Physical Chemistry
Dónal Leech
donal.leech@nuigalway.ie
Ext 3563
Room C205, Physical Chemistry
Notes for downloading (powerpoint and word)
http://www.nuigalway.ie/chem/Donal/Teaching.htm
Chemistry
Physical
Sciences
Biological
Sciences
Organisms
Sub-atomic
Organs
Atoms
Tissues
Chemistry
Materials
Cells
Atmosphere
DNA
Stellar
Molecules
Molecular Bonds
Sciences Forces
Physical Chemistry
Establishes and develops the principles that
are used to explain and interpret the
observations made in chemistry
Bulk

Thermodynamics 
Equilibrium

Individual
 Rates
 Chemical reactions
Quantum
mechanics &
spectroscopy
Structure
ENERGY

Change
Textbook
• Brown, LeMay, Bursten
• Chemistry: The Central Science, 9th or 10th
Edition
Companion Web-site
http://www.prenhall.com/brown
Aqueous Reactions
Chapter 4
Reactions in Solution
The most important substance on earth is water.
In chemistry, water is necessary for many reactions to take place.
Table salt (NaCl) when put into water dissolves into its ions, Na+ and Cl-.
Water is the solvent and NaCl, is the solute.
The mixture is an aqueous solution.
The Water Molecule allows many
substances to be dissolved in them.
One side of the water molecule is
negatively charged, and the other side is
positively charged. Water is a polar
molecule.
Electricity and Solutions
A useful characteristic of solutions is the ability to
conduct electricity. To determine if a solution has the
ability to conduct electricity, an electrical conductivity
apparatus is used. An electrical conductivity apparatus
is basically a battery and light bulb setup which lights
up when electricity is conducted through the solution.
Electrolytes are substances that produce ions upon
dissolving. There are two ways to provide these
mobile ions for conducting purposes.
1. Dissociation of Ionic Compounds
2. Ionisation of Polar Covalent Molecular Substances
Common Ions
Electrolytes
1.
Dissociation of Ionic Compounds:
Ionic compounds are made of cations (+) and anions (-).
(Note see tables 2.4 and 2.5 for formulas and charges of common ions)
These ions are locked into position in their crystal structure and are not
able to move around.
In water, the water molecules, are attracted to the ions. The ions are said to
be dissociated, and able to carry electrical particles to conduct current.
K3PO4 + H2O  3K+ (aq) + PO4-3 (aq)
Such substances are said to be electrolytes.
Salts that are completely soluble in water are strong electrolytes.
Salts that are slightly soluble are weak electrolytes at best.
The strength of an electrolyte is measured by its ability to conduct
electrical current.
See permanganate and NaCl dissolution
NaCl Dissolution
Electrolytes
2.
Ionisation of Polar Covalent Molecular Substances
Polar molecular substances are substances whose atoms are covalently
bonded. Each molecule has a net molecular dipole moment and thus a
positive and a negative end.
Polar water molecules can line up around polar molecule. If this dipoledipole interaction can overcome the dissociation energy of a bond the
molecule will fragment with bonding electrons going with the most
electronegative atom in the broken bond, creating ions.
(Electronegativity is the electron attracting ability of an atom)
Such polar molecular compounds are called electrolytes.
An example of a strong electrolyte is any of the strong acids, such as HBr.
H-Br + H2O  H3O+ (aq) + Br- (aq)
Electrolytes
Some polar molecular substances have such strong covalent bonding that water
is only able to overcome these stronger dissociation energies in a portion of the
molecules.
CH3COOH + H2O ⇌H3O++CH3COOFor example,a weak acid such as ethanoic acid, CH3-COO-H, dissolves in
water with only a small percentage of the molecules being ionized.
Non-electrolytes are substances that do not
produce ions when they dissolve. Sugar-sucrose
This results when polar molecular substances
are large enough and their covalent bonding is
strong enough so that water is not able to break
any of the covalent bonds during the solvation
process. As a result, the neutral molecules are
solvated (separated by solvent water molecules)
without any ionization occurring.
Precipitation reactions
Occur when pairs of oppositely charged ions attract each
other so strongly that they form an insoluble solid
(precipitate) that drops out of solution, removing material, and
therefore driving the reaction along.
AgNO3 (aq.)+ NaCl(aq.)  AgCl (s.)+ NaNO3 (aq.)
To predict whether a reaction will occur we must know the
solubilities of the potential products in a reaction.
Solubility definition: Amount of substance that can be dissolved
in a given quantity of solvent.
Solubility Guidelines
Table 4.1
Guidelines:
1.
