1. 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

2. Chemistry Biological Sciences  Organisms  Organs  Tissues  Cells  DNA Physical Sciences  Sub-atomic  Atoms  Materials  Atmosphere  Stellar Chemistry Molecular Sciences  Molecules  Bonds  Forces

3. Physical Chemistry Establishes and develops the principles that are used to explain and interpret the observations made in chemistry Bulk  Individual  Rates Equilibrium  Structure  Change Thermodynamics   Chemical reactions Quantum mechanics & spectroscopy ENERGY

4. Textbook Brown, LeMay, Bursten Chemistry: The Central Science, 9 th or 10 th Edition Companion Web-site http://www.prenhall.com/brown

5. Aqueous Reactions Chapter 4

6. 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.

7. 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

8. Common Ions

9. 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. K 3 PO 4 + H 2 O  3K + (aq) + PO 4 -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

10. NaCl Dissolution

11. 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 + H 2 O  H 3 O + (aq) + Br - (aq) Electrolytes

12. 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. For example,a weak acid such as ethanoic acid, CH 3 -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. CH 3 COOH + H 2 O ⇌ H 3 O + +CH 3 COO -

13. 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. AgNO 3 (aq.) + NaCl (aq.)  AgCl (s.) + NaNO 3 (aq.)

14. Solubility Guidelines Table 4.1 Guidelines: 1.Compounds containing the group 1 (Li, Na, K etc) or NH 4 + cations are most likely soluble. 2.Compounds containing the Halide anions, Cl -, Br - and I - (except Ag +, Hg 2 2+ and Pb 2+ compounds), the nitrate (NO 3 - ), acetate (ethanoate) or sulphate anions (except Sr 2+, Ba 2+, Hg 2 2+ and Pb 2+ sulphates) are most likely soluble. 3.Compounds containing sulfide (except NH 4 +, group 1 and heavy group 2), carbonate (except NH 4 + and group 1), phosphate (except NH 4 + and group 1) and hydroxyl (except NH 4 +, group 1 and heavy group 2) anions are most likely insoluble. Periodic Table: http://www.webelements.com/http://www.webelements.com/

15. Table 4.1

16. Predictions Give the chemical formula for the following, and then classify as soluble or insoluble: Sodium carbonate Lead sulfate Ammonium phosphate Na 2 CO 3 soluble PbSO 4 insoluble (NH 4 ) 3 PO 4 soluble

17. Metathesis (transposition) Reactions involving exchange of partners AX + BY  AY + BX AgNO 3 (aq.) + KCl(aq.)  AgCl(s) + KNO 3 (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

18. Molecular and ionic equations Sometimes convenient to identify whether dissolved substances are present as ions Molecular equation AgNO 3 (aq.) + KCl(aq.)  AgCl(s) + KNO 3 (aq.) Ionic equation Ag + (aq.) + NO 3 − (aq.) + K + (aq.) + Cl − (aq.)  AgCl(s) + K + (aq.) + NO 3 − (aq.) NOTE K +, NO 3 − are SPECTATOR ions Net ionic equation Ag + (aq.) + Cl − (aq.)  AgCl(s)

19. 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

21. ArrheniusArrhenius 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, H 3 O + ) 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

22. 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 CH 3 COOH + H 2 O ⇌ H 3 O + +CH 3 COO - See student activities

23. Strong acids (memorise) dissociate completely in water HClO 4, HClO 3, HCl, HBr, HI, HNO 3 and H 2 SO 4 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 CH 3 COOH + H 2 O  H 3 O + +CH 3 COO -

24. Acid-Base Reactions Acid/Base reactions are reactions that involve the neutralisation of an acid through the use of a base. HCl + NaOH  NaCl + H 2 O 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 -  H 2 O 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)milk of magnesia

25. 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 gain of electron(s) for reduction Thus redox reactions are electron transfer reactions. 2Na  2Na + + 2e - Cl 2 + 2e -  2Cl - 2Na + Cl 2  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

26. 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, O 2 2- 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: H 2 S, Na 2 SO 3, NH 4 Cl, KMnO 4,Na 2 S 2 O 3

27. Redox Reactions Oxidation-increase in oxidation numberReduction-decrease in oxidation number example.: rusting of iron. 4Fe(s) + 3O 2 (g)  2Fe 2 O 3 (s) Identify substance that is oxidised, then identify substance that is reduced. Identify oxidising and reducing agents. 00+3 -2

28. 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 H 2 O). 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.

29. 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 H 2 O if you can. Try these: Cr 2 O 7 2- + Fe 2+  Cr 3+ + Fe 3+ in acid SO 3 2- + MnO 4 -  SO 4 2- + MnO 2 in base Using the ion-electron (half-reaction) method

30. Redox Titrations in Analyses Take the example of 2.00g of iron ore converted with acid to Fe 2+ (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 Ce 4+ that would have been required to titrate the Fe 2+ solution

31. Solution Concentration Molarity: moles of solute in a litre of solution Example: prepare 250mL of 1.00 M solution of CuSO 4. Molarity=moles/litre moles = litres x Molarity = 0.25L x 1.00 M = 0.25moles 1mole CuSO 4 : Cu(63.5g)+S(32g)+4O(4x16g)=159.5g 0.25moles = 39.9g

32. 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 Na 2 SO 4  2Na + + SO 4 2- 1.0M Na 2 SO 4 gives a solution containing 2.0M of sodium ions and 1.0M of sulphate ions.

33. 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.

34. 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 -  H 2 O 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 V e (mL) x (1L/1000mL) = 0.001 mol therefore:V e (mL) = 0.001 mol  (0.1mol/L x 1L/1000mL) = 10 mL