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Chemical Reactions

Chemical Reactions. Chapter 4, part 2: Reactions in Aqueous Solution. The dissolving process. When a solid is put into a liquid, solute-solute attractions compete with solute-solvent & solvent-solvent attractions. Solubility. Let’s assume our solvent is water . . .

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Chemical Reactions

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  1. ChemicalReactions Chapter 4, part 2: Reactions in Aqueous Solution

  2. The dissolving process • When a solid is put into a liquid, solute-solute attractions compete with solute-solvent & solvent-solvent attractions

  3. Solubility • Let’s assume our solvent is water . . . • If solute-water attractions > solute-solute & water-water attractions, solute particles are pulled out one by one into the water: • The solute is SOLUBLE in water

  4. Solubility • But if solute-water attractions < solute-solute & water-water attractions, solute particles remain together: • The solute is INSOLUBLE in water

  5. Solution conductivity • Solution conductivity depends on type of solute particles

  6. Solution conductivity • Ionic solutes (salts) made of cations & anions • Ions DISSOCIATE (separate) during dissolving • Molecular solutes made of molecules • each solute particle that moves into solution is identical

  7. Solution conductivity • Solutions of ionic solutes contain independent mobile ions • Solution conducts electricity • Solute is an ELECTROLYTE • Solute also conducts when melted, but not when solid (ions can’t move)

  8. Solution conductivity • Solutions of most molecular solutes contain independent neutral molecules • Solution does not conduct electricity • Solute is a NONELECTROLYTE • Solute also does not conduct when melted or solid

  9. Solution conductivity • Some molecular compounds (acids & bases) react with water to produce ions, as if they dissociated • A few acids & bases do this very well, producing lots of ions = strong electrolytes, strong acids & bases • Most acids & bases do this weakly, producing a few ions = weak electrolytes, weak acids & bases

  10. Strong acids & bases STRONG BASES STRONG ACIDS All other acids & bases are WEAK

  11. Predicting Electrolytes • All soluble salts and strong acids & bases are strong electrolytes • Weak acids & bases are weak electrolytes • All other molecular compounds are nonelectrolytes

  12. Precipitation reactions

  13. Precipitation reactions spectator ions precipitate

  14. Predicting Precipitation Reactions • To predict whether a precipitate will form, you need to know which compounds are soluble (no ppt) and which are insoluble (ppt forms) • Memorize the guidelines in Table 4.1 on page 154 in your text! • Add this guideline: CrO42– acts like SO42–

  15. Solubility guidelines: soluble • Compounds of these ions are generally soluble and do NOT form precipitates: • Alkali metals (group 1A, except Li1+) • Ammonium (NH41+) • Nitrates (NO31–) and acetates (C2H3O21–) • Chlorides, bromides, iodides, except Pb2+, Ag1+, Hg22+ • Sulfates (SO42–) & chromates (CrO42–), except Sr2+, Ba2+, Pb2+, and Hg22+

  16. Solubility guidelines: insoluble • Compounds of these ions are generally insoluble and DO form precipitates: • Hydroxides (OH1–) and Sufides (S2–) • Alkali metals (group 1A) and ammonium (NH41+) are soluble • Sulfides of group 2A metals are generally soluble • Hydroxides of Ca2+, Sr2+, and Ba2+ are soluble • Carbonates (CO32–) and phosphates (PO43–)

  17. Net ionic equations • AgNO3 (aq) + NaCl (aq)  • AgNO3 (aq) + NaCl (aq)  AgCl + NaNO3 • AgNO3 (aq) + NaCl (aq)  AgCl (s) + NaNO3 (aq) • The (aq) substances are dissociated: Ag1+ (aq) + NO31– (aq) + Na1+ (aq) + Cl1– (aq)  AgCl (s) + Na1+ (aq) + NO31– (aq)

  18. Net ionic equations Na1+ and NO31– are spectator ions Ag1+ (aq) + NO31– (aq) + Na1+ (aq) + Cl1– (aq)  AgCl (s) + Na1+ (aq) + NO31– (aq) • The net ionic equation omits spectators: Ag1+ (aq) + Cl1– (aq)  AgCl (s)

  19. Is that really a spectator? • An ion is a spectator if and only if it is in exactly the same form in the products and reactants: Na2CO3 (aq) + BaCl2 (aq)  BaCO3 (s) + 2 NaCl (aq) CO32– (aq) + Ba2+ (aq)  BaCO3 (s) Na2CO3 (s) + 2 HCl (aq)  2 NaCl (aq) + H2O + CO2 (g) Na2CO3 (s) + 2 H1+ (aq)  2 Na1+ (aq) + H2O + CO2 (g) Only Cl1– is a spectator Na1+ is not a spectator because it was (s), then (aq)

