Chapter 4 Outline. Electrolytic Properties

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1 General Properties of Aqueous Solutions Solution = a homogeneous mixture of two or more substances Solvent = substance present in greatest quantity Solute = the other substance(s) present in a solution Aqueous solution = a solution in which the dissolving medium is water Electrolytic Properties

2 Electrolyte = a substance whose aqueous solutions contain ions (ionic substances mainly) Nonelectrolyte = a substance that does not form ions in solution (molecular substances) Ionic Compounds in Water Solid ionic substances = tightly packed crystals of ions In water ionic compounds may dissociate into component ions Water is an effective solvent for ionic compounds because the O is δ - and the H is δ + (hydration). Solution formation Coulombic attractions between solute and solvent pull solvent particles away from one another. Solute particles are pulled away from lattice structure. Solvation = solvent particles surround and stabilize the attractive forces of the solute particles. Solvation prevents recombining of ions and precipitation.

3 If ions are more attracted to each other than to the solvent, the substance will not dissolve or dissolve only slightly. Molecular Compounds in Water Covalent bonds are stronger than attractions between molecules and solvents so the molecule stays intact Some molecules will ionize when placed in water Acids are the most important of these Hydrogen ions form H + (aq) leaving the remaining portion of the molecule as an anion.

4 Strong & Weak Electrolytes Strong electrolytes = solutes that in solution, exist completely or nearly completely as ions; reactions would be written using a single arrow Weak electrolytes = solutes that mostly remain as molecules, with only a small fraction ionizing Solution consists of intact molecules uniformly dispersed in solvent Some molecular species DO ionize in solution: especially ACIDS Weak acids form an equilibrium between ions and undissociated molecules; use a double headed arrow when writing reactions

5 4.2 Precipitation Reactions Precipitate = an insoluble solid formed by a reaction in solution Occur when certain pairs of oppositely charged ions attract each other very strongly Solubility Guidelines for Ionic Compounds Solubility = the amount of a substance that can be dissolved in a given quantity of solvent at a certain temperature Solubility ranges vary greatly Miscible = dissolve in each other in any proportion Insoluble = solubility less than 0.01 mol/liter o Intermolecular forces are too strong for solvent particles to separate solute particles Solvent Solute Attraction Non-polar Polar Low (solute sticks to itself) Polar Non-polar Low (solvent sticks to itself) Non-polar Non-polar Low (but soluble) Polar Polar High (partial charges attract) Polar Ionic High (partial + full charge attract)

6 Memorize the key solubility rules! (p. 121) Exchange (Metathesis) Reactions aka: double displacement reactions General format: AB + CD AD + CB Example: BaCl2(aq) + MgSO4(aq) BaSO4(s) + MgCl2(aq) Ionic Equations Molecular equation: shows complete chemical formulas of reactants and products Complete ionic equation: shows all soluble strong electrolytes as ions Spectator ion: ion that appears in identical forms on both sides of equation; present, but plays no direct role Ba 2+ (aq) + 2Cl - (aq) + Mg 2+ (aq) + SO4 2- (aq) BaSO4(s) + Mg 2+ (aq) + 2Cl - (aq) Spectators

7 Net ionic equation: omits all spectator ions; only shows ions & molecules directly involved in reaction (Very important on AP exam!) Ba 2+ (aq) + SO4 2- (aq) BaSO4(s)

8 4.3 Acids, Bases & Neutralization Reactions Acids = substances that ionize in aqueous solutions to form hydrogen ions (H + ) proton donor hydronium ion, H3O + monoprotic vs. diprotic acids monoprotic = producing 1 H + (HCl, HNO3, HI) diprotic = producing 2 H + (H2SO4, H2CO3) o Polyprotic acids ionize in steps. o Each step yields 1 H +. H2CO3(aq) + H2O(aq) H3O + (aq) + HCO3 - (aq) HCO3 - (aq) + H2O(aq) H3O + (aq) + CO3 2- (aq) o Each step has its own equilibrium position. o In H2SO4, only first H + ionization is complete o The H + ionized in an acid comes from a highly electronegative bond o H listed first in formula or contains -COOH Base = substances that accept (react with) hydrogen ions; produce OH- ions when dissolved in water Ionic hydroxides donate OH - directly to solution

9 Other bases may remove H + from water Strong and Weak Acids and Bases Strong acids/bases are strong electrolytes Ionize completely in solution Weak acids/bases are weak electrolytes Ionize partially in solution Most acids are weak Most common weak base = NH3 Memorize the formulas of common strong acids & bases Strong Acids Hydroiodic Acid HI Hydrobromic Acid HBr Perchloric Acid HClO4 Hydrochloric Acid HCl Sulfuric Acid H2SO4* Nitric Acid HNO3 Chloric Acid HClO3 *Only 1st H+

10 Group 1 metal hydroxides Strong Bases Heavy Group 2 metal hydroxides LiOH NaOH KOH RbOH CsOH Ca(OH)2 Sr(OH)2 Ba(OH)2 Identifying Strong and Weak Electrolytes Graphic representations of acid strength Neutralization Reactions & Salts Operational Definitions Acids: sour taste, turn litmus paper red, react with metals to release H2(g), Bases: bitter taste, turn litmus paper blue, turn phenolphthalein pink

