Transcript Document 7642264

Chapter 4
Types of Chemical Reactions and
Solution Stiochiometry
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Preview
the contents of this chapter will introduce
you to the following topics:
 Water, Nature of aqueous solutions,
types of electrolytes, and dilution
 Types of chemical reactions:
precipitation, acid base reactions and
oxidation-reduction reaction
 Stoichiometry of reactions and
balancing the chemical equations
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4.1 Water, the Common Solvent
Water is one of the most important substances on earth:
– Cooling – engine, nuclear power plants, and many.
– Transportation …etc.
Water dissolve many different substances,
e.g. salts, sugar, and many other
To understand this process, we need to consider the nature of water:
– As molecule, is H2O
– Shape is V-shape with angle 105o
– Band type of each O – H is covalent band and polar [electrons
are not equirdlntty shared]
– Polarity is polar with +ve charges of hydrogen and -ve on
oxygen.
This Polarity of water gives it the greatest ability to dissolve
compounds.
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4.1 Water, the Common Solvent
• This Polarity of water gives it the greatest ability to dissolve
compounds. Figure 4.2 shows schematic ionic solid dissolving in
water. This process is called "Hydration".
Note:
• The stronger the ion-water attraction, the higher the solubility.
• Therefore, not all the solid have the same solubility [chapter 11].
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4.1 Water, the Common Solvent
Water also dissolve non-ionic substances e.g. alcohols – "polar" and
compatible structure to water: "like – dissolve – like"
Figure 4.3
many substances don't dissolve in water e.g. animal fats "non-polar"
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4.2 The Nature of Aqueous Solutions:
Strong and Weak Electrolytes & non–Electrolyte
 Solution – composite of solute (substance to be dissolved) +
Solvent (Major substance e.g. water).
 Solute:



dissolves in water (or other “solvent”)
changes phase (if different from the solvent)
is present in lesser amount (if the same phase as the solvent)
 Solvent:


retains its phase (if different from the solute)
is present in greater amount (if the same phase as the solute)
Solution
Solvent
Solute
Soft drink (l)
H2O
Sugar, CO2
Air (g)
N2
O2, Ar, CH4
Soft Solder (s)
Pb
Sn
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4.2 The Nature of Aqueous Solutions:
Strong and Weak Electrolytes & non–Electrolyte
 One major property for characterizing aqueous solutions is its "Electrical
conductivity" or its ability to conduct an electric current. Three solutions are
observed: non electrolytes, weak, electrolytes and strong electrolyte. figure
4.4
 The basis for conductivity properties of solutions was first correctly identified
by Svante Arrhenius (1859 – 1927)
nonelectrolyte
weak electrolyte
strong electrolyte
An electrolyte is a substance that, when dissolved in
water, results in a solution that can conduct electricity.
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4.2 The Nature of Aqueous Solutions:
Strong and Weak Electrolytes & non–Electrolyte
 Arrhenius postulated that the electric current depend
directly on the number of ions present.
Strong Electrolyte – 100% dissociation
NaCl (s)
H 2O
Na+ (aq) + Cl- (aq)
Weak Electrolyte – not completely dissociated
CH3COOH
CH3COO- (aq) + H+ (aq)
Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s)
H 2O
C6H12O6 (aq)
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4.3 The Composition of Solutions:
To perform stoichiomctric calculations at any chemical reactions
you must know two things:
– The nature of the reaction – exact forms of chemical in solutions.
– The amounts of chemicals – "concentration"
concentration of a solution can be described in many different
ways, % , molar, molal, mole. faction .. etc.
We will consider here the unit "molar"(M) or molarity: "the method
used to prepare molar solutions".
Molarity (M) = moles of solute
per volume of solution in liters:
M  molarity 
moles of solute
liters of solution
6 moles of HCl
3 M HCl 
2 liters of solution
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4.3 The Composition of Solutions:
Notes:
 For ionic systems the solution prepared will contain the number of
moles prepared but for ionic species its different story
e.g. 1.0M NaCl contains 1.0 mole NaCl or more accurate 1.0 mole of
Na+ and 1.0 mole of Cl-.
