1. Unit 7 Reactions in Solution Chem II Objectives  Describe the driving force for a chemical reaction.  Use generalizations to predict the products of simple reactions.  Describe reactions in solutions by writing molecular, complete ionic and net ionic equations.  Identify the key characteristics of the reactions between strong acids and strong bases  Identify the general characteristics of a reaction between a metal and a nonmetal.  Classify a reaction as one of five basic types: synthesis, decomposition, single displacement, double displacement or combustion. 1

2. Will a Reaction Occur? What causes reactions to occur? Reactants → Products Driving forces that make reactants go in the direction of the arrow include: – Formation of a solid – Formation of a gas – Formation of water – Transfer of electrons 2

3. Reactions which form solids One driving force for a chemical reaction is the formation of a solid, a process called precipitation. The solid formed is called a precipitate The reaction is called a precipitation reaction Example: Potassium chromate + barium nitrate ↑ Write equation for products 3

4. Precipitation Reactions To predict the identity of the precipitate, consider the possible products First look at the nature of each reactant in an aqueous solution Barium nitrate contains Ba 2+ and NO 3 - ions Typically, when a solid containing ions dissolves in water, the ions separate, moving around independently 4

5. Precipitation Reactions Ba(NO 3 ) 2 (aq) does not contain Ba(NO 3 ) 2 units The solution contains Ba 2+ and NO 3 - ions There are two NO 3 - ions for every Ba 2+ ion When each unit of a substance that dissolves in water produces separated ions, it is called a strong electrolyte Ba(NO 3 ) 2 is a strong electrolyte as is K 2 CrO 4 5

6. Ionic Compounds When ionic compounds dissolve, the resulting solution contains ions. We usually write these reactants as Ba(NO 3 ) 2 (aq) + K 2 CrO 4 (aq) → Products It is more appropriately shown as Ba 2+ (aq) + 2NO 3 - (aq) + 2K + (aq) + CrO 4 2- (aq) → Products The mixed solution contains four types of ions. 6

7. What products form? A solid compound must have a net zero charge Must contain cations and anions K + and Ba 2+ could not combine to form the solid Most ionic materials contain only two types of ions – one cation and one anion Possible compounds are K 2 CrO 4 Ba(NO 3 ) 2 K(NO 3 ) BaCrO 4 7

8. Predicting Products Need more information to determine which of these is more likely to be the precipitate K(NO 3 ) BaCrO 4 Potassium nitrate is a white solid. The chromate ion is yellow. So, the likely precipitate is BaCrO 4 and the unbalanced equation is Ba(NO 3 ) 2 (aq) + K 2 CrO 4 (aq) →KNO 3 (aq) + BaCrO 4 (s) 8

9. Solubility Vocabulary A soluble solid readily dissolves in water. An insoluble solid, or slightly soluble solid, means only a small amount of the solid dissolves in water. A salt means an ionic compound. 9

10. Solubility Rules Table 7.1, Z page 178 1.Most nitrate (NO 3 - ) salts are soluble. 2.Most salts of Na +, K + and NH 4 + are soluble. 3.Most chloride salts are soluble. Notable exceptions are AgCl, PbCl 2 and Hg 2 Cl 2. 4.Most sulfate salts are soluble. Notable exceptions are BaSO 4, PbSO 4 and CaSO 4. 5.Most hydroxide compounds are insoluble. Notable exceptions are NaOH, KOH, Ba(OH) 2 and Ca(OH) 2. 6.Most sulfide (S 2- ), carbonate (CO 3 2- ) and phosphate (PO 4 3- ) salts are insoluble. 10

11. Predicting Products - Z page 180 Step 1 – Write the reactants as they actually exist before any reaction occurs. Step 2 – Consider the various solids that could form. To do this, simply exchange the anions of the added salts. Step 3 – Use the solubility rules to decide whether a solid forms and to predict the identity of the solid. 11

12. Predicting Products Guided example When an aqueous solution of silver nitrate is added to an aqueous solution of potassium chloride, a white precipitate is formed. Identify the solid and write a balanced equation for the reaction. 12

