Transcript Chapter Four - Salina USD 305

1
Chapter Four
Chemical Reactions in
Aqueous Solutions
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Four
Today…
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Chapter One
Unit 1 Test Results
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Chapter One
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Quick Review
• How do we know a
chemical reaction has
occurred? What do we
observe?
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Chapter Four
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Evidence of a Chemical Rxn
1.
2.
3.
4.
Color Change
Precipitate (solid) forms
Gas Evolved
Heat/Light Given Off
– Endothermic – heat absorbed (gets
cold)
– Exothermic – heat released (gets hot)
• More about this in Unit 5 - Thermodynamics
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Chapter Four
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Review of Basics
• When writing out a balanced equation, it is
necessary to indicate what state the
substance is in (s, l, g, or aq).
• For elements, you can look at the PT to
determine their standard state.
– Examples:
• Mercury
• Fluorine
• Iron
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Chapter Four
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Review of Basics
• There are 7 diatomic elements.
• They are only diatomic when
they are alone. They are all
gases, except Br2, which is a
liquid, as indicated on the PT.
• Remember the Super 7?
–H2, N2, O2, F2, Cl2, Br2, I2
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Chapter Four
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Review of Basics
• Soluble means the substance dissolves
in solution (water).
– Label soluble substances aqueous (aq)
• Insoluble means the substance does not
dissolve in solution.
– Label insoluble substances solid (s)
• To determine if a substance is soluble
or insoluble, use a solubility table.
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Chapter Four
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Solubility Chart
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Chapter Four
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Other Basic Tips
• Acids and bases are aqueous
• Water is almost always a liquid
• Write water as HOH when it is
a reactant
• Some common gases that aren’t
diatomic are CO2, CO, SO3,
SO2, H2S, and NH3
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Chapter Four
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Other Basic Tips
• If a reaction indicates that a catalyst was
present, it is not part of the actual reaction.
Instead, you write the catalyst above the
arrow.
• Catalysts are not used up in the reaction.
They are just there to speed it up.
• If a reaction is heated, you draw a triangle
above the arrow. Heat is used as a catalyst.
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Chapter Four
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Synthesis Reactions
• Occur when two or more
reactants combine to form a
single product.
–In Chem 1 we called this type
“Dating”
• There are several common types
of synthesis reactions.
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Chapter Four
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Synthesis Reactions
Type 1: A metal combines with a
nonmetal to form a binary salt.
Example: A piece of lithium metal
is dropped into a container of
nitrogen gas.
6Li (s) + N2 (g) → 2Li3N (aq)
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Synthesis Reactions
Type 2: Metallic oxides and water
form bases (metallic hydroxides).
Example: Solid sodium oxide is added to
water
Na2O (s) + HOH (l) → 2NaOH (aq)
Example: Solid magnesium oxide is added to
water.
MgO (s) + HOH (l) → Mg(OH)2 (aq)
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Synthesis Reactions
Type 3: Nonmetallic oxides and
water form acids. The nonmetal
retains its oxidation number.
Example: Carbon dioxide is bubbled in water.
CO2 (g) + H2O (l) → H2CO3 (aq)
Example: Dinitrogen pentoxide is bubbled in
water.
N2O5 (g) + H2O (l) → 2HNO3 (aq)
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Chapter Four
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Synthesis Reactions
Type 4: Metallic oxides and
nonmetallic oxides form salts.
Example: Solid sodium oxide is added to
carbon dioxide.
Na2O (s) + CO2 (g) → Na2CO3 (aq)
Example: Solid calcium oxide is added to
sulfur trioxide.
CaO (s) + SO3 (g) → CaSO4 (aq)
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Chapter Four
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Try It Out 1
• Solid barium oxide is added to distilled
water (p. 142)
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Chapter Four
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Try It Out 2
• Sulfur trioxide gas is added to excess water
(p. 146)
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Chapter Four
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Decomposition Reactions
• Occur when a single reactant is
broken down into two or more
products.
– In Chem 1 we called this “The Break Up”
• There are several common
types of decomposition
reactions.
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Chapter Four
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Decomposition Reactions
Type 1: Metallic carbonates
decompose into metallic oxides
and carbon dioxide
Example: A sample of magnesium
carbonate is heated.
