Acid - Base Equilibria AP Chapter 16

Acids and Bases Arrhenius acids have properties that are due to the presence of the hydronium ion (H + ( aq )) They turn litmus red. Arrhenius bases have properties that are due to the presence of the hydroxide ion (OH - ( aq )). They turn litmus blue.

Brønsted-Lowery Acids and Bases Bronsted-Lowery definitions state that acid-base reactions involve the transfer of hydronium ions (H + ) from one substance to another. A hydronium ion is simply a proton with no surrounding valence electrons. Remember – acids donate and bases accept!

Hydronium ion

Proton transfer reactions The polar water molecule promotes the ionization of acids in water solution by accepting a proton to form H 3 O +.

Brønsted-Lowery definitions ACID BASE Acids donate protons Bases accept protons

The Hydronium ion The hydronium ion (H 3 O + ) is a hydrated proton. When water accepts a proton from an acid, the product is a hydronium ion. Hydronium ions are represented by either the H 3 O + ( aq ) or H + ( aq )

Conjugate Acid-Base Pairs Conjugate acid-base pairs are two substances in an aqueous solution whose formulas differ by an H +. The acid is the more positive species having an extra H.

Amphoteric An amphoteric substance is a substance that can act as either an acid or a base.

Relative Strengths of Acids and Bases A strong acid completely transfers its protons to water, leaving no undissociated molecules in water. It totally dissociates in water.

Relative Strengths, continued A weak acid only partially dissociates in water, and exists as a mixture of acid molecules and their constituent ions. The conjugate base of a weak acid is a weak base.

Relative Strengths, continued A substance with negligible acidity, such as CH4, contains hydrogen, but does not demonstrate any acidic behavior in water. It’s conjugate base is a strong base, reacting completely with water.

Acid-Base Equilibrium In every acid-base reaction the position of equilibrium favors transfer of the proton from the stronger acid to the stronger base to form the weaker acid and the weaker base.

Autoionization of Water Water has the ability to act as either an acid or a base.

The ion product of water Because the autoionization of water is an equilibrium process, there is an equilibrium-constant expression: K c = [H 3 O + ][OH - ]

Ion-Product Constant for H 2 O K w = [H 3 O + ][OH - ] = 1.0 x can also be written = [H + ][OH - ] = 1.0 x

Acid and Base Ionization Constants The acid ionization constant ( Ka ) is the equilibrium constant for the ionization of a weak acid in water. The base ionization constant ( Kb ) is the equilibrium constant for a weak base. For any conjugate acid-base pair, Kw = Ka x Kb.

pH Scale pH = -log[H + ] Neutral solution: pH = -log(1.0 x ) = -(-7.00) = 7.00 The pH decreases as the [H+] increases.

The pH Scale

Calculating the pH of a Basic Solution Calculate the pH of a basic solution, where the [OH - ] > 1.0 x M. Suppose [OH - ] = 2.0 x10 -3 M. Calculate the H + value for this solution. [H + ] === 5.0 x M Kw 1.0 x [OH - ] 2.0 x pH = -log(5.0 x ) = 11.30

pH and pOH pH and pOH = 14.00

Measuring pH pH meter Acid-base indicators (less precise) Methyl orange Litmus phenolphthalein Etc.

Strong Acids and Bases The seven most common strong acids include 6 monoprotic acids (HCl, HBr, HI, HNO 3, HClO 3, and HClO 4 ) and one diprotic acid, H 2 SO 4. HNO 3 ( aq ) + HOH( l ) → H 3 O + ( aq ) + NO 3 - ( aq ) HNO 3 ( aq ) → H + ( aq ) + NO 3 - ( aq )

Calculating the pH of a strong acid What is the pH of a M solution of HClO 4 ? pH = -log(0.040) = 1.40

Strong Bases Common strong bases are the ionic hydroxides of alkali metals and the heavy alkaline earth metals. The cations of these metals have negligible acidity.

Weak Acids Weak acids are only partially ionized (or dissociated.) They are weak electrolytes. HA( aq ) + HOH( l ) ↔ H 3 O + ( aq ) + A - ( aq ) Ka = The larger the value of Ka, the stronger the acid. [H 3 O + ][A - ] [HA]

Polyprotic Acids Polyprotic acids have more than one ionizable proton, such as H 2 SO 3. These acids have acid-dissociation constants that decrease in magnitude in the order K a1 >K a2 >K a3. Because nearly all the H + (aq) in a polyprotic solution comes from the first dissociation, the pH can usually be estimated using only K a1.

Weak Bases Weak bases include NH 3, amines and the anions of weak acids. K b = the dissociation constant for the base. The relationship between the strength of an acid and the strength of its conjugate base is expressed by the equation K a x K b = K w

Using K b to Calculate OH - NH 3 ( aq ) + HOH( l ) ↔ NH 4 + ( aq ) + OH - ( aq ) K b = [NH 4 + ][OH - ] [NH 3 ] = 1.8 x Reference the problem example on page 691.

Hydrolysis Acid-base properties of salts can be attributed to the behavior of their respective cations and anions. The reaction with water, with a resulting change in pH, is called hydrolysis. Cations of alkali metals and alkaline earth metals and anions of strong acids don’t hydrolyze. Salt + water = acid + base

Acid-Base Behavior and Chemical Structure A molecule containing H will transfer a proton only of the H-X bond is polarized:

Bond Strength Strong bonds do not dissociate as easily as weaker bonds, so they are less likely to form acidic ions in solution. Since HF has such a strong bond due to the electronegativities, it does not dissociate readily and is therefore a weak acid.

Oxyacids Oxyacids are acids in which OH groups and possible additional oxygen atoms are present. What determines whether it is an acid or a base? Generally, as the electronegativity of the attached element increases, so will the acidity of the substance.

Oxyacids The strength of an acid will increase as additional electronegative atoms bond to the central atom. Electronegativity

Acid strength increases as the number of oxygen atoms attached to the central atom increases.

Carboxylic Acids Acids that contain carboxyl groups are called carboxylic acids. These form the largest category of organic acids.

Lewis Acids and Bases The Lewis concept of acids emphasizes the shared electron pair rather than the proton. A Lewis acid is an electron-pair acceptor. A Lewis base is an electron-pair donor. This concept is more general than the Brønsted-Lowery definition – it explains why many hydrated metal cations can form acidic aqueous solutions.

Acidity of Metal Cations The acidity of a hydrated metal cation depends on the cation charge and size.