AP CHEMISTRY

 Acids ◦ Sour, can corrode metals, cause certain dyes to change colors  Bases ◦ Bitter taste, feel slippery, usually used in cleaning products  Arrhenius Acid ◦ substances when dissolved in water, increase the concentration of H + ions  Arrhenius Base ◦ Substances, when dissolved in water, increase the concentration of OH - ions

 Arrhenius definition is restricted to aqueous solutions  Brønsted-Lowery Acid ◦ Substance (molecule or ion) that can donate a proton to another substance  Brønsted-Lowery Base ◦ Substance that can accept a proton

 In any acid-base equilibrium, both the forward and reverse reactions involve protong transfers.  An acid and base in an equation that differ only in the presence or absence of a proton are called conjugate acid-base pairs. ◦ Every acid has a conjugate base and every base has a conjugate acid

 What is the conjugate base of each of these acids? ◦ HClO 4, PH 4 +, HNO 2 ◦ ClO 4 -, PH 3, NO 2 -  What is the conjugate acid of each of these bases? ◦ CN -, H 2 O, HCO 3 - ◦ HCN, H 3 O +, H 2 CO 3  Identify the conjugate base pairs. NH 3(aq) + H 2 O (l) ⇆ NH 4 + (aq) + OH - (aq)

 Some acids/bases are better proton donors/acceptors than others.  The stronger the acid, the weaker its conjugate base (and the stronger the base, the weaker the conjugate base.  Strengths of acids and bases ◦ Strong Acids completely transfer their protons to water. (Their conjugate base have negligible tendency to be protonated.) ◦ Weak Acids only partly dissociate and therefore exist as a mixture of acid molecules and their ions (Conjugate base is a weak base as well.) ◦ Negligible Acidity contain hydrogen but do not demonstrate any acidic behavior in water. (Conjugate bases are strong bases, reacting completely in water)

 Relative Strengths of Conjugate Acid-Base pairs

 Think of proton-transfer reactions as being governed by the relative abilities of two bases to abstract protons HC 2 H 3 O 2(aq) + H 2 O (l) ⇆ H 3 O + (aq) + C 2 H 3 O 2 - (aq)  C 2 H 3 O 2 - is a stronger base than H 2 O and therefore abstracts the proton from H 3 O +.  In every acid-base reaction the position of the equilibrium favors transfer of the proton to the stronger base.

 One of the most important chemical properties of water is its ability to act as either a Brønsted-Lowery acid or base.  This is called the autoionization of water ◦ Reaction is rapid and no individual molecules remains ionized for long (only 2 in every 10 9 molecules is ionized at any moment)

 Because water ionizes, we can write its equilibrium constant.  Keq = [H 3 O + ] [OH - ] ◦ Kw = 1.0 x (at 25 ⁰ C) ◦ Ion-product constant ◦ Commit this to memory!!  Important because this equilibrium applies to any dilute aqueous solution and can be used to calculate the [H + ] or [OH - ] ◦ The product of [H + ] and [OH - ] = 1.0 x ◦ [H + ] = [OH - ] is said to be neutral. ◦ [H + ] exceeds [OH - ] is acidic ◦ [OH - ] exceeds [H + ] is basic  [H 3 O + ] [OH - ] = 1.0 x

 The molar [H+] is very small in aqueous solutions so we express [H + ] in terms of pH, which is the negative logarithm in base 10 of [H + ].  pH = -log[H + ]

 What is the pH for a solution that has [H + ] = 5.6 x M?  pH = -log(5.6 x )  = 5.25  A sample of apple juice has a pH of Calculate the [H + ].  pH = -log[H + ] = 3.76  [H+] = = 1.7 x M

 pOH = -log[OH-]  pH + pOH = 14 (at 25˚C)

 7 Common Strong Acids ◦ 6 Monoprotic: HCl, HBr, HI, HNO 3, HClO 3, and HClO 4 ◦ 1 Diprotic: H 2 SO 4  Strong acids completely dissociate into their ions ◦ HNO 3(aq) + H 2 O(l) H 3 O + (aq) + NO 3 - (aq)  [H + ] = the original concentration of acid. ◦ If 0.20 M HNO 3 was used, [H + ] = 0.20 M ◦ pH = -log(0.20) = 0.7 ◦ (diprotic acid is a little more complex)

