1. FACTORS (driving forces) THAT CAUSE REACTION TO OCCUR: Formation of a solid Formation of water Transfer of electrons Formation of a gas *Review polyatomic ions; these will be needed to write chemical equations in this chapter.

2. Precipitation: Reactions in which solid forms One of the reaction products is insoluble Pb(NO 3 ) 2 (aq) + 2 KI(aq)  2 KNO 3 (aq) + PbI 2 (s) Many reactions are done by mixing aqueous solutions of electrolytes (solutions containing ions) together. - When this is done, often a reaction will take place as a result of the cations and anions in the two solutions exchanging. - If the ion exchange results in forming a compound that is insoluble in water, it will come out of solution as a precipitate. The reaction is known as a precipitation reaction.

3. No Precipitate Formation Results in No Reaction Occurring KI(aq) + NaCl(aq)  KCl(aq) + NaI(aq) If all ions are still present, then there is no reaction

4. 1.Most (NO 3 - ) salts are soluble. 2.Most salts of Na +, K +, and NH 4 + are soluble. 3.Most chloride salts are soluble. Exceptions: AgCl, PbCl 2, and Hg 2 Cl 2 4.Most sulfate salts are soluble. Exceptions: BaSO 4, PbSO 4, and CaSO 4 5.Most hydroxide compounds are only slightly soluble. Exceptions: NaOH and KOH. Ba(OH) 2 and Ca(OH) 2 are only slightly soluble. 6.Most sulfides (S 2- ), carbonate (CO 3 2- ), and phosphate (PO 4 3- ) salts are only slightly soluble. –note: slightly soluble  insoluble Solubility Rules: See Table 7.1

5. Process for Predicting the Products of a Precipitation Reaction 1.Determine what ions are present in each aqueous reactant. 2.Exchange the Ions as if a reaction occurred: –(+) ion from one reactant with (-) ion from another. 3.Balance the charges of the newly combined ions to get a formula for each product. 4.Balance the Equation: –count all of the atoms in the reactants and products and make sure the same number and kinds of atoms are in both reactants and products. 5.Determine the Solubility of Each Product in Water: –use the Solubility Rules –if a product is insoluble or slightly soluble, it will precipitate. –if neither product will precipitate, there is no reaction.

6. Example - When an aqueous solution of sodium carbonate is added to an aqueous solution of copper(II) chloride, a white solid forms 1.Write the formulas of the reactants Na 2 CO 3 (aq) + CuCl 2 (aq)  2.Determine the ions present when each reactant dissociates (Na + + CO 3 2- ) + (Cu +2 + Cl - )  3.Exchange the Ions (Na + + CO 3 2- ) + (Cu +2 + Cl - )  (Na + + Cl - ) + (Cu +2 + CO 3 2- ) 4.Write the formulas of the products –cross charges and reduce Na 2 CO 3 (aq) + CuCl 2 (aq)  NaCl + CuCO 3 5.Balance the Equation Na 2 CO 3 (aq) + CuCl 2 (aq)  NaCl + CuCO 3 6.Determine the solubility of each product NaCl is soluble and CuCO 3 is insoluble 7.Write an (s) after the insoluble products and a (aq) after the soluble products Na 2 CO 3 (aq) + CuCl 2 (aq)  NaCl(aq) + CuCO 3 (s)

7. Additional Examples 1.In the following reaction a white solid forms, identify the solid and write the balanced equation: Ba(NO 3 ) 2 (aq) + Na 2 SO 4 (aq)  2.In the following reaction a yellow solid forms, identify the solid and write the balanced equation: Pb(NO 3 ) 2 (aq) + KI(aq)  3.Predict what will happen when the following solutions are mixed and write the balanced equation: Ba(NO 3 ) 2 (aq) + K 3 PO 4 (aq)  4.Predict what will happen when the following solutions are mixed and write the balanced equation: Na 2 SO 4 (aq) + KCl(aq) 

8. Reactions in Aqueous Solutions Equations which describe the chemicals put into the water and the product molecules are called molecular equations: 2 KOH(aq) + Mg(NO 3 ) 2 (aq) →2 KNO 3 (aq) + Mg(OH) 2 (s) Equations which describe the actual dissolved species are called complete ionic equations: –aqueous electrolytes are written as ions soluble salts, strong acids, strong bases –insoluble substances and nonelectrolytes are written as molecules solids, liquids and gases are not dissolved, therefore they are written as molecules: 2K +1 (aq) + 2OH -1 (aq) + Mg +2 (aq) + 2NO 3 -1 (aq)  K +1 (aq) + 2NO 3 -1 (aq) + Mg(OH) 2 (s) Net Ionic Equation: 2OH -1 (aq) + Mg +2 (aq)  Mg(OH) 2 (s)

9. Conventional (Molecular) Equation HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq) Pros: Actual starting materials Easier to balance equations Easier for stoichiometry problems Cons: Does not accurately describe rxn, reactants or products Less useful with equilibrium problems

