CHEM 1412 Concept Review: Acids, Bases, & Acid-Base Equilibrium

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Definitions for Acids and Bases

	Arrhenius Definitions	Bronsted-Lowry Definitions	Lewis Definitions
For an ACID	A substance that produces H+ ions when dissolved in solution.	A substance that donates a proton (H+) as part of a chemical reaction.	A substance that accepts a pair of electrons as part of a chemical reaction.
For a BASE	A substance that produces OH- ions when dissolved in solution.	A substance that accepts a proton (H+) as part of a chemical reaction.	A substance that donates a pair of electrons as part of a chemical reaction.

Note:  We will largely be focusing on acids and bases from the perspective of the Bronsted-Lowry and Arrhenius definitions in this chapter.

 

 

The Auto-Ionization Equilibrium of Water

	Arrhenius framework	Bronsted-Lowry framework
Chemical Equation	$H_2O\left(l\right)\longleftrightarrow H^+\left(aq\right)+OH^-\left(aq\right)$	$2H_2O\left(l\right)\longleftrightarrow H_3O^+\left(aq\right)+OH^-\left(aq\right)$
Equilibrium Constant	$K_w=\left[H^+\right]\left[OH^-\right]=1.0\times10^{-14}$	$K_w=\left[H_3O^+\right]\left[OH^-\right]=1.0\times10^{-14}$

 

Relative Acid-Conjugate Base / Base-Conjugate Acid Strength Chart

As you can see from above, the stronger the acid is, the weaker its conjugate base will be and vice-versa.

 

 

Bronsted-Lowry Acid-Base/Conjugate Acid-Conjugate Base Reactions

Note: The conjugate acid and the conjugate base are the acid and base for the REVERSE reaction, and are therefore always products.

 

 

Weak Acid Dissociation Equilibrium (K_a)

	Arrhenius framework	Bronsted-Lowry framework
Generic Chemical Equation	$HA\left(aq\right)\longleftrightarrow H^+\left(aq\right)+A^-\left(aq\right)$	$HA\left(aq\right)+H_2O\left(l\right)\longleftrightarrow H_3O^+\left(aq\right)+A^-\left(aq\right)$
Equilibrium Constant Expression	$K_a=\frac{\left[H^+\right]\left[A^-\right]}{\left[HA\right]}$	$K_a=\frac{\left[H_3O^+\right]\left[A^-\right]}{\left[HA\right]}$

Note:  The equilibrium reaction equation associated with K_a is always the reaction of a weak acid with WATER to form H₃O⁺ and the conjugate base of the weak acid.

 

 

Percent Ionization of an Weak Acid:

$Percent\:Ionization=\frac{\left[H_3O^+\right]_{equilibrium}}{\left[HA\right]_{initial}}\times100\%$

([H₃O⁺] can be determined using an ICE chart and Ka)

 

 

Weak Base Dissociation Equilibrium (K_b)

Generic Chemical Equation	$B\left(aq\right)+H_2O\left(l\right)\longleftrightarrow BH^+\left(aq\right)+OH^-\left(aq\right)$
Equilibrium Constant Expression	$K_b=\frac{\left[BH^+\right]\left[OH^-\right]}{\left[B\right]}$

 

 

What is pH?

In pH, the “p” means “power of”, therefore the pH is the “power of hydrogen”, which refers to the magnitude of the negative exponent of the hydrogen ion concentration.  Mathematically, we determine the “power of hydrogen” by taking the negative logarithm of the hydrogen ion concentration, therefore mathematically “p” means “-log”.  This same relationship holds for other quantities as well.  Thus, we have the following relationships:                

pH and Related Equations

Logarithmic Form		Exponential Form
$pH=-\log\left[H_3O^+\right]$	AND	$\left[H_3O^+\right]=10^{-pH}$
$pOH=-\log\left[OH^-\right]$	AND	$\left[OH^-\right]=10^{-pOH}$
$pK_a=-\log\left(K_a\right)$	AND	$K_a=10^{-pK_a}$
$pK_b=-\log\left(K_b\right)$	AND	$K_b=10^{-pK_b}$

NOTE:  H⁺ may be used interchangeably with H₃O⁺ in the above equations.

Because the pH represents the magnitude of the negative exponent, the larger the pH, the fewer hydrogen ions are present in solution, making the solution less acidic.  The equilibrium maintained in aqueous solutions between [H₃O⁺] and [OH^-] leads to the following relationships between pH, pOH, pK_a, and pK_b:

$pH+pOH=14.00$             $pK_a+pK_b=14.00$             $K_a\times K_b=K_w=1.0\times10^{-14}$