 Common Soluble Strong Bases ◦ Alkali metal hydroxides and heavier alkaline earth metal hydroxides  LiOH, NaOH, KOH, Ca(OH) 2, Sr(OH) 2, and Ba(OH) 2  Strong Bases also completely dissociate so calculating pH is also straightforward. ◦ What is pH of M solution of NaOH?  pOH = -log(0.028) = 1.55  pH = – 1.55 = OR  [H + ] = 1.0 x = 2.35 x  pH = -log(3.57 x ) = 12.45

 Weak acids only partially ionize in aqueous solutions  We can use the equilibrium constant for the ionization reaction to express the extent to which an acid ionizes.  Ionization reaction can be written in two ways: ◦ HA (aq) + H 2 O (l) ⇌ H (aq) + A - (aq) ◦ HA (aq) ⇌ H + (aq) + A - (aq)

 HA (aq) + H 2 O (l) ⇌ H (aq) + A - (aq)  HA (aq) ⇌ H + (aq) + A - (aq)  K a = [H ][A - ] or [H + ][A - ] [HA]  K a is acid-dissociation constant  The larger the value of K a, the stronger the acid

 A 0.10 M solution of formic acid (HCHO 2 ) has a pH of 2.38 at 25 ⁰ C. Calculate the K a for formic acid at this temperature. ◦ HCHO 2(aq) ⇌ H + (aq) + CHO 2 - (aq) ◦ K a = [H + ][CHO 2 - ] [HCHO 2 ] ◦ pH = -log[H + ] = 2.38 ◦ [H + ] = = 4.2 x M ◦ x M ≈ 0.10 ◦ K a = [H + ][CHO 2 - ] = (4.2 x M)(4.2 x M) = 1.8 x [HCHO 2 ] 0.10 HCHO 2 H+H+ CHO 2 - I0.10 M00 C-4.2 x M+4.2 x M E x M 4.2 x M

 Know the value of K a and the initial concentration of the weak acid, we can calculate the [H + ] in a solution of a weak acid.  Calculate the pH of a 0.30 M solution of acetic acid. 1.Write the ionization for acetic acid: HC 2 H 3 O 2 ⇆ H + (aq) + C 2 H 3 O 2 - (aq) 2.Write the acid-dissociation constant expression and calculate the value: K a = [H+][C 2 H 3 O 2 -] = 1.8 x 10-5 (from table 16.2 page 628) [HC 2 H 3 O 2 ] 3.Use an ICE box to express equilibrium concentrations

4.Substitute equilibrium concentrations into acid dissociation constant expression (0.30 – x ≈ 0.30) K a = [H + ][C 2 H 3 O 2 - ] = (x)(x) = 1.8 x [HC 2 H 3 O 2 ] Solve for x x 2 = (0.30)(1.8 x ) = 5.4 x x = 2.3 x Solve for [H + ] [H + ] = x = 2.3 x M 7.Solve for pH pH = -log(2.3 x ) = 2.64 HC 2 H 3 O 2 H+H+ C2H3O2-C2H3O2- I0.30 M00 C-x M+x M E(0.30 –x) Mx Mx Mx Mx M

 Percent Ionization = [H + ] equilibrium x 100 [Acid] initial  Example  Percent Ionization of HC 2 H 3 O 2  = M x 100 = 0.77% 0.30 M

 In polyprotic acids, H atoms ionize in successive steps: ◦ H 2 SO 4(aq) ⇌ H + + HSO 3 - (aq) Ka = 1.71 x ◦ HSO 3 - (aq) ⇌ H + (aq) + SO 3 2- (aq) Ka = 6.4 x  It is always easier to remove the first proton from a polyprotic acid than the second.  As long as successive K a values differ by a factor of 10 3 or more, it is possible to obtain a satisfactory estimate of the pH of polyprotic acid solutions by considering K a1

 The most commonly encountered weak base is NH 3 NH 3(aq) + H 2 0 ⇌ NH 4 + (aq) + OH - (aq)  Base-dissociation constant  The constant K b always refers to the equilibrium in which a base reacts with H 2 O to form the corresponding conjugate acid and OH -  Table 16.4 on page 636 provides formulas and K b values for several weak bases in water.

