Acids and Bases Notes

1. Acids and Bases Notes http://www.nearingzero.net/screen_res/nz065.jpg

2. I. Strength of Acids and Bases • A. Bases: • Strong Bases: metal hydroxides of Group I and II metals (except Be) that are soluble in water and dissociate (separates into ions) completely in dilute aqueous solutions Weak Bases: a molecular substance that ionizes only slightly in water to produce an alkaline (basic) solution (ex. NH3)

3. Johnny was a chemist, a chemist he's no more for what he thought was H2O was H2SO4. B. Acids: Strong Acids: an acid that ionizes (separates into ions) completely or very nearly completely in aqueous solutions (will not be on Ka chart). Weak Acids: an acid that ionizes only slightly in dilute aqueous solutions (will be on the Ka chart).

4. Why do white bears dissolve in water? Because they're polar. 1. Binary or hydrohalic acids – HF, HCl, HBr, HI, etc. “hydro____ic acid” are usually strong acids HF and H2S are weak hydrohalic acid. Although the H-F bond is very polar, the bond is so strong (due to the small F atom) that the acid does not completely ionize.

5. 2. Oxyacids – contain a polyatomic ion a. Most common form (MCF) “ic” ending– strong acids (contain 2 oxygen per hydrogen) HNO3 – nitric from nitrate H3PO4 - phosphoric from phosphate H2SO4 - sulfuric from sulfate HClO3 - chloric from chlorate

6. b. Acids with l less oxygen than the MCF “ous” ending- weaker acids HNO2 – nitrous from nitrite H3PO3 - phosphorous from phosphite H2SO3 - sulfurous from sulfite HClO2 - chlorous from chlorite c. Acids with 2 less oxygen than the MCF “hypo___ous” – very weak acids HNO - hyponitrous H3PO2 - hypophosphorus HClO - hypochorous

7. d. Acids with 1 more oxygen than the MCF “per______ic” – very strong acids HClO4 – perchloric acid HNO4 - pernitric acid e. Organic acids – have carboxyl group -COOH - usually weak acids HC2H3O2 - acetic acid C7H5COOH - benzoic acid

8. Q: What's the most important thing to learn in chemistry?A: Never lick the spoon II. Characteristics of Acids and Bases (page 453-458) There are 3 different definitions for acids and bases and I have summarized them all here. Keep in mind that there are some exceptions and this is just a guide to follow when trying to determine if a substance is an acid or a base.

9. Usually begin with H Sour taste Litmus paper turns red pH paper 1-6 or pH meter Phenolphthalein - colorless Feels like water React with metals toproduce H2 Electrolyte pH<7 Usually contain OH Bitter taste Litmus paper turns blue pH paper 8-14 or pH meter Phenolphthalein – pink Feels slippery Electrolyte pH>7 Acids Bases

10. Arrhenius Acid: donates (or produces) hydronium ions (H3O+) in water or hydrogen ions (H+) in water Bronsted-Lowry Acid: donates a proton (H+) in water, H3O+ has an extra H+, if it donated it to another molecule it would be H2O (page 467) HNO3 + H2O  H+ + NO3- HNO3 + H2O  H3O+ + NO3- HCl + H2O  H+ + Cl- HCl + H2O  H3O+ + Cl- Arrhenius Base: donates (or produces) hydroxide ions (OH-) in water Bronsted – Lowry Base: accepts a proton in water, OH- needs an extra H+ if it accepts one from another molecule it would be H2O (page 468) KOH + H2O  K+ + OH- NH3 + H2O  NH4+ + OH- Acids Bases

11. Lewis Acid: Not all acids contain H, any atom, ion, or molecule that accepts an electron pair from a covalent bond is an acid Lewis Base: Not all bases contain OH, any atom, ion, or molecule that donates an electron pair to form a covalent bond it a base Acids Bases Q. What do you do when you find a dead chemist? A. Barium. http://www.gis.net/~sjp3/ps.html

12. Conjugates: HF + H2O  F - + H3O+ Acid Base Conjugate Conjugate Base Acid Here HF donated a proton (H+) to the water and the water accepted the proton (H+). HF is referred to as the acid and water is referred to as the base. The fluoride ion, F- is referred to as the conjugate base of HF. F- can accept a proton (H+) to be stable. The hydronium ion, H3O+ is referred to as the conjugate acid of water. H3O+ can donate a proton (H+) to be stable.

13. Example: Determine the acid, base, conjugate acid, and conjugate base in each of the following equations: HCl + H2O  Cl- + H3O+ Acid Base Conjugate Conjugate   Base Acid  H2SO4 + H2O  HSO4- + H3O+ Acid Base Conjugate Conjugate   Base Acid NH3 + H2O  OH- + NH4+ Base Acid Conjugate Base Conjugate Acid

14. 2. What is the conjugate base of the following substances? • H2O ________________ • NH4+________________ • HNO2_______________ • HC2H3O2_________________ 3. What is the conjugate acid of the following substances? • HCO3-__________________ • H2O____________ • HPO42-____________ • NH3___________

15. III. pH Scale[ ] brackets mean concentration The pH scale indicates the hydronium ion concentration, [H3O+] or molarity, of a solution. (In other words how many H3O+ ions are in a solution. If there are a lot we assume it is an acid, if there are very few it is a base.) Q: How do you make a 24-molar solution? A: Put your artificial teeth in water.

16. Where does the pH scale come from? [H+] pH

17. pOH Scale The pOH scale indicates the hydroxide ion concentration, [OH-] or molarity, of a solution. (In other words how many OH- ions are in the solution. If there are a lot we assume it is a base, if there are very few it is an acid.) Two chemists meet for the first time at a symposium. One is American, one is British. The British chemist asks the American chemist, "So what do you do for research?" The American responds, "Oh, I work with aerosols." The British chemist responds, "Yes, sometimes my colleagues get on my nerves also."