10. Total Ionic Equation H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) H 2 O (l) + Na + (aq) + Cl - (aq) Pros: Accurately describes rxn, reactants and products Cons: Cannot determine starting material Less useful with stoichiometry Difficult to balance equation Spectator ions on each side Long

11. Net Ionic Equation H + (aq) + OH - (aq) H 2 O (l) Pros: Indicates chemical change Concise Same as total ionic, but w/o spectator ions Cons: Cannot determine starting material The net ionic equation tells the chemistry that occurs, and nothing else (that is its purpose)

12. Writing Net Ionic Equations Conventional (Molecular) Equation Total Ionic Equation Net Ionic Equation HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq) H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) H 2 O (l) + Na + (aq) + Cl - (aq) H + (aq) + OH - (aq) H 2 O (l)

13. Writing Net Ionic Equations Na 2 SO 4(aq) + Pb(NO 3 ) 2(aq) PbSO 4(s) + 2 NaNO 3(aq) Conventional (Molecular) Equation Total Ionic Equation 2 Na + (aq) + SO 4 2- (aq) + Pb 2+ (aq) + 2 NO 3 - (aq) PbSO 4(s) + 2 Na + (aq) + SO 4 2- (aq) Net Ionic Equation SO 4 2- (aq) + Pb 2+ (aq) PbSO 4(s)

14. Reaction that form Water Acids and Bases Everyday acids and bases we use: AcidsBases Soft DrinksDetergents FruitsDrain-cleaning products Where have we heard and used an ‘acid’ or ‘base’ in this class?

15. What Are Acids and Bases? Every aqueous solution has mostly water molecules, but also has both hydronium ions (H + ) and hydroxide ions (OH - ) H+H+ OH - H2OH2O H2OH2O H2OH2O H2OH2O H+H+ H2OH2O H2OH2O H2OH2O H2OH2O H+H+ H2OH2O H2OH2O H2OH2O H2OH2O H+H+ H+H+ An acidic solution is one where there are more hydronium ions (H + ) than hydroxide ions (OH - ) A basic solution is one where there are more hydroxide ions (OH - ) than hydronium ions (H + ) This is the Arrhenius theory

16. Acidic, Basic, or Neutral? H+H+ OH - H2OH2O H2OH2O H2OH2O H2OH2O H+H+ H2OH2O H2OH2O H2OH2O H2OH2O H+H+ H2OH2O H2OH2O H2OH2O H2OH2O H+H+ H+H+ Neutral [H + ] ~ [OH - ] Acidic [H + ] > [OH - ] Basic [OH - ] > [H + ] or [H + ] < [OH - ]

17. Acid, Base Strength Vinegar and sulfuric acid are both acids One we consume, one will kill us Baking soda and drain cleaners are both bases One we put in the refrigerator, the other will make us sick Acids and bases come in a variety of strengths - some are strong and others are weak

18. Examples of Strong Acids Recognize the following strong acids: *Nitric acid (HNO 3 ) *Sulfuric acid (H 2 SO 4 ) *Hydrochloric acid (HCl) Hydrobromic acid (HBr) Hydroiodic acid (HI) Chloric acid (HClO 3 ) *Perchloric acid (HClO 4 )

19. Strong Acid: very conductive solution as it ionizes

20. Weak Acid (or base); barely lights up as there are few ions in solution

21. Example of Weak Acids: Most organic acids (acetic, oxalic, citric, fatty acids) HF Hydrofluoric acid H 2 CO 3 Carbonic acid (soda pop) H 3 PO 4 Phosphoric acid (also in Coke) H 3 BO 3 Boric acid (used in eye washes)

22. Acid-Base Reactions Acid-Base reactions are also called neutralization reactions because the acid and base neutralize each other’s properties to form a salt and water. In the reaction of an acid with a base, the H + from the acid combines with the OH - from the base to make water. The cation from the base combines with the anion from the acid to make the salt acid + base  salt + water 2 HNO 3 (aq) + Ca(OH) 2 (aq)  Ca(NO 3 ) 2 (aq) + 2 H 2 O(l) the net ionic equation for an Acid-Base reaction is H +1 (aq) + OH -1 (aq)  H 2 O(l) –as long as the salt that forms is soluble in water

23. Oxidation/Reduction Atoms undergo changes in oxidation state. Atoms that lose electrons become oxidized and atoms that gain electrons become reduced. They always accompany each other: 2Mg(s) + O 2 → 2MgO(s) you cannot have one without the other

24. Reactions of Metals with Nonmetals Reactions of Metals with Nonmetals (Oxidation-Reduction ) Metals react with nonmetals to form ionic compounds –ionic compounds are solids at room temperature. The metal loses electrons and becomes a cation –the metal undergoes oxidation. The nonmetal gains electrons and becomes an anion –the nonmetal undergoes reduction. In the reaction, electrons are transferred from the metal to the nonmetal: 2 Na(s) + Cl 2 (g)  2 NaCl(s)