 Weak Bases fall into 2 general categories 1.Neutral substances that have an atom with a nonbonding pair of electrons that can serve as a proton acceptor. 2.Anions of weak acids ClO - (aq) + H 2 O (l) ⇌ HClO (aq) + OH - K b = 3.33 x 10 -7

 The product of the acid-dissociation constant for and acid and the base-dissociation constant for its conjugate base is the ion- product for water ◦ (K a )(K b ) = K w AcidKaKa BaseKbKb HNO 3 (strong acid)NO 3 - Negligible HF6.8 x F-F- 1.5 x HC 2 H 3 O x C2H3O2-C2H3O x H 2 CO x HCO x NH x NH x HCO x CO x OH - NegligibleO 2- (Strong base)

 Sometimes acid- and base- dissociation constants are expressed as pK a or pK b ◦ pK a = -log K a ◦ pK b = -log K b ◦ pK a + pK b = pK w = at 25˚C

 Salts completely dissociate in water  Many ions are able to react with water to generate H + or OH - ions. ◦ This type of reaction is called hydrolysis ◦ The pH of an aqueous salt solution can be predicted by considering the ions of which the salt is composed.

 In general, an anion, X-, in solution can be considered the conjugate base of an acid. ◦ Whether or not an anion will react with water to produce OH- depends on the strength of the acid the anion would create. ◦ To identify the acid, add a proton to the anion ◦ X - + a proton = HX ◦ If the acid is a strong acid, the anion will have a negligible tendency to abstract a proton from water. ◦ If the acid is a weak acid, X - will react to a small extent. He pH would be higher (more basic) ◦ If the ion has an ionizable proton like HSO 3 -, it is amphoteric. The behavior is determined by the magnitude of K a and K b

 Most metal ions can react with water to decrease the pH of an aqueous solution ◦ Ions of alkali metals and of the heavier alkaline earth metals do not react with water and therefore do not affect pH. (same cations that make strong bases)

1. An anion that is conjugate base of a strong acid will not affect the pH of solution 2. An anion that is conjugate base of a weak acid will cause an increase in pH 3. A cation that is the conjugate acid of a weak base will cause a decrease in pH 4. With the exception of group 1A and heavy 2B, metal ions will decrease pH 5. When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the ion with the largest ionization constant will have the greatest affect on pH

 Predict whether the salt Na 2 HPO 4 will form an acidic or basic solution on dissolving in water. ◦ Na 2 HPO 4 ⇆ 2Na + + HPO 4 2- ◦ HPO 4 2- can act like an acid or base  1. HPO 4 2- ⇆ H+ + PO 4 3- (acid)  2. HPO H 2 O ⇆ H 2 PO OH - (base)  The reaction with the larger ionization constant will determine whether it is acidic or basic  K a1 = 4.2 x (Table 16.3)  (K a2 )(K b2 ) = K w  (6.2 x )(K b2 ) = 1.0 x  K b = 1.6 x  Since K b > K a, reaction is basic

 Factors that Affect Acid Strength 1.A molecule containing H will transfer a proton only if the H-X bond is polarized in the following way  In ionic hydrides (NaH) H has a negative charge and acts as a proton acceptor 2.Strong bonds (Table 8.4) are more difficult to dissociate than weak bonds 3.The greater the stability of the acid’s conjugate base, the stronger the base

 H-X strength is the most important factor in determining acid strength in binary acids  The H-X strength decreases as element X increases in size

 Acids in which OH groups and possibly additional oxygen atoms are bound to a central atom are called oxyacids.  Consider -Y-O-H ◦ When Y is a metal, they are sources of OH - ions and behave as bases ◦ When Y is a nonmetal, the OH bond is more polar so it will donate a H + and behave as an acid

 For oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom. ◦ Example: Strength of H-O-Y, increases as electronegativity of Y increases  For oxyacids that have the same central atom Y, acid strength increases as the number of oxygen atoms attached to Y increases. ◦ Example: the strength of the oxyacids of chlorine increases from HClO to HClO 2 to HClO 3 to HClO 4

 Another group of acids are those with a carboxyl group, often written COOH ◦ Example: acetic acid HC 2 H 3 O 2 can also be written CH 3 COOH  Other examples:  Formic AcidBenzoic Acid

 The acid strength of carboxylic acids increase as the number of electronegative atoms in the acid increase. ◦ Example: trifluoroacetic acid (CF 3 COOH) is stronger than acetic acid (CH 3 COOH)

 For a substance to be a proton acceptor (base), it must have an unshared pair of electrons for binding a proton.  G.N. Lewis noticed this and proposed definitions for acid and base ◦ A Lewis acid is an electron-pair acceptor ◦ A Lewis base is an electron-pair donor ◦ **Not using this definition in AP Chem.**