  20. Examples • Indicate whether a ppt forms and if so, complete the reaction as a balanced net ionic equation: • AlCl3 (aq) + KOH (aq)  • K2SO4 (aq) + FeBr3 (aq)  • CaI2 (aq) + Pb(NO3)2 (aq)  • Na3PO4 (aq) + AlCl3 (aq)  • Al2(SO4)3 (aq) + BaCl2 (aq)  • (NH4)2CO3 (aq) + Pb(NO3)2 (aq) 

  21. Acids • Acids produce H3O1+ in aqueous solution: • Strong acids are molecular compounds that react completely with water to produce H3O1+: HCl (g) + H2O  H3O1+ (aq) + Cl1– (aq) • For convenience, we often show acids as simply dissociating to produce H1+: HCl (g)  H1+ (aq) + Cl1– (aq) • There are only 6 strong acids (memorize them, pg 161)

  22. Acids • Acids produce H3O1+ in aqueous solution: • Weak acids are molecular compounds that react incompletely with water: HC2H3O2 (aq) + H2O  H3O1+ (aq) + C2H3O21– (aq) HC2H3O2 (aq)  H1+ (aq) + C2H3O21– (aq) • All acids that are not strong are weak

  23. Acids • H1+ is a proton • The reaction with water is proton transfer: HCl (g) + H2O  H3O1+ (aq) + Cl1– (aq) HC2H3O2 (aq) + H2O  H3O1+ (aq) + C2H3O21– (aq) • Acids with one H1+ to transfer are monoprotic acids HCl HNO3 HC2H3O2 • Acids with more than one H1+ to transfer are polyprotic acids H2SO4 H3PO4 H2C3H2O4

  24. Acids • Polyprotic acids produce H3O1+ in steps: H2SO4 (aq) + H2O  H3O1+ (aq) + HSO41– (aq) HSO41– (aq) + H2O  H3O1+ (aq) + SO42– (aq) • For H2SO4 the first step is strong and the second weak H2C2O4 (aq) + H2O  H3O1+ (aq) + HC2O41– (aq) HC2O41– (aq) + H2O  H3O1+ (aq) + C2O42– (aq) • For all other polyprotic acids, the first step is weak and the second step is weaker

  25. Bases • Bases produce OH1– in aqueous solution: • Strong bases are ionic hydroxide compounds that are completely dissociated in water: NaOH (s)  Na1+ (aq) + OH1– (aq) • The strong bases are the hydroxides of group 1A and 2A metals (memorize them) • Weak bases are molecular compounds that react incompletely with water to produce OH1–: NH3 (aq) + H2O (l)  NH41+ (aq) + OH1– (aq) • Amines (–NH2 ) are weak bases

  26. Neutralization • Acids and bases neutralize each other • The H1+ from the acid is transferred to the base: HCl (aq) + NaOH (aq)  HOH (l) + NaCl (aq) H1+ (aq) + OH1– (aq)  HOH (l) • The base is not always an OH1– compound: HCl (aq) + NH3 (aq)  NH41+ (aq) + Cl1– (aq) H1+ (aq) + NH3 (aq)  NH41+ (aq)

  27. Neutralization and net ionic equations • It is important to recognize strong acids and bases when writing net ionic equations HCl (aq) + NaOH (aq)  H2O (l) + NaCl (aq) H1+ (aq) + OH1– (aq)  H2O (l) HC2H3O2 (aq) + NaOH (aq)  H2O (l) + NaC2H3O2 (aq) HC2H3O2 (aq) + OH1– (aq)  H2O (l) + C2H3O21– (aq) The weak acid is not significantly ionized, so acetate is not a spectator (even if it is aq)

  28. Examples • Write the molecular and net ionic equations for • HNO3 + NaOH  • HF + KOH  • HC2H3O2 + NH3 • H2SO4 + Ba(OH)2 • H2C2O4 + NaOH  • HCHO2 + Ca(OH)2

  29. Gas forming reactions • Some neutralizations produce a gas: • CO32– + 2 H1+ H2CO3 • H2CO3 is unstable and decomposes immediately • H2CO3 CO2 (g) + H2O • The overall reaction is • CO32– + 2 H1+ CO2 (g) + H2O

  30. burning sulfur smell rotten egg smell ammonia Gas forming reactions • Memorize these gas-formers (pg 166): • SO32– + 2 H1+ SO2 (g) + H2O • HSO31– + H1+ SO2 (g) + H2O • CO32– + 2 H1+ CO2 (g) + H2O • HCO31– + H1+ CO2 (g) + H2O • S2– + 2 H1+ H2S (g) • NH41+ + OH1– NH3 (g) + H2O

  31. Redox • Redox (oxidation-reduction) reactions are those in which electrons are transferred • The loss of electrons is oxidation • The gain of electrons is reduction • LEO says GER

