Reactions in Aqueous Solutions

7.1 Predicting Whether a Reaction Will Occur? Driving Forces Formation of solid Formation of water Transfer of electrons Formation of a gas

7.2 Reactions in Which a Solid Forms? Driving forces Precipitation – formation of solid E.g Ba(NO 3 ) 2 (aq) + K 2 CrO 4 (aq)  yellow solid yellow sol.colorless

What happens When an Ionic Compound Dissolves in Water? Ba(NO 3 ) 2 (aq) => barium nitrate (white solid) has been dissolved in water Containing Ba 2+ (aq) and 2 NO 3 - (aq) ions Strong electrolytes: A unit of substance that dissolves in water produces separate ions E.g K 2 CrO 4 (aq) contains K +, K+ and CrO 4 2- ions => strong electrolyte

Strong Electrolytes Ba(NO 3 ) 2 (aq) + K 2 CrO 4 (aq)  yellow solid

How to Decide What Products Form NO 3 - CrO 4 2- K+K+K+K+ KNO 3 Ba 2+ Ba(NO 3 ) 2 Write out all possible formula that can be formed

Using Solubility Rules General Rules for Solubility of ionic compounds (salts) in water at 25 o C Most nitrate (NO 3 - ) salts are soluble Most salts of Na +, K + and NH 4 + are soluble Most chloride, bromide, iodide (Cl -, Br - and I - ) salts are soluble. Notable exceptions are AgCl, PbCl 2 and Hg 2 Cl 2. Most sulfate salts are soluble. Notable exceptions are BaSO4, PbSO 4 and CaSO 4 Most hydroxide ( - OH) compounds are slightly soluble. The important exception are NaOH and KOH. Ba(OH) 2 and Ca(OH) 2 are only moderately soluble Most sulfide (S 2- ), carbonate (CO 3 2- ) and phosphate (PO 4 3- ) salts are slightly soluble * The term insoluble and slightly soluble really mean the same thing

Identifying Precipitates in Reactions Where a Solid Forms Step 1: Write the reactants as they actually exist before any reaction occurs Step 2:Consider the various solids that could form. Step 3: use solubility rules to decide whether a solid forms E.gAgNO 3 (aq) + KCl(aq)  white solid

Examples Using solubility rules to predict the product of Reactions KNO 3 (aq) + BaCl 2 (g)  Na 2 S(aq) + Cu(NO 3 ) 2 (aq)  KOH(aq) + Fe(NO 3 ) 2 (aq) 

Describing Reactions in Aqueous Solutions Molecular equation: shows overall reaction but not necessary the actual forms of the reactants and products in solution Complete ionic equation: represents all reactants and products that are strong electrolytes as ions. All reactants and products are included. The net ionic equation includes only those components that undergoes a change. Spectator ions are not included

Examples  For each of the following reactions, write molecular equation, the complete ionic equation and the net ionic equation  Aqueous sodium chloride is added to aqueous silver nitrate to form solid silver chloride plus aqueous sodium nitrate  Aqueous nickel (II) nitrate is added to aqueous potassium carbonate to form solid nickel (II) carbonate and aqueous potassium nitrate

7.4 Reactions That Form Water: Acids and Bases Arrhenius acids: a substance that produces H + ions (protons) when it dissolved in water Strong acids: strong electrolytes Common strong acids: HCl, HBr, HI, H 2 SO 4, HClO 4, HNO 3 dissolved in H 2 O HCl(aq)  H + (aq) + Cl - (aq)

7.4 Reactions That Form Water: Acids and Bases Arrhenius Bases: substance that produces –OH ion (hydroxide ions) in water Strong bases: Strong electrolytes Common strong bases: KOH, LiOH, NaOH, Ba(OH) 2 and Sr(OH) 2 dissolved in H 2 O NaOH(aq)  Na + (aq) + - OH(aq)

Writing Equations for Acid-Base Reactions Nitric acid is a strong acid. Write the molecular, complete ionic and net ion equations for the reaction of aqueous nitric acid and aqueous potassium hydroxide. Hydrobromic acid is a strong acid. Write the molecular, complete ionic and net ion equations for the reaction of aqueous hydrobromic acid and aqueous barium hydroxide

7.5 Reactions of Metals with Nonmetals (Oxidation Reduction) Oxidation – Reduction or Redox reaction: process of transferring electrons from one to the other Step in determination Redox Reaction Determine the charges of each atom Oxidation: loss of electron(s); charge becomes more positive Reduction: gain of electron(s); charge becomes less positive E.g2 Na(s) + Cl 2 (g)  2 NaCl(s)

Examples For each of the following reactions, show how electrons are gained and lost 2 Al(s) + 3 I 2 (g)  2 AlI 3 (s) 2 Cs(s) + F 2 (g)  2 CsF(s) 2SO 2 (g) + O 2 (g)  2SO 3 (g)

7.6 Ways to Classify Reactions Consider the driving forces Formation of solid Precipitation reaction Double replacement Formation of water Acid – base reaction Transfer of electrons Oxidation – Reduction reaction

7.7 Other Ways to Classify Reactions Combustion Organic Compound + O 2 (g) + heat  CO 2 (g) + H 2 O(g) Combination Element + element  compound Decomposition Compound  element + element

Examples Classify each of the following reactions in as many ways as possible 2 K(s) + Cl 2 (g)  2 KCl(s) BaCl 2 (aq) + Na 2 SO4(aq)  BaSO 4 (s) + 2 NaCl(aq) HNO 3 (aq) + NaOH(aq)  H 2 O(l) + NaNO 3 (aq) 2C 2 H 2 (g) + 5O 2 (g)  4 CO 2 (g) + 2H 2 O(l) PbO 2 (s)  Pb(s) + O 2 (g) Fe 2 O 3 (s) + 2 Al(s)  Al 2 O 3 (s) + 2 Fe(s)