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Precipitation Reactions

Precipitation Reactions. Example: Write the complete balanced molecular equation for the reaction between Na 3 PO 4 (aq) and CaCl 2 (aq). Write the complete ionic equation and the net ionic equation for the reaction. Ions: Na + PO 4 3- Ca 2+ Cl -. Possible Products: NaCl

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Precipitation Reactions

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  1. Precipitation Reactions Example: Write the complete balanced molecular equation for the reaction between Na3PO4 (aq) and CaCl2 (aq). Write the complete ionic equation and the net ionic equation for the reaction. Ions: Na+ PO43- Ca2+ Cl- Possible Products: NaCl Ca3(PO4)2

  2. Precipitation Reactions Solubilities: NaCl Ca3(PO4)2 (aq) soluble (s) insoluble

  3. Precipitation Reactions Unbalanced Molecular Eq’n: Na3PO4(aq) + CaCl2(aq) Ca3(PO4)2(s) + NaCl (aq) Balanced Molecular Eq’n: 2 Na3PO4(aq) + 3 CaCl2(aq)Ca3(PO4)2(s) + 6 NaCl(aq)

  4. Precipitation Reactions • Complete Ionic Equation: 6 Na+ (aq)+ 2PO43-(aq)+ 3 Ca2+(aq)+ 6Cl- (aq)Ca3(PO4)2 (s)+ 6 Na+ (aq) + 6 Cl- (aq) • Net Ionic Equation: 3 Ca2+(aq) + 2 PO43-(aq) Ca3(PO4)2(s)

  5. Precipitation Reactions Example: Write the equation for the reaction between NH4Cl (aq) and CaSO4 (aq). Ions: NH4+ Cl-, Ca2+ SO42- Possible Products: (NH4)2SO4 CaCl2

  6. Precipitation Reactions Solubility of Products: (NH4)2SO4 CaCl2 (aq) soluble (aq) soluble

  7. Precipitation Reactions NH4Cl (aq) + CaSO4 (aq) No Reaction • If both products are soluble (aq), then NO REACTION is presumed to occur between the ionic compounds involved. • The reaction mixture simple contains a mixture of soluble ions. • On your exam, you should be able to indicate whether a reaction occurs or not. • If it occurs, write the formulas for the products and balance. • If no reaction occurs, write the words, “No Reaction”

  8. Acids & Bases • Acids: • substances that ionize in aqueous solution to form one or more hydrogen ions (H+) • increase the concentration of H+ ions in solution • Acids are sometimes called proton donors. • H+ is often called a proton

  9. Acids & Bases • Examples of Acids: • HCl hydrochloric acid • HNO3 nitric acid • HC2H3O2 acetic acid • H2SO4 sulfuric acid • H3PO4 phosphoric acid Note: Acids can form different numbers of H+ ions!

  10. Acids & Bases • Monoprotic acids • have one H in the formula • form a single H+ ion when they ionize HNO3 (aq) H+(aq) + NO3- (aq)

  11. Acids & Bases • Diprotic acids • have two H’s in the formula • can form two H+ ion when they ionize completely H2SO4 (aq) 2H+(aq) + SO42- (aq) • Polyprotic acids: • Have two or more H’s in the formula • Form two or more H+ ions when they ionize completely

  12. Acids & Bases • Bases: • substances that accept (react with) H+ ions. • any substance that increases the OH- concentration when added to water • Examples: • Hydroxide ion (OH-) OH- (aq) + H+ (aq)  H2O (l)

  13. Acids & Bases • Examples (cont) • Common hydroxide containing bases: • NaOH, KOH, Ca(OH)2 • Note: These are strong electrolytes! • NaOH (aq)  Na+ (aq) + OH- (aq)

  14. Acids & Bases • Examples (cont): • Ammonia (NH3) • Does not contain OH- • Accepts H+ ion from water and increases the OH- concentration in the water NH3(aq) + H2O(l) NH4+ (aq) + OH- (aq) NH3 is a weak electrolyte!!

