1. Preparation for AP Exam Free Response Question #4 Writing Reactions

2. In this question, you will be given just the reactants of a reaction, described in words. You will have to write and balance the entire reaction (predicting products) and also answer a question (????) about each reaction. This section was different a couple years ago, but it is still the hardest section for chemists and students. We’ll do our best to prepare! You need to be familiar with the various reaction types to help predict the products: 5 types of reactions - Double Replacement -Lewis Acid-Base -Organic -Redox -Complex Ion Formation

3. Precipitation Reactions (“double replacement”) Description two ionic compounds (or acids/bases) mix and form two new compounds. one of the new combinations of ions prevents the ions from reforming the original compounds either: - because it forms a stable precipitate - one product breaks apart forming gas, which leaves the solution

4. Preparation & Hints Write water as HOH and acids with hydrogens out front and recognize that it is H + and OH - Memorize: special metal & polyatomic ions and charges solubility rules to predict precipitates strong acids and bases so that you recognize the weak ones (write strong acids and bases as ions, weak in molecular form) the double replacement products that break into gases  you can think of or write chemicals in different ways for different reactions. Ammonia may be NH 3 (aq) for complex ions or NH 4 OH for double replacement or acid-base reactions

5. A potassium chromate solution and barium nitrate solution are combined. molecular equation: K 2 CrO 4 (aq) + Ba(NO 3 ) 2 (aq)  product complete ionic equation : 2K + + CrO 4 2- + Ba 2+ + 2NO 3 -  2K + + CrO 4 2- + Ba 2+ + 2NO 3 - Net ionic equation: CrO 4 2- (aq) + Ba 2+ (aq)  BaCrO 4 (s) Always write the net ionic equation! All strong electrolytes should be shown as free ions (eliminate spectators) and weak electrolytes shown as precipitates (solids). You need to balance equations but you don’t need to include the phases of the reactants and products. Solutions: split into ions (unless insoluble) Solids, crystals, chunks: these are not ions, even if they are things that should be soluble A carbonic acid solution is combined with a calcium chloride solution. H 2 CO 3 (aq) + Ca 2+ ---> 2 H + (aq) + CaCO 3 (s)

6. Neutralization Solutions of ammonia and sulfuric acid are mixed. strA + wkB = salt + HOH + H + H 2 SO 4 + NH 3  H + + SO 4 2- + NH 4 OH  NH 4 + +SO 4 2- + H 2 O H + + NH 4 OH  NH 4 + + H 2 O wkA + strB = salt + HOH + OH - strA + strB  salt + HOH (Net ionic always: H + + OH - =  H 2 O ) wkA + wkB = depends anhydrides Solid sodium oxide reacts with sulfuric acid Na 2 O + H 2 SO 4  Na 2 SO 4 + H 2 O Na 2 O + H +  Na + + H 2 O Sulfur trioxide gas is bubbled through a sodium hydroxide solution 2NaOH + SO 3  Na 2 SO 4 + H 2 O OH - + SO 3  SO 4 2- + H 2 O

7. Know the strong acids & bases!

8. Gas-Forming Reactions Sulfuric acid erodes a limestone statue (solid calcium carbonate) CaCO 3 (s) + H 2 SO 4 (aq) ---> CaSO 4 (s) + H 2 CO 3 (aq) Carbonic acid is unstable: H 2 CO 3 (aq) ---> CO 2 + water CaCO 3 (s) + H + (aq) ---> Ca 2+ + CO 2 + H 2 O

10. Description Two species come together and share electrons in a coordinate covalent bond. The species who donates the electrons is the Lewis base and the accepting species is the Lewis acid. Preparation & Hints “Have pair will share” = Lewis base; e- acceptors are Lewis acids Aluminum ions and Fe 3+ ions make a solution acidic because they are Lewis acids, drawing e- away from O in H 2 O, making the O-H bond more polar, causing H + to pull off H 2 O molecules Watch for: NO 3 -, NO 2 -, CO 3 2-, SO 3 2-, SO 4 2- ions in solids that are heated in a vacuum (not air), this is not combustion. A gas (ex. CO 2 ) and an anhydride (ex. CaO) are formed. Two uncombined elements can form a salt. A Lewis acid and base can also form a salt: CO 2 + CaO  CaCO 3 When metals and metal hydrides react with water, they usually form bases and hydrogen gas. If a gas compound is bubbled through a solution, water is acting like a Lewis base or acid. Lewis Acid-Base Reactions