Compounds containing the group 1 (Li, Na, K etc) or NH4+
cations are most likely soluble.
2.
Compounds containing the Halide anions, Cl-, Br- and I- (except
Ag+, Hg22+ and Pb2+ compounds), the nitrate (NO3-), acetate
(ethanoate) or sulphate anions (except Sr2+, Ba2+, Hg22+ and
Pb2+ sulphates) are most likely soluble.
3.
Compounds containing sulfide (except NH4+, group 1 and
heavy group 2), carbonate (except NH4+ and group 1),
phosphate (except NH4+ and group 1) and hydroxyl (except
NH4+, group 1 and heavy group 2) anions are most likely
insoluble.
Periodic Table: http://www.webelements.com/
Table 4.1
Predictions
• Give the chemical formula for the
following, and then classify as soluble or
insoluble:
• Sodium carbonate Na2CO3 soluble
PbSO4 insoluble
• Lead sulfate
• Ammonium phosphate (NH4)3PO4 soluble
Metathesis (transposition)
• Reactions involving exchange of partners
AX + BY  AY + BX
AgNO3(aq.) + KCl(aq.)  AgCl(s) + KNO3(aq.)
STEPS
Determine ions present as reactants
Combine cation of one reactant with anoin of the
other
Balance equation
Try:
barium chloride mixed with potassium sulfate
Molecular and ionic equations
• Sometimes convenient to identify whether
dissolved substances are present as ions
• Molecular equation
AgNO3(aq.) + KCl(aq.)  AgCl(s) + KNO3(aq.)
• Ionic equation
Ag+(aq.) + NO3−(aq.) + K+(aq.) + Cl−(aq.)  AgCl(s) +
K+(aq.) + NO3−(aq.)
NOTE K+, NO3−are SPECTATOR ions
• Net ionic equation
Ag+(aq.) + Cl−(aq.)  AgCl(s)
Acids and Bases
The properties of acids include the following:
•
Taste sour (but don't taste them!!)
•
Their water solutions conduct electrical current (electrolytes)
•
They react with bases to form salts and water
•
Turns Blue Litmus Paper to Red
The properties of bases include the following:
•
Have a slippery feel between the fingers
•
Have a bitter taste (but don't taste them!!)
•
React with acids to form salts and water
•
Turns Red Litmus Blue
•
Their water solutions conduct electrical current (electrolytes)
Acids and Bases
Arrhenius in 1884 discovered that acids give off H+ ions and
allow for a good flow of electricity through a solution.
Arrhenius also discovered that bases give off OH- ions and
OH- ions also allow for a good flow of electricity through the
solution.
Traditionally Svante Arrhenius defined:
Acid released Hydrogen ion (as Hydronium ions,
H3O+) in water solution.
Base produced Hydroxide ion in water solution.
The limitations on these definitions were:
1. The need for water
2. The need for a protic acid
3. The need for Hydroxide bases
Bronsted/Lowry acids and bases
Bronsted and Lowry defined these two terms the following:
Acid-Proton donor Base-Proton acceptor
These definitions are not as restrictive as Arrhenius’ definitions.
1. No need for water although it can be present, it need not be.
2. Bases do not have to be Hydroxide compounds. Ammonia as a base!
However, one restriction still remaining is the need for a protic acid.
(see Lewis theory later)
Each Bronsted acid is coupled to a
conjugate base to constitute a
CONJUGATE ACID-BASE PAIR
CH3COOH + H2O ⇌H3O++CH3COO-
See student activities
Acid and Base Strength
Strong acids (memorise) dissociate completely in water
HClO4, HClO3, HCl, HBr, HI, HNO3 and H2SO4
Strong bases are the metal hydroxides of Group 1 and
heavy Group 2
E.g. LiOH, NaOH, KOH, Ba(OH)2 etc
Weak acids and bases are not completely ionised in solution
CH3COOH + H2O H3O++CH3COO-
Acid-Base Reactions
Acid/Base reactions are reactions that involve the neutralisation of an acid through
the use of a base.
HCl + NaOH NaCl + H2O
In this reaction, the Na+ and the Cl- are called spectator ions because they play no
role in the overall outcome of the reaction. The only thing that reacts is the H+
(from the HCl) and the OH- (from the NaOH). So the reaction that actually takes
place is:
H+ + OH- H2O
If in the end, the OH- was the limiting reagent and there are H+'s still left in the
solution then the solution is acidic, but if the H+ was the limiting reagent and OH-'s
were left in the solution then the solution is basic.
Example application: antacids (milk of magnesia invented by an Irishman, James
Murray)
Oxidation-Reduction (REDOX) reactions
Originally oxidation was assigned to the combination of an element with oxygen to
give an oxide and reduction was the reverse.