MgCO3 (s) → MgO (s) + CO2 (g)
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Decomposition Reactions
Type 2: Metallic chlorates
decompose into metallic chlorides
and oxygen.
Example: A sample of magnesium
chlorate is heated.
Mg(ClO3)2 (aq) → MgCl2 (aq) + 3O2 (g)
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Decomposition Reactions
Type 3: Ammonium carbonate
decomposes into ammonia, water,
and carbon dioxide.
Example: A sample of ammonium
carbonate is heated.
(NH4)2CO3 (aq)→ 2NH3 (g) + H2O (l) + CO2 (g)
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Decomposition Reactions
Type 4: Sulfurous acid
decomposes into sulfur dioxide
and water.
Example: A sample of ammonium
carbonate is heated.
H2SO3 (aq) → H2O (l) + SO2 (g)
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Chapter Four
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Decomposition Reactions
Type 5: Carbonic acid
decomposes into carbon dioxide
and water.
Example: A sample of carbonic acid
is heated.
H2CO3 (aq) → H2O (l) + CO2 (g)
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Decomposition Reactions
Type 6: A binary compound may
break down to produce two
elements.
Example: Molten sodium chloride is
electrolyzed.
2NaCl (l) → 2Na (s) + Cl2 (g)
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Chapter Four
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Decomposition Reactions
Type 7: Hydrogen peroxide
decomposes into water and
oxygen.
Example: 2H2O2 (aq) → 2H2O (l) + O2 (g)
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Chapter Four
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Decomposition Reactions
Type 8: Ammonium hydroxide
decomposes into ammonia and
water.
Example: NH4OH (aq) → NH3 (g) + HOH (l)
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Chapter Four
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Try It Out 1
• Solid calcium chlorate is heated in the
presence of a magnesium dioxide catalyst
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Chapter Four
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Try It Out 2
• A sample of lithium carbonate is heated
strongly
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Chapter Four
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STOP
• Complete Worksheet #1
by next class.
Prentice Hall © 2005
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Four
Today…
• Turn in:
• Get Out WS#1
• Our Plan:
• Questions on WS#1
• Quick Review
• Notes – Single & Double Replacement
• Begin WS#2
• Homework (Write in Planner):
• WS#2 Due Friday
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Chapter One
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Quick Review
• Write formulas for the reactants and predicted products for
the chemical reactions that follow:
1. Solid calcium carbonate is strongly heated in a test tube.
2. Solid lithium chlorate is heated in the presence of a
manganese dioxide catalyst.
3. Excess chlorine gas is passed over hot iron filings.
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Chapter One
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Single Replacement
• Reactions that involve an element
replacing one part of a compound.
• The products include the displaced
element and a new compound.
– In Chem 1 we called this “Cheating”
• An element can only replace
another element that is less active
than itself.
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Chapter Four
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Single Replacement
General Activity Series for Metals:
Most Active
Li
Ca
Na
Mg
Al
Zn
Fe
Pb
H2
Cu
Ag
Pt
Au
Least Active
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Chapter Four
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Single Replacement
General Activity Series for Nonmetals:
Most Active
F2
Cl2
Br2
I2
Least Active
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Chapter Four
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Remember the Trend?
• Or you can just look at the Periodic
Table
– The most reactive metals are on the
bottom left (Francium)
– The most reactive nonmetals are on the
top right (Fluorine)
– Hydrogen is the tricky one.
– Use the pink sheet for electronegativity
potentials otherwise.
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Chapter Four
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For metals, the
more negative the
reduction potential,
the more reactive.
For nonmetals, the
more positive the
reduction potential,
the more reactive.
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Chapter Four
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Single Replacement
Type 1: Active metals replace less
active metals from their
compounds in aqueous solution.
Example: Magnesium turnings are
added to a solution of iron (III)
chloride.
3Mg (s) + 2FeCl3 (aq)→ 2Fe (s) + 3MgCl2 (aq)
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Single Replacement
Type 2: Active metals replace
hydrogen in water.
Example: Sodium is added to water.
2Na (s) + 2HOH (l)→ 2NaOH (aq) + H2 (g)
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Single Replacement
Type 3: Active metals replace
hydrogen in acids.
Example: Lithium is added to
hydrochloric acid.
