more equilibrium Dr. Fred Omega Garces Chemistry, Miramar College 1
Acids-Bases Characteristics Acids (Properties) Taste Sour Dehydrate Substances Neutralizes bases Dissolves metals Examples: Juices: TJ, OJ, AJ Wine Banana Coffee Vitamin C Soda Base (Properties) Taste Bitter Denatures Proteins Neutralizes acids Turns metal g hydroxides Examples: Milk of Magnesia Lime water Lye, Drain Ammonia blood Soap 2
Practice Naming Acids HNO 2 HI HF Nitrous Acid Hydroiodic Acid Hydrofluoric Acid Oxy-anions Add H+ Oxy-acids H 3 PO 4 H 3 PO 3 Phosphoric Acid Phosphorous Acid Add H+ HClO 2 HClO Chlorous Acid Hypochlorous Acid Add H+ H 2 CO 3 HCN Carbonic Acid Hydrocyanic Acid Add H+ HC 2 H 3 O 2 Acetic Acid H 2 Cr 2 O 7 Dichromic Acid HClO 4 Perchloric Acid 3
Arrhenius Definition Svante Arrhenius (1859-1927) Acid - Increases H + (H 3 O + ) concentration Base - Increases OH - concentration Arrhenius acids and bases are limited to aqueous solutions. Examples: Acids are substances that are able to ionize to form hydrogen ion and thereby increase the concentration of H + (aq) ions in aqueous solutions. HNO 3 (aq) g H + (aq) + NO 3 - (aq) Bases are substances that accept (react with) H + ions. Hydroxide ions, OH -, are basic because they readily react with H + ions to form water: H + (aq) + OH - (aq) g H 2 O (l) 4
New Definition: Bronsted-Lowry Acids-Bases Bronsted - Lowry definition Acid - Proton H + (H 3 O + ) donor Base - Proton H + (H 3 O + ) acceptor. example: acids: HCl (aq) H + (aq) + Cl - (aq) Bases: NH 3 (aq) + H 2 O (l) NH + 4 (aq) + OH - (aq) HCl (aq) + NH 3 (aq) NH + 4 (aq) + Cl - (aq In an acid - base reaction, H + & OH - always combine together to form water and an ionic compound (a salt): HCl (aq) + NaOH (aq) g H 2 O (l) + NaCl (aq) 5
Factors Affecting Acid/Base Strength What determines the strength of acids and Base? Dissociation property- Electrolyte - Substances which dissociate in water. Strong electrolyte completely dissociates to ions Weak electrolyte undergoes partial dissociation. Acid Examples : HClO 4 + H 2 O g H 3 O + + ClO - 4 100% Dissociation (Strong Acid) HNO 3 + H 2 O g H 3 O + + NO 3-100% Dissociation (Strong Acid) H 2 S + H 2 O D H 3 O + + HS - Less 100% Dissociation (Weak Acid) Base Examples : NaOH + H 2 O g OH - + Na + 100% Dissociation (Strong Base) Ca(OH) 2 + H 2 O g 2OH - + Ca +2 100% Dissociation (Strong Base) NH 3 + H 2 O D OH - + NH + 4 Less 100% Dissociation (Weak Base) 6
Strong Acids Strong Acids Earlier we described acids and bases as electrolyte; that is, these substances ionizes (break up to ions) in solution. Strong electrolyte are substances which completely dissociates (100%). SA-Strong acids completely dissociates to H + and anion. HA g H + + A - Strong Acids; HX, H 2 SO 4, HNO 3, HClO 4 HClO 3 : All others are considered weak. HX: HCl, HBr, HI (all Halogens, except F ) H 2 SO 4 Sulfuric Acid HNO 3 Nitric Acid HClO 4 Perchloric Acid HClO 3 Chloric Acid Weak acids, incompletely dissociation in water. i.e., CH 3 COOH (acetic acid) only 1 in 100 dissociate 7
Strong Acids; indicative by acid dissociation constant, Ka The K a values shown are the acid dissociation constant for monoprotic acids. The K a is the equilibrium constant for the ionization of the acid, HA D H + + A. Strong acids do not have meaningful K a values and are left out from this table. 8
Strong Bases Strong Bases Substances which completely dissociate into cation and hydroxide. MOH g M + + OH - Strong Bases, M(OH) n ; NaOH, KOH, LiOH, Ca(OH) 2, Ba(OH) 2, Sr(OH) 2 M n O; Li 2 O, Na 2 O, K 2 O All others are considered weak. Bases: 3 ways of OH - formation Extraction of H + from H 2 O: i.e.., reaction of NH 3 in water: NH 3 (aq) + H 2 O (l) D NH + 4 (aq) + OH - (aq) Extraction of H + from H 2 O by charge i.e.., reaction of CO 2-3 in water: CO 2-3 (aq) + H 2 O (l) D HCO - 3 (aq) + OH - (aq) Dissociation of ionic substance to OH - i.e.., dissociation of NaOH in water: NaOH (aq) D Na + (aq) + OH - (aq) 9
Chemistry of Acids and Bases Acids Base Strong Acid Weak Acid 10
