Chapter 4
Reactions in Aqueous Solutions
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Many chemical and almost all biological
reactions occur in the aqueous medium
Substances (solutes) that dissolve in water
(solvent) can be divided into two
categories:
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Electrolytes
Non-Electrolytes
Three Major Types of Reactions
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Precipitation Reaction – the product an
insoluble substance separates from the
solution
Acid/Base Reactions – A proton transfer
from an acid to a base
Oxidation/Reduction (Redox) “the bane of
the AP Test” – Electrons are transferred
from a reducing agent to an oxidizing agent
Solution Stoichiometry
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Quantitative studies with known
concentrations (Molarity) of solutions
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Gravimetric Analysis
Titrations
General Properties of a Solution
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Solution – a homogenous mixture of two or
more substances
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Solution may be gaseous (air), solid (alloy) or
Liquid (salt water)
In this chapter we will deal only with
aqueous solutions
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Most common Solvent - Water
Electrolytes versus Nonelectrolytes
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Electrolytes – Ionic compounds that
completely or partially dissociate in solution
with the ability to pass electric current in
solution
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Acids/Bases will ionize in solution, therefore
electricity can be conducted
Non-Electrolytes – Molecular compounds
that do not dissociate in solution, therefore
no electric current can be pass
Ionic Compounds in Solution
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Water is a great solvent for ionic
compounds because it is polar, the positive
end attracts the Negative Ion and vice
versa
Acids and Bases as Electrolytes
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Some acid/bases competely dissociate in solution
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HCl
HNO3
H2SO4
Ba(OH)2
NaOH
While others only partially dissociate
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CH3COOH
HF
HNO2
HN3
Writing Partial Dissociation Equations
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Partial dissociation equations are written
with a double arrow, indicating a reversible
reaction
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Write partial dissociation for CH3COOH
Precipitation Reactions
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A double replacement reaction (metathesis)
in which a product is insoluble
Solubility Rules
In water at 25 Degrees
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All common compounds of Group I and ammonium ions are
soluble.
All nitrates, acetates, and chlorates are soluble.
All binary compounds of the halogens (other than F) with
metals are soluble, except those of Ag, Hg(I), and Pb.
All sulfates are soluble, except those of barium, strontium,
calcium, lead, silver, and mercury (I). The latter three are
slightly soluble.
Except for rule 1, carbonates, hydroxides, oxides, silicates,
and phosphates are insoluble.
Sulfides are insoluble except for calcium, barium, strontium,
magnesium, sodium, potassium, and ammonium.
Soluble or Insoluble at 25 Degrees
Celsius in Water
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PbSO4
BaCO3
Li3PO4
FeS
Ca(OH)2
Co(NO3)3
Net Ionic Equations
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Write the correctly balanced equation and
decide on state of each product
Write free state of all ions and insoluble
product
Cancel out spectator ions – anyone not part
of the reaction
Check charges and balancing in net ionic
Practice Net Ionic
Predict, Balance and write net ionic
1)
2)
3)
4)
Lead Nitrate and Potassium Iodide
Barium Chloride and Sodium Sulfate
Potassium Phosphate and Calcium
Nitrate
Aluminum Nitrate and Sodium Hydroxide
Acid – Base Reactions
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Acids react with metal such as Zn, Mg and
Fe to produce hydrogen gas
Acids react with carbonates and
bicarbonates to produce carbon dioxide
gas, water and the salt
Bronsted Acid and Bases
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Bronsted Acid is a proton donor
Bronsted Base is a proton acceptor
HCl (aq)
H+ (aq) + Cl-(aq)
In water the H+ attracts to the water molecule
producing the hydronium ion
Monoprotic Acids
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Each unit of acid yields one hydrogen ion
upon ionization
Diprotic Acids
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Each unit of the acid gives up two hydrogen
ions in two separate steps (they strip)
Triprotic Acids
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Yield three hydrogen ions in three separate
steps (they strip)
Bronsted Acid is a proton donor
Bronsted Base is a proton acceptor
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Classify each of the following as an
Bronsted acid or Bronsted base, explain
your reasoning based on the definition
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HBr
-2
SO 4
HI
HCO 3
NO2
Neutralization Reaction
Acid and Base will form Salt and
Water
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Write the net ionic for the following
Hydrochloric acid and Sodium Hydroxide
Sulfuric acid and Aluminum Hydroxide
Acid – Base Reactions Leading to
Formation of a Gas
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Certain Salts – Carbonates, bicarbonates,
sulfites and sulfides react with acids to form
gaseous products
Oxidation Numbers
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Oxidation Reaction – refers to half reaction that
involves loss of electrons
Reduction reaction – refers to a half reaction that
involves the gain of electrons
The extent of oxidation in a redox reaction must
be equal to the extent of reduction; that is the
number of electrons lost by a reducing agent must
be equal to the number of electrons gained by an
oxidizing agent
The half-reactions of a redox reaction
or oxidation-reduction reaction
Oxidation Number
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The number of charges the atom would
have in a molecule if electrons are transfer
completely
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The convention is that the cation is written first in a formula, followed by the
anion.
