Aqueous Reactions and Solution Stoichiometry

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Chapter 4 Aqueous Reactions and Solution
Stoichiometry
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Properties of Aqueous Solutions
Definitions Solvent Solute
Electrolytic Properties Aqueous solutions
(solutions in which water is the solvent) have
the potential to . The ability of the
solution to conduct depends on the number of ions
in solution.
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The structure of water (with partial positive and
negative regions) promotes interaction with ions
and molecules.
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The ability of water to dissolve substances is
related to the structure of the water molecule.
This side will be attracted by and groups with
a d- charge.
d
d-
d
This side will be attracted by and groups with
a d charge
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What happens when NaCl (table salt) is placed in
water?
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There are three types of solute, which conduct
electricity in different ways, indicating
different chemical properties and solubilities.
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Transport of ions through solution causes flow
of current, which can be measured either visually
(qualitatively) with a lightbulb or
quantitatively with a voltmeter.
Depending on the electrolytic strength of the
solution, a lightbulb will glow brightly,
faintly, or not at all.
You will see this later on in Chapter 20
Electrochemistry!
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Most strong electrolytes ionize (or dissociate)
completely when dissolved in water.
NaCl(s) H2O ?
HClO4(l) H2O ?
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But some strong electrolytes are only somewhat
soluble in H2O
eg BaSO4 KClO4 Ca(OH)2
In these cases, there may be solid material
(undissolved) present, but the solution will
conduct electricity.
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Some compounds only ionize partially when they
dissolve in water. The ionized and unionized
forms exist in equilibrium with each other. For
example acetic acid
CH3COOH H2O CH3COO- H3O
Solution will contain all species in equation.
Such compounds are electrolytes. And
acetic acid is a acid.
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Most molecular compounds (two or more
!) dissolve in water, but do not ionize.
Sugar is a good example, as is methanol, CH3OH.
Most acids are an exception!
No ions in solution no electric charge no
conductivity. These compounds are
non-electrolytes.
Would a lightbulb glow?
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Water Solubility Guidelines
All common ionic compounds of alkali metals
(Group 1A) and ammonium (NH4) are soluble in H2O.
Ionic compounds containing the following are
soluble ALL nitrates (NO3-) and acetates
(CH3COO-) Cl-, Br-, and I- EXCEPT with Ag,
Hg22, or Pb2 SO42- EXCEPT with Ag, Hg22, or
Pb2
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Ionic compounds containing the following are
insoluble
Carbonates (CO32-), chromates (CrO42-), oxalates
(C2O42-) and phosphates (PO43-) EXCEPT with Group
1 elements and NH4 Sulphides (S2-) EXCEPT with
Group 1 and Group 2 elements, and NH4
Hydroxides (OH-) and oxides (O2-) EXCEPT with
Group 1 and 2 elements (although Ca(OH)2 and
Sr(OH)2 are only sparingly soluble)
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Precipitation Reactions
When two solutions are mixed and a solid is
formed, the solid is called a
.
The solubility guidelines allows predictions of
whether mixtures of ionic solutions will ppt or
not.
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A precipitation reaction is an example of an
or reaction, and involves
the swapping of ions in solution.
AX BY ? AY BX
Mixing a solution of AX(aq) and BY(aq) initially
gives us a solution with the ions A(aq),
B(aq), X-(aq) and Y-(aq) in solution.
We know that AX(s) and BY(s) are soluble.
A precipitate will form if and/or
is insoluble.
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Consider mixing calcium chloride and sodium
carbonate.
Precipitation reactions can be written as a
molecular equation,
All reactants and products are shown as compounds.
as a complete ionic equation,
Molecular equation rewritten to show all ions in
solution (except for solid).
or as a net ionic equation.
Spectator ions (those not participating in
reaction) are left out. Seen on both sides of
the complete ionic equation. If all ions cancel
out NO REACTION!
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Acid-Base Reactions
Acids ionize in solution to produce hydronium
ions (H)aq, which are actually protons due to H
atoms only consisting of 1p and 1e-.
Hydronium ions can either be written as H or as
an aquated water molecule, H3O. Both indicate a
substance acting as an acid.
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Acids with one acidic proton are called
monoprotic HCl hydrochloric HBr hydrobromic
HI hydroiodic HNO3 nitric HNO2 nitrous HCl
O4 perchloric HClO3 chloric HClO2 chlorous
HClO hypochlorous
When dissolved in water, reactions look like
this HA(l) H2O(l) Strength of acid
determines electrolytic nature, weak acids have a
, and are weak electrolytes. Why isnt HF
listed as a strong monoprotic acid?
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Acids with two acidic protons are called
diprotic H2SO4 sulphuric H2SO3 sulphurous
In most cases, the second (or third) proton will
come off, too, but will be in equilibrium with
the protonated species.
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are substances that add OH- to
solutions.
Group 1A and 2A hydroxides are common strong
bases and react with water as follows
NaOH(s) H2O(l)
Bases can also be compounds that do NOT contain
OH-, like ammonia (NH3), a weak base.
NH3(l) H2O(l)
Brönsted-Lowry definition
Water is acting as an acid (a proton donor)
towards NH3. NH3 is the proton acceptor.
Think ahead Would NH4(aq) act as an acid or
base?
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Strong and Weak Acids and Bases Strong acids and
bases are strong electrolytes. They are
in solution. Weak acids and
bases are weak electrolytes. They are
in solution.
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Neutralization Reactions and Salts
Neutralization occurs when acids and bases are
mixed HCl(aq) NaOH(aq) ? Notice we form a
salt (NaCl) and water.
Salt
For di- or polyprotic acids, neutralization
requires equal moles of base molecules as there
are acidic protons. H2SO4(aq) NaOH(aq) ?
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Acid-Base Reactions with Gas Formation
Sulphides, carbonates and bicarbonates (hydrogen
carbonates) are usually basic compounds which
react with H and neutralize acids. A gas is
evolved in these reactions.
2HCl(aq) Na2S(aq) ? 2HNO3(aq) K2CO3(aq)
? HClO4(aq) NaHCO3(aq) ?
Acid-base reactions with carbonate or bicarbonate
will always produce CO2 and H2O as products (as
well as a salt).
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Kitchen Chemistry