2Li (s) + 2HCl (aq) → 2LiCl (aq) + H2 (g)
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Single Replacement
Type 4: Active nonmetals replace
less active nonmetals from their
compounds in aqueous solution.
Example: Chlorine gas is bubbled
into a solution of potassium iodide.
Cl2 (g) + 2KI (aq) → I2 (g) + 2KCl (aq)
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Chapter Four
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Single Replacement
Type 5: If a less reactive element is combined
with a more reactive element in compound form,
their will be no resulting reaction.
Example: Chlorine gas is bubbled into a solution of
potassium iodide.
Cl2 (g) + KF (aq) → No Reaction
Example: Zinc is added to a solution of sodium
chloride.
Zn (s) + NaCl (aq) → No Reaction
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Chapter Four
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Try It Out 1
• A piece of aluminum metal is added to a
solution of gold nitrate. (p. 146)
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Chapter Four
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Try It Out 2
• Liquid bromine is shaken with a potassium
iodide solution. (p. 152)
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Chapter Four
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Try It Out 3
• Small chunks of potassium are added to
water. (p. 164)
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Chapter Four
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Try It Out 4
• Small strips of platinum are placed in
hydrochloric acid.
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Chapter Four
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Double Replacement
• Reactions between two
compounds in aqueous solution
where the cations and anions
appear to “switch partners”.
AX + BY → AY + BX
• In Chem 1 we called this
“Swapping”
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Double Replacement
• All double replacement
reactions (aka metathesis)
must have a “driving force”
or reason why the reaction
will occur or “go to
completion”.
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Chapter Four
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Double Replacement
• “Driving Force” for reactions:
1. Formation of a precipitate
2. Formation of a gas
3. Formation of primarily
molecular species
(nonelectrolytes, water, weak
acids)
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Double Replacement
• If one of these “driving forces” is NOT
present, then the reaction does not go to
completion.
• This type of reaction is indicated by a
double arrow
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Chapter Four
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Reactions that Form Precipitates
• There are limits to the amount of a solute
that will dissolve in a given amount of
water.
• If the maximum concentration of solute is
less than about 0.01 M, we refer to the
solute as insoluble in water.
• When a chemical reaction forms such a
solute, the insoluble solute comes out of
solution and is called a precipitate.
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Chapter Four
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Silver Iodide Precipitation
A solution containing
silver ions and nitrate
ions, when added to …
… a precipitate of
silver iodide.
… a solution
containing potassium
ions and iodide ions,
forms …
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What is the net ionic
equation for the
reaction that has
occurred here? (Hint:
what species actually
reacted?)
General Chemistry 4th edition, Hill, Petrucci, McCreary, Perry
Chapter Four
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Double Replacement (Precipitate)
• In order to predict double
replacement reactions yielding
precipitates, one must
memorize the solubility rules.
• I will let you use a solubility
chart on your test!
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Chapter Four
54
Solubility Rules
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Chapter Four
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Example 1
• Predict and balance the following double
replacement reactions based on the
solubility of the products. Use the
abbreviations (aq) and (s) for the reactant
and products. All reactants are aqueous.
1. Solutions of manganese (II) sulfate and
ammonium sulfide are mixed.
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Chapter Four
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Example 2
2. Solutions of sodium iodide and lead (II)
nitrate are combined.
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Try it Out 1
• Solutions of sodium carbonate and
iron (III) nitrate combine.
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Try it Out 2
• Solutions of lead (II) acetate and
calcium chloride are combined.
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Chapter Four
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Double Replacement (Gas)
• Common Gases formed: H2S, CO2,
SO2, NH3
• Reactions that produce three of the
gases (CO2, SO2, and NH3) involve
the initial formation of a substance
that breaks down to give the gas
and HOH.
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Chapter Four
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Double Replacement (Gas)
Common Gases
H2S
CO2
SO2
NH3
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Any sulfide (salt of S2-) plus any acid form H2S (g)
and a salt.
Any carbonate (salt of CO32-) plus any acid form CO2
(g), HOH, and a salt
Any sulfite (salt of SO32-) plus any acid form SO2 (g),
HOH, and a salt.
Any ammonium salt (salt of NH4+) plus any soluble
strong hydroxide react upon heating to form NH3 (g),
HOH, and a salt.
