Transcript CHEM1310 Lecture

Chapter 4
Types of Chemical Reactions and Solution Stoichiometry
4.1 Water, the Liquid of Life
4.2 Aqueous Solutions: Strong and Weak Electrolytes
4.3 The Composition of Solutions
4.4 Types of Chemical Reactions
4.5 Precipitation Reactions
4.6 Describing Reactions in Solution
4.7 Selective Precipitation
4.8 Stoichiometry of Precipitation Reactions
4.9 Acid-Base Reactions
4.10 Oxidation-Reduction Reactions
4.11 Balancing Oxidation-Reduction Equations (skip)
4.12 Simple Oxidation-Reduction Titrations (skip)
Water, the Liquid of Life (part 1)
Water is polar. In a water
molecule, each oxygen atom is
partially negative. Each
hydrogen atom is partially
positive. This phenomenon is
called charge separation.
.
Oxygen is an electron hog. It does not
share electrons well with other atoms.
Oxygen takes more than its share of
electron density, that is why it has a
partial negative charge. The hydrogen
atoms are stripped of much of their
electron density and so carry a partial
positive charge.
In liquid water and especially in solid water, the
molecules interact strongly with each other,
with preferred orientations.
Water, the Liquid of Life (part 2)
• Water has all sort of strange and unusual
properties.
1) Its density decreases when it freezes (ice
floats).
2) It has a high boiling point and high heat of
vaporization.
3) It has high surface tension.
4) It dissolves many salts (like sodium chloride)
and polar molecules (like ethanol).
5) It does not dissolve non-polar substances (oil
and water don’t mix).
6) It has high heat capacity
7) In the presence of amphipaths, it readily
forms compartments (like in cells).
8) Where there is liquid water, there is life (I
bet).
Water, the Liquid of Life (part 3)
Water dissolves NaCl.
NaCl dissolves in water. To dissolve is to mix
at the level of individual ions and molecules.
A solution is a homogeneous mixture of two or more substances. A solution may exist in
any phase.
Water is the solvent
The solvent is the component of a solution that is present in the greatest amount.
NaCl is the solute
The solute is the substance that is dissolved in a solution.
Pairs of liquids that mix in any proportion are miscible.
Liquids that do not mix are immiscible.
Ice
Classification of Molecules
1) Non-polar: No charge separation, every atom is
neutral. Examples: graphite = Cbig, hexane =
C8H18, or olive oil.
2) Polar: With charge separation, some atoms
carry partial charge.
Examples: water = H2O, ammonia = NH3,
fructose = C6H12O6.
3) Ionic Compounds: With Formal (unit) charge.
Ionic Compounds (Salts)
Examples:
Na+Cl-, Na+Br-, Mg2+Cl-2, NH4+NO3-, CH3COO-Na+
Salts contain charged (unit charge, not partial
charge) atoms or groups of atoms.
Cations: Na+, K+, Mg2+ Ca2+, NH4+
Anions: Cl-, F-, SO42-, PO43-, CH3COO-
Properties of Ionic Compounds:
All ionic compounds are solids in the range of temperature in
which water is a liquid. They have rigid lattices in which
strong electrostatic forces (ionic bonds) hold the constituent ions
in place.
The high melting points of ionic compounds indicate that a good
deal of energy must be supplied to destroy the lattice and produce
a liquid (molten) form in which the ions move more freely.
Solubility in Water: Water molecules are polar.
When water interacts with an ionic compound such
as NH4NO3, or NaCl, the partially negative ends
of some of the water molecules attract the positive
ions (cations) while the positive poles of others
attract the negative ions (anions).
Ionic Compounds in Water:
An ionic solute is said to dissociate into ions upon dissociation. When a
salt dissolves in water ionic bonds are broken. When a salt precipitates
from water, ionic bonds are formed
NaCl (s) + H2O (l) → Na+ (aq) + Cl– (aq)
Water molecules interact favorably with both cations and
anions.
Water “solvates” ions. The negative pole of water interacts
with cations. The positive pole of water interacts with
anions.
Lies my Chemistry Teacher told me:
"Like dissolves like”
substances with similar intermolecular attractive forces tend to be soluble in
one another.
Water is polar: polar molecules like fructose dissolve in water.
Hexane is non-polar: non-polar molecules do not dissolve in water.
But DMSO and acetone are polar. Hexane (non-polar) dissolves in DMSO
(polar) and acetone (polar).