18. [OH-] pOH

19. Example: • Lemon juice (citric acid) pH = 2, pOH = 12 • Pure water pH = 7, pOH = 7 • Milk of magnesia pH = 10, pOH = 4 • The last words of a chemist: • And now for the taste test. • 2. I wonder if this is hot? • 3. And now a little bit of this... • 4. And now shake it a bit.

20. 4. Calculations Involving pH, pOH, [H3O+], and [OH-] of strong Acids and Bases 1st: determine which ion will be produced, either OH or H3O+ (Acids produce H3O+ and bases produce OH-). 2nd: use formula to determine pH or pOH. 3rd: check to see if answer is reasonable. pH = -log [H3O+] pOH = -log [OH-] pOH + pH = 14

21. Example Problems: 1.      What is the pH of a 0.001M NaOH solution? 1st step: Hydroxide will be produced and the [OH-] = 0.001M 2nd step: pOH = -log [0.001] pOH = 3 pH = 14-3 = 11 The answer to the problem was "log(1+x)". A student copied the answer from the good student next to him, but didn't want to make it obvious that he was cheating, so he changed the answer slightly, to "timber(1+x)."

22. 2. What is the pH of a 3.4X10-5M HCl solution? 3. What is the pH of a solution if the pOH = 5? 4. What is the pH of a 10 liter KOH solution if 5.611 grams of KOH were used to prepare the solution? 5. What is the pOH of a 1.1X10-5M HNO3 solution? 6. If the pH of a KOH solution is 10.75, what is the molar concentration of the solution? What is the pOH? What is the [H+]?

23. The pH of a strong acid cannot be greater than 7. If the acid concentration [H3O+] is less than 1.0X10-7, the water becomes the important source of [H3O+] or [H+] and the pH is 7.00. Just remember to check if you answer is reasonable! 7.      What is the pH of a 2.5X10-10M HCl solution? 8.      What is the pH of a 1.0X10-11M HNO3 solution?

24. 3. Amphoterism and WaterWater is an amphoteric substance - a substance that can act as an acid or as a base.Autoionization of water: H2O + H2O  H3O+ + OH-

25. Ion product constant for water (Kw) Kw = [H30+][OH-] or Kw = [H+][OH-] Because at 25°C [H+] = l.0 X l0-7 M and [OH-] = l.0 X l0-7 M the Kw = 1.0 X l0-14 mol2/L2 No matter what an aqueous solution contains, at 25°C [H+] [OH-] = l.0 x l0-14 Kw varies with temperature. Neutral solution [H+] = [OH-] Acidic solution [H+] > [OH-] Basic solution [H+] < [OH-]

26. Calculations Involving Weak Acids • Weak acid strength is compared by the Ka values of the acids. The smaller the Ka, the weaker the acid. Strong acids do not have Ka values because strong acids completely ionize in water so there would be almost no unionized acid

27. Calculations Involving Weak Bases • Weak base strength is compared by the Kb values of the bases. The smaller the Kb, the weaker the base. Strong bases do not have Kb values because strong bases completely ionize in water so there would be almost no unionized base

28. C. Calculations Involving pH of Weak Acids - RICE • Calculating the pH of weak acids involves setting up an equation. Always start by writing the equation, setting up the acid equilibrium expression (Ka), defining initial concentrations, changes, and final concentrations in terms of x, substituting values and variables into the Ka expression and solving for x. The RICE method will be used.

29. C. Calculations Involving pH of Weak Acids - RICE

30. Example: Calculate the pH of a 3.5 M solution of acetic acid.

31. K = [products]coef [reactants]coef So K= [H30+ ][C2H302-] HC2H302 K for acteic acid is 1.8 X 10-5. Substituting we get: 1.8 X 10-5 = [x][x] 3.5 – x Because of sig figs we can neglect the “-x”

32. Practice problem 2 • Find the pH of a l.50 M solution benzoic acid (C6H5COOH) solution. Ka = 6.4 X 10-5

33. Practice problem 3 • Find the pH of a l5.0 M ammonia solution.

34. IV. Acid Rain Many industrial processes produce gases such as NO, NO2, CO2, SO2, and SO3. These compounds can dissolve in atmospheric water to produce acidic solutions that fall to the ground in the form of rain or snow. Marble found in many buildings and statues is composed of calcium carbonate, when acid snow or rain falls on these structures a great deal of damage is caused. (page 475)

35. V. Neutralization Reactions • Neutralization Reaction– the reaction of an acid with a base to produce water and a salt. (This occurs when H3O+ and OH- ions are supplied in equal numbers by the reactants.) • Salt - an ionic compound composed of a cation from a base and an anion from an acid. • Example: HCl + NaOH  HOH + NaCl

36. Practice • H2SO4 + KOH  K2SO4 + H2O • HCl + NaOH  NaCl + H2O 3. phosphoric acid and ammonium hydroxide

37. VI. Titrations • Titration – A neutralization reaction of an acid by a base or vice versa; it is usually used to find the concentration, molarity, of an unknown acid or base. The concentration of the other is known. • Standard solution – solution of known concentration (in buret). The titrant is added to a solution of unknown concentration until the substance being analyzed is just consumed (stoichiometric point or equivalence point). • _____________- change color at the end point of a titration. Examples: phenolphthalein and methyl orange • ____________________ point-the point at which the two solutions used in a titration are present in chemically equivalent amounts.