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Chapter Four
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Double Replacement (Gas)
• Example 1: The reaction of Na2SO3 and HCl
produces H2SO3:
Na2SO3 (aq) + 2HCl (aq) → H2SO3 (aq) + 2NaCl (aq)
• Bubbling is observed in this reaction because the
H2SO3 (sulfurous acid) is unstable and
immediately decomposes to give HOH and SO2
gas:
H2SO3 (aq) → HOH (l) + SO2 (g)
• The molecular equation for the overall or
complete reaction, therefore, is:
Na2SO3 (aq) + 2HCl (aq) → HOH (l) + SO2 (g) + 2NaCl (aq)
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Chapter Four
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Double Replacement (Gas)
• Example 2: A typical reaction of a carbonate
and an acid is:
K2CO3 (aq) + 2HNO3 (aq) → HOH (l) + CO2 (g) + 2KNO3 (aq)
• Bubbling is also observed in this reaction.
Theoretically H2CO3, carbonic acid, is formed,
but the acid is unstable and immediately
decomposes to form carbon dioxide gas and
water according to the following equation:
H2CO3 (aq) → HOH (l) + CO2 (g)
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Chapter Four
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Double Replacement (Gas)
• Example 3: Ammonium salts and soluble bases
react as follows (particularly when the solution is
warmed):
NH4Cl (aq) + NaOH (aq) → NH3 (g) + HOH (l) + NaCl (aq)
• The odor of ammonia gas is noted and moist blue
litmus paper held near the mouth of the container
will turn blue. Theoretically NH4OH, ammonium
hydroxide, is produced (also known as ammonia
water). The compound is unstable and
decomposes into ammonia gas and water:
NH4OH (aq) → NH3 (g) + HOH (l)
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Double Replacement (Gas)
• Example 4: The odor of rotten
eggs and bubbling are noted when
an acid is added to a sulfide. A
typical reaction producing
hydrogen sulfide gas is:
FeS (s) + 2HCl (aq) → FeCl2 (aq) + H2S (g)
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Chapter Four
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Try It Out 1
• Dilute hydrochloric acid is added to a
solution of potassium sulfite. (p. 152)
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Chapter Four
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Try It Out 2
• Concentrated hydrochloric acid is added to
solid manganese (II) sulfide. (p. 150)
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Chapter Four
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Today…
• Turn in:
• Get Out WS#2
• Our Plan:
• Questions on WS#2
• Quick Review/Video
• Notes – Aqueous Reactions, Net Ionic Equations, &
Math with Ion Concentration
• Begin WS#3
• Homework (Write in Planner):
• WS#3 Due Monday
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Chapter Four
68
A Note about phases
• A product can only be aqueous if there is an
aqueous or liquid reactant. That’s why
some of the ones on the first WS that you
labeled aqueous were solid.
• Examples:
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Chapter Four
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Quick Discussion of Weak Electrolytes
• One of the driving forces for a reaction is
the production of weak electrolytes (water
or a weak acid).
• For a list of strong acids see page 22 in
booklet. Everything else is weak.
• The only thing to note is that if you form a
weak acid, it is aqueous but you would have
a forward arrow, not a double arrow.
• See examples on the board.
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Chapter Four
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Quick Review
• Solutions of sodium fluoride and dilute
hydrochloric acid are combined.
• Dilute acetic acid solution is added to
solid magnesium carbonate.
• A saturated solution of calcium
hydroxide was added to a solution of
magnesium chloride.
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Chapter Four
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Crash Course in Precipitation
• http://www.youtube.com/watch?v=IIu16dy
3ThI
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Chapter Four
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Arrhenius’s Theory of
Electrolytic Dissociation
• Why do some solutions conduct electricity?
• An early hypothesis was that electricity
produced ions in solution, and those ions
allowed the electricity to flow.
• Arrhenius’s theory:
– Certain substances dissociate into cations and
anions when dissolved in water.
– The ions already present in solution allow
electricity to flow.
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Electrolytic
Properties
of Aqueous
Solutions
• Electrolytes
dissociate to
produce ions.
The more the electrolyte dissociates, the more ions it produces.
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Chapter Four
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Types of Electrolytes
• A strong electrolyte dissociates completely.