13. Predicting Products When an aqueous solution of silver nitrate is added to an aqueous solution of potassium chloride, a white precipitate is formed. AgNO 3 (aq) + KCl (aq) → AgCl (?) + KNO 3 (?) Solubility rule 1 – NO 3 - is soluble Solubility rule 2 – K + is soluble Solubility rule 3 – AgCl is insoluble AgNO 3 (aq) + KCl (aq) → AgCl (s) + KNO 3 (aq) Equation is balanced 13

14. Predicting Products In class examples 1.KNO 3 (aq) and BaCl 2 (aq) 2.Na 2 SO 4 (aq) and Pb(NO 3 ) 2 (aq) 3.KOH (aq) and Fe(NO 3 ) 3 (aq) 14

15. Predicting Products Homework Page 202, problems 18 and 22 15

16. Reactions in Aqueous Solutions Returning to this equation, note that it is called a molecular equation. Ba(NO 3 ) 2 (aq) + K 2 CrO 4 (aq) →2KNO 3 (aq) + BaCrO 4 (s) It shows the complete formulas of all products and reactants. This does not give a clear picture of what occurs in solution. 16

17. Reactions in Aqueous Solutions Aqueous solutions of barium nitrate, potassium chromate and potassium nitrate contain the individual ions, not the complete molecules implied by the molecular equation. The complete ionic equation better represents the actual forms of reactants and products in solution. Ba 2+ (aq) + 2NO 3 - (aq) + 2K + (aq) + CrO 4 2- (aq) → BaCrO 4 (s) + 2K + (aq) + 2NO 3 - (aq) 17

18. Reactions in Aqueous Solutions The complete ionic equation reveals that only some of the ions participate in the reaction. Ions which do not participate in the reaction, such as K + and NO 3 -, are called spectator ions. The ions that participate in the reaction are Ba 2+ (aq) + CrO 4 2- (aq) → BaCrO 4 (s) This equation, called the net ionic equation, includes only those components that are directly involved on the reaction. 18

19. Molecular, Complete Ionic and Net Ionic Equations In class Z page 184 examples 7.3a and 7.3b 19

20. Equations 7.3a) Molecular equation NaCl (aq) + AgNO 3 (aq) → AgCl (s) + NaNO 3 (aq) Complete ionic equation Na + (aq) + Cl - (aq) + Ag + (aq) + NO 3 - (aq) → AgCl (s) + Na + (aq) + NO 3 - (aq) Net ionic equation Cl - (aq) + Ag + (aq) → AgCl (s) 20

21. Equations 7.3b ) Molecular equation 3KOH(aq) +Fe(NO 3 ) 3 (aq) → Fe(OH) 3 (s) + 3KNO 3 (aq) Complete ionic equation 3K + (aq) + 3OH - (aq) + Fe 3+ (aq) + NO 3 - (aq) → Fe(OH) 3 (s) + 3K + (aq) + 3NO 3 - (aq) Net ionic equation 3OH - (aq) + Fe 3+ (aq) → Fe(OH) 3 (s) 21

22. Molecular, Complete Ionic and Net Ionic Equations Homework Page 202 and 203, problem 26 22

23. Reactions that Form Water Acids – first associated with the sour taste of citrus fruits. Dervied from Latin word acidus meaning sour Essential nature of acids discovered by Svante Arrhenius in the late 1800’s Arrhenius proposed that an acid is a substance that produces H + ions when it is dissolved in water Studies show that when HCl, HNO 3 and H 2 SO 4 are placed in water nearly every molecule dissociates to give ions Because these substances are strong electrolytes that produce H + ions they are called strong acids 23

24. Reactions that Form Water Bases – characterized by bitter taste and slippery feel Arrhenius found that aqueous solutions that exhibit basic behavior always contain hydroxide ions (OH - ). He defined a base as a substance that produces OH - in water. Most common base in the lab is NaOH which dissolves in water to form Na + and OH - ions 24

25. Reactions that Form Water When strong acids and strong bases are mixed, the fundamental change that occurs is that H + and OH - ions react to form water H + (aq) + OH - (aq) → H 2 O (l) The tendency to form water is one of the driving forces for chemical reactions 25

26. Reactions that Form Water The reaction between hydrochloric acid and sodium chloride shown as a molecular equation is: HCl (aq) + NaOH (aq) → NaCl (aq) + H 2 O (l) 26

