Key Points for

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…1- Bromate Chlorate Iodate Nitrate Acetate Azide Cyanide Hydroxide Permanganate Thiocyanate

2- Carbonate Chromate Sulfate Oxalate Dichromat

e Peroxide

3-ArsenateBoratePhosphateCitrate

1+AmmoniumHydronium

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 NamesDerive 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 EquationsDissociate 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 PredictionsOne 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.

Solid lithium metal is added to water.

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.

Solid potassium chlorate is strongly heated.

RedoxSeveral 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 O2 6 CO2 +

8H2O or 2C3H7OH +9 O2 6 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