Chapter 15 Aqueous Equilibria: Acids and Bases 15.1 -15.3: Definition of Acids and Bases Properties of Base
Properties of Acid • Taste bitter • Sour taste alkaloids = plant product that is alkaline • Ability to dissolve many metals often poisonous • Ability to neutralize bases • Feel slippery • Change blue litmus paper to red • Ability to turn red litmus paper blue
• Ability to neutralize acids
Definition of Arhenius acids and bases Acid: any compounds that produce H+ when dissolve in water Base: any compounds that produce –OH when dissolve in water
Base: any compounds that produce –OH when dissolve in water Problems:
• It does not explain why molecular substances, such as NH3, dissolve in water to form basic solutions, even though they do not contain OH– ions.
• It does not explain how some ionic compounds, such as Na2CO3 or Na2O, dissolve in water to form basic solutions, even though they do not contain OH– ions.
• It does not explain why molecular substances, such as CO2, dissolve in water to form acidic solutions, even though they do not contain H+ ions. • It does not explain acid–base reactions that take place outside aqueous solution. This theory is limited to only “those compounds”
Dang1 The Brønsted-Lowry Theory: a better definition for acids and bases in aqueous solution Brønsted-Lowry Acid: A substance that can transfer hydrogen ions, H+. In other words, a proton donor loses a proton Conjugate Acid-Base Pairs: Chemical + - HA(aq) + H2O(l) H3O (aq) + A (aq) species whose formulas differ only by one hydrogen ion, H+ gains a proton HA and A-
+ H2O and H3O
+ In the acid example, notice the formation of H3O -- this species is named the hydronium ion. It lets you know that the solution is acidic!
+ - HCl(aq) + H2O(l) H3O (aq) + Cl (aq) (neutral compound)
Brønsted-Lowry Base: A substance that can accept hydrogen ions, H+ . In other words, a proton acceptor Loses a proton Conjugate Acid-Base pairs - + B(aq) + H2O(l) OH(aq) + BH (aq) B and BH+
gains a proton - H2O and OH
- + NH3(aq) + H2O(l) OH(aq) + NH4 (aq)
Amphoteric Substances
• Amphoteric substances can act as either an acid or a base because they have both a transferable H+ and an atom with lone pair electrons.
• Water acts as base, accepting H+ from HCl.
– + HCl(aq) + H2O(l) → Cl (aq) + H3O (aq)
+ • Water acts as acid, donating H to NH3.
+ – NH3(aq) + H2O(l) NH4 (aq) + OH (aq)
Examples: Write a balanced equation for the dissociation of each of the following Bronsted-Lowry acids in water
H2CO3(aq)
Dang2 - HSO3 (aq)
Examples: Write a balanced equation for the dissociation of each of the following Bronsted-Lowry bases in water
2- CO3 (aq)
CH3NH2(aq)
Example: In the following equation, label each species as an acid (A) or a base (B) on the left. Identify the conjugated acid (CA) or base on the right (CB) - - a. NO2 (aq) + HF(aq) HNO2(aq) + F (aq)
+ - b. HClO4(aq) + H2O(l) H3O (aq) + ClO4 (aq)
15.4 Acid Strength and the Acid Ionization Constant, Ka Strength is determined by the position of the "dissociation" equilibrium
Acid Strength and Base Strength
Strong acids/strong bases/strong electrolyte 1. dissociates completely in water 2. all the base molecules form H+ or OH– ions, either Dangthrough3 dissociation or reaction with water, 3. have very large dissociation or K values • Strong acids donate practically all their H’s. 100% ionized in water Strong electrolyte + • [H3O ] = [strong acid]ini • Has no Ka values (why?)
Other acids are considered “weaker acids”
Weak acids donate a small fraction of their H’s.
Most of the weak acid molecules do not donate H to water.
Much less than 1% ionized in water
+ [H3O ] << [weak acid]
1. dissociate only to a slight extent in water 2. only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water 2. dissociation constant is very small Do Not confuse concentration with strength!
Strengths of Acids and Bases
Dang4 • Commonly, acid (HA) or base (B) strength is measured by determining the equilibrium constant of a substance’s reaction with water. - + HA + H2O A + H3O + − B + H2O HB + OH
The stronger the acid the weaker its conjugate base, the converse is also true. • The stronger an acid is at donating H, the weaker the conjugate base is at accepting H.