25. Oxidation-Reduction Reactions Any reaction that has an element that is uncombined on one side and combined on the other is a redox reaction: –uncombined = free element –2 CO + O 2  2 CO 2 –2 N 2 O 5  4 NO 2 + O 2 –3 C + Fe 2 O 3  3 CO + 2 Fe –Mg + Cl 2  MgCl 2 Any reaction where a cation changes charge is redox: –CuCl + FeCl 3  FeCl 2 + CuCl 2 –SnCl 2 + F 2  SnCl 2 F 2

26. Your Car Battery

27. Rust is the result of an Oxidation/Reduction Reaction

28. Gas Evolving Reactions “Reactions that form a gas” Some reactions form a gas directly from the ion exchange: K 2 S(aq) + H 2 SO 4 (aq)  K 2 SO 4 (aq) + H 2 S(g) Other reactions form a gas by the decomposition of one of the ion exchange products into a gas and water: K 2 SO 3 (aq) + H 2 SO 4 (aq)  K 2 SO 4 (aq) + H 2 SO 3 (aq) H 2 SO 3  H 2 O(l) + SO 2 (g)

29. Compounds that Undergo Gas Evolving Reactions Reactant Type Reacting With Ion Exchange Product Decom- pose? Gas Formed Example metal n S, metal HS acidH2SH2SnoH2SH2S K 2 S(aq) + 2HCl(aq)  2KCl(aq) + H 2 S(g) metal n CO 3, metal HCO 3 acidH 2 CO 3 yesCO 2 K 2 CO 3 (aq) + 2HCl(aq)  2KCl(aq) + CO 2 (g) + H 2 O(l) metal n SO 3 metal HSO 3 acidH 2 SO 3 yesSO 2 K 2 SO 3 (aq) + 2HCl(aq)  2KCl(aq) + SO 2 (g) + H 2 O(l) (NH 4 ) n anionbaseNH 4 OHyesNH 3 KOH(aq) + NH 4 Cl(aq)  KCl(aq) + NH 3 (g) + H 2 O(l)

30. Types of Chemical Reactions: 1. Combination or Synthesis Reaction Two or more elements combine to form one molecule A + XAX 2 Cu + SCu 2 S 2 CO + O 2 2 CO 2 2 Mg + O 2 2 MgO HgI 2 + 2 KI K 2 HgI 4

31. 2.Decomposition Reactions One compound breaks down into simpler substances AXA + X 2 H 2 O2 H 2 + O 2

32. 3.Complete Oxidation (Burning or Combustion) Burning of C, H, O - containing molecules in the presence of oxygen C x H y O z + O 2 CO 2 + H 2 O C 3 H 8 + O 2 3 CO 2 + 4 H 2 O Propane + oxygen carbon dioxide + water

33. The complete oxidation of hydrocarbons by burning in the presence of oxygen gas always yields CO 2 and H 2 O as products! The combustion of glucose C 6 H 12 O 6 + 6O 2 → 6CO 2 + 6H 2 O Methane gas burns to produce carbon dioxide gas and gaseous water whenever something burns it combines with O 2 (g) CH 4 (g) + O 2 (g) → CO 2 (g) + H 2 O(g) H H C H H OO + O O C + O HH

34. Combustion Products To predict the products of a combustion reaction, combine each element in the other reactant with oxygen If the ReactantThe Combustion Product is contains CCO 2 (g) contains HH 2 O(g) contains SSO 2 (g) contains NNO(g) or NO 2 (g) contains metalM 2 O n (s)

35. 4.Double-Replacement Prec. Rxns Ions exchange with oppositely-charged ions to form a precipitate AX + BY AY + BX NaCl + AgNO 3 NaNO 3 + AgCl(s)

36. 5.Double-Replacement Neutralization Rxn The acid and base neutralize each other in a neutralization reaction and water is formed. An ionic compound called a salt is also formed HX + MOH H 2 O + MX acid base water salt HCl + NaOH H 2 O + NaCl

37. Double Displacement Reactions Two ionic compounds exchange ions during a reaction. It may be followed by decomposition of one of the products to produce a gas. X  Y  (aq) + A  B  (aq)  XB + AY Precipitation, acid-base and gas-evolving reactions are all examples of double displacement reactions.

38. Single-replacement reduction-oxidation reactions (redox) - reaction looks as if one element is replacing another A + BX AX +B Cu + AgNO 3 CuNO 3 + Ag 6.Redox Rxns What is reduction? What is oxidation? It is the gaining or losing of electrons - more later

39. Displacement of Copper by Zinc

40. Classifying Reaction Type Reaction Type Combination Decomposition Complete oxidation Single-replacement redox Double replacement Precipitation Neutralization Equation Type A + X AX AX A + X C x H y O z + O 2 CO 2 + H 2 O A + BX AX + B AX + BY AY + BX HX + MOH H 2 O + MX

41. Classifying Reaction Type Classify the following as one of the six reaction types Mg + N 2 C 3 H 7 CHO + O 2 HgO Zn(OH) 2 + H 2 SO 4 Combo, oxidation, decomp, 2 repl. neut.