  15. Acids & Bases • Strong Acid: • an acid that is a strong electrolyte • ionizes completely in solution • Weak Acid: • an acid that is a weak electrolyte • an acid that does not ionize completely

  16. Acids & Bases • Strong acids: • Know the names and formulas of the 7 common strong acids: • HCl (aq) hydrochloric acid • HBr (aq) hydrobromic acid • HI (aq) hydroiodic acid • HClO3 chloric acid • HClO4 perchloric acid • HNO3 nitric acid • H2SO4 sulfuric acid

  17. Acids & Bases • Examples of Weak Acids • HF (aq) hydrofluoric acid • H3PO4 phosphoric acid • HC2H3O2 acetic acid

  18. Acids & Bases • Strong Base: • a base that is a strong electrolyte • ionizes completely in solution • Weak Base: • a base that is a weak electrolyte • does not ionize completely in solution

  19. Acids & Bases • Strong Bases:Know the names and formulas of the strong bases • Alkali metal (1A) hydroxides • LiOH lithium hydroxide • NaOH sodium hydroxide • KOH potassium hydroxide • RbOH rubidium hydroxide • CsOH cesium hydroxide

  20. Acids & Bases • Strong bases to know (con’t): • Heavy alkaline earth metal (2A) hydroxides • Ca(OH)2 calcium hydroxide • Sr(OH)2 strontium hydroxide • Ba(OH)2 barium hydroxide

  21. Acids & Bases • Examples of Weak Bases: • ammonia (NH3) • sodium bicarbonate (NaHCO3) • baking soda • a component of Alka-Seltzer

  22. Acid-Base Reactions • Reactions between acids and bases are called neutralization reactions. • The products of these reactions have very different properties than the reactants. HCl (aq) + NaOH (aq) H2O (l) + NaCl (aq) Sharp sour bitter slippery salt

  23. Acid-Base Reactions • Salt: • any ionic compound whose cation comes from a base and whose anion comes from an acid • An ionic compound that is neither an acid nor a base • In general, acid + metal hydroxide  a salt + water

  24. Acid-Base Reactions • Neutralization reactions are a type of metathesis reaction. • To predict the products: • identify the ions present • exchange anions • write the correct formulas for the products • write a balanced equation

  25. Acid-Base Reactions Example: Write the balanced equation for the reaction between HBr (aq) and Ca(OH)2 (aq). Ions: H+ Br- Ca2+ OH- Possible Products: H-OH = H2O CaBr2 2HBr (aq) + Ca(OH)2(aq) CaBr2(aq) + 2H2O(l)

  26. Acid-Base Reactions • Notice that you can also write complete and net ionic equations for acid-base reactions: • Molecular equation: 2HBr (aq) + Ca(OH)2(aq) CaBr2(aq) + 2H2O (l) • Complete ionic equation: 2 H+ (aq) + 2Br-(aq) + Ca2+(aq) + 2 OH-(aq) Ca2+(aq) + 2 Br-(aq) + 2 H20 (l)

  27. Acid-Base Reactions • Net ionic equation: H+(aq) + OH-(aq) H2O (l) • Note:This is the net ionic equation between any strong acid and strong base. • Complete ionic equation: 2 H+ (aq) + 2Br-(aq) + Ca2+(aq) + 2 OH-(aq) Ca2+(aq) + 2 Br-(aq) + 2 H20 (l)

  28. Acid-Base Reactions Example: Write the balanced equation for the reaction between Mg(OH)2 (s) and HCl (aq). Ions: Mg2+ OH- H+ Cl- Products: MgCl2 (aq) H-OH = H2O (l)

  29. Acid-Base Reactions Molecular Equation: Mg(OH)2(s) + 2 HCl (aq) MgCl2(aq) + 2 H2O (l) Complete Ionic Equation: Mg(OH)2(s) + 2 H+(aq) + 2 Cl-(aq) Mg2+(aq) + 2 Cl-(aq) + 2 H2O (l) Net ionic equation: Mg(OH)2(s) + 2 H+(aq) Mg2+(aq) + 2 H2O (l)

  30. Acid-Base Reactions • There are many bases that do not contain OH- • Na2S • NaCN • NaHCO3 • These bases react with acids to form gaseous products.