For example, in NaH, the H is H-; in HCl, the H is H+.
The oxidation number of a free element is always 0.
The atoms in He and N2, for example, have oxidation numbers of 0.
The oxidation number of a monatomic ion equals the charge of the ion.
For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3.
The usual oxidation number of hydrogen is +1.
The oxidation number of hydrogen is -1 in compounds containing elements that are
less electronegative than hydrogen, as in CaH2.
The oxidation number of oxygen in compounds is usually -2.
Exceptions include OF2, since F is more electronegative than O, and BaO2, due to
the structure of the peroxide ion, which is [O-O]2-.
The oxidation number of a Group IA element in a compound is +1.
The oxidation number of a Group IIA element in a compound is +2.
The oxidation number of a Group VIIA element in a compound is -1, except
when that element is combined with one having a higher electronegativity.
The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in
HOCl.
The sum of the oxidation numbers of all of the atoms in a neutral compound is
0.
The sum of the oxidation numbers in a polyatomic ion is equal to the charge of
the ion.
For example, the sum of the oxidation numbers for SO42- is -2.
Assign oxidation numbers to all the
elements in the following compounds
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Na2O
HNO2
-2
Cr2O7
PF3
MnO4
Arrange the following species in order
of increasing oxidation number of the
sulfur atoms
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H2S
SO2
SO3
S8
H2SO4
-2
S
HS
Concentration
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Molarity = moles of solute/liters of solution
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Example: How many grams of potassium
dichromate are required to prepare a 125ml
solution whose concentration is 1.83M
Concentration
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In a biochemical assay a chemist needs a
to add 4.07g of glucose to a reaction
mixture. Calculate the volume in milliliters
the volume of a 3.16M glucose she should
use
Dilution of Solutions
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The procedure of making a less
concentrated solution from a high
concentration solution
MiVi = MfVf
Dilution Problem
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Describe hou you would prepare 2.50 * 102
ml of a 2.25M H2SO4 solution, starting with
a 7.41 M stock solution of H2SO4
Dilution Problem #2
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How would you prepare a 200ml of a .866M
KOH solution, starting with 5.07M stock
solution
Acid – Base Titrations
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In a titration a solution of an accurately
known concentration, called the standard is
added gradually to another solution of
unknown until reaction is neutralized
(equivalence point)
Indicators are used to color the reaction
when it is complete
Titration Problem
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In a titration experiment, a student finds
that 25.46ml of a NaOH solution is needed
to neutralize .6092g of KHP. What is the
concentration of the NaOH solution?
Titration Problem #2
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How many milliliters of a .836M NaOH
solution is needed to neutralized 25ml of a
.335M of H2SO4?
Solution Stoichiometry
When sodium chloride reacts with silver
nitrate, silver chloride precipitates. What
mass of silver chloride is produced from
150ml 3M of silver nitrate?
1.
When Magnesium chloride reacts with
silver nitrate, silver chloride precipitates.
What mass of silver chloride is produced
from 4.5M in 250ml of silver nitrate? What
is the name of the other product of the
reaction?