key points for. review nomenclature know symbols/names positive monatomic ions’ names are...
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Key Points for
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Review Nomenclature Know symbols/names Positive monatomic ions’
names are _____________. Negative monatomic
ions/names end in __________. ____________ ion is named
first.
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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 ______________
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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
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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.
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Acid NamesDerive from negative monatomic or polyatomic ion
Monatomic ion plus hydrogen ________________
-ate becomes ______ -ite becomes ______
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General Information Uncombined elements have an oxidation number of zero.
Learn polyatomic ions and their charges!!
Think about reasonable possibilities for given reactants.
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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
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Net Ionic EquationsDissociate all ions when in
solution. (Is it soluble???)Cancel ions that appear on both sides of an equation (spectator ions).
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Solutions of nickel (II) nitrate and cesium hydroxide are mixed.
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Equal volumes of equimolar sulfuric acid and sodium hydroxide are mixed.
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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.
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Reaction PredictionsOne of a few general types:
Double Replacement or Metathesis
Single Replacement (also redox) Redox Synthesis/Decomposition Complex Ions Combustion
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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.
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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.
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Precipitates
A solution of barium chloride is combined with a solution of potassium sulfate.
Solutions of cobalt II chloride and lithium sulfite are combined.
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Both
Sulfuric acid solution in combined with solid calcium hydroxide.
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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.
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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.
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Solid lithium metal is added to water.
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Aqueous solutions of oxalic acid (H2C2O4) and excess potassium hydroxide are mixed.
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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
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Synthesis
Be familiar with diatomics Sulfur and phosphorus
occur as S8 and P4. Nonmetal oxides + water
make acids. Metal oxides + water make
bases.
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Synthesis
Sulfur is burned in oxygen.
Sulfur (VI) oxide is added to water
Calcium oxide is added to water.
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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
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Decomposition Solid potassium chlorate is
heated.
Solid aluminum oxide is heated.
An electric current is passed through molten sodium chloride.
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Solid potassium chlorate is strongly heated.
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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.
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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!
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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
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Oxidizing Agents
Cause oxidation in something else while being reduced themselves
Electron acceptors Nonmetal atoms Positive ions Permanganates, peroxides and
dichromates
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Autooxidation
Some species such as peroxide are self-oxidzining
One species is both oxidized and reduced
Ex: H2O2 H2O + O2
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Acidic or Basic Conditions See page 833 Reactions may be different
under acidic or basic conditions
Use the reduction potential table for reference
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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
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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.
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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
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SUMMARY Two uncombined reactants—synthesis
Single reactant—decomposition Water as a reactant—metals & metal oxides produce bases; nonmetals and nonmetal oxides produce acids.
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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
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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
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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.
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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)
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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+
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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)
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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
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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