11 Neutralization reaction = reaction of an acid and a base Acid + metal hydroxide water + salt o Salt = any ionic compound whose cation comes from a base and whose anion comes from an acid = any ionic compound that is neither an acid nor a base NaOH(aq) + HCl(aq) NaCl(aq) + H2O(l) Na + (aq) + OH - (aq) + H + (aq) + Cl - (aq) Na + (aq) + Cl - (aq) + H2O(l) H + (aq) + OH - (aq) H2O(l) Mg(OH)2(s) + 2HCl(aq) MgCl2(aq) + 2H2O(l) Mg(OH)2(s) + 2H + (aq) + 2Cl - (aq) Mg 2+ (aq) + 2Cl - (aq) + 2H2O(l) Mg(OH)2(s) + 2H + (aq) Mg 2+ (aq) + 2H2O(l) Acid-base reactions with gas formation Sulfide ion hydrogen sulfide gas (H2S) 2HCl(aq) + Na2S(aq) 2NaCl(aq) + H2S(g) 2H + (aq) + 2Cl - (aq) + 2Na + (aq) + S 2- (aq) 2Na + (aq) + 2Cl - (aq) + H2S(g) 2H + (aq) + S 2- (aq) H2S(g)

12 Carbonates, hydrogen carbonates carbon dioxide (because carbonic acid is unstable and decomposes) NaHCO3(aq) + HCl(aq) NaCl(aq) + H2O(l) + CO2(g) Na + (aq) + HCO3 - (aq) + H + (aq) + Cl - (aq) Na + (aq) + Cl - (aq) + H2O(l) +CO2(g) HCO3 - (s) + H + (aq) H2O(l) + CO2(g) BE ABLE TO WRITE THESE AS NET IONIC EQUATIONS!

13 4.4 Oxidation-Reduction Reactions (REDOX) = reactions in which electrons are transferred between reactants Oxidation = loss of electrons (gets more positive) Called oxidation due to early study of reactions of metals with oxygen The more reactive the metal, the faster it is oxidized Reduction = gain of electrons (gets more negative) Keep track of oxidation/reduction by using oxidation numbers The oxidation number of atoms in their elemental form is zero. For monatomic ions, the oxidation number equals the charge on the ion. The oxidation number of oxygen is usually -2 in compounds; the major exception is the peroxide ion, which has an oxidation number of -1. The oxidation number of hydrogen is usually +1 when bonded to nonmetals and -1 when bonded to metals. The oxidation number of fluorine is -1 in all compounds. The other halogens have an oxidation

14 number of -1 in binary compounds but may have other oxidation numbers when oxygen is present. The sum of all the oxidation numbers in a neutral molecule is zero. The sum of all the oxidation numbers in a polyatomic ion is the charge of the polyatomic ion. Oxidation of Metals by Acids and Salts Single Displacement Reactions: A + BC AC + B Be able to write these as net ionic equations!

15 Reactions of metals w/ acids to produce H2 gas o The oxidation number of the metal increases indicating the atom has lost electrons (oxidized). o The oxidation number of the H atoms decrease indicating the atoms have gained electrons (reduction). o The oxidation number of the remaining species remains the same and is a spectator ion in the redox reaction. Net ionic equations for these reactions would be: Mg(s) + 2H + (aq) Mg 2+ (a) + H2(g) Zn(s) + 2H + (aq) Zn 2+ (a) + H2(g) Oxidation of metal by an aqueous salt Cu(s) + 2AgNO3(aq) 2Ag(s) + Cu(NO3)2(aq) Whenever one substance is oxidized, another must be reduced.

16 The Activity Series Used to predict if a single displacement reaction will occur Active metals (at top of series) are easily oxidized; the noble metals at the bottom of the series have low reactivity A metal can replace any metal below it in the series in a single displacement reaction A similar series exists for halogen replacements You will not have (or need) an activity series for the AP Chemistry exam

17 4.5 Concentrations of Solutions Concentration = amount of solute dissolved in a given amount of solvent Molarity (M) Molarity = Moles of solute Liters of solution Concentrations of Electrolytes When ionic compounds dissolve, the relative concentration of the ions depends on the chemical formula (SUBSCRIPTS); so concentration of a particular ion in solution may equal or be greater than the solution concentration MgCl2(s) Mg 2+ (aq) + 2Cl - (aq) 0.10 M 0.10 M 0.20 M Na2CO3(s) 2Na + (aq) + CO3 2- (aq) 0.50 M 1.0 M 0.50 M Fe2(SO4)3(s) 2Fe 3+ (aq) + 3SO4 2- (aq) 0.05 M 0.10 M 0.15 M Dilution Stock solution = sol n stored at a particular concentration Diluting stock sol n to a desired concentration If only adding solvent to a solution, moles of solute remains constant M stock = Moles stock Liters stock M dilute = Moles dilute Liters dilute M stock Liters stock = Moles stock Moles dilute = M dilute Liters dilute M= stock Liters stock = M= dilute Liters dilute =

18 4.6 Solution Stoichiometry & Chemical Analysis Stoichiometry = comparing chemical species by way of a balanced equation These comparisons must be done on a mole: mole basis Molarity may be used to find moles Titration = combining a sample of the solution with a reagent solution of known concentration (standard solution) Equivalence point = when stoichiometrically correct number of moles of each reagent is present End point = point where the indicator changes Indicator = weak acid or base that undergoes color change Indicator chosen so that endpoint is close to equivalence point Types of titrations: acid-base precipitation redox Problem solving is the key!