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Learn symbols, names, and charges of ions, as well as the rules for naming polyatomic ions and acids. Understand the types of reactions and how to predict them. Practice balancing equations and writing net ionic equations.
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Key Points for Reaction Predictions
Review Nomenclature • Know symbols/names • Positive monatomic ions’ names are _____________. • Negative monatomic ions/names end in __________. • ____________ ion is named first.
Polyatomic Ions • Know “home base” ion (-ate ion) and charge. Others can be figured out from there. • “Home base” ion plus one oxygen ________________ • “Home base” ion minus one oxygen ________________ • “Home base” ion minus two oxygens ______________
You Gotta Know… • 3- • Arsenate • Borate • Phosphate • Citrate 1- • Bromate • Chlorate • Iodate • Nitrate • Acetate • Azide • Cyanide • Hydroxide • Permanganate • Thiocyanate 2- • Carbonate • Chromate • Sulfate • Oxalate • Dichromate • Peroxide • 1+ • Ammonium • Hydronium
Other Hints: • Bi- or hydrogen in the name means an extra hydrogen is present—Add H and add +1 to the charge. • Thio- means a sulfur replaces an oxygen—Add S & remove O; charge is the same.
Acid Names • Derive from negative monatomic or polyatomic ion • Monatomic ion plus hydrogen ________________ • -ate becomes ______ • -ite becomes ______
General Information • Uncombined elements have an oxidation number of zero. • Learn polyatomic ions and their charges!! • Think about reasonable possibilities for given reactants.
Equations • Word: Sodium chloride + silver nitrate sodium nitrate + silver chloride • Formula: NaCl + AgNO3 NaNO3 + AgCl • Complete Ionic: Na + + Cl- + Ag+ + NO3- AgCl + Na+ + NO3- • Net Ionic: Cl- + Ag+ AgCl
Net Ionic Equations • Dissociate all ions when in solution. (Is it soluble???) • Cancel ions that appear on both sides of an equation (spectator ions).
Solutions of nickel (II) nitrate and cesium hydroxide are mixed.
Equal volumes of equimolar sulfuric acid and sodium hydroxide are mixed.
General Information • Will always appear on the exam • Three equations to write & three questions to answer • Must write a net ionic equation—NO SPECTATOR IONS • Balance the equations—atoms & charges. • All reactions will occur. Don’t worry about activity series.
Reaction Predictions One of a few general types: • Double Replacement or Metathesis • Single Replacement (also redox) • Redox • Synthesis/Decomposition • Complex Ions • Combustion
Double Replacement or Metathesis Starting point: two compounds, often in solution • Look for an acid/base reaction • Look for the formation of an insoluble compound (precipitate) • Both can happen simultaneously.
Acid/Base • A solution of hydrochloric acid is combined with a solution of sodium hydroxide. • A solution of hydrofluoric acid is combined with solid zinc hydroxide.
Precipitates • A solution of barium chloride is combined with a solution of potassium sulfate. • Solutions of cobalt II chloride and lithium sulfite are combined.
Both • Sulfuric acid solution in combined with solid calcium hydroxide.
Single Replacement Starting point: one compound (often in solution) and one uncombined element • Like will replace like in a reaction. • Oxidation numbers must change since an uncombined element’s oxidation # is always zero.
Single replacement • Fluorine gas is bubbled through a solution of potassium chloride. • A piece of solid aluminum is placed in a solution of copper II sulfate.
Aqueous solutions of oxalic acid (H2C2O4) and excess potassium hydroxide are mixed.
Synthesis/Decomposition • Synthesis—two reactants combine • Synthesis—starting point will be two separate substances—possibly elements • Decomposition—one reactant breaks apart • Starting point will be only one reactant—look for heat or electricity
Synthesis • Be familiar with diatomics • Sulfur and phosphorus occur as S8 and P4. • Nonmetal oxides + water make acids. • Metal oxides + water make bases.
Synthesis • Sulfur is burned in oxygen. • Sulfur (VI) oxide is added to water • Calcium oxide is added to water.
Decomposition • Know some special types: • Metal carbonates metal oxides & CO2 • Metal hydroxides metal oxides and water • Metal chlorates metal chlorides and oxygen gas • Hydrogen peroxide H2O & O2
Decomposition • Solid potassium chlorate is heated. • Solid aluminum oxide is heated. • An electric current is passed through molten sodium chloride.
