Reaction Prediction

1. Reaction Prediction Reactions in Solution Redox-Reactions Original PowerPoint obtained from: archbishopspalding.org

2. Types of Reactions In solution: Double Replacement Acid-Base Precipitation Redox Single Replacement (sometimes in solution) Oxidation-Reduction in Acid Solution Oxidation-Reduction in Basic Solution Combustion (a type of redox reaction) Synthesis (also a type of redox reaction) Decomposition (yet another type of redox reaction!)

3. Reactions between ions in solution Neutralization is an example of a reaction between ions in solution. When ions react, we might observe the formation of a precipitate or a gas. AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Na2CO3 (aq) + 2HNO3 (aq) 2NaNO3 (aq) +H2O (l) CO2 (g) However, not all ions will react in solution. KNO3 (aq) + NaCl (aq) No reaction Solubility rules can help predict reactions.

4. Some simple solubility rules All acids are soluble. All Na+, K+ and NH4+ salts are soluble. All nitrate, acetate, and perchlorate salts are soluble. All chlorides except AgCl and Hg2Cl2 are soluble. PbCl2 is slightly soluble. All sulfates are soluble except PbSO4, Hg2SO4, SrSO4 and BaSO4. Ag2SO4 and CaSO4 are slightly soluble. All sulfides are insoluble except those of the Group IA (1), IIA (2) and ammonium sulfide. All hydroxides are insoluble except those of the group IA(1) and Ba(OH)2. Sr(OH)2 and Ca(OH)2 are slightly soluble.

5. Ionic equations When ionic substances dissolve in water, they dissociate into ions. AgNO3 Ag+ + NO3- KCl K+ + Cl- When a reaction occurs, only some of the ions are actually involved in the reaction. Ag+ + NO3- + K+ + Cl- AgCl(s) + K+ + NO3-

6. Ionic equations To help make the reaction easier to see, we commonly list only the species actually involved in the reaction. Complete ionic equation Ag+ + NO3- + K+ + Cl- AgCl(s) + K+ + NO3- Net ionic equation Ag+ + Cl- AgCl(s) NO3- and K+ are referred to as spectator ions. On Question 4 � ALL reactions should be shown in net ionic form N.B. � Not all of the reactions will occur in solution, therefore, the �net ionic equation� won�t contain any ions Example - Combustion

7. Single replacement reaction Where one element displaces another in a chemical compound. The �displacer� must be more reactive than the �displacee�. H2 + CuO Cu + H2O In this example, hydrogen replaces copper. This type of reaction always involves oxidation and reduction (REDOX). In this case there are no spectator ions. When single replacement reactions occur in solution, a spectator ion will usually exist.

8. Single Replacement Reactions in Solution Example: CuSO4 (aq) + Mg (s) ? MgSO4 (aq) + Cu (s) Any species in an aqueous state should be written as dissociated ions Complete Ionic Equation: Cu2+(aq) + SO42-(aq) + Mg(s) ? Mg2+(aq) + SO42-(aq) + Cu(s) Net Ionic Equation: Cu2+(aq) + SO42-(aq) + Mg(s) ? Mg2+(aq) + SO42-(aq) + Cu(s) Cu2+(aq) + Mg(s) ? Mg2+(aq) + Cu(s) Another Question #4 hint � Equations need not be balanced! Only the species need to be correct. Also, states do not have to be shown. The reactant states are usually indicated or implied by the wording of the question, although they don�t have to be written

9. Single replacement reactions If various metals are in water, we observe that some are more reactive than others. 2Na (s) + 2H2O (l) 2NaOH (aq) + H2 (g) (fast) Ca (s) + 2H2O (l) Ca(OH)2 (s) + H2 (g) (slow) Mg (s) + H2O (l) no reaction This indicates that the order of reactivity of these metals towards water is Na > Ca > Mg We can show the reactivity of metals towards water and acids using an activity series.

10. Activity series of metals

11. Activity series of metals -various metals in HCl

12. Reactivity of nonmetals

13. Decomposition Often occurs when a solid is heated. Several relevant patterns: Chlorates- Decompose to form the metal chloride and oxygen gas Hydrogen carbonates (bicarbonates)- Decompose to form the carbonate, water vapor, and carbon dioxide gas Carbonates- Decompose to form the metal oxide and carbon dioxide gas Hydrogen peroxide Decomposes (in light or with heat) to form liquid water and oxygen gas

14. Synthesis Reactions often occur between metals and nonmetals to form ionic compounds. The metal is oxidized. The nonmetal is reduced. Syntheses don�t occur in solution, so species will not be written as dissociated ions. Example: Na (s) + Cl2 (g) ? NaCl (s)

15. Oxidation number or state When dealing with simple ions, this is easy to determine. It is simply the charge on the ion. Examples Group IA (1) +1 Group IIA (2) +2 Group VIIA (17) -1 Oxygen -2 usually Hydrogen +1 if bonded to nonmetal Hydrogen -1 if bonded to metal

16. Oxidation number For elements in their elemental state, the oxidation number is also pretty straightforward. Since all of the atoms are the same, the electrons are shared equally so the oxidation number is zero. Examples The atoms in N2, Na, P4, H2 and O2 all have oxidation numbers of zero. The bonds in the molecules above are covalent (the electrons are shared) and nonpolar (shared equally)

17. Oxidation numbers With polar covalent bonds electrons are shared but not equally. We will see a little later that some covalent bonds are more polar than others. For electrons that are shared in these compounds, we assign the shared electrons to the most electronegative element. We are just acting as though the electronegativity difference was large enough for the transfer of electrons to occur.

