Advanced Chemical Reactions

1. Advanced Chemical Reactions Reaction rates, Equilibrium, Acids/Bases, Redox Reactions

2. Entropy (S) • Measure of disorder or randomness in a system • Natural tendency for system to increase entropy (more random) • EXAMPLE – Diffusion • As molecules are dispersed, entropy increases • Continued dispersal leads to a uniform solution

3. Spontaneous reaction • Remember, things tend towards an increase in entropy • Spontaneous reaction favors the products (exothermic) and releases free energy • C + O2 CO2 • Exothermic • Solid  gas increases entropy • Gibbs free energy – max amt of E that can be used in another process

4. 2nd law of thermodynamics • Entropy never decreases in a system and instead will increase over time • UNLESS you change the surroundings • Spraying air freshener • Spray it into a collapsible box

5. Kinetics • Study of reaction rates (rate at which a chemical reaction takes place) • Measured by: • Rate of formation of products • Rate of disappearance of reactants • Changes in concentration of reactants or products

6. Factors that Affect Rxn Rate • Concentration • Pressure • Temperature • Surface Area • All of the above have a DIRECT relationship

7. When do chemical reactions occur? • When reactants collide • Normally, molecules bounce off each other b/c of electron clouds repulsion • BUT, if those molecules have a LARGE amount of energy, they can overcome the repulsion and react • Molecules also must collide in the right orientation

8. Activation Energy (Ea) • Energy required to start a chemical reaction • A nudge, a spark • Potential E

9. Activation Energy • Activated complex – “speed bump” of the reaction – point at which it could go either way • H2O + CO2 H2CO3  H+ + HCO3-

10. Catalyst • Another factor that affects reaction rate • Speeds the reaction by lowering the activation energy • Not used up by reaction

12. Equilibrium • Two basic categories for reactions • Completion reactions – 1-way (combustion, decomp, rusting) • Reversible reactions – products can re-form original reactants • Reversible reactions often use 2 arrows b/c reactions occur at the same time

13. Equilibrium • Chemical equilibrium is DYNAMIC, not STATIC

14. Equilibrium • Chemical equilibrium – reactions in which the forward and reverse reaction rates are equal

15. Equilibrium Constant • Every reaction has a condition of equilibrium at a given temperature • That means that 2 reactants will react to form products until a state is reached where the amounts of products and reactants no longer change • CO2 in a half-filled, sealed soda bottle • Things will stay that way until the system is somehow altered

16. Equilibrium Constant • Equilibrium constant, Keq – a number that expresses the necessary concentrations of reactants and products for the reaction to be at equilibrium • aA + bB  cC + dD • Keq = [C]c [D]d [A]a [B]b • If Keq >1, the reaction favors the products • If Keq <1, the reaction favors the reactants

17. Equilibrium Constant • Calculate the Keq of the following equation CO2 (g) + H2 (g)  CO (g) + H2O (g) If the [CO2] = 1.5 M, [ H2 ] = 1.5 M, [ CO ] = 0.6 M, [ H2O] = 0.6 M • Keq= [CO]1 [H2O]1 = [0.6] [0.6] = 0.16 [CO2]1 [H2]1 [1.5] [1.5] • So this reaction favors the….

18. Le Chatelier’s Principle • When a system at equilibrium is disturbed, the system adjusts in a way to reduce the change. • Chemical equilibria responds to 3 kinds of stress or change • Change in concentration • Change in temperature • Change in pressure

19. Changes in Concentration • Increasing concentration of reactant will make the rate of the forward reaction faster than the reverse • Called a shift right • Continues until new equilibrium • H3O+ + HCO3 2H2O + CO2 • Increasing concentration of product leads to shift left

20. Changes in Temperature • Remember that endothermic & exothermic are opposites • Increasing the temp adds E so the endothermic will go faster to use it • If it is exothermic forward, increasing the temp favors the reactants • If it is endothermic forward, increasing the temp favors the products

21. Changes in Pressure • Only affects gases • Imagine volume has been decreased, increasing the pressure • Immediate effect is increase in concentration of both product & reactant • According to principle, system will adjust to decrease the pressure

22. Changes in Pressure • A pressure increase favors the reaction that produces fewer molecules (stoichiometry) • 2NOCl   2 NO + Cl2 • H2O + CO   H2 + CO2

23. Acids & Bases

24. Characteristics • Acids – sour taste, conduct electricity well, react with many metals, generate hydronium ions (H3O+), turn litmus paper red • Bases – bitter taste, slippery feel, varying solubility, generate hydroxide ions (OH-), turn litmus paper blue

25. STRONG vs WEAK • Strong acids & bases COMPLETELY dissociate or ionize in water (one way reaction) • HNO3 + H2O  H3O+ + NO3- • NaOH  Na+ + OH- • Weak acids & bases only partially dissociate (reversible reaction) • HOCl + H2O  H3O+ + ClO- • NH3 + H2O  NH4+ + OH-

26. Arrhenius • Acid – ionizes to form an H3O+ ion when added to water • Base – generate OH- when dissolved in water

27. Bronsted-Lowry • Acid – donates a proton (H+) to another substance • Base – accepts a proton (H+) • NH3 + H2O  NH4+ + OH- • H2O is the Bronsted-Lowry acid & NH3 is the Bronsted-Lowry base • Always reactants

