakly ionized in water.

Reactions in Water NH3(aq) + H2O (l) NH4+(aq) + OH-(aq) kf NH3(aq) + H2O (l) NH4+(aq) + OH-(aq) Forward direction (strong electrolytes) Reverse direction (weak electrolytes) When kr is larger than kf, then the compound is weakly ionized. kr

Simple Rules for Solubility 1. Most nitrate (NO3-) salts are soluble. 2. Most alkali (group 1A) salts and NH4+ are soluble. 3. Most Cl-, Br-, and I- salts are soluble (NOT Ag+, Pb2+, Hg22+) 4. Most sulfate salts are soluble (NOT BaSO4, PbSO4, HgSO4, CaSO4) 5. Most OH- salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble) 6. Most S2-, CO32-, CrO42-, PO43- salts are only slightly soluble.

Example A precipitation reaction forming PbI2

Describing Reactions in Solution 1. Molecular equation (reactants and products as compounds) AgNO3(aq) + NaCl(aq) ® AgCl(s) + NaNO3(aq) 2. Complete ionic equation (all strong electrolytes shown as ions) Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) ® AgCl(s) + Na+(aq) + NO3-(aq)

Describing Reactions in Solution (continued) 3. Net ionic equation (show only components that actually react) Ag+(aq) + Cl-(aq) ® AgCl(s) Na+ and NO3- are spectator ions.

Examples Molecular equation: Complete ionic: Pb(NO3)2 + Na2SO4 PbSO4 + 2 NaNO3 Complete ionic: Pb2+ + 2 NO3- + 2 Na+ +SO42- PbSO4 (s) + 2 Na+ + 2 NO3- Net ionic: Pb2+ (aq) + SO42-(aq) PbSO4 (s)

Types of Solution Reactions Precipitation reactions AgNO3(aq) + NaCl(aq) ® AgCl(s) + NaNO3(aq) Acid-base reactions NaOH(aq) + HCl(aq) ® NaCl(aq) + H2O(l) Oxidation-reduction reactions Fe2O3(s) + Al(s) ® Fe(l) + Al2O3(s)

Acid-Base Concepts Two concepts of acid-base theory includes: The Arrhenius concept: an acid increases the concentration of H+, and a base increases the concentration of OH- Arrhenius acids/bases are only defined for water as a solvent. The Brønsted-Lowry concept: an acid donates a proton, and a base accepts a proton. The Bronsted-Lowry definition covers a wider range of compounds than the Arrhenius concept. 2

hydrochloric and sulfuric acid Acids Strong acids - dissociate completely to produce H+ in solution hydrochloric and sulfuric acid Weak acids - dissociate to a slight extent to give H+ in solution acetic and formic acid

Bases Strong bases - react completely with water to give OH- ions. sodium hydroxide Weak bases - react only slightly with water to give OH- ions. ammonia

Acid-Base Reactions Involving Gases Molecular equation: HNO3(aq) + KOH(aq) H2O(l) + KNO3(aq) Total ionic: H+ + NO3- + K+ + OH- H2O(l) + K+ + NO3- Net ionic: H+(aq) + OH-(aq) H2O (l)

Performing Calculations for Acid-Base Reactions 1. List initial species and predict reaction. 2. Write balanced net ionic reaction. 3. Calculate moles of reactants. 4. Determine limiting reactant. 5. Calculate moles of required reactant/product. 6. Convert to grams or volume, as required.

Molarity Molarity (M) = moles of solute per volume of solution in liters: M = Molarity = 3 M solution of HCl = 3 moles HCl/liter soln moles of solute liter of solution

Concentration Example (Dilution) How would one make 1 liter of a 3 M solution of HCl from a solution of concentrated HCl? (conc. HCl is 12 M) use: M1 * V1 = M2* V2 or Mconc * Vconc = MHCl * VHCl Vconc = = 0.25 liters = 250 ml 3M * 1 liter 12 M

Common Terms of Solution Concentration Stock - routinely used solutions prepared in concentrated form. Concentrated - relatively large ratio of solute to solvent. (5.0 M NaCl) Dilute - relatively small ratio of solute to solvent. (0.01 M NaCl)