11. Sulfur dioxide gas is bubbled through water SO 2 (g) + H 2 O(l)  H 2 SO 3 (aq) Nonmetal oxides (CO 2 SO 2 SO 3 NO 2 P 4 O 10 ) form acids in water A calcium oxide solution is prepared. CaO(s) + H 2 O(l) --> Ca(OH) 2 (aq) Metal oxides (MgO, CaO, Na 2 O) form bases in water

12. Organic/Combustion Reactions Description Involves hydrocarbons, commonly combustion but quite often other reactions Preparation and Hints Combustion produces carbon dioxide and water If something reacts with an alkene or alkane, it results in substitution for hydrogen Combustion or burned: O 2 is a reactant

13. Methane reacts with chlorine gas CH 4 + Cl 2  CH 3 Cl + HCl cis-1,2 dichloroethene reacts with hydrogen gas C H Cl H C + H-H  C H Cl HC H H Methane gas is combusted in air products depend on how much oxygen is available

14. Alcohols: *Esterification with carb. acids *Oxidation (they lose e- when lose a H) to form aldehydes, ketones, or carb acids Esters: Hydrolysis to yield carb. acids and alcohols Ketones: Reduction to yield an alcohol Aldehydes: *Carbonyl groups (ketones AND aldehydes) react with PCl 5 to make acid chlorides *Oxidation to form carb. acids Amines: React with acids to form ammonium salts; condensation to form amides

15. Complex Ion Formation Description A transition metal (and some other metals) reacts and gets surrounded by ligands, resulting in a charged compound formed of multiple ions (ex Cu(NH 3 ) 4 2+ ) Ligand: has lone pair to act as lewis base and bond with metal ion( lewis acid) Preparation and Hints A ligand can be a polar molecule (H 2 O, NH 3 ) or a negative anion: Cl -, OH -, SCN -, CN -, S 2 O 3 2- ‘Excess’ and ‘concentrated’ may indicate complex ions, not precipitates: AgNO 3 + HCl forms a white precipitate but concentrated HCl results in AgCl 2 - formation and the solution clears. Watch for FeSCN 2+

16. Central Metal IonCoordination number Ag + 2 Cu 2+ 4 Hg 2+ 4Generally, you place twice Cd 2+ 4the number of ligands around Sn 2+ 4 or 6the central ion as the charge Fe 3+ 6on that central ion. Al 3+ 4 This is not always correct but it will do. Silver chloride dissolves in ammonia solution AgCl(s) + 2NH 3  [Ag(NH 3 ) 2 ] + (aq) + Cl - Aluminum hydroxide dissolves in a concentrated NaOH solution Al(OH) 3 (s) + OH -  Al(OH) 4 - (aq)

17. Redox Reactions (“Oxidation-Reduction”) Description An oxidizing species accepts electrons from a reducing species. The oxidation number of one element increasing as the other decreases. Preparation & Hints Memorize: -the strong oxidizers (ions with lots of oxygens) and what they turn in to (ex HNO 3 may be an acid or an oxidizer forming NO 2, ) - the strong reducers and what they turn in to Look for: - free neutral element in the reaction. -battery reactions. Metal with the greatest E⁰ will reduce - key words like ‘acidified solution’ or an acid included in the reactants. The H + will form H 2 O with oxygens from the oxidizer LeO says GeR When the hydrides of an alkali metal, Ca, Ba, or Sr dissolve in water, hydroxides form and H 2 gas is released (redox) When a hydroxide dissolves, hydroxides form but no gas is released (not redox)

18. Redox Reactions Mg + 2H +  Mg 2+ + H 2 Combination: Combination: sulfur reacts with oxygen S + O 2  SO 2 Decomposition: Decomposition: potassium chlorate undergoes decomposition 2KClO 3 2KCl + 3O 2 00 +4-2 +1+5-2+10 Single Replacement: Single Replacement: Magnesium metal reacts with hydrochloric acid Disproportionation Reaction: Disproportionation Reaction: Element is simultaneously oxidized and reduced Chlorine gas is bubbled through a strongly basic solution Cl 2 + 2OH - ClO - + Cl - + H 2 O 0 +1