– A strong electrolyte is present in solution almost
exclusively as ions.
– Strong electrolyte solutions are good conductors.
• A nonelectrolyte does not dissociate.
– A nonelectrolyte is present in solution almost exclusively
as molecules.
– Nonelectrolyte solutions do not conduct electricity.
• A weak electrolyte dissociates partially.
– Weak electrolyte solutions are poor conductors.
– Different weak electrolytes dissociate to different extents.
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Is it a strong electrolyte, a weak
electrolyte, or a nonelectrolyte?
• Strong electrolytes include:
– Strong acids (HCl, HBr, HI, HNO3, H2SO4, HClO4)
– Strong bases (IA and IIA hydroxides)
– Most water-soluble ionic compounds
• Weak electrolytes include:
– Weak acids and weak bases
– A few ionic compounds
How do we tell
whether an acid
(or base) is
weak?
• Nonelectrolytes include:
– Most molecular compounds
– Most organic compounds (most of them are molecular)
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Chapter Four
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Aqueous Solutions & Ionic Equations
• On the AP Exam, you do not
have to write out complete
molecular equations like we
have been doing.
• Instead, you have to write ionic
equations.
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Chapter Four
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Overall (Total) Ionic Equation
• Formulas of the reactants and products are
written to show the predominant form of
each substance as it exists in aqueous
solution.
• Soluble salts, strong acids, and strong
bases are written as separated ions.
Everything else is written as a molecule.
• We will memorize strong acids and strong
bases later in the year. For now, a list will
be provided.
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Chapter Four
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Memorize Later…
Strong Acids
HClO4
H2SO4
HI
HBr
HCl
HNO3
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Strong Bases
LiOH
RbOH
NaOH
CsOH
KOH
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
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Chapter Four
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Overall Ionic Equation Example 1
Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)
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Overall Ionic Equation Example 2
CaCO3 (aq) + 2HCl (aq) → CaCl2 (aq) + HOH (l) + CO2 (g)
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Try it Out
(NH4)2S (aq) + 2LiOH (aq) → Li2S (aq) + 2NH3 (g) + 2HOH (l)
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Chapter Four
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Net Ionic Equation
• Net ionic equations are written to show only
the species that react or undergo change in
aqueous solutions.
• The net ionic equation is obtained by
eliminating the spectator ions from the
overall equation.
• All that is left are the ions that have changed
chemically.
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Net Ionic Equation Example 1
Cd(NO3)2 (aq) + Na2S (aq) → CdS (s) + 2NaNO3 (aq)
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Example 2
• For each of these word equations, predict the product
and write an overall and net ionic equation.
• Copper (I) nitrate is combined with silver chloride.
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Example 3
• Magnesium hydroxide is combined with lead (II)
sulfate.
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Try It Out
• Mercury (I) nitrate is combined with sodium sulfide.
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Ion Concentrations in Solution
• “Trick” question …
What is the concentration of Na2SO4 in a solution
prepared by diluting 0.010 mol Na2SO4 to 1.00 L?
• The answer is:
… zero …
• WHY??
• And … how do we describe the concentration of
this solution?
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Calculating Ion Concentrations in
Solution
• In 0.010 M Na2SO4:
– two moles of Na+ ions are formed for each
mole of Na2SO4 in solution, so [Na+] =
0.020 M.
– one mole of SO42– ion is formed for each
mole of Na2SO4 in solution, so [SO42–] =
0.010 M.
• An ion can have only one concentration
in a solution, even if the ion has two or
more sources.
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Chapter Four
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Example 4.1
Calculate the molarity of each ion in an aqueous solution
that is 0.00384 M Na2SO4 and 0.00202 M NaCl. In
addition, calculate the total ion concentration of the
solution.
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Try It Out!
• Exercise 4.1A: Seawater is essentially 0.438 M NaCl and
0.0512 M MgCl2, together with several other minor
solutes. What are the molarities of Na+1, Mg+2, and Cl-1 in
seawater?
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Example 4.6
One cup (about 240 g) of a certain clear chicken broth
yields 4.302 g AgCl when excess AgNO3(aq) is added to
it. Assuming that all the Cl– is derived from NaCl, what is
the mass of NaCl in the sample of broth?
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Chapter Four
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Stop!
