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Lecture 13

Acids and Bases

Shaun Williams, PhD

What Are Acids and Bases?

Arrhenius Model

Hydronium Ion

  • A fundamental problem with the Arrhenius model is the treatment of the behavior of the hydrogen ion, \(\chem{H^+}\)
  • Hydrogen ions are better represented as hydronium ions, \(\chem{H_3O^+}\), in solution
A representation of an oxygen atom with three hydrogen atoms around it.

Brønsted-Lowry Theory

Conjugate Acid-Base Pairs

\[ \underbrace{\chem{HCl(g)}}_{\text{acid}} + \underbrace{\chem{H_2O(l)}}_{\text{base}} \rightarrow \underbrace{\chem{H_3O^+(aq)}}_{\text{ conjugate acid}} + \underbrace{\chem{Cl^-(aq)}}_{\text{conjugate base}} \]

Amphoteric Substances

Acidic Hydrogen Atoms

  • If an acid has more than one hydrogen atom, we need to determine which hydrogen atoms are acidic.
  • Typically, in oxoacids, any hydrogen atoms bonded to oxygen atoms are acidic.
Representations showing acetic acid with its one acidic hydrogen atom and phosphorous acid with its two acidic hydrogen atoms.

Strong and Weak Acids and Bases

Strong Acids and Bases

\(\chem{HCl}\) - Example of a Strong Acid

\[ \chem{HCl(aq)+H_2O(l) \rightarrow H_3O^+(aq)+Cl^-(aq)} \]

A picture of gaseous hydrogen chloride bubbling into water, dissolving and breaking up into ions.

Common Strong Acids

Formula Name
\(\chem{HCl}\) hydrochloric acid
\(\chem{HBr}\) hydrobromic acid
\(\chem{HI}\) hydroiodic acid
\(\chem{HNO_3}\) nitric acid
\(\chem{HClO_3}\) chloric acid
\(\chem{HClO_4}\) perchloric acid
\(\chem{H_2SO_4}\) sulfuric acid

\(\chem{NaOH}\) - Example of a Strong Base

A picture of solid sodium hydroxide going into water and dissolving and breaking up into its ions.

Common Strong Bases

Formula Name
\(\chem{LiOH}\) lithium hydroxide
\(\chem{NaOH}\) sodium hydroxide
\(\chem{KOH}\) potassium hydroxide
\(\chem{Mg(OH)_2}\) magnesium hydroxide
\(\chem{Ca(OH)_2}\) calcium hydroxide
\(\chem{Ba(OH)_2}\) barium hydroxide

Weak Acids and Bases

\(\chem{HC_2H_3O_2}\) - Example of a Weak Acid

A picture of acetic acid in water with very few of the acetic acid molecules dissociated into ions.

\(\chem{NH_3}\) - Example of a Weak Base

A picture of ammonia in water with very few of the ammonia molecules existing as ammonium ions.

Some Common Weak Acids

Formula Name Occurance
\(\chem{HC_2H_3O_2}\) Acetic acid Vinegar, sour wine
\(\chem{H_2CO_3}\) Carbonic acid Soda, blood
\(\chem{H_3C_6H_5O_7}\) Citric acid Fruit, soda
\(\chem{HF}\) Hydrofluoric acid Glass etching
\(\chem{HOCl}\) Hypochlorous acid Sanitize pool and drinking water
\(\chem{HC_3H_5O_3}\) Lactic acid Milk
\(\chem{HC_4H_4O_5}\) Malic acid Fruit
\(\chem{H_2C_2O_4}\) Oxalic acid Nuts, cocoa, parsley
\(\chem{H_3PO_4}\) Phosphoric acid Soda, blood
\(\chem{H_2C_4H_4O_6}\) Tartaric acid Candy, wine, grapes

Some Common Weak Base

Formula Name Occurance
\(\chem{NH_3}\) Ammonia Glass cleaners
\(\chem{CaCO_3}\) Calcium carbonate Antacids, minerals
\(\chem{Ca(ClO)_2}\) Calcium hypochlorite Chlorine source for swimming pools
\(\chem{CH_3NH_2}\) Methylamine Herring brine
\(\chem{(CH_3)_3N}\) Trimethylamine Rotting fish