Volumetric Flask

Making solutions from solid solutes

Volumetric Flask

Add liquid up to the volume indicator mark (ie., the meniscus)

Meniscus is U-shaped in polar liquids, Inverted U for non-polar liquids

Volumetric and Gravimetric Analysis (Titrations) Methods used to accurately determine the concentration of a solution. Gravimetric analysis involves precipitating a solute, and weighing the amount of precipitate formed (requires stoichiometric calculations). Volumetric analysis involves using a solution of known concentration that reacts predictably and completely with the solution of unknown concentration. (ie., How many ml of known solution are required to completely react with a solution of unknown concentration?)

Gravimetric Analysis

Acid-Base Indicators Organic dyes that change color depending upon the amount of H+ (ie., H3O+) ions present in the solution.

Red Cabbage Extracts Are Excellent Acid-Base Indicators

Volumetric Analysis: very similar the dilution calculations It took 20 ml of a 3 M solution of NaOH to neutralize 50 ml of HCl. What is the concentration of the HCl solution? use: M1 * V1 = M2* V2 or MNaOH * VNaOH = MHCl * VHCl MHCl = = 1.2 M 3 M * 20 ml 50 ml

Key Titration Terms Titrant - solution of known concentration used in titration Analyte - substance being analyzed Equivalence point - enough titrant added to react exactly with the analyte Endpoint - the indicator changes color so you can tell the equivalence point has been reached.

Titration Apparatus Bruet Titrant Erlenmeyer Flask Analyte

Final Volume of Titrant used Indicator changes color at the endpoint

Titration Curves Endpoint

pH curves and the buffer region 0 , because [base] = [acid] buffer region pH = pKa, at midpoint of buffer region 2

Solution Examples NOTE: Precipitation reactions are generally used for Gravimetric Analysis. (ie., the sample is weighed to determine the amount of a substance in a sample). A QUANTATIVE ANALYSIS METHOD Problems: 76, 78, 80, 84, 105, 115

Oxidation-Reduction Reactions A type of chemical reaction where an exchange of electrons occurs. Oxidation – a loss of electrons (e-’s) Reduction – a gain of electrons Oxidizing Agent – receives electrons (is reduced) Cu2+ (aq) + 2 e- Cu (s) Reducing Agent – donates electrons (is oxidized) Fe (s) Fe2+ (aq) + 2 e-

Single Displacement Reaction Cu (s) 2Fe (s) + 3CuSO4 Fe2(SO4)3 + 3Cu (s) oxidized reduced

Types of Redox Reactions Combination Reactions 2 Na (s) + Cl2 (g) 2 NaCl (s) Decomposition Reactions 2 HgO (s) 2 Hg (l) + O2 (g) Single/Double Displacement Reactions Zn (s) + 2 HCl (aq) ZnCl2 (aq) + H2 (g) Combustion Reactions 2C4H10 (g) + 13O2 (g) 8CO2 + 10H2O (g)

Combination Reaction

Combustion Reaction

Decomposition Reaction

Rules for Assigning Oxidation States 1. Oxidation state of an atom in an element = 0 2. Oxidation state of monatomic element = charge 3. Oxygen = -2 in covalent compounds (except in peroxides where it = -1) 4. H = +1 in covalent compounds 5. Fluorine = -1 in compounds 6. Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion

Balancing by Half-Reaction Method (p. 625; sec. 19.1) Write separate reduction, oxidation reactions; 2. For each half-reaction: - Balance atoms other than H or O Balance O using H2O Balance H using H+ Make sure charges are balanced (use electrons to balance charges) Use multipliers to balance e-’s

Balancing by Half-Reaction Method (continued) 3. If necessary, multiply by integer to equalize electron count. 4. Add half-reactions. 5. Check that elements and charges are balanced.

Half-Reaction Method - Balancing in Base 1. Balance as in acid. 2. Add OH- that equals H+ ions (both sides!) 3. Form water by combining H+, OH-. 4. Check elements and charges for balance.

Redox Reactions: Examples Oxidation Numbers: 51, 53 Balancing Redox Equations: 60 Calculations involving Redox reactions: 134, 137