Reality:
Water is special. It is a very polar solvent, but that is
only part of the story. Non-polar substances are
driven out of water by “the hydrophobic effect”,
which is a property of water, and of water only. The
molecular interactions between water molecules
drive non-polar solutes out of solution. The
interactions between solute molecules are not the
issue. Hexane does not care whether it is sitting
next to other hexanes or next to water molecules.
Water is unstable when it is directly adjacent to
hexane.
Electrolytes and Non-Electrolytes:
Electrolytes: Mainly ionic compounds such as potassium
sulfate (K2SO4) and sodium chloride (NaCl),
which, as a solute, increase the electrical
conductance of water when they dissolve.
Non-Electrolytes: Do not increase the conductance of
water when they dissolve.
Strong electrolytes: Increase
conductivity more
than weak
electrolytes.
The Composition of Solutions
The Molarity (M) of a solute in a solution is the amount
of the solute (n in moles) divided by the total volume (V
in liters) of the solution.
Example: Obtain 180 g of fructose. Add it to a volumetric
flask. Add water, with stirring, until the total volume is 1
L. What is the molarity?
MWt = 6x12 + 12 + 6x16 = 180 g/mol
n = 180 g (180 g/mol)-1 = 1 mol
M = 1 mol/1L = 1 = mol L-1
How to make a mistake: If you add 180 g of fructose to 1 L of water,
the final volume will be greater than 1 L, and the concentration will be
< 1 M.
Calculate the molarity (M in mol / L) of a solution
obtained by dissolving 10.0 g of Al(NO3)3 in enough
water to make 250.0 mL of solution.
nAl(NO3)3 = mass Al(NO3)3 g x [mol weight g/mol]-1
Vsolution = given
MAl(NO3)3 = M (M = mol/L
1. Precipitation Reactions (of ionic
compounds)
–
–
Types of
Chemical
Reactions
Ionic Equations and Net Ionic Equations
Predicting Precipitation Reactions
2. Acids and Bases (move protons)
–
–
–
–
–
–
Arrhenius Acids and Bases Theory
Strong and Weak acid
Naming Acids
Weak Bases
Modifying the Arrhenius model
Acid-Base Titrations
3. Oxidation-Reduction Reactions (move
electrons)
–
–
–
Oxidizing and Reducing Agents
Oxidation Number
Types of Redox Reactions
» Combination and Decomposition
» Oxygenation
» Hydrogenation
» Displacement Reactions
» Disproportionation
Predicting Dissolution and Precipitation Reactions
CHEM1310 McKelvy Lecture
Ionic Equations (precipitation equation)
What happens when you add solutions of barium chloride and potassium
sulfate?
BaCl2(aq) + K2SO4(aq) → ???
1) Identify all the possible ions, rewrite the equation with explicit ions.
Ba2+(aq) + 2 Cl-(aq) + 2 K+(aq) + SO42-(aq) → ???
2) Look at a table of solubilities. Notice that barium sulfate is insoluble. Write the
‘complete ionic equation.
Ba2+(aq) + 2 Cl-(aq) + 2 K+(aq) + SO42-(aq) →
BaSO4(s) + 2 K+(aq) + 2 Cl-(aq)
3) Identify the ‘spectator ions’ (chloride and potassium).
4) Write out the “Net Ionic Equation”.
BaCl2(aq) + K2SO4(aq) → BaSO4(s) + 2 KCl(aq)
Write a net ionic equation to represent the formation of
the precipitate observed when aqueous solutions of
CaCl2 and NaF are mixed. Identify the precipitate and
the spectator ions.
CaCl2(aq) + NaF(aq) → ???
Rewrite with explicit ions (complete ionic equation.
Ca2+(aq) + 2Cl-(aq) + Na+(aq) + F-(aq) → ???
Check the solubility table. Note that CaF2 is insoluble. Balance the equation.
Ca2+(aq) + 2 Cl-(aq) + 2 Na+(aq) + 2 F-(aq) →
CaF2(s) + 2 Na+(aq) + 2 Cl-(aq)
Na+ and Cl- are spectator ions. CaF2 is the precipitate.