27. Reactions that Form Water Because HCl and NaOH exist as completely separated ions in water, the complete ionic equation can be written as: H + (aq) + Cl - (aq) + Na + (aq) + OH - (aq) → Na + (aq) + Cl - (aq) + H 2 O (l) 27

28. Reactions that Form Water Notice that Cl - and Na + are spectator ions so the net ionic equation is: H + (aq) + OH - (aq) → H 2 O (l) 28

29. Acid-Base Reactions In class Z Page 187 – Example 7.4 29

30. Acid-Base Reactions In class Page 187 – Example 7.4H Molecular equation: HNO 3 (aq) + KOH (aq) → H 2 O (l) + KNO 3 (aq) Complete ionic equation: H + (aq) + NO 3 - (aq) + K + (aq) + OH - (aq) → H 2 O (l) + NO 3 - (aq) + K + (aq) Net ionic equation: H + (aq) + OH - (aq) → H 2 O (l) 30

31. Acid-Base Reactions Note two things when looking at the reaction between HCl and NaOH and the reaction between HNO 3 and KOH: 1.The net ionic equation is the same H + (aq) + OH - (aq) → H 2 O (l) 2.Besides water, the second product formed is an ionic compound called a salt. The salt may precipitate or remain in solution, depending on solubility. How could you recover the salt if it is soluble? If it is insoluble? 31

32. Acid-Base Reactions Homework Page 203 problems 39a, 39c, 39d, 40a, 40c, 40d 32

33. Oxidation-Reduction Reactions In a reaction between a metal and a non- metal, electrons are transferred from the metal to the non-metal. Consider the reaction between sodium metal and chlorine gas 2Na (s) + Cl 2 (g) → 2NaCl (s) The product, sodium chloride, consists of Na + ions and Cl - ions. 33

34. Oxidation-Reduction Reactions The sodium in the product has a positive charge because it transferred an electron to the chlorine. The chlorine has a negative charge because it gained an electron from the sodium. A reaction in which electrons are transferred is called an oxidation-reduction reaction. 34

35. Oxidation-Reduction Reactions Consider the reaction between magnesium metal and oxygen. 2Mg(s) + O 2 (g) → 2 MgO(s) Each magnesium atom loses 2 electrons. Mg → Mg 2+ + 2e - Each oxygen atom gains 2 electrons O + 2e - → O 2- 35

36. Oxidation-Reduction Reactions Page 190 - Example 7.5a Show how electrons are gained and lost. 2 Al (s) + 3 I 2 (g) → 2AlI 3 (s) 36

37. Oxidation-Reduction Reactions Page 190 - Example 7.5a Show how electrons are gained and lost. 2 Al (s) + 3 I 2 (g) → 2AlI 3 (s) Al → Al 3+ + 3e - I + e - → I - 37

38. Oxidation-Reduction Reactions Page 191 – Self check exercises 7.3a 2 Na (s) + Br 2 (l) → 2 NaBr (s) 7.3b 2 Ca (s) + O 2 (g) → 2 CaO (s) 38

39. Oxidation-Reduction Reactions Page 191 – Self check exercises 7.3a 2 Na (s) + Br 2 (l) → 2 NaBr (s) Na → Na + + e - Br + e - → Br - 7.3b 2 Ca (s) + O 2 (g) → 2 CaO (s) Ca → Ca 2+ + 2e - O + 2e - → O 2- 39

40. Oxidation-Reduction Reactions 1.When a metal reacts with a nonmetal, an ionic compound is formed. The ions are formed when the metal transfers one or more electrons to the nonmetal, the metal becoming a cation and the nonmetal becoming an anion. Therefore, a metal- nonmetal reaction can always be assumed to be an oxdiation-reduction reaction, which involves the transfer of electrons. 40

41. Oxidation-Reduction Reactions 2.Two nonmetals can also undergo an oxidation-reduction reaction. At this point we can recognize these cases only by looking for O 2 as a reactant or product. When two nonmetals react, the compound formed is not ionic. 41

42. Classifying Reactions Formation of a solid (precipitation reaction) Formation of water (acid-base reaction) Transfer of electrons (oxidation-reduction reaction) -single replacement -double replacement Other types of reactions Combustion reaction Synthesis or Combination reaction Decomposition reaction 42