28. Conjugate • Conjugate Acid – Formed when a base gains a proton (H+) • Conjugate Base – Formed when an acid loses a proton (H+) • NH3 + H2O  NH4+ + OH- • NH4+ is the conjugate acid & OH- is the conjugate base • Always products

29. Amphoteric • Can act as an acid or a base depending on what it is combined with

30. Water • Can act as a Bronsted-Lowry acid or base • H2O + H2O   H3O+ + OH- • Called the self-ionization of water • Results in equal concentrations of H3O+ and OH- in pure water • [H3O+] = [OH-] = 1.00 x 10-7 M

31. Water • [H3O+] x [OH-] = 1.00 x 10-7 x 1.00 x 10-7 = 1.00 x 10-14 • Found to be true for other aqueous solutions at equilibrium • [H3O+] x [OH-] = 1.00 x 10-14 • Also abbreviated as Kw

32. Acids & Bases • Have proportional amounts of H3O+ & OH- • [H3O+] x [OH-] = 1.00 x 10-14 OH- H3O+ H3O+ OH- OH- H3O+ ACID NEUTRAL BASE

33. Calculating [H3O+] & [OH-] using Kw • [H3O+] x [OH-] = 1.00 x 10-14 • If [H3O+] = 1.00 x 10-2, what is [OH-]? • [OH-] = 1.00 x 10-12 • If [H3O+] = 1.00 x 10-5, what is [OH-]? • [OH-] = 1.00 x 10-9

34. The pH Scale • 1909 – Soren Sorenson – negative exponents are annoying… • So let’s just look at the exponents! • Logarithm – power to which 10 must be raised to equal that number • log 100 = 2 because 100 = 102 • log 0.001 = -3 because 0.001 = 10-3

35. Logarithm Practice • log 10,000 = • log 0.01 = • log 10 = • log 0.000001 = • log 1 =

36. The pH Scale • Represents the “power” of “Hydrogen” • pH = - log [H3O+] • What is the pH of a 0.00010 M solution of HNO3? • pH = - log [1.0 x 10-4] = -(-4) = 4

37. pH Practice • What is the pH of a 0.2 M solution of a strong acid? • pH = - log [.2] • pH = 0.70

38. What if it’s a base? • [H3O+] x [OH-] = 1.00 x 10-14 • pH + pOH = 14 • You can calculate [H3O+] by 1.00 x 10-14 / [OH-] • Then you can calculate pH

39. Example • What is the pH of a 0.0136 M solution of KOH, a strong base? • [H3O+] = 1.00 x 10-14 / 0.0136 • [H3O+] = 7.35 x 10-13 • pH = -log [H3O+] • pH = - log [7.35 x 10-13] • pH = 12.13

40. Example • Lemonade has a hydronium ion concentration of 0.0050 moles/L. What is it’s pH? • pH = -log [H3O+] • pH = 2.3 • What is it’s pOH?

41. Neutralization • Reaction of H3O+ & OH- to form water molecules and often a salt • H3O+ & OH-  2H2O • Neutral means [H3O+] = [OH-] • HCl + NaOH  H2O + NaCl • Common way to deal with acid & base spills • Baking soda = NaHCO3,Ammonia = NH3

42. Indicators • Change color at a certain pH level • Red cabbage juice – changes to blue between 3 & 4 and to green at 8/9 • Litmus paper – red or blue • Phenolphthalein – turns bright pink in the presence of a base

43. Titration • Used to determine the unknown concentration of a known reactant • Uses an indicator to show the equivalence point • For strong acid/strong base… • Equivalence point is where [H3O+] = [OH-] or where moles of acid = moles of base • Often uses phenolphthalein

44. Titration curve

45. Titration Set-up

46. Oxidation-Reduction Reactions • Remember that electronegativity is a measure of how tightly atoms hold on to their electrons • Atoms with large electronegativity differences form ionic bonds by electron transfers • 2Na + Cl2 2NaCl • Can be written as 2Na + Cl2  2Na+Cl-

47. Oxidation-Reduction Reactions • Oxidation = Loss of electrons • Na  Na+ • Reduction = Gain of electrons • Cl2 2 Cl- • These 2 reactions happen together • Oxidation-Reduction or REDOX • OIL RIG

48. How do we know if an atom is oxidized or reduced? • Use “oxidation” numbers • The number of electrons that must be added or removed to convert the atom to elemental or neutral form • In other words, it’s the charge the atom would have if it were an ion

49. Determining Oxidation Numbers 1. Look at the equation 2. Assign known oxidation numbers 3. Calculate unknowns & verify - Sum of all atoms in a molecule is zero - Sum of all atoms in a polyatomic is equal to the charge on that ion

50. Assigning known oxidation #s • Uncombined = 0 O2 • Monatomic ion = ion charge Zn 2+ • Flourine = -1 (most electronegative) • Group 1 = +1 K • Group 2 = +2 Ca • Binary compounds – most electronegative element = ion charge CaCl2