Relative Strengths of Weak Acids

Acid Ionization Constants

Weak Acids and \(K_a\) Values

Acid \(K_a\) Conjugate Base
Strongest Acids \(\chem{HF}\) \(6.3 \times 10^{-4}\) \(\chem{F^-}\) Weakest Bases
\(\chem{HNO_2}\) \(5.6 \times 10^{-4}\) \(\chem{NO_2^-}\)
\(\chem{HCO_2H}\) \(1.8 \times 10^{-4}\) \(\chem{HCO_2^-}\)
\(\chem{HC_2H_3O_2}\) \(1.8 \times 10^{-5}\) \(\chem{C_2H_3O_2^-}\)
\(\chem{HOCl}\) \(4.0 \times 10^{-8}\) \(\chem{OCl^-}\)
\(\chem{NH_4^+}\) \(5.6 \times 10^{-10}\) \(\chem{NH_3}\)
Weakest Acids \(\chem{HCN}\) \(6.2 \times 10^{-10}\) \(\chem{CN^-}\) Strongest Bases

Acidic Hydrogen Atoms

  • An acid that contains more than one acidic hydrogen and can thus donate more than one \(\chem{H^+}\) ion
  • The acid donates one \(\chem{H^+}\) ion at a time in steps
  • The \(K_a\) values for polyprotic acids are often labeled to indicate the particular step in the overall ionization process (\(K_{a1}\), \(K_{a2}\), \(K_{a3}\), etc.)
Images of phosphoric acid, carbonic acid, and sulfuric acid.

\(K_a\) for Polyprotic Acids

Name Formula \(K_{a1}\) \(K_{a2}\) \(K_{a3}\)
Carbonic acid \(\chem{H_2CO_3}\) \(4.5 \times 10^{-7}\) \(4.7 \times 10^{-11}\) \(\)
Citric acid \(\chem{H_3C_6H_5O_7}\) \(7.4 \times 10^{-4}\) \(1.7 \times 10^{-5}\) \(4.0 \times 10^{-7}\)
Hydrosulfuric acid \(\chem{H_2S}\) \(8.9 \times 10^{-8}\) \(1.0 \times 10^{-19}\) \(\)
Oxalic acid \(\chem{H_2C_2O_4}\) \(5.6 \times 10^{-2}\) \(1.5 \times 10^{-4}\) \(\)
Phosphoric acid \(\chem{H_3PO_4}\) \(6.9 \times 10^{-3}\) \(6.2 \times 10^{-8}\) \(4.8 \times 10^{-13}\)
Sulfuric acid \(\chem{H_2SO_4}\) Strong \(1.0 \times 10^{-2}\) \(\)
Tartaric acid \(\chem{H_2C_4H_4O_6}\) \(1.0 \times 10^{-3}\) \(4.3 \times 10^{-5}\) \(\)

Acidic, Basic, and Neutral Solutions

Ion-Product Constant of Water

Definitions of Neutral, Acidic, and Basic Aqueous Solutions

Type of solution Relative Concentration \(\chem{ \left[ H_3O^+ \right]}\) \(\chem{ \left[ OH^- \right]}\) \(K_w\)
Neutral \(\chem{ \left[ H_3O^+ \right] = \left[ OH^- \right] }\) \(1.0\times 10^{-7}\) \(1.0\times 10^{-7}\) \(1.0\times 10^{-14}\)
Acidic \(\chem{ \left[ H_3O^+ \right] \gt \left[ OH^- \right] }\) \(\gt 1.0\times 10^{-7}\) \(\lt 1.0\times 10^{-7}\) \(1.0\times 10^{-14}\)
Basic \(\chem{ \left[ H_3O^+ \right] \lt \left[ OH^- \right] }\) \(\lt 1.0\times 10^{-7}\) \(\gt 1.0\times 10^{-7}\) \(1.0\times 10^{-14}\)

The pH Scale

  • The pH of a solution is the negative logarithm (base 10) of the \(\chem{H_3O^+}\) concentration: \[ pH=-\log \left[ \chem{H_3O^+} \right] \]
  • It is convenient to express the acidity of aqueous solutions on a pH scale (shown at right).
An imagine comparing the pH scale to the concentration of hydronium ions.

Calculating \(pOH\)

Calculating Concentrations from \(pH\) and \(pOH\)

Calculate Concentrations from pH and pOH Graphically

A graphical representation of converting between pH, pOH, the concentration of hydronium ion, and the concentration of hydroxide ions.

Measuring pH

  • pH meters and \(pH\) indicators are often used to determine the \(pH\) of a solution.
A photograph of a pH sensor in a solution. The pH meter is reading 5.48 at a temperature of 16.3 Celsius.

A chart showing the pH ranges where a variety of pH indicators change colors.

Buffered Solutions

A graphic showing that in a buffer with equal parts carbonic acid and bicarbonate, if we add an acid we lose some bicarbonate and get carbonic acid and if we add a base we lose some carbonic acid and get bicarbonate ions.

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