Summary: Three Types of Equations Used to
Describe Reactions in Solution
• Molecular Equation
BaCl2(aq) + K2SO4(aq) → BaSO4(s) + 2 KCl(aq)
• Complete Ionic Equation
Ba2+(aq) + 2 Cl-(aq) + 2 K+(aq) + SO42-(aq) →
BaSO4(s) + 2 K+(aq) + 2 Cl-(aq)
• Net Ionic Equation
Ba2+(aq) + SO42-(aq) → BaSO4(s)
Acids and Bases and Their Reactions
(more in Chapter 7)
Acid Base Theory
1. Arrhenius Acids and Bases Acids are
•
•
H+ donors
Bases are OH- donors
2. Arrhenius Broadened Definition
•
•
Acids increase H+ concentration or [H+]
increases
Bases increase OH- concentration or [OH-]
increases
3. Brønsted-Lowry Acids and Bases (1923)
•
•
Acids donate H+
Bases accept H+
Arrhenius
1903
Nobel Prize
Acids and Bases
Neutralization reaction:
HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) →
H2O(l) + Na+(aq) + Cl-(aq)
HCl (the acid) has donated a proton (H+) to OHOH- (the base) has accepted a proton from HCl
Neutralization occurs at the equivalence point, where
nbase = nacid
number of moles of acid = number of moles of base
The titration of an acid with a base.
Base (Titrant):
Standard Solution
Indicator
End point =
Equivalence point =
Stoichiometric point
Example Equivalence Point Calculation
Compute the molarity of a solution of sodium hydroxide if 26 ml
of solution must be added to a solution containing 0.53 g of
KH5C8O4 (KHP, potassium hydrogen phthalate) to reach the
end-point.
If you see “end-point” or “equivalence point”, think nacid = nbase
1) nacid = 0.53 g (39 + 5 + 96 + 76 g/mol)-1 =2.5 mmol
2) nbase = 2.5 mmol NaOH
[NaOH] = 2.5 mmol / 26 mL
= 0.10 mmol/ml
= 0.10 mol/L
= 0.10 M
[x] is molarity of x
Arrhenius Acids and Bases:
Strong acids: Hydrochloric acid (HCl)
Sulfuric acid (H2SO4)
Perchloric acid (HClO4)
Hydroiodic acid (HI)
Hydrobromic acid (HBr)
Nitric acid (HNO3)
Chloric acid (HClO3)
Strong bases: Sodium hydroxide (NaOH)
Potassium hydroxide (KOH)
Strong acids and bases are strong electrolytes.
Strong acids and bases dissociate to ions completely in water.
Weak Acids
Formic acid (HCOOH)
Acetic acid (CH3COOH)
Phosphoric acid (H3PO4)
Hydrofluoric acid (HF) Oxalic acid (H2C2O4)
CH3COOH(aq) → CH3COO−(aq) + H+(aq)
(partial dissociation of acetic acid)
Neutralization Reaction (weak or strong acid plus strong base)
CH3COOH(aq) + NaOH(aq) → CH3COO− (aq) + H2O(l)
(molecular eq.)
CH3COOH(aq) + Na+(aq) + OH-(aq) →
Na+(aq) + CH3COO−(aq) + H2O(l)
(complete ionic eq.)
Weak acids partially dissociate in water. Some on the acid(aq) remains.
Weak acids are weak electrolytes.
Weak Bases
Ammonia (NH3)
acetate CH3COO−
NH3(aq) + H2O(l) → NH4+(aq) + OH-(aq)
CH3COO-(aq) + H2O(l) → CH3COOH(aq) + OH-(aq)
Weak acids partially dissociate in water.
Weak base are weak electrolytes.
Oxidation-Reduction Reactions
Redox: an extensive and important class of reactions that is
characterized by the transfer of electrons.
2 Mg(s) + O2(g) → 2 MgO(s)
Magnesium is oxidized: it gives up electrons as the oxidation
state of its atoms increases from zero to +2.
Oxygen is reduced: it gains electrons as the oxidation state
decreases from zero to -2 (i.e., becomes more negative).
Term
Oxidation
Reduction
Oxidizing Agent, does
the oxidizing
Reducing Agent, does
the reducing
Substance Oxidized
Substance Reduced
Oxidation
Electron
Number Change Change
Increase
Loss of
Electrons
Decrease
Gain of
Electrons
Decrease
Picks Up
electrons
Increase
Supplies
Electrons
Increase
Loses
Electrons
Decrease
Gains
Electrons
•
Oxidation States: Examples
1. NaCl
2. CrO3
3. TlCl3
4. Mn3N2
5. O2
6. H2
7. CH4
8. CO2
9. C6H12O6
10. MnO4-
Oxidation-Reduction reactions
Oxidation Number
Change
Electron Change
Oxidizing Agent, does the oxidizing
Decrease
Picks Up electrons
Reducing Agent, does the reducing
Increase
Supplies Electrons
Substance Oxidized
Increase
Loses Electrons
Substance Reduced
Decrease
Gains Electrons
Term
O2
+
H2
→
H2O
Cl2
+
Na
→
NaCl
H+
+
Mg
→
Mg++ + H2
A summary of an
oxidation-reduction
process, in which M
is oxidized and X is
reduced.