• Baking powder vs. Baking soda

Baking soda NaHCO3(s) If baking something
acidic (has yogurt, citrus, honey), use this to
make batter or dough rise!
Basic!
Baking powder two acids, some NaHCO3(s) and
cornstarch If baking a neutral cake, using
baking powder does everything for youtry
dissolving some in warm water it should fizz if
active.
Acidic AND Basic!
Self-rising flour contains baking powder!
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Oxidation and Reduction
Oxidation-Reduction (Redox) reactions involve
between reactants to form different products.
Electrons must be balanced, so if oxidation takes
place, reduction must also.
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Oxidation Numbers
Keeping track of electrons oxidation numbers are
assigned to all species in a redox reaction.

• Atoms in elemental form, oxidation number is
  zero.
• (Cl2, H2, P4, Ne are all zero)
• Monoatomic ion, the oxidation number is the
  charge on the ion.
• (Na 1 Al3 3 Cl- -1)
• Oxidation number of O is usually -2. But in
  peroxides (like H2O2 and Na2O2) it has an
  oxidation number of -1.
• Oxidation number of H is 1 when bonded to
  nonmetals and -1 when bonded to metals.
• (1 in H2O, NH3 and CH4 -1 in NaH, CaH2 and
  AlH3)
• The oxidation number of F is -1
• The sum of the oxidation numbers for the molecule
  is the charge on the molecule (zero for a neutral
  molecule).

