A subset of chemical reactions Reactions in solution A subset of chemical reactions

Learning objectives Define solution, solute, and solvent Distinguish among strong, weak and non-electrolyte Identify strong acids and strong bases Apply solubility rules to prediction of precipitate formation Classify types of chemical reaction Predict course of reaction based on activity series Define oxidation and reduction Identify oxidizing and reducing agent in reactions Determine oxidation numbers in ions and compounds

Solution A homogeneous mixture of two or more substances May or may not contain electrolytes Electrolytes- substances that conduct electricity when dissolved (acids, bases, and ionic compounds)

Electrolytes and ionic compounds All ionic compounds are electrolytes when dissolved in water Not all ionic compounds are soluble How do we tell? Rules to predict solubility Covalent molecular compounds* are non-electrolytes – no ions produced *Except acids and bases

Dissociation and ionization: same or different? Ionic compounds dissociate in water Ions already exist in the solid Acids or bases* ionize in water A pure acid or base contains no ions *Except strong bases like NaOH, Ca(OH)2 are ionic

When the weak are made strong Strong electrolytes are characterized by their nearly complete dissociation in water Weak electrolytes dissociate to a much smaller extent.

Strong, weak or non electrolyte? All soluble (ionic) salts are strong electrolytes Strong acids and bases are strong electrolytes Weak acids and bases are weak electrolytes Insoluble compounds are non-electrolytes Molecular compounds are non-electrolytes (except acids/bases)

Classifying chemical reactions Single displacement reactions Double displacement (metathesis)/ (partner exchange) reactions (in solution) Acid-base reactions Oxidation-reduction reactions Combination reactions Decomposition reactions

Single replacement (displacement) Element displaces another element from compound (redox)

Metathesis (double displacement) reactions involve changing partners AX + BY = AY + BX Driven by removal of ions from solution One example: Precipitation reactions Two solutions react to form an insoluble solid (precipitate)

Precipitation reactions Does one of the possible cation-anion combinations produce an insoluble salt? Initial compounds are all soluble Use solubility rules to investigate If yes, a precipitate is produced

Writing Equations for Rxns in Solution Molecular equation: normal equation 2KCl (aq) + Pb(NO3)2 (aq) 2KNO3 (aq) + PbCl2 (s) Ionic equation: shows all ions separately 2K+ (aq) + 2Cl- (aq) + Pb2+ (aq) + 2NO3– (aq)  2K+ (aq) + 2NO3 – (aq) + PbCl2 (s) Net ionic equation: does not show spectator ions 2Cl- (aq) + Pb2+ (aq)  PbCl2 (s) Spectator ion – any ion that is present both before and after the reaction ALL EQUATIONS STILL NEED TO BE BALANCED!

Writing Equations for Rxns in Solution Molecular equation: normal equation 2KCl (aq) + Pb(NO3)2 (aq) 2KNO3 (aq) + PbCl2 (s) Ionic equation: shows all ions separately 2K+ (aq) + 2Cl- (aq) + Pb2+ (aq) + 2NO3– (aq)  2K+ (aq) + 2NO3 – (aq) + PbCl2 (s) Net ionic equation: does not show spectator ions 2Cl- (aq) + Pb2+ (aq)  PbCl2 (s) Spectator ion – any ion that is present both before and after the reaction ALL EQUATIONS STILL NEED TO BE BALANCED!

Solubility rules– you will have to memorize these for the AP exam!

For your worksheet Write all answers on another sheet of paper. For Part 1: show ionic and net ionic equations For Part 2: Step 1: Finish the equation by switching the anions K3PO4(aq) + Al(NO3)3(aq)  KNO3 + AlPO4 make sure you write the new compounds with the correct subscripts! (use charges) Step 2: Figure out which product is soluble (aq), and which is insoluble (s) by using solubility rules Step 3: Write the ionic and net ionic equations

Precipitation Lab You must show me your pre-lab questions before you may begin. IMPORTANT: We are using nickel nitrate (Ni(NO3)2) instead of lead II nitrate (Pb(NO3)2)

Day 4 – Acid-Base Reactions

Know your acids The six strong acids All other acids are weak HCl, HBr, HI (but not HF) HNO3 (but not HNO2) H2SO4 (but not H2SO3) HClO4 (maybe HClO3) All other acids are weak

