The Chemistry of Acids and Bases To play the movies and simulations included, view the presentation in Slide Show Mode.

Acid and Bases

Acid and Bases

Some Properties of Acids Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen ion attached to a water molecule) Taste sour Corrode metals Electrolytes React with bases to form a salt and water pH is less than 7 Turns blue litmus paper to red “Blue to Red A-CID”

Some Properties of Bases Produce OH- ions in water Taste bitter, chalky Are electrolytes Feel soapy, slippery React with acids to form salts and water pH greater than 7 Turns red litmus paper to blue “Basic Blue”

Some Common Bases NaOH sodium hydroxide lye KOH potassium hydroxide liquid soap Ba(OH)2 barium hydroxide stabilizer for plastics Mg(OH)2 magnesium hydroxide “MOM” Milk of magnesia Al(OH)3 aluminum hydroxide Maalox (antacid)

Acid/Base definitions Definition #1: Arrhenius (traditional) Acids – produce H+ ions (or hydronium ions H3O+) acids can be classified as monoprotic, diprotic or triprotic Bases – produce OH- ions problem: some bases don’t have hydroxide ions Nice, simple definitions but has limitations… restricted to aqueous solutions couldn’t explain basic properties of NH3

Acid/Base Definitions Definition #2: Brønsted – Lowry Acids – proton donor Bases – proton acceptor A “proton” is really just a hydrogen atom that has lost it’s electron!

A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor conjugate acid conjugate base base acid

Conjugate Pairs

Bronsted-Lowry A-B Because many A/B reactions are reversible, it is useful to identify conjugates (opposites) that will react in the reverse reaction conjugate A-B pairs = 2 formulas in an equation whose formulas differ by a H+ strong A-B are not equilibrium expressions, but all other A-B are reversible

Acids & Base Definitions Definition #3 – Lewis Lewis acid - accepts an electron pair Lewis base - donates an electron pair

Lewis Acid/Base Reaction

Lewis Acid-Base Interactions in Biology The heme group in hemoglobin can interact with O2 and CO. The Fe ion in hemoglobin is a Lewis acid O2 and CO can act as Lewis bases Heme group

Brønsted-Lowry vs. Lewis All B/L bases are Lewis bases BUT, by definition, a B/L base cannot donate its electrons to anything but a proton (H+) While B/L is most useful for our purposes, Lewis allows us to treat a wider variety of reactions (even if no H+ transfer occurs) as A/B reactions

Strong Acids/Bases The strength of an acid (or base) is determined by the amount of IONIZATION (Not pH value). HCl, HBr, HI, HNO3, H2SO4 HClO3 and HClO4 are the strong acids.

Strong Acids/Bases The strong bases are: alkali-metal hydroxides and alkaline earth metal hydroxides. Ex. NaOH, KOH, CsOH, Ca(OH)2, etc. Note: The stronger the acid/base, the weaker its conjugate pair.

The pH scale is a way of expressing the strength of acids and bases The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion. Under 7 = acid 7 = neutral Over 7 = base

pH of Common Substances

pH and acidity Acidity or Acid Strength depends on Hydronium Ion Concentration [H3O+] The pH system is a logarithmic representation of the Hydrogen Ion concentration (or OH-) as a means of avoiding using large numbers and powers. pH = - log [H3O+] pOH = - log [OH-] In pure water [H3O+] = 1 x 10-7 mol / L (at 25oC)  pH = - log(1 x 10-7) = - (0 - 7) = 7 pH range of solutions: 0 - 14 pH < 7 (Acidic) [H3O+] > 1 x 10-7 m / L pH > 7 (Basic) [H3O+] < 1 x 10-7 m / L

pH calculations – Solving for H+ If the pH of Coke is 3.12, [H+] = ??? Because pH = - log [H+] then - pH = log [H+] Take antilog (10x) of both sides and get 10-pH = [H+] [H+] = 10-3.12 = 7.6 x 10-4 M

pOH Since acids and bases are opposites, pH and pOH are opposites! pOH does not really exist, but it is useful for changing bases to pH. pOH looks at the perspective of a base pOH = - log [OH-] Since pH and pOH are on opposite ends, pH + pOH = 14

pH [H+] [OH-] pOH

pH indicators Indicators are dyes that can be added that will change color in the presence of an acid or base. Some indicators only work in a specific range of pH Once the drops are added, the sample is ruined Some dyes are natural, like radish skin or red cabbage

Indicators

Wait, water can go both ways? amphoteric substances can behave as either an acid or base depending on what they react with. water and anions with protons (H+) attached are most common amphoterics