18. Example Assign the oxidation states for all elements in water. The electronegativities are: H = 2.2, O = 3.5 The electrons from both hydrogen are assigned to the oxygen. Oxidation numbers: O = -2 H = +1

19. Rules for assigningoxidation numbers The sum of the oxidation number of all atoms must equal the net charge of the species. In compounds: Group IA are +1. Group IIA are +2. B and Al are +3, and F is -1. Hydrogen is +1 except when combined with a metal. Then it is -1. Oxygen is -2 except for peroxides (-1) and superoxides (- �). Elements in their elemental state have an oxidation number of zero.

20. Oxidation numbers Many elements have more than one possible oxidation number. Often, it is possible to determine the oxidation number of those elements in a compound simply by looking at what you do know. Follow the previous rules and then assign an oxidation number that insures that the overall compound has no net charge

21. Example Find the oxidation state for all elements in:

22. Example See what you do know and find the difference. HNO3

24. Oxidation numbers and the periodic table Some observed trends in compounds. Metals have positive oxidation numbers. Transition metals typically have more than one oxidation number. Nonmetals and semimetals have both positive and negative oxidation numbers. No element exists in a compound with an oxidation number greater than +8. The most negative oxidation numbers equals 8 - the group number

25. Oxidation number and nomenclature Stock system For metals with several possible oxidation numbers, use Roman numeral in the name. FeSO4 iron(II) sulfate Fe2(SO4)3 iron (III) sulfate Cu2O copper(I) oxide CuO copper(II) oxide PbCl2 lead(II) chloride PbCl4 lead(IV) chloride

26. Oxidation number and nomenclature Inorganic oxygen-containing acids and anions Oxo acids and oxo anions rely on a modification of the name to indicate the oxidation number. Acids Anions per ________ic per ________ate ________ic ________ate ________ous ________ite hypo________ous hypo ________ite

27. Oxidation number and nomenclature Examples Cl oxidation number Formula Name +7 HClO4 Perchloric acid +5 HClO3 Chloric acid +3 HClO2 Chlorous acid +1 HClO Hypochlorous acid +7 NaClO4 Sodium perchlorate +5 NaClO3 Sodium chlorate +3 NaClO2 Sodium chlorite +1 NaClO Sodium hypochlorite

28. Identifying oxidation-reduction reactions. Oxidation-Reduction - REDOX A chemical reaction where there is a net change in the oxidation number of one or more species. Both an oxidation and a reduction must occur during the reaction.

29. REDOX reactions Oxidation An increase in oxidation number. Reduction A decrease in oxidation number. If the oxidation number of any element changes in the course of a reaction, the reaction is oxidation-reduction. Example. 2 Fe(NO3)3 (aq) + Zn(s) 2 Fe(NO3)2 (aq) + Zn(NO3)2 (aq)

30. Example 2Fe(NO3)3 (aq) + Zn(s) 2Fe(NO3)2 (aq) + Zn(NO3)2 (aq)

31. Balancing REDOX equations Many REDOX equations can be balanced by inspection. H2S (g) + H2O2 (aq) S (s) + 2 H2O (l) However, others are more difficult. 2KMnO4 (aq) + H2O2 (l) + 3H2SO4 (aq) 2MnSO4 (aq) + K2SO4 (aq) + 3O2 (g) + 4H2O (l)

32. Balancing REDOX equations Half-Reaction method. With this approach, the reaction is broken into two parts. Oxidation half-reaction. The portion of the reaction where electrons are lost. A An+ + ne- Reduction half-reaction. The portion of the reaction where electrons are gained. me- + B Bm-

33. Balancing REDOX equations The goal is then to make sure that the same number of electrons are being produced and consumed. (m) ( A An+ + ne- ) (n) (me- + B Bm- ) nB + mA mAn+ + nBm+ When properly balanced, the electrons will cancel out.

34. Half reactions Example. Half-reactions can be of the �net ionic� form. Balance the following: Fe3+ + Zn (s) Fe2+ + Zn2+ 2 ( Fe3+ + e- Fe2+ ) (reduction) Zn(s) Zn2+ + 2e- (oxidation) 2Fe3+ + Zn (s) 2Fe2+ + Zn2+

35. Half reactions Another Example Determine the balanced equation for the reaction of Fe2+ with Cr2O72- in an acidic solution. Fe2+ + Cr2O72- Fe3+ + Cr3+ The two half-reactions would be: Fe2+ Fe3+ Cr2O72- Cr3+

36. Half reactions First, balance each half-reaction for all elements except hydrogen and oxygen. Fe2+ Fe3+ Cr2O72- 2Cr3+ Next, balance each half-reaction with respect to oxygen by adding an appropriate number of H2O. Fe2+ Fe3+ Cr2O72- 2Cr3+ + 7H2O

37. Half reactions Remember that this reaction occurs in an acid solution so we can add H+ as needed. Fe2+ Fe3+ 14H+ + Cr2O72- 2Cr3+ + 7H2O Now we need to know how many electrons are produced or consumed and place them in our half-reactions. For iron, one e- is produced. For dichromate, six e- are consumed.