• Complete Worksheet #4 by
Monday.
–Let’s look at it together.
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Chapter Four
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Today…
• Turn in:
• Get Out WS#3
• Our Plan:
• Questions on WS#3
• Quick Review
• Notes – Acid/Base Reactions & Titrations
• Begin WS#4
• Homework (Write in Planner):
• WS#4 Due Wednesday
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Chapter Four
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Quick Review
• Write the molecular and net ionic
equation for the following:
1. Chlorine gas is bubbled into a
solution of potassium iodide.
2. Aqueous solutions of potassium
chromate and silver nitrate react.
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Acid-Base Reactions
• We will cover acid/base reactions in
great detail at the end of the school
year. For now, we’re just going to
cover very simple neutralization
reactions.
• Neutralization is the transfer of
PROTONS from an acid to a base.
• In neutralization, an acid and a base
form a salt and water.
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Acid-Base Reactions
• When writing net ionic equations, keep in mind
which acids are strong (written in ionic form)
and which are weak (written in molecular
form).
• Check the solubility rules of the salt produced.
If it is soluble, it is written in ionic form; if it is
insoluble it is written in molecular form.
• The salt produced always consists of a cation
from the base and an anion from the acid.
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Remember, these are in your notes…
Strong Acids
HClO4
H2SO4
HI
HBr
HCl
HNO3
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Strong Bases
LiOH
RbOH
NaOH
CsOH
KOH
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
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Acid-Base Reactions
• Example 1: Hydrogen sulfide
gas is bubbled through excess
potassium hydroxide solution.
H2S (g) + 2KOH (aq) → K2S (aq) + 2HOH (l)
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Acid-Base Reactions
• Polyprotic acids (more than one hydrogen like
H2SO4) can be tricky. If the base is in excess,
all hydrogen ions will react with strong base to
produce water. See Example 2.
• If however, this same reaction were described
in terms of mixing equal numbers of moles,
then the coefficients for both reactants would
be one (the same number of H and OH must be
given away). See Example 3.
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Acid-Base Reactions
• Example 2: Dilute sulfuric acid is
reacted with excess sodium
hydroxide.
H2SO4 (aq) + 2NaOH (aq) → Na2SO4 (aq) + 2HOH (l)
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Acid-Base Reactions
• Example 3: Equal number of
moles of sulfuric acid and sodium
hydroxide react.
H2SO4 (aq) + NaOH (aq) → NaHSO4 (aq) + HOH (l)
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Acid-Base Reactions
• As the following example demonstrates, it
is important to take into account the
quantity (concentration and amount) of each
reactant.
• Example 4: Equal volumes of 0.1 M
phosphoric acid and 0.2 M sodium
hydroxide are reacted together.
H3PO4 (aq) + 2NaOH (aq) → Na2HPO4 (aq) + 2HOH (l)
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Try It Out 1
• Equal molar volumes of sodium hydroxide
and hydrobromic acid are mixed together.
Write the molecular equation and the net
ionic equation.
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Chapter Four
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Try It Out 2
• Equal volumes of 0.1 M phosphoric acid
and potassium hydroxide solutions are
mixed.
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Chapter Four
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Titrations
• A titration is a procedure where one
reactant is carefully added to another
reactant until the two have combined in
their exact stoichiometric proportions.
• There are different types of titrations,
but no matter the type, both reactants
are fully consumed at the end of the
titration.
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Titrations
• The purpose of a titration is to find the number of
moles, grams, concentration, or percentage of an
unknown.
• The unknown is called the ANALYTE.
• This is done by measuring the precise volume
and concentration of a known solution
(TITRANT) needed to react completely with the
analyte.
• Stoichiometry is then used to determine whatever
quantity you’re looking for.
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Acid-Base Titrations
• Acid-base titrations are the most common type
• You conducted this type of titration in Chem 1
(remember the antacid lab?)
• There are 3 things that you need for a titration:
– A way to accurately measure the titrant
(BURET)
– A way to know your reaction is complete
(INDICATOR or you can look at a
TITRATION CURVE)
– A solution whose concentration you know
(STANDARD SOLUTION)
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Acid-Base Titrations
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Acid-Base Titrations
• Equivalence point – in an acid-base titration
it is the point when the titrant completely
neutralizes the analyte.