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Example
Determine the oxidation state of all elements in
ammonium thiosulphate (NH4)2(S2O3)
(NH4)2(S2O3)
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occurs when a metal is
attacked by something in its environment.
When a metal is corroded by an acid, it loses
electrons to form cations (in salt form) and
hydrogen is released. The reaction is shown as a
balanced net ionic equation. Zn(s) 2H(aq) ?
Zn2(aq) H2(g)
Zinc is Loses electrons
Hydrogen is Gains electrons
How many electrons were transferred in the
reaction?
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Balancing Redox Reactions, Chapter 20
Usually, redox reactions are shown as net ionic
equations which may or may not be balanced.
Are the following net ionic equations
balanced? Ag(aq) Li(s) ? Ag(s)
Li(aq) Fe(s) Na(aq) ? Fe2(aq) Na(s)
Look at the electrons that are transferred in
half-reactions.
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Another example of metal oxidation by
acids 2Al(s) 3H2SO4(aq) ? Al2(SO4)3(aq)
3H2(g) During the reaction, 2H(aq) is reduced
to H2(g) and Al is oxidized from 0 to 3.
Some metals can also be oxidized by the salts of
other metals Fe(s) Ni(NO3)2(aq) ? Fe(NO3)2(aq)
Ni(s) Notice that the Fe is oxidized to Fe2
and the Ni2 is reduced to Ni.
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Some metals are easily oxidized whereas others
are not. For example, Na metal corrodes rapidly
in air, but Au does not.
Activity series a list of metals arranged in
decreasing ease of oxidation.
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The higher the metal in the activity series, the
that metal. Any metal can be oxidized by the
ions of elements below it.
Wed expect Cu2 to oxidize Fe to Fe2 or Fe to
reduce Cu2 to Cu. (If one happens, the other
MUST!!)
Elements below hydrogen coins and jewelry
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Concentrations of Solutions
Nomenclature
Solution Solute Water as solvent
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Molarity
Units mol L-1 (or mol l-1 or mol/L) mol
dm-3 M
Note 1 dm3 and 1 cm3
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Example
What is the concentration of a solution made by
dissolving 3.50 g glucose (C6H12O6) in water and
adjusting the volume to 100.0 ml? Data molar
mass glucose 180.16 g/mol
n m/MW
c n/V
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Example
What is the concentration of Na ions in a 0.250
mM solution of sodium phosphate?
When Na3PO4 dissolves Na3PO4 H2O ?
Hence Na 0.750 mM
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Example

• Explain how to prepare 500 mL of a 0.25 M
  solution of NaCl.

Data molar mass of NaCl 58.44 g/mol
We know that Molarity moles of solute
liters of solution
Therefore, after rearranging, moles
This mass would be placed in a 500 mL volumetric
flask and water would be added to the mark.
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Dilutions
Example
What volume of a 1.00 M CuSO4 solution must be
diluted to 250 ml in order to produce a solution
which is 0.100 M in CuSO4?
Several approaches!
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What volume of a 1.00 M CuSO4 solution must be
diluted to 250 ml in order to produce a solution
which is 0.100 M in CuSO4?
Using moles
How many moles CuSO4 in target (dilute) solution?
n c x V
What volume of stock (concentrated) solution
contains this number of moles?
V n/c 25.0 cm3 or 25.0 mL
Moles taken from stock solution moles needed in
dilute solution Only VOLUME changes! USE THIS
APPROACH IN YOUR LAB!! Quick calculation c1V1
c2V2 (Use to check your result in lab.)
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What volume of a 1.00 M CuSO4 solution must be
diluted to 250 ml in
order to produce a solution which is 0.100 M in
CuSO4?
Using c1V1 c2V2
What do we know? What are we calculating? c1
1.00 M V1 ?c2 0.100 MV2 250 mL
Think of it as before and after (with all 1s
being conc and 2s being dilute).
25.0 mL
Big benefit equation works for ANY type of
dilution, as long as the same units are on both
sides of the equation!
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For very dilute solutions
There are 103 mg in 1 g There are 103 g in 1 kg
Hence there are mg in 1 kg
Define this as 1 part per million (1 ppm)
1 ppm 1 mg solute/1 kg solution 1 mg solute/1
liter solution
And even smaller1 part per billion (ppb) 1 µg
solute/1 liter solution
c1V1 c2V2 even works for ppm and ppb dilutions!
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Solution Stoichiometry and Chemical Analysis
Gravimetric analysis

• produce precipitate
• collect dry
• weigh

Example calculate amount of SO42- in solution

• treat with BaCl2(aq)
• precipitate sulphate as BaSO4(s)
• collect, dry, and weigh
• 1 mol BaSO4 is equivalent to 1 mol SO42- in
  solution

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Titrations
Determine unknown concentrations of analytes
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Titrations
Known volume conc of solution 1 reacted
with known volume unknown conc of solution 2

• calculate unknown conc.

Equivalence point
End point of reaction signaled with an indicator
(color change) or a quantitative measurement
(voltage). Achieving a useful endpoint can be
tricky and takes patience go SLOWLY in lab!!
Equivalence point and end point should be !!
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Titration Example
It takes 28.75 ml of a 0.400 M solution of NaOH
to completely neutralize 25.00 ml of a H2SO4
solution. Calculate H2SO4.
2NaOH(aq) H2SO4(aq) ?
28.78 mL 0.400 M
25.00 mL ?? M

• cn/V so n cV

• equation stoichiometry

• c n/V
• 0.230 mol dm-3 or 0.230 M

Please note You cannot directly apply c1V1
c2V2 to this problem due to stoic!!
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END OF CHAPTER 4