Water: Acid-Base Properties Auto-ionization (Self-Ionization) of water Why does water have a ph of 7? Water is Amphoteric (it reacts with itself) 2 in 1 billion water molecules self-ionizes. H 2 O (l) + H 2 O (l) + Energy K w D H 3 O (aq) + OH - (aq) K eq = K w Endothermic reaction K w (ion-product constant) = 1 10-14 see that, [H 3 O + ] = 1 10-7 M [OH - ] = 1 10-7 M The equation K w = [H 3 O + ][OH - ] is valid in pure water and in any aqueous solution. K w is temperature-dependent, the autoionization rxn is endothermic, so K w increases with temperature. C Kw 10 0.29 10-14 15 0.45 10-14 20 0.68 10-14 25 1.01 10-14 30 1.47 10-14 50 5.48 10-14 11
K w & [H 3 O + ] How does K w dictate concentration of H 3 O + and OH -? K w = 1 10 14 # = H 3 O + & # & $ % ' ( $ % OH '( at 25 C For pure water, If [H 3 O] + or [OH] - is concentration known, the other can be determine through the mass action expression and a iδe table. 2H 2 O (l)! OH - + H 3 O + i Excess 0 0 Δ - 2x + x + x e excess x x K w = 1 10-14 = x 2 ph g 7.0 1 10-7 = x = [H 3 O] + = [OH] - ph calculation to be discussed. 12
Consequences of Auto-ionization A change in [H 3 O + ] causes an inverse change in [OH - ] & vice versa. Higher [H 3 O + ] g lower [OH - ] Higher [OH - ] g lower [H 3 O + ] Both ions are present in all aqueous systems. Acidic solution g [H 3 O + ] > [OH - ] Neutral solutions g [H 3 O + ] = [OH - ] Basic Solutions g [OH - ] > [H 3 O + ] Kw = [H 3 O + ] [OH - ] or pkw = ph + poh (as shown later) 13
For pure water [H 3 O + ] = 1 10-7 M Since 1 10-7 M = 0.0000001M Concentration is so small, it is much more convenient to use a scale called p-h (power of hydrogen or potential of hydrogen) ph Calculation 1 10-7 M g ph = 7 ph = - log [H 3 O + ] or [H 3 O + ] = 1 10 -[ph] Example: 1) What is the ph of a hydrobromic acid solution with a molar concentration of 1.67 10-5 M. (3-Sig figures) 2) Double check your answer by taking the ph answer from #1, and calculate the molar concentration of the H 3 O + for the hydrobromic acid solution. Calculate: ph for [H 3 O + ] = 1.67 10-5 M [H 3 O + ] of ph = 4.777 Calculator Sequence: 1.67 EE 5 +/- log +/- 4.78 +/- 10 x Answer : 4.777 (3-Sig figs) 1.67 10-5 14
ph and the Concentration of Acids ph Scale Conc [H 3 O + ] Exp [H 3 O + ] ph poh 1M 1 10 0 0 14 0.1M 1 10-1 1 13 0.01M 1 10-2 2 12 0.001M 1 10-3 3 11 0.0001M 1 10-4 4 10 0.000001M 1 10-5 5 9 0.0000001M 1 10-6 6 8 0.00000001M 1 10-7 7 7 0.000000001M 1 10-8 8 6 0.0000000001M 1 10-9 9 5 0.00000000001M 1 10-10 10 4 0.000000000001M 1 10-11 11 3 0.0000000000001M 1 10-12 12 2 0.00000000000001M 1 10-13 13 1 0.000000000000001M 1 10-14 14 0 15
K w, ph and poh ph measures the concentration of [H 3 O + ] K w = 1 10-14 = [H 3 O + ] [OH - ] pk w = ph + poh 14 = ph + poh Example: What is the ph of a potassium hydroxide solution with a molar concentration of 3.0 10-4 M (2 sig figs) Answer: Potassium hydroxide, KOH is a strong base, therefore, the molar concentration of KOH (3.0 10-4 M) is equal to the [OH-] concentration or [OH-] = 3.0 10-4 M poh = -log (3.0 10-4 ) = 3.52 ph + poh = 14.00 g ph = 10.48 16
Determining ph, poh, [OH - ], [H 3 O + ] Use this chart to determine acid and base concentration at (25 C) K w = 1.00 10-14 & pk w = 14 [H 3 O + ] [OH - ] = K w & ph + poh = pk w K w / [H 3 O + ] [OH - ] [H 3 O + ] [OH - ] = K w -log [H 3 O + ] K w / -log 10 -ph 14.0 - ph ph + poh = pk w = 14 14.0 - poh 10 -poh 17
Kw, ph and poh Example 2 ph measures the concentration of [H 3 O + ] Kw = 1 10-14 = [H 3 O + ] [OH - ] pkw = ph + poh 14 = ph + poh Example: What is the ph of a solution having a [OH - ] = 3.0 10-13 M Way of Multiplying Way of adding / subtracting Using the formula: Using the formula: 1 10-14 = [H 3 O + ] [OH - ] 14 = ph + poh 1 10-14 = [H 3 O + ] [OH - ] 14 = ph + poh 1 10-14 =[H 3 O + ] [3.0 10-13 ] 14 = ph + 12.52 [H 3 O + ] =3.33 10-2 M ph = 1.48 ph = 1.48 18
Acid Base in Water Summary Pure water has a low conductivity because it autoionizes to a small extent. This process is described by an equilibrium reaction whose equilibrium constant is the ion-product constant for water, K w (1 10-14 at 25 C). Thus, [H 3 O + ] and [OH - ] are inversely related. In acidic solution, [H 3 O + ] is greater than [OH-], the reverse is true in basic solution, and the two are equal in neutral solution. To express small values of [H 3 O + ] more simply, we use the ph scale (ph = -log [H 3 O + ]). A high ph represents a low [H 3 O + ]. Similarly, poh = -log [OH - ], and pk = - log K. At 25 C, in acidic solutions, ph < 7.00, in basic solutions, ph > 7.00; and in neutral solutions, ph = 7.0. The sum of the ph and poh equals pk w (14.00 at 25 C) 19