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Chapter 7

Chapter 7. 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

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Chapter 7

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  1. Chapter 7 Reactions in Aqueous Solutions

  2. 7.1 Predicting Whether a Reaction Will Occur? • Driving Forces • Formation of solid • Formation of water • Transfer of electrons • Formation of a gas

  3. 7.2 Reactions in Which a Solid Forms? • Driving forces • Precipitation – formation of solid E.g Ba(NO3)2(aq) + K2CrO4(aq)  yellow solid yellow sol. colorless

  4. What happens When an Ionic Compound Dissolves in Water? • Ba(NO3)2(aq) => barium nitrate (white solid) has been dissolved in water • Containing Ba2+ (aq) and 2 NO3-(aq) ions • Strong electrolytes: A unit of substance that dissolves in water produces separate ions • E.g K2CrO4(aq) contains K+, K+ and CrO42- ions • => strong electrolyte

  5. Strong Electrolytes • Ba(NO3)2(aq) + K2CrO4(aq)  yellow solid

  6. Write out all possible formula that can be formed How to Decide What Products Form

  7. Using Solubility Rules • General Rules for Solubility of ionic compounds (salts) in water at 25oC • Most nitrate (NO3-) salts are soluble • Most salts of Na+, K+ and NH4+ are soluble • Most chloride, bromide, iodide (Cl-, Br- and I-) salts are soluble. Notable exceptions are AgCl, PbCl2 and Hg2Cl2. • Most sulfate salts are soluble. Notable exceptions are BaSO4, PbSO4 and CaSO4 • 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 (S2-), carbonate (CO32-) and phosphate (PO43-) salts are slightly soluble * The term insoluble and slightly soluble really mean the same thing

  8. 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.g AgNO3(aq) + KCl(aq)  white solid

  9. Examples • Using solubility rules to predict the product of Reactions • KNO3(aq) + BaCl2(g)  • Na2S(aq) + Cu(NO3)2(aq)  • KOH(aq) + Fe(NO3)2(aq) 

  10. 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

  11. 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

  12. 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, H2SO4, HClO4, HNO3 dissolved in H2O HCl(aq) -------------- H+(aq) + Cl-(aq)

  13. 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 H2O NaOH(aq) ----------------- Na+(aq) + -OH(aq)

  14. 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

  15. 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.g 2 Na(s) + Cl2(g)  2 NaCl(s)

  16. Examples • For each of the following reactions, show how electrons are gained and lost • 2 Al(s) + 3 I2(g)  2 AlI3(s) • 2 Cs(s) + F2(g)  2 CsF(s) • 2SO2(g) + O2(g)  2SO3(g)

  17. 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

  18. 7.7 Other Ways to Classify Reactions • Combustion • Organic Compound + O2(g) + heat  CO2(g) + H2O(g) • Combination • Element + element  compound • Decomposition • Compound  element + element

  19. Examples • Classify each of the following reactions in as many ways as possible • 2 K(s) + Cl2(g)  2 KCl(s) • BaCl2 (aq) + Na2SO4(aq)  BaSO4 (s) + 2 NaCl(aq) • HNO3(aq) + NaOH(aq)  H2O(l) + NaNO3(aq) • 2C2H2 (g) + 5O2 (g)  4 CO2 (g) + 2H2O(l) • PbO2 (s)  Pb(s) + O2 (g) • Fe2O3(s) + 2 Al(s)  Al2O3(s) + 2 Fe(s)

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