Recognizing acids Mineral acids: HCl, HNO3 etc. Conventionally H appears first in the formula All strong acids are mineral May be strong or weak Organic acids: CH3COOH etc Harder to spot Sometimes written with H in front – HCH3CO2 Always weak Presence of –OH (-SH): necessary but not sufficient Not all –OH are acidic (CH3OH is not an acid)

Recognizing bases Mineral bases usually distinguished by OH groups – all strong NaOH, Ca(OH)2 Ammonia, NH3, is an exception – is weak Organic bases do not contain –(OH) – all weak

Neutralization Combine acid with base: ACID + BASE = SALT + WATER HCl(aq) + NaOH(aq) = H2O(l) + NaCl(aq) Mg(OH)2(s) + 2HCl(aq) = MgCl2(aq) + 2H2O(l) Salt contains anion of acid and cation of base: HCl + NaOH = NaCl + H2O HCl + KOH = KCl + H2O HNO3 + KOH = KNO3 + H2O 2HCl + Ca(OH)2 = CaCl2 + 2H2O HCN + NaOH = NaCN + H2O

Acid-base reaction with gas formation Tums... HCl(aq) + NaHCO3(aq) = NaCl(aq) + H2CO3(aq) H2CO3 is unstable: H2CO3(aq) = H2O(l) + CO2(g) Bad egg gas: 2HCl + Na2S = H2S(g) + 2NaCl(aq)

Oxidation - reduction Oxidation is loss of electrons Reduction is gain of electrons Oxidation is always accompanied by reduction The total number of electrons is kept constant Oxidizing agents oxidize and are themselves reduced Reducing agents reduce and are themselves oxidized

Oxidation numbers Oxidation number is the number of electrons gained or lost by the element in making a compound Metals are typically considered more 'cation-like' and would possess positive oxidation numbers, while nonmetals are considered more 'anion-like' and would possess negative oxidation numbers.

Predicting oxidation numbers Oxidation number of atoms in element is zero Oxidation number of element in monatomic ion equals charge Sum of oxidation numbers in compound is zero Sum of oxidation numbers in polyatomic ion equals charge F has ON –1 H has ON +1; except in metal hydrides where it is –1 Oxygen is usually –2. Exceptions: O is –1 in hydrogen peroxide, and other peroxides O is –1/2 in superoxides KO2 In OF2 O is +2

Position of element in periodic table determines oxidation number G1A is +1 G2A is +2 G3A is +3 (some rare exceptions) G5A are –3 in compounds with metals, H or with NH4+ Exceptions are compounds with elements to right (e.g. NO2, PF5); in which case use rules 3 and 4. G6A below O (S, Se etc.) are –2 in binary compounds with metals, H or NH4+ When combined with O or lighter halogen (e.g. SeO2, SF6) use rules 3 and 4. G7A elements are –1 in binary compounds with metals, H or NH4+ or with a heavier halogen (e.g. Cl in BrCl3) When combined with O or a lighter halogen, use rules 3 and 4 (e.g. Br in BrCl3 or Cl in ClO4-).

Identifying reagents Those elements that tend to give up electrons (metals) are typically categorized as reducing agents and those that tend to accept electrons (nonmetals) are referred to as oxidizing agents.

Identify redox by change in oxidation numbers Reducing agent increases its oxidation number (Na) Oxidizing agent decreases its oxidation number (H in H2O)

Nuggets of redox processes Where there is oxidation there is always reduction Oxidizing agent Reducing agent Is itself reduced Is itself oxidized Gains electrons Loses electrons Causes oxidation Causes reduction

Iron reduces Cu2+ to Cu Iron reduces Cu2+ ions to Cu Cu does not reduce Fe2+

Applying activity series to metals in acids Mg is higher than H in activity series – forms H2 Cu is lower than H in activity series – no H2 produced

Element can be oxidizer and reducer depending on relative positions in activity series Fe reduces Cu2+ Cu reduces Ag+ (lower activity) Fe2+ is reduced by Zn (higher activity)

Combination reactions Element + element  compound (redox) Metal + nonmetal  binary ionic compound Nonmetal + nonmetal  binary covalent compound Compound + element  compound (redox) Compound + compound  compound

Decomposition reactions Compound  element + element (redox) Compound  element + compound (redox) Compound  compound + compound

Production of a gas If product is a gas that has low solubility in water, reaction produces gas Any carbonate with an acid for example