  31. Acid-Base Reactions Examples: • Na2S(aq) + 2HCl (aq)H2S(g) + 2 NaCl (aq) • HCl (aq) + NaHCO3(aq)  NaCl (aq) + H2CO3(aq) but: H2CO3(aq) H2O (l) + CO2 (g) so: HCl (aq) + NaHCO3(aq)  NaCl (aq) + H2O (l) + CO2 (g)

  32. Oxidation-Reduction Reactions • Precipitation Reactions: • ions combine to form insoluble products • Neutralization Reactions: • H+ ions and OH- ions combine to form H2O • Oxidation-Reduction (Redox) Reactions: • Atoms or ions give or accept electrons

  33. Redox Reactions • Corrosion of your car battery terminal is caused by a reaction between the metal terminal, oxygen, and the battery acid, H2SO4. • This reaction is a redox reaction.

  34. Redox Reactions • Oxidation-Reduction Reactions (Redox Reactions) • reactions that involve the transfer of electrons between two reactants • an element in one reactant is oxidized while an element in another reactant is reduced Mg (s) + 2 H+ (aq) Mg2+(aq) + H2(g) oxidized reduced

  35. Redox Reactions • Oxidation: • the loss of electrons • chemical species becomes more positively charged • the gain of oxygen • the loss of hydrogen

  36. Redox Reactions • Reduction: • the gain of electrons • the chemical species becomes more negatively charged • the gain of hydrogen • the loss of oxygen

  37. Redox Reactions • LEO: • Lose Electrons Oxidation • GER: • Gain Electrons Reduction GER LEO LEO says GER

  38. Redox Reactions • Oil : • OxidationInvolves Lossof e- • Rig : • ReductionInvolves Gainof e- Oil Rig

  39. Redox Reactions • Electrons are not explicitly shown in chemical equations. • Oxidation Numbers are used to keep track of electrons gained and lost during redox reactions. • Oxidation number • a hypothetical number assigned to an individual atom present in a compound using a set of rules. • May be positive, negative, or zero

  40. Rules for Oxidation Numbers • Oxidation numbers are always reported for individual atoms or ions not groups of atoms or ions!!!!!!!!!!! • For an atom in its elemental form, the oxidation number is always zero. • H2: oxidation # = 0 for each H atom • Cu: oxidation number = 0 • Cl2: oxidation # = 0 for each Cl atom

  41. Rules for Oxidation Numbers • For any monoatomic ion, oxidation # = charge on ion • K+ oxidation # = +1 • Cl- oxidation # = -1 • S2- oxidation # = -2

  42. Rules for Oxidation Numbers • Group 1A MetalCations: • Always +1 • Group 2A Metal Cations: • Always +2 • Hydrogen (H) • +1 when bonded to nonmetals • -1 when bonded to metals

  43. Rules for Oxidation Numbers • Oxygen (O) • -1 in peroxides (O22-) • -2 in all other compounds • Fluorine (F) • always -1

  44. Rules for Oxidation Numbers • The sum of the oxidation numbers of all atoms in any chemical species (ion or neutral compound) is equal to the charge on that chemical species • H2O: 1 + 1 + -2 = 0 • MgCl2: 2 + -1 + -1 = 0 • MnO4- : 7 + -2 + -2 + -2 + -2 = -1

  45. Oxidation Numbers • For many compounds, you will be able to directly apply the rules to determine the oxidation number of all atoms except for one. • Use the last two rules to determine the oxidation number of that last element.

  46. Oxidation Numbers Example: Determine the oxidation state of all elements in SO3. Is it elemental? No Are any monoatomic ions present? No Which elements have rules? O = -2 Set up an equation to find the remaining oxidation number. S + 3(-2) = 0  S = +6

  47. Oxidation Numbers Example: Determine the oxidation number of Mn and O in MnO4-. Is it elemental? No Are any monoatomic ions present? No Which elements have rules? O = -2 Set up an equation to find the remaining oxidation number. Mn + 4(-2) = -1 so Mn = +7

  48. Oxidation Numbers Example: Determine the oxidation state of all elements in NaNO3 Is it elemental? No Are any monoatomic ions present? Na+ Which elements have rules? Na = +1, O = -2 Set up an equation to find the remaining oxidation number. 1 + N + 3(-2) = 0  N = +5

  49. Oxidation Numbers Example: Determine the oxidation number of P in HPO42-

  50. Oxidation Numbers Example: Determine the oxidation state of all elements in Cr2O72-.

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