• Higher oxidation number = stronger oxyacid
– H2SO4 > H2SO3; HNO3 > HNO2
• Cation stronger acid than neutral molecule; neutral stronger acid than anion
+ − + − – H3O > H2O > OH ; NH4 > NH3 > NH2
• The farther the equilibrium position lies toward the products, the stronger the acid or base. • The position of equilibrium depends on the strength of attraction between the base form and the H+. Stronger attraction means stronger base or weaker acid
Trend in base strength opposite
Examples Dang5 If you mix equal amount concentrations of reactants and products, decide which species (reactants or products) are favored at the completion of the reaction. Why?
◦ + - H2SO4(aq) + NH3(aq) NH4 (aq) + HSO4 (aq)
◦ - - HF(aq) + NO3 (aq) HNO3(aq) + F (aq)
15.5 Autoionization of water Fredrich Kohlrausch, around 1900, found that no matter how pure water is, it still conducts a minute amount of electric current. This proves that water self-ionizes. • Since the water molecule is amphoteric, it may dissociate with itself to a slight extent. • Only about 2 in a billion water molecules are ionized at any instant!
+ − H2O(l) + H2O(l) H3O (aq) + OH (aq)
The equilibrium expression used here is referred to as the autoionization constant for water, Kw o + - -14 At 25 C Kw = [H3O ] [ OH] = 1.0 x 10
+ - Knowing this value will help to calculate [H3O ] or [ OH] at various situations
Acidic: Neutral Basic + + + [H3O ] > [OH ] [H3O ] = [OH ] [H3O ] < [OH ]
o −13 + - Example: At 60 C, At 60°C, the value of Kw is 1.0 × 10 . Calculate the [H3O ] and [ OH] in a neutral solution
The pH Scale Used to designate the [H+] in most aqueous solutions where [H+] is small. pH = −log [H+] pOH = − log [-OH] pH + pOH = 14.00
Acidic: pH < 7 Neutral: pH = 7 Basic: pH > 7
Dang6 The pH Scale and Significant figures + The hydronium ion concentration for lemon juice is approximately 0.0025 M. What is the pH when [H3O ] = 0.0025 M? 2 significant figures
pH = log(0.0025) = 2.60
2 decimal places
- + -5 Examples: Calculate the [ OH] in a solution with [H3O ] = 7.5 x 10 M. Is this solution basic, acidic or neutral?
Calculate the pH of a solution with [-OH] = 8.2 x 10-10M
+ – Calculate the concentration of H3O and OH for a solution with a pH of 8.37
15.6 Measuring pH Acid-Base Indicator: A substance that changes color in a specific pH range. Indicators exhibit pH-dependent color changes because they are weak acids and have different colors in their acid (HIn) and conjugate base (In) forms.
- HIn(aq) + H2O(l) H3O+(aq) + In(aq)
• Indicators are chemicals that change color depending on the solution’s acidity or basicity.
• Many vegetable dyes are indicators.
– Anthocyanins
• Litmus Dang7 – From Spanish moss
– Red in acid, blue in base
• Phenolphthalein
– Found in laxatives
Red in base, colorless in acid
pH of strong acids and strong bases
A strong monoprotic acds – 100% dissociated in aqueous solution (HClO4, HCl, HBr, HI, H2SO4 HNO3 etc…) ◦ Contains a single dissociable proton + - HA(aq) + H2O(l) H3O (aq) + A (aq)
◦ + pH = - log H3O
◦ + - [H3O ] = [A ] = initial concentration of the acid ◦ Undissociated [HA] = 0 Strong Bases Alkali metal hydroxide, MOH ◦ Water-soluble ionic solids ◦ Exits in aqueous solution as alkali metal cations and hydroxide anions ◦ Calculate pH from [-OH]
Alkaline earth metal hydroxide, M(OH)2 where M= Mg, Ca, Sr, Ba ◦ Less soluble than alkali hydroxide, therefore lower [-OH]
Dang8 The pH in Solutions of Strong Acids and Strong Bases - Example: Calculate the [ OH] and pH of 0.30 M HNO3
What is the pH of a 0.0250 M solution of Ca(OH)2?
The pH of an unknown acid solution is 2.45. What is the concentration of the acid solution?
Dang9