Redox Several starting points: • Key compounds or ions: MnO4-, H2O2, Cr2O72-, HNO3 • Metals with multiple oxidation states: Sn2+, Sn4+, Cr2+, Cr3+, Cr6+, etc. • Acidic or basic conditions • Sometimes halogens, i.e. I-, IO-, IO2-, etc.
Redox • In redox, one reactant gains electrons and one reactant loses. • LEO says GER: • Losing electrons = oxidation • Gaining electrons = reduction • One process cannot occur without the other!
Reducing Agents • Cause reduction in something else while being oxidized themselves • Electron donors (any species that loses electrons) • Metal atoms • Negative ions • Positive metal ions that may still be able to lose more electrons
Oxidizing Agents • Cause oxidation in something else while being reduced themselves • Electron acceptors • Nonmetal atoms • Positive ions • Permanganates, peroxides and dichromates
Autooxidation • Some species such as peroxide are self-oxidzining • One species is both oxidized and reduced • Ex: H2O2 H2O + O2
Acidic or Basic Conditions • See page 833 • Reactions may be different under acidic or basic conditions • Use the reduction potential table for reference
To Predict: • Single replacement—follow normal pattern • Synthesis—follow normal pattern • All others: • Think about the most stable state of a substance (i.e. K vs K+) • Use the standard reduction potential table to help
An acidic solution of potassium dichromate is added to a solution of iron II nitrate • Find acidified dichromate as a reactant (oxidizing agent—It is reduced.) • Find iron II as a product. (It is oxidized.) • Reverse the one that is oxidized (iron II). • Combine and balance.
A strip of copper is immersed in dilute nitric acid • Copper begins with ox. # of zero; can only form positive ions; must be oxidized; reverse reaction. • NO3- + 4H+ NO + 2H2O
SUMMARY • Two uncombined reactants—synthesis • Single reactant—decomposition • Water as a reactant—metals & metal oxides produce bases; nonmetals and nonmetal oxides produce acids.
SUMMARY • Acid & base reactants (including salts)—neutralize to salt and water • Two salt solutions—look for a precipitate or a gas produced • Combustion of hydrocarbon—produce CO2 & H2O • Solid metal placed in solution—single replacement/redox rxn
SUMMARY • Transition metal with ammonia (NH3), hydroxide (OH-), cyanide (CN-), or thiocyanate (SCN-)—form complex ions in which ions attach to metals • Doesn’t matter how many—just get charge correct
Special Notes • If a product would be carbonic acid (H2CO3), it will break down into CO2 and H2O. • If a product would be ammonium hydroxide, it would break down into NH3 and H2O. • Ammonium carbonate decomposes into CO2, NH3 & H2O.
EXAMPLES • A piece of aluminum metal is added to a solution of silver nitrate. • Al + 3Ag+ Al3+ + 3Ag • A piece of solid bismuth is strongly heated in oxygen. • 4Bi + 3O2 2Bi2O3 (or another oxide)
EXAMPLES • An excess of sodium hydroxide solution is added to a solution of magnesium nitrate • Mg2+ + 2OH- Mg(OH)2 • Solid lithium hydride is added to water • LiH + H2O LiOH + H2 • A concentrate solution of ammonia is added to a solution of zinc nitrate • Zn2+ + 4NH3 Zn(NH3)42+
Examples • Concentrate hydrochloric acid is added to solid manganese II sulfide. • 2 H+ + MnS H2S + Mn2+ • Excess chlorine gas is passed over hot iron filings. • 3Cl2 + 2Fe 2FeCl3 (possibly FeCl2)
Examples • Solid ammonium carbonate is heated • (NH4)2CO3 CO2 + NH3 + H2O • Equal volumes of 0.1 M sulfuric acid and 0.1 M potassium hydroxide are mixed • H+ + OH- H2O
EXAMPLES • Propanol is burned completely in air • 2CH3CH2CH2OH +9 O26 CO2 + 8H2O or 2C3H7OH +9 O26 CO2 + 8 H2O • A solid sample of magnesium carbonate is heated strongly. • MgCO3 MgO + CO2 • Ethene gas is bubbled through a solution of bromine. • C2H4 + Br2 C2H4Br2