Today, a much broader definition is given:
loss of electron(s) for oxidation
2Na 2Na+ + 2e-
gain of electron(s) for reduction
Cl2 + 2e-  2Cl-
Thus redox reactions are electron transfer reactions.
2Na + Cl2 2Na+ + 2Cl-
In more complex reactions a bookkeeping system, oxidation
numbers, is used to keep track of electron transfers.
A redox reaction is therefore a reaction in which changes in
oxidation numbers occur.
See student activities
Oxidation Numbers
Rules for assigning oxidation numbers:
1. An atom in its elemental state, 0.
2. An atom in a simple monoatomic ion, charge on the ion. Group 1, +1 etc.
3. Non-metals usually have negative oxidation numbers.
In its compounds O, -2, except for peroxides, O22- ion, -1
In its compounds H, +1 bonded to non-metals, -1 bonded to metals
In its compounds F, -1. Other halogens mostly -1, but can have
positive oxidation numbers when combined with oxygen
4. The sum of all the oxidation numbers in a molecule or a polyatomic ion,
charge on the particle.
Try these:
H2S,
Na2SO3,
NH4Cl,
KMnO4,
Na2S2O3
Redox Reactions
Oxidation-increase in oxidation number
example.:
Reduction-decrease in oxidation number
rusting of iron.
4Fe(s) +
0
3O2(g) 
0
2Fe2O3(s)
+3 -2
Identify substance that is oxidised, then identify substance that is reduced.
Identify oxidising and reducing agents.
Balancing redox reactions (Chapter 20)
Using the ion-electron (half-reaction) method
In acidic solutions
•
Find oxidised and reduced species.
•
Divide chemical equation into two half-reactions.
•
Balance atoms (excluding H and O).
•
Balance O (by adding H2O).
•
Balance H (by adding H+).
•
Balance charge (by adding electrons).
•
Make electron gain equivalent to electron loss, then add the half-reactions.
•
Cancel similar species on both sides of the chemical equation.
Balancing redox reactions
Using the ion-electron (half-reaction) method
In basic solutions
8.
Add the same number of hydroxyl ions as there are protons to both sides
of the chemical equation.
9.
Combine protons and hydroxyls to give water molecules.
10. Cancel H2O if you can.
Try these:
Cr2O72- + Fe2+ Cr3+ + Fe3+
in acid
SO32- + MnO4-  SO42- + MnO2
in base
Redox Titrations in Analyses
Take the example of 2.00g of iron ore converted with acid to
Fe2+(aq.). Titrated solution required 27.45mL of 0.100 M
potassium permanganate. What is the %Fe in the ore?
For the same sample, evaluate the volume of 0.100 M Ce4+ that
would have been required to titrate the Fe2+ solution
Solution Concentration
Molarity: moles of solute in a litre of solution
Example: prepare 250mL of 1.00 M solution of
CuSO4.
Molarity=moles/litre
moles = litres x Molarity = 0.25L x 1.00 M = 0.25moles
1mole CuSO4: Cu(63.5g)+S(32g)+4O(4x16g)=159.5g
0.25moles = 39.9g
Electrolyte Concentration
When an ionic compound dissolves, the
relative concentration of the ions depends
on the chemical formula.
NaCl Na+ + Cl1.0M NaCl gives a solution containing1.0M of its ions
Na2SO4  2Na+ + SO421.0M Na2SO4 gives a solution containing 2.0M of sodium
ions and 1.0M of sulphate ions.
Acid-Base Titrations
• Titration is the process of mixing acids and
bases to analyse one of the solutions. For
example, if you were given an unknown
acidic solution and a 1 M NaOH solution,
titration could be used to determine what
the concentration of the other solution was.
Simple Acid-Base Titrations
The goal of titration is to determine the equivalence point. The equivalence point
is the point in which all the H+ and the OH- ions have been used to produce water.
Titration also usually involves an indicator. An indicator is a liquid that turns a
specific colour at a specific pH. (Different indicators change colours at different
pH's). Indicators are chosen to allow a colour change at the equivalence point.
Titration of a strong acid with a strong base
50.00mL of 0.020M HCl with 0.100M NaOH
H+ + OH- H2O
at equivalence pt.:
nb mol HCl = nb mol NaOH
HCl:
0.02mol/L x 50 mL x (1L/1000mL) = 0.001 mol
NaOH: 0.1mol/L x Ve(mL) x (1L/1000mL) = 0.001 mol
therefore:
Ve(mL) = 0.001 mol  (0.1mol/L x 1L/1000mL) = 10 mL