 Molarity can be used to determine number of moles per certain
volumes where:
Moles = Liters of solution x Molarity.
Example 4.3
Gives the concentration of each type of ion in the solutions of
0.50M Co(NO3)2
Example 4.4
Calculate the number of moles of Cl- ions in 1.75 L of 1.0x10-3M
ZnCl2
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4.3 The Composition of Solutions:
Standard Solution:
Solution used in chemical analysis. it has accurately known concentration
Dilution:
it is the procedure to get low concentrated solution (diluted) from a high
concentrated one.
Moles of solute after dilution = moles of solute before dilution.
(M x V) after = (M x V) before.
Example 4.7
What volume of 16 M sulfuric Acid must be used to
prepare 1.5 L of a 0.10M H2sO4 solution?
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4.4 Types of Solution Reactions
Millions of possible chemical reaction needs system for
grouping them into classes. The commonly used by chemists:
•Precipitation reactions
AgNO3(aq) + NaCl(aq)  AgCl(s) +
NaNO3(aq)
•Acid-base reactions
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
•Oxidation-reduction reactions
Fe2O3(s) + Al(s)  Fe(s) + Al2O3(s)
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4.5 Precipitation Reactions
Precipitate – insoluble solid that separates from solution
precipitate
Pb(NO3)2 (aq) + 2NaI (aq)
PbI2 (s) + 2NaNO3 (aq)
molecular equation
Pb2+ + 2NO3- + 2Na+ + 2I-
PbI2 (s) + 2Na+ + 2NO3-
ionic equation
Pb2+ + 2IPbI2
PbI2 (s)
net ionic equation
Na+ and NO3- are spectator ions
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4.5 Precipitation Reactions
Simple Rules for Solubility of Aq. Solutions
1.
2.
3.
4.
5.
6.
Most nitrate (NO3) salts are soluble.
Most alkali (group 1A) salts and NH4+ are soluble.
Most Cl, Br, and I salts are soluble (NOT Ag+, Pb2+, Hg22+)
Most sulfate salts are soluble
(Except BaSO4, PbSO4, HgSO4, CaSO4)
Most OH salts are only slightly soluble
(Except NaOH, KOH are soluble)
Most S2, CO32, CrO42, PO43 salts are only slightly soluble, i.e.,
Not soluble.
Exercise 4.8:
Using the solubility rules in table 4.1, predict what will
happen when the following pairs of solutions are mixed.
a. KNO3(aq) & BaCl2(aq)
b. Na2SO4(aq) & Pb(NO3)2(aq)
c. KOH(aq) & Fe(NO3)(aq)
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4.6 Describing Reactions in Solution
Writing Net Ionic Equations
1.
Write the balanced molecular equation.
2.
Write the ionic equation showing the strong electrolytes
3.
Determine precipitate from solubility rules
4.
Cancel the spectator ions on both sides of the ionic equation
Write the net ionic equation for the reaction of silver
nitrate with sodium chloride.
AgNO3 (aq) + NaCl (aq)
AgCl (s) + NaNO3 (aq)
Ag+ + NO3- + Na+ + Cl-
AgCl (s) + Na+ + NO3-
Ag+ + Cl-
AgCl (s)
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4.7 Stoichiometry of Precipitation Reactions
The procedures for calculating quantities of reactants and
products involved in chemical reaction. The following steps
summarized the procedure:
Step 1: Identify the present in the combined solution, and
determine what reaction occurs.
Step 2: write the balanced net ionic equation.
Step 3: calculate the moles of reactants.
Step 4: determine which reactant is limiting.
Step 5: calculate the moles of product or product as required.
Step 6: convert to grams or other units, as requires.