Autoionization of Water H2O + H2O OH- + H3O+ @ equilibrium Keq = [H3O+] [OH-] = so Keq becomes Kw and equals 1.00 x 10-14 at 25 oC In a neutral solution [H3O+] = [OH-]

Weak Acids/Bases Only partially ionize (less than 100%) in water. Acids (symbol=HA) @ equilibrium Keq = [H3O+] [OH-]/[HA] = so Keq becomes Ka The bigger the Ka the stronger the acid Polyprotic acids have more than 1 H+ available to ionize Have multiple Ka values (Ka1 Ka2 Ka3) Always easier to remove 1st H+ than 2nd If Ka2 is different than Ka1 by 103 (or more) consider only Ka1 to get pH

Weak Acids/Bases Bases (symbol=B) Relationship between Ka and Kb 2 general categories Contain neutral substances that have an atom with non-bonding electrons Ex. NH3  NH4+ Anions of weak acids Ex. ClO-  HClO @ equilibrium Keq = [HB+] [OH-]/[B] = so Keq becomes Kb Relationship between Ka and Kb Ka x Kb = Kw pKa + pKb = pKw = 14.00

Equilibrium Constants for Weak Acids Weak acid has Ka < 1 Leads to small [H3O+] and a pH of 2 - 7

Equilibrium Constants for Weak Bases Weak base has Kb < 1 Leads to small [OH-] and a pH of 7-12

Relation of Ka, Kb, [H3O+] and pH

Equilibria Involving A Weak Acid You have 1.00 M HA. Calc. the equilibrium concs. of HA, H3O+, A-, and the pH. HA + H2O H3O+ + A- Step 1. Define equilibrium concs. in ICE table [HA] [H3O+] [A-] initial change equilib 1.00 0 0 -x +x +x 1.00-x x x

Equilibria Involving A Weak Acid Step 2. Solve equilibrium (Ka) expression This is a quadratic. Solve using quadratic formula. because Ka for most weak acids is less than 10-3, 1-x is about equal to 1, so Ka = x2/1.00

Equilibria Involving A Weak Acid So the Ka expression becomes Now we can more easily solve this approximate expression.

Equilibria Involving A Weak Acid Step 3. Calculate the pH x = [H3O+] = [A-] = 4.2 x 10-3 M pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37

Equilibria Involving A Weak Base You have 0.010 M of B. Calc. the pH. B + H2O OH- + HB+ Kb = 1.8 x 10-5 Step 1. Define equilibrium concs. in ICE table [B] [HB+] [OH-] initial change equilib 0.010 0 0 -x +x +x 0.010 - x x x

Equilibria Involving A Weak Base Step 2. Solve the equilibrium expression b/c Kb for most weak bases is less than 10-3, 0.010-x is about equal to 0.010, so Kb = x2/0.010 x = [OH-] = [HB+] = 4.2 x 10-4 M and [B] = 0.010 - 4.2 x 10-4 ≈ 0.010 M

Equilibria Involving A Weak Base Step 3. Calculate pH [OH-] = 4.2 x 10-4 M so pOH = - log [OH-] = 3.37 Because pH + pOH = 14, pH = 10.63

A-B properties of salt solutions for the most part anions are slightly basic (because they attract protons) and cations are slightly acidic (because they can donate protons) Anions Look at strength of the acid to which it is a conjugate If acid is strong, anion won’t really take H+ from water If acid is weak, anion will react to a small extent ( pH) If amphoteric need to look at magnitudes of Ka and Kb Cations Polyatomic cations can be the conjugate acid of a weak base Usually react to make H3O+ (  pH)

A/B Behavior & Chemical Structure Polarity if H is the positive end = acid if H is negative end (NaH) = base if non-polar, neither acid or base Strength of bond Look at what H is bonded to How easy is it to dissociate into ions? Stability of conjugate: the greater the stability of the conjugate, the stronger the acid

A/B Behavior & Chemical Structure Binary Acids 1. HX bond strength is most important 2. acidity increases down a group - larger elements - longer and weaker bonds 3. in same row, look at bond polarity - across from L to Right increases EN - Across from L to R increases acidity

A/B Behavior & Chemical Structure Oxyacids (have O in compound) 1. Have same # of OH groups and the same # of O atoms, acid strength  with  in EN of center atom 2. Have same center atom (Y) acid strength  as the # of O atoms attached to Y  3. in a series of oxyacids, acidity  with as the oxidation # of the center atom 

A/B Behavior & Chemical Structure Carboxylic Acids (have –COOH in compound) H on OH group is ionized the other O atom draws elctron density away from OH bond increasing polarity the conj. base (COO-) can exhibit resonance making it more stable Acid strength of COOH group increases as the # of electronegative atoms in the acid increase.