Reaction Prediction

Reaction Prediction Reactions in Solution Redox-Reactions

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!)

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.

Some simple solubility rules • All group IA and ammonium salts are soluble. • All nitrate, acetate, and perchlorate salts are soluble. • All chlorides, bromides, and iodides are soluble except Ag+, Hg2+2 and Pb+2. PbCl2 is slightly soluble, more so in hot water than cold. (For our purposes here, slightly soluble will be the same as insoluble.) • 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, 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. • Most carbonates, chromates, and phosphates are only slightly soluble.

H2O H2O Ionic equations • When ionic substances dissolve in water, they dissociate into ions. • AgNO3 Ag++ NO3- • KClK+ + Cl- • When a reaction occurs, only some of the ions are actually involved in the reaction. • Ag++ NO3- +K+ + Cl- AgCl(s) + K+ + NO3-

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 AP exam – ALL reactions should be shown in net ionic form • However, not all of the reactions will occur in solution, therefore, the “net ionic equation” won’t contain any ions • Example - Combustion

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.

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)

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.

Activity series of metals potassium sodium React vigorously with cold water  calcium React slowly with cold water magnesium aluminum zinc chromium React very slowly with steam but are quite reactive in acid increasing reactivity iron nickel tin lead React moderately with acid copper silver platinum gold Unreactive in acid

Activity series of metals -various metals in HCl Iron Zinc Magnesium

Increased reactivity C N O F P S Cl Se Br Increased reactivity I Reactivity of nonmetals

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

More Decomposition • Metal Oxides (heated) • produce the metal + oxygen gas. (This is how oxygen was discovered.) • Acids (heated) • produce non-metal oxide plus water • The non-metal oxide is called an acid anhydride since it makes an acid when combined with water (think synthesis reaction) • Bases (heated) • produce metal oxide plus water • The metal oxide is called a basic anhydride since it makes a base when combined with water (again, think synthesis reaction). • Electrolysis • Uses electricity to break a compound into its elements. The compound must be in solution or in liquid form. Solid salts don’t conduct electricity.

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)

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

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 Increased oxygen and Oxidation number

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

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. Mg (s) + Cl2 (g) MgCl2 (s) Here the oxidation number of Mg has changed from zero to +2. Cl has changed from zero to -1.

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)

Example +2 +2 +3 0 • 2Fe(NO3)3 (aq) + Zn(s) 2Fe(NO3)2 (aq) + Zn(NO3)2 (aq) Fe3+ is reduced to Fe2+ Zn is oxidized to Zn2+ NO3- is a spectator ion.

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)

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-

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.

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+

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+ H+

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

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.

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.

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.

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.

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

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

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

Possible Shapes of [Cu(NH3)4]2+ Square planar Tetrahedral http://www.elecuter.co.uk/Scinet/chemistry/ss/dblock.php

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

Structure of [Ag(NH3)2]+ H3N Ag NH3+ Linear http://www.elecuter.co.uk/Scinet/chemistry/ss/dblock.php

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

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)3or [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

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.

Structure similarity to [Fe(SCN)6]3- [Fe(CN)6]3- Octahedral http://www.elecuter.co.uk/Scinet/chemistry/ss/dblock.php

Reaction Prediction