• If you’re using an indicator, you choose one
that changes color close to the
neutralization point.
• If you are looking at a titration curve, it is
the straight line on a graph of volume
versus pH.
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Acid-Base Titrations
• When the indicator changes color
in a titration, you have reached the
endpoint.
• At the endpoint the titration is
stopped and the volume of titrant is
recorded.
• You want the equivalence point
and endpoint to be the same!
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Chapter Four
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Titration Curve
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Calculation Example 1
• Example 4.9: What volume (mL) of 0.210 M
NaOH is required to neutralize 20.00 mL of 0.103
M HCl in an acid-base titration?
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Chapter Four
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Calculation Example 2
• Example 4.10: A 10.00 mL sample of an aqueous solution
of calcium hydroxide is neutralized by 23.30 mL of
0.020000 M HNO3 (aq). What is the molarity of the
calcium hydroxide solution?
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Chapter Four
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Stop
• Complete Worksheet #4 by next
class and complete the pre-lab
for Investigation 4.
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Chapter Four
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Today…
• Turn in:
• Get Out WS#4
• Our Plan:
• Questions on WS#4
• Investigation 4
• Homework (Write in Planner):
• Lab Report Due Monday
• Post results (Average concentration of
H3PO4) to Edmodo by Friday.
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Chapter Four
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Quick Review
• Equal volumes of 0.1 M sulfuric acid and
0.2 M potassium hydroxide are mixed.
Write the balanced molecular and net ionic
equation. (p. 148)
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Chapter Four
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Today…
• Turn in:
• Questions on Lab Report so far?
• Our Plan:
• Notes – Oxidation Numbers & Redox
Titrations
• Begin WS#5 (10 minutes)
• Sample Test (Pep Assembly Today)
• Homework (Write in Planner):
• WS#5 Due Monday & Lab Report due Mon
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Chapter Four
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Redox Reactions
• Balancing Redox Equations is
complicated. We will spend an entire
week at the end of the year working on
it.
• For now, you just have to identify
substances that are oxidized and those
that are reduced.
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Chapter Four
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Reactions Involving
Oxidation and Reduction
• Oxidation: Loss of electrons
• Reduction: Gain of electrons
• Both oxidation and reduction must occur
simultaneously.
– A species that loses electrons must lose them to
something else (something that gains them).
– A species that gains electrons must gain them from
something else (something that loses them).
• To remember, use the phrase LEO GER or
OIL RIG.
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A Redox Reaction:
Electrons are transferred
from Mg metal to Cu2+
ions and …
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Mg + Cu2+ → Mg2+ + Cu
… the products
are Cu metal
and Mg2+ ions.
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Chapter Four
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Oxidation Numbers
• An oxidation number is the charge on an ion, or a
hypothetical charge assigned to an atom in a
molecule or polyatomic ion.
• Examples: in NaCl, the oxidation number of Na is
+1, that of Cl is –1 (the actual charge).
• In CO2 (a molecular compound, no ions) the
oxidation number of oxygen is –2, because oxygen
as an ion would be expected to have a -2 charge.
• The carbon in CO2 has an oxidation number of +4
(Why?)
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Rules for Assigning Oxidation Numbers
1. For the atoms in a neutral species—an
isolated atom, a molecule, or a formula unit—
the sum of all the oxidation numbers is 0.
2. For the atoms in an ion, the sum of the
oxidation numbers is equal to the charge on
the ion.
3. In compounds, the group 1A metals all have
an oxidation number of +1 and the group 2A
metals all have an oxidation number of +2.
4. In compounds, the oxidation number of
fluorine is –1.
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Rules for Assigning Oxidation Numbers
5. In compounds, hydrogen has an oxidation
number of +1.
6. In most compounds, oxygen has an oxidation
number of –2.
7. In binary compounds with metals, group 7A
elements have an oxidation number of –1, group
6A elements have an oxidation number of –2,
and group 5A elements have an oxidation
number of –3.
8. Elements in their standard state (solid, liquid, or
gas – look at PT) have an oxidation number of
0.
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Chapter Four
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Identifying Oxidation–Reduction
Reactions
• In a redox reaction, the oxidation number of a
species changes during the reaction.
