Physical Chemistry

Dónal Leechdonal.leech@nuigalway.ie

Ext 3563Room C205, Physical Chemistry

Notes for downloading (powerpoint and word)

http://www.nuigalway.ie/chem/Donal/Teaching.htm

Chemistry

Biological Sciences

Organisms

Organs

Tissues

Cells

DNA

Physical SciencesSub-atomic

Atoms

Materials

Atmosphere

Stellar

Chemistry

Molecular Sciences

Molecules

Bonds

Forces

Physical ChemistryEstablishes and develops the principles that are used to explain and interpret the observations made in chemistry

Bulk Individual Rates

Equilibrium Structure Change

Thermodynamics Chemical reactionsQuantum mechanics & spectroscopy

ENERGY

Textbook

• Brown, LeMay, Bursten

• Chemistry: The Central Science, 9th or 10th Edition

Companion Web-site

http://www.prenhall.com/brown

Aqueous ReactionsChapter 4

Reactions in SolutionThe 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 SolutionsA 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

Electrolytes1. 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

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 dipole-dipole 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

ElectrolytesSome 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.

For 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.

CH3COOH + H2O ⇌H3O++CH3COO-

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.

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.

AgNO3 (aq.)+ NaCl(aq.) AgCl (s.)+ NaNO3 (aq.)

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+, Hg2

2+ 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

• Lead sulfate

• Ammonium phosphate

Na2CO3 soluble

PbSO4 insoluble

(NH4)3PO4 soluble

Metathesis (transposition)

• Reactions involving exchange of partnersAX + 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 equationAgNO3(aq.) + KCl(aq.) AgCl(s) + KNO3(aq.)

• Ionic equationAg+(aq.) + NO3

−(aq.) + K+(aq.) + Cl−(aq.) AgCl(s) + K+(aq.) + NO3

−(aq.)

NOTE K+, NO3−are SPECTATOR ions

• Net ionic equationAg+(aq.) + Cl−(aq.) AgCl(s)

 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

Acids and BasesArrhenius 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 basesBronsted 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

Strong acids (memorise) dissociate completely in water

HClO4, HClO3, HCl, HBr, HI, HNO3 and H2SO4

Acid and Base Strength

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 + H2OIn 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) reactionsOriginally 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

gain of electron(s) for reduction

Thus redox reactions are electron transfer reactions.

2Na 2Na+ + 2e-

Cl2 + 2e- 2Cl-

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 oxygen4. 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 ReactionsOxidation-increase in oxidation number Reduction-decrease in oxidation number

example.: rusting of iron.

4Fe(s) + 3O2(g) 2Fe2O3(s)

Identify substance that is oxidised, then identify substance that is reduced.

Identify oxidising and reducing agents.

0 0 +3 -2

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

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

Using the ion-electron (half-reaction) method

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+ + Cl- 1.0M NaCl gives a solution containing1.0M of its ions

Na2SO4 2Na+ + SO42-

1.0M 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 TitrationsThe 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 NaOHH+ + 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