Example 4.11
When aqueous solutions of Na2SO4 and Pb(NO3)2 are mixed,
PbSO4 precipitates. Calculate the mass of PbSO4 formed
when 1.25L of 0.0500M Pb(NO3)2 and 2.00L of 0.0250M
Na2SO4 are mixed.
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4.8 Acid – Base Reactions:
Acids
Have a sour taste.
Taste of vinegar is due to acetic acid.
Citrus fruits contain citric acid.
React with certain metals to produce hydrogen gas.
React with carbonates and bicarbonates to produce carbon dioxide gas
Bases
Have a bitter taste.
Feel slippery. Many soaps contain bases.
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4.8 Acid – Base Reactions:
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
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4.8 Acid – Base Reactions:
A Brønsted acid is a proton donor
A Brønsted base is a proton acceptor
base
acid
acid
base
A Brønsted acid must contain at least one
ionizable proton!
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4.8 Acid – Base Reactions:
They are also called neutralization reaction
acid + base
salt + water
Describing Reactions in Solution
HCl (aq) + NaOH (aq)
H+ + Cl- + Na+ + OHH+ + OH-
NaCl (aq) + H2O
Na+ + Cl- + H2O
H2O
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4.8 Acid – Base Reactions:
They are also called neutralization reaction
x acid + y base
salt + water
Main reaction is titration, the key terms are:
Titrant - solution of known concentration used in titration
Analyte - substance being analyzed
Equivalence point - enough titrant added to react exactly with
the analyte
Endpoint - the indicator changes color so you can tell the
equivalence point has been reached.
The neutralization reaction calculation:
a. Write the correct balanced acid–base reaction.
b. Use the following equation:
y. (M.V)acid = x. (M.V)base
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4.8 Acid – Base Reactions:
In a titration a solution of accurately known
concentration is added gradually added to another
solution of unknown concentration until the chemical
reaction between the two solutions is complete.
Equivalence point – the point at which the reaction is complete
Indicator – substance that changes color at (or near) the
equivalence point
Slowly add base
to unknown acid
UNTIL
the indicator
changes color
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4.7
4.9 Oxidation-Reduction Reactions:
(electron transfer reactions)
2Mg (s) + O2 (g)
2Mg
2MgO (s)
2Mg2+ + 4e-
Oxidation half-reaction (lose e-)
O2 + 4e2O2Reduction half-reaction (gain e-)
2Mg + O2 + 4e2Mg + O2
2Mg2+ + 2O2- + 4e2MgO
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4.9 Oxidation-Reduction Reactions:
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4.9 Oxidation-Reduction Reactions:
Rules for Assigning Oxidation States
1. Oxidation state of an atom in an element = 0
2. Oxidation state of monatomic element ions = charge
3. Oxygen =-2 in covalent compounds
(except in peroxides where it = -1)
4. H = +1 in covalent compounds
5. Fluorine = -1 in compounds
6. Sum of oxidation states = 0 in compounds Sum of oxidation states
= charge of the ion
Oxidation numbers of
the elements in the
following ?
IF7
F = -1
7x(-1) + ? = 0
I = +7
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4.10 Balancing Oxidation – Reduction Equations
Balancing by Half-Reaction Method (Acidic)
1. Write separate reduction, oxidation reactions
2. For each half-reaction:
•
•
•
•
Balance elements (except H, O)
Balance O using H2O
Balance H using H+
Balance charge using electrons
3. If necessary, multiply by integer to equalize electron
count
4. Add half-reactions
5. Check that elements and charges are balanced
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4.10 Balancing Oxidation – Reduction Equations
Half-Reaction Method - Balancing in Base
1.
2.
3.
4.
Balance as in acid.
Add OH- that equals H+ ions (both
sides!)
Form water by combining H+, OHCheck elements and charges for balance
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4.10 Balancing Oxidation – Reduction Equations
MnO4- + ClO2- → MnO2 + ClO4-
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