• Oxidation occurs when the oxidation number
increases (species loses electrons).
• Reduction occurs when the oxidation number
decreases (species gains electrons).
• If any species is oxidized or reduced in a reaction,
that reaction is a redox reaction.
• Examples of redox reactions: displacement of an
element by another element; combustion;
incorporation of an element into a compound, etc.
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LEO - GER
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Chapter Four
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Example 4.7
What are the oxidation numbers assigned to the atoms of
each element in
(a) KClO4
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(b) Cr2O72–
(c) CaH2
(d) Na2O2
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(e) Fe3O4
Chapter Four
128
Example 1
• Which is oxidized and which is reduced in
the following equation:
Fe2O3 (s) + 3CO (g) → 2Fe (s) + 3CO2 (g)
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Example 2:
• Which is reduced/oxidized?
MnO2 + 4H+1 + 2Cl-1 → Mn+2 + 2H2O + Cl2
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Oxidation Numbers of Nonmetals
• The maximum
oxidation number of
a nonmetal is equal
to the group number.
– For nitrogen, +5.
– For sulfur, +6.
– For chlorine, +7.
• The minimum
oxidation number is
equal to the group
number – 8.
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Chapter Four
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Activity Series
of Some Metals
In the activity series, any metal
above another can displace that
other metal.
Mg metal can
react with …
… Cu2+ ions to
form Cu metal.
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Will lead metal react
with Fe3+ ions?
Will iron metal
dissolve in an acid to
produce H2 gas?
Chapter Four
132
Applications of Oxidation
and Reduction
• Everyday life: to clean (bleach) our clothes,
sanitize our swimming pools (“chlorine”),
and to whiten teeth (peroxide).
• In foods and nutrition: redox reactions
“burn” the foods we eat; antioxidants react
with undesirable free radicals.
• In industry: to produce iron, steel, other
metals, and consumer goods.
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Chapter Four
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Redox Titrations
• In a redox titration, one reactant (often the
titrant) is an oxidizing agent and the other
reactant is a reducing agent.
• Permanganate ion, usually from KMnO4, is
one of the most commonly used oxidizing
agents in the chemical laboratory and makes
an excellent titrant.
• The math is very similar to that used in
acid-base titrations.
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Chapter Four
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Calculation Example 1
• Example 4.12: A 0.2865 g sample of an iron ore is dissolved in
acid, and the iron is converted entirely to Fe+2 (aq). To titrate the
resulting solution, 0.02645 L of 0.02250 M KMnO4 (aq) is
required. What is the mass percent of iron in the ore?
5Fe+2 (aq) + MnO4-1 (aq) + 8 H+1 (aq) → 5Fe+3 (aq) + Mn+2 (aq) + 4H2O (l)
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Calculation Example 2
• Example 4.12A: Suppose the titration in Example 4.12 was carried
out with 0.02250 M K2Cr2O7 (aq) rather than KMnO4 (aq). What
volume of K2Cr2O7 (aq) would be required?
6Fe+2 (aq) + Cr2O7-2 (aq) + 14H+1 (aq) → 6Fe+3 (aq) + 2Cr+3 (aq) + 7H2O (l)
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Stop!
• Complete Worksheet #5
by next class period and
finish your titration lab
report!
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Chapter Four
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Today…
• Turn in:
• Get out WS#5 to Check
• Put Lab Report in basket (rubric on top)
• Our Plan:
• Questions on WS#5
• Worksheet Race
• Crash Course Video
• Begin Report for Investigation 8
• Homework (Write in Planner):
• Lab Report Due Friday
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Chapter Four
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Crash Course Water &
Solutions
• Good introduction to redox lab:
• http://www.youtube.com/watch?v=AN4Kif
V12DA
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Chapter Four
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Today…
• Turn in:
• Nothing
• Our Plan:
• Investigation 8
• Test Review
• Homework (Write in Planner):
– Test Next Class
– Breakfast Club Friday at 6:30 am!
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Chapter Four
141
Today…
• Turn in:
• Lab Report – rubric on top
• Our Plan:
• Questions on Test Review
• Unit 2 Test
• Homework (Write in Planner):
• Watch the Alkanes Notes online and
complete the booklet to p. 7
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Chapter Four
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Do you feel like this?
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Chapter Four