  32. Oxidation states • The oxidation state (O.S.) or oxidation number is a convenient but artificial way to describe the electron environment around an atom • It is related to the number of electrons gained, lost, or apparently used in forming compounds • Oxidation states are assigned using the rules on page 169 of your text (memorize these in order)

  33. Assigning oxidation states 1. The O.S. of each atom in an element is zero 2. The O.S. of a monoatomic ion is equal to its charge 3. The total of the O.S. of all atoms in any species (formula unit, molecule or ion) equals the charge on that species 4. In compounds, metals always have a positive O.S. • Group 1A metals are always O.S. +1 and Group 2A metals are always O.S. +2 5. For nonmetals in compounds, • the O.S. of fluorine is –1. • the O.S. of hydrogen is +1. • the O.S. of oxygen is –2. 6. In binary compounds, the O.S. of a Group 7A element is –1, Group 6A element –2, and Group 5A element –3.

  34. Examples • What is the oxidation state of each element in • S8 Cr2O72– Cl2O KO2 • What is the oxidation state of each element in • S2O32– Hg2Cl2 KMnO4 H2CO

  35. Redox • In a redox reaction, atoms change (O.S.) • The element oxidized loses electrons and its O.S. becomes more positive • The element reduced gains electrons and its O.S. becomes more negative

  36. Examples • Which of these reactions is a redox reaction? Identify the species oxidized and reduced. HCl (aq) + NaOH (aq)  H2O + NaCl (aq) 2 Pb(NO3)2 (s)  2 PbO (s) + 4 NO2 (g) + O2 (g) NH4Cl (s) + NaOH (aq)  NH3 (g) + H2O + NaCl (aq) • Identify the species oxidized and reduced: 5 VO2+ (aq) + MnO41– (aq) + H2O  5 VO21+ (aq) + Mn2+ (aq) + 2 H1+ (aq)

  37. Agents of Oxidation and Reduction • An agent makes something happen • An oxidizing agent makes oxidation happen by being reduced • A reducing agent makes reduction happen by being oxidized • The “agent” is the entire species in which the oxidized or reduced atom appears

  38. Examples • Identify the elements oxidized & reduced and the oxidizing & reducing agents in • 2 NO2 (g) + 7 H2 (g)  2 NH3 (g) + 4 H2O (g) • 5 H2O2 (aq) + 2 MnO41– (aq) + 6 H1+ (aq)  8 H2O + 2 Mn2+ (aq) + 5 O2 (g) • S2O32– (aq) + 4 Cl2 (aq) + 5 H2O  2 HSO41– (aq) + 8 H1+ (aq) + 8 Cl1– (aq) • 6 Fe2+ (aq) + 14 H1+ (aq) + Cr2O72– (aq)  6 Fe3+ (aq) + 2 Cr3+ (aq) + 7 H2O • S2O32– (aq) + 2 H1+ (aq)  S (s) + SO2 (g) + H2O

  39. Reactions • You now know how to write & balance 4 types of reactions • combustion • precipitation • acid-base neutralization • gas-forming • and how to recognize redox reactions

  40. Reaction quizzes • Reaction quizzes (RQ) will replace NQ • I give you the names of the reactants • You write the formulas of the reactant and the formulas of the products • Cross out spectator ions • Write the balanced net ionic equation • You need not include state symbols such as (aq) or (g)

  41. Example • Solutions of silver nitrate and potassium chloride are mixed • AgNO3 (aq) + KCl (aq)  • AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq) • cross out K1+ and NO31– as spectators • Ag1+ + Cl1– AgCl • check: it is already balanced

  42. Example • Hydrochloric acid solution is added to solid sodium hydrogen carbonate • HCl (aq) + NaHCO3 (s)  • HCl (aq) + NaHCO3 (s)  H2O + CO2 (g) + NaCl (aq) • cross out Cl1– as a spectator • do NOT cross out Na1+ because NaHCO3 is solid • H1+ + NaHCO3 H2O + CO2 + Na1+ • check: it is already balanced

  43. Example • Ethoxyethane burns in air • C2H5OC2H5 + O2 • C2H5OC2H5 + O2 CO2 + H2O • there are no spectators in combustion • balance • C2H5OC2H5 + 6 O2 4 CO2 + 5 H2O

  44. Example • Solutions of nitrous acid and sodium hydroxide are mixed • HNO2 (aq) + NaOH (aq)  • HNO2 (aq) + NaOH (aq)  H2O + NaNO2 (aq) • cross out Na1+ as a spectator • do NOT cross out NO21– because HNO2 is a weak acid and is not significantly dissociated in solution! • HNO2 + OH1– H2O + NO21– • check: it is balanced

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