Reactions in Aqueous Solutions Chapter 4

General properties

Solution

Electrolyte

Hydration

HCl(l) H+(aq) + Cl−(aq) CH3COOH(aq) H+(aq) + CH3COO−(aq)

Review of Concepts The diagrams here show three compounds AB2 (a), AC2 (b), and AD2 (c) dissolved in water. Which is the strongest electrolyte and which is the weakest? (For simplicity, water molecules are not shown.)

Precipitate reactions

Double-displacement reaction

Solubility Rules

Equations Pb2+(aq) + 2I−(aq) PbI2(s) Pb(NO3)2(aq) + 2 KI(aq) PbI2(s) + 2KNO3(aq) Pb2+(aq) + 2NO3−(aq) + 2K+(aq) + 2I−(aq) PbI2(s) + 2K+(aq) + 2NO3(aq) Pb2+(aq) + 2I−(aq) PbI2(s)

Example 4.2 page 125 in textbook K3PO4(aq) + Ca(NO3)2(aq) Example 4.2 page 125 in textbook

Review of Concepts Which of the diagrams here acuratly describes the reaction between Ca(NO3)2(aq) and Na2CO3(aq)? For simplicity, only the Ca2+ (yellow) and CO32− (blue) ions are shown.

Acid-base reactions

General Properties Acid base Sour taste Color changes in plant dyes React with metals to produce H2 gas React with carbonates and bicarbonates to produce CO2 gas Aqueous acid solutions conduct electricity Taste bitter Feel slippery Color changes in plant dyes Aqueous base solutions conduct electricity

Brønsted Acid and Bases Proton donor Monoprotic Diprotic Triprotic Proton acceptor

Strong acids Strong bases HI HBr HClO4 HCl H2SO4 HNO3 NaOH KOH LiOH RbOH CsOH                     Ca(OH)2 Ba(OH)2 Sr(OH)2 Strong acids/bases are strong electrolytes and will completely dissociate in water.

Review of Concepts Which of the following diagrams best represents a weak acid? Very weak acid? Strong acid? The proton exists in water as the hydronium ion. All acids are monoprotic. (For simplicity, water molecules are not shown.)

Acid-Base Neutralization Reaction between an acid and a base Generally aqueous solutions result in water and a salt Ex: HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) *this is a strong acid and strong base so they completely dissociate and the net ionic equation is H+(aq) + OH−(aq) H2O(l) Ex: HCN(aq) + NaOH(aq) NaCN(aq) + H2O(l) *this is a weak acid and strong base so the acid does not completely ionize in water. When writing the ionic and net ionic equations you cannot break the weak acid apart! The net ionic equation is HCN(aq) + OH−(aq) CN−(aq) + H2O(l)

Gas formation Certain salts react with acids to produce gaseous products HNO3 breaks down into H2O(l) + NO2(g) + NO(g) H2CO3 breaks down into H2O(l) + CO2(g) H2SO3 breaks down into H2O(l) + SO2(g) NH4OH breaks down into H2O(l) + NH3(g) H2S(g) CO2(g) H2(g) If you get one of these as a product in your molecular equation, they immediately breakdown as above Gasses do not ionize

Double Replacement Rxns Review Driving Force How do you recognize it? Precipitate You must memorize the solubility rules. Any compound formed from two ions can be recognized as soluble (written as separate ions) or as a precipitate (written as a molecule). Gas formed You must memorize the combinations that decompose into gases (there are 4). You must also memorize the gases that form. For example, when you H2SO3 as a product, you must know it decomposes into H2O and SO2 gas. Weak electrolyte You must memorize the short list of strong acids and strong bases so you will recognize all the weak acids and bases that dissolve, but do not dissociate into ions. The weak base ammonia, NH3, is in this category. It exits in water as NH3(aq) and only slightly forms the ions NH4+ + OH−

Oxidation-reduction reactions

Half-reaction OIL RIG Oxidation reaction Reduction reaction Reaction that involves the loss of electrons Contains reducing agent-donates electrons Involves the gain of electrons Contains oxidizing agent-accepts electrons

Oxidation Number Charge of the atom would have in a molecule if electrons were transferred completely Rules Uncombined elements = 0 Neutral compounds sum = 0 Ion = ion charge (polyatomic ions sum to charge) Exceptions Hydrogen +1 w/ nonmetals, −1 w/ metals Oxygen −2 except w/ fluorine (+2), in peroxides (−1) Fluorine ALWAYS −1

More common oxidation numbers are in red.

Example 2 Mg(s) + O2(g)  2 MgO(s)

Types of Redox Reactions Combination 2 Al(s) + 3 Br2(l) 2 AlBr3(s) Decomposition 2 NaH(s) 2 Na(s) + H2(g) Combustion C2H8(g) + 5 O2(g) 3 CO2(g) + 4H2O(l) Displacement Hydrogen Ca(s) + 2 H2O(l) Ca(OH)2(s) + H2(g) Metal Zn(s) + 2 HCl(aq) ZnCl2(aq) + H2(g) Halogen Cl2(g) +2 KBr(aq) 2 KCl(aq) + Br2(l) Disproportionation 2 H2O2(aq) 2H2O(l) + O2(g)

Activity Series For Halogens: F2 > Cl2 > Br2 > I2

Elements most likely to undergo disproportionation

Concentration n M = V Molarity = moles of solute liters of solution Ex: 1M KCl solution KCl(s) K+(aq) + Cl−(aq) H2O Ex: 1M Ba(NO3)2 solution Ba(NO3)2(s) Ba2+(aq) + 2 NO3−(aq) H2O

Example How many grams of potassium dichromate (K2Cr2O7) are required to prepare a 250 mL solution whose concentration is 2.16M?

Dilutions MiVi = MfVf

Example Describe how you would prepare 5.00x102 mL of a 1.75M H2SO4 solution, starting with an 8.16M stock solution of H2SO4.

Review of Concepts What is the final concentration of a 0.6M NaCl solution if its volume is doubled and the number of moles of solute is tripled?

Quantitative analysis Gravimetric analysis Titrations Acid-base redox

Gravimetric Analysis

Example A 0.5662 g sample of an ionic compound containing chloride ions and an unknown metal is dissolved in water and treated with and excess of AgNO3. if 1.0882 g of AgCl precipitate forms, what is the percent by mass of Cl in the original compound?

Acid-base titrations

Example How many mL of a 0.610M NaOH solution are needed to neutralize 20.0 mL of a 0.245M H2SO4 solution?

Redox titrations

Example A 16.42 mL volume of 0.1327M KMnO4 solution is needed to oxidize 25.00 mL of a FeSO4 solution in an acidic medium. What is the concentration of the FeSO4 solution in molarity? The net ionic equation is 5Fe2+ + MnO4− + 8H+  Mn2+ + 5Fe3+ + 4H2O