38. Half-reactions Fe2+ Fe3+ + e- 6e- + 14 H+ + Cr2O72- 2Cr3+ + 7H2O We need the same number of electrons produced and consumed so: 6Fe2+ 6Fe3+ + 6e- 6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O As our final step, we need to combine the half-reactions and cancel out the electrons.

39. Half-reactions 6Fe2+ + 14H+ + Cr2O72- 6Fe3+ + 2Cr3+ + 7H2O In this reaction, Fe2+ is oxidized and the dichromate ion is reduced. This reaction is used for the determination of iron by titration.

40. Disproportionation reactions In some reactions, the same species is both oxidized and reduced. Examples 2H2O2 (l) 2H2O (l) + O2 (g) Notice this is a decomposition reaction! Remember � decomposition is one sub-type of redox. 3Br2 (aq) +6OH- (aq) BrO3-(aq) +5Br-(aq) +3H2O(l) For this to occur, the species must be in an intermediate oxidation state. Both a higher and lower oxidation state must exist.

41. Complex Ion Formation Transition metal ions for complex ions when they combine with a �ligand� Common metals: Fe, Co, Ni, Cr, Cu, Zn, Ag, Al Sometimes form/dissolve with the addition of acid or base Possible ligands: NH3 CN- OH- SCN- General rule: Look at the charge on the metal ion and double to determine the number of ligands to use � maximum = 6 Typically, one coordination/complex reaction choice out of the eight. Key phrases: �excess concentrated ammonia� �excess cyanide solution� �excess �fill-in-the-blank- hydroxide solution� �thiocyanate� Names of substances that are obviously complex ions Complex ion formulae are always placed in square brackets with the net charge outside Be careful of the charge of the metal ions and ligands to arrive at the correct net charge

42. Example #1a � Formation of ammine complexes Copper(II) chloride solution is combined with an excess of concentrated ammonia solution Cu2+ + NH3 ? [Cu(NH3)4]2+ Notes: There are four ammonia ligands on the Cu2+ The charge of the complex ion is the same as the Cu2+, as the ammonia ligand is electrically neutral Chloride is a spectator ion

43. Example #1b � Formation of ammine complexes The ammonia ligand can �kick out� a hydroxide ion to form a complex ion, almost like a single replacement reaction An excess of concentrated ammonia solution is added to freshly precipitated copper(II) hydroxide. NH3 + Cu(OH)2 ? [Cu(NH3)4]2+ + OH- Notes: Number of ligands in complex ion Charge of complex ion Reaction is electrically balanced

44. Possible Shapes of [Cu(NH3)4]2+

45. Example #1c � Dissolution of ammine complexes Ammine complexes for via the addition of a base, therefore, they can be dissolved by adding strong (but not necessarily concentrated) acid. Remember your list of strong acids & strong bases! A solution of diamminesilver(I) chloride is treated with dilute nitric acid. [Ag(NH3)2]+ + Cl- + H+ ? AgCl + NH4+ Notes: Nomenclature for complex ion The complex ion chloride is soluble, as most chlorides are The silver chloride product is written as a compound as it is one of the insoluble chlorides The ammonium ion product balances the reaction electrically * �Concentrated ammonia� is really ammonium hydroxide The nitrate ion from the acid is a spectator

46. Structure of [Ag(NH3)2]+

47. Example #2 � Formation of cyanide complexes Excess sodium cyanide solution is added to a solution of silver nitrate CN- + Ag+ ? [Ag(CN)2]- Notes: If balanced, there would be two cyanide ions in the reactants and the reaction would be electrically balanced, as well Notice the number of cyanide ligands The sodium and nitrate ions are spectators Because we recognize that cyanide typically forms complexes we do not write: CN- + Ag+ ? AgCN

48. Example #3 � Formation of hydroxo complexes Excess potassium hydroxide solution is added to a solution of aluminum nitrate There are two acceptable products: The complex ion and the aluminum hydroxide precipitate OH- + Al3+ ? Al(OH)3 or [Al(OH)6]3+ Excess potassium hydroxide is added to a precipitate of aluminum hydroxide in water Now our only possible product is the complex ion We may have a maximum of six ligands, but no fewer than four

49. Exammple #4 � Formation of thiocyanato complexes A solution of ammonium thiocyanate is added to a solution of iron(III) chloride SCN- + Fe3+ ? [Fe(SCN)6]3- Notes: Again we write a complex ion product, because we recognize the SCN- to be a ligand-forming ion The best bet is always to add double the charge of ligands, even though sometimes fewer is OK, but never exceed six.

50. Structure similarity to [Fe(SCN)6]3-