Acids and Bases in Chemistry

Acids and Bases 1

• Two definitions of acids and bases are examined: Brønsted-Lowry and Lewis.
• Brønsted-Lowry theory defines an acid as a proton (H⁺) donor and a base as a proton acceptor.
• An example of a Brønsted-Lowry acid-base reaction is shown with water and ammonia.
• In this reaction, water acts as an acid and ammonia acts as a base.
• Another example of a Brønsted-Lowry reaction is given with sulphuric acid and water.
• The Brønsted-Lowry theory has a disadvantage: it doesn't explain highly acidic solutions resulting from substances that aren't proton donors, such as FeCl₃.
• Lewis theory defines a Lewis acid as a lone pair acceptor and a Lewis base as a lone pair donor.
• An example illustrates a Lewis acid-base reaction where ammonia donates a lone pair to boron trifluoride.
• Another example shows the reaction between hydroxide and hydrogen chloride, where hydroxide acts as a Lewis base and hydrogen chloride acts as a Lewis acid.
• Lewis definitions are more comprehensive than previous definitions of acids and bases.
• These definitions are widely used in organic chemistry.
• Aqueous solutions of FeCl₃ are acidic because the hydrogen atoms in water are attracted to the metal cation, weakening their bonds with oxygen, which can then be abstracted by water molecules to form more H₃O⁺ ions.

Brønsted-Lowry Theory - Revision Slide

• In the Brønsted-Lowry theory example, H₂O donates a proton (H⁺ ion) to NH₃, making water a Brønsted-Lowry acid.
• In turn, NH₃ accepts a proton, making it a Brønsted-Lowry base.
• The reverse reaction shows a conjugate acid-base pair. NH₄⁺ donates a proton (H⁺ ion) to hydroxide (OH⁻), making NH₄⁺ a Brønsted-Lowry acid and OH⁻ a Brønsted-Lowry base.
• In summary, conjugate acid-base pairs are formed by an acid and a base in an equilibrium reaction.

Strong and Weak Acids

• A strong acid dissociates completely in water.
• The equilibrium lies to the right for strong acids.
• Strong acids have a very large equilibrium constant (HCl, for example).
• The conjugate base of a strong acid is very weak.

Weak Acids

• Weak acids dissociate only partially in water.
• The equilibrium for weak acids lies towards the left.
• Weak acids have a very small equilibrium constant, for example acetic acid.
• The conjugate base of weak acids is relatively strong.

Strengths of Acids and Their Conjugate Bases

• A table displays the relative strengths of various acids and their conjugate bases.
• The table is ordered from strong acids to weak acids.
• The stronger the acid, the weaker its conjugate base.
• The weaker the acid, the stronger its conjugate base.

Dissociation of Water

• Water is amphoteric, acting as both an acid and a base.
• Water as a base accepts protons (H⁺ ions), for example, when reacting with hydrochloric acid.
• Water as an acid donates protons (H⁺ ions), for example, when reacting with ammonia.
• A reaction between two water molecules illustrates how one water molecule acts as an acid and the other acts as a base, forming hydronium (H₃O⁺) and hydroxide (OH⁻) ions.

Water Dissociation

• Distilled water is a weak electrolyte.
• The concentration of hydronium ions ([H₃O⁺]) and hydroxide ions ([OH⁻]) in distilled water at 25 °C is 1.0 × 10⁻⁷ M.
• Two molecules in a billion dissociate.

Ion-Product Constant

• The ion-product constant (K_w) of water is a measure of the equilibrium between water molecules, hydronium (H₃O⁺) and hydroxide (OH⁻).
• K_w = [H₃O⁺][OH⁻] = 1.0 × 10⁻¹⁴ at 25ºC
• The concentrations of [H₃O⁺] and [OH⁻] in pure water are equal, making it neutral.
• Solutions with greater [H₃O⁺] than [OH⁻] are acidic, and basic solutions have a higher [OH⁻] concentration than [H₃O⁺].

The pH Scale

• pH is a measure of the acidity or basicity of a solution.
• pH = -log[H₃O⁺]
• Acidic solutions have pH values less than 7.
• Neutral solutions have a pH of 7.
• Basic solutions have pH values greater than 7.
• Values for pH can be negative or greater than 14, which depends on temperature.

pH of Some Common Substances

• A chart shows the pH values of common substances.

Measuring pH

• pH can be measured using an electronic meter, which is highly accurate.
• The electronic meter requires calibration with solutions of known pH like distilled water or buffer solutions.
• Acid-base indicators can also approximate pH.
• Universal indicators (paper strips or liquid) are another way to measure pH, offering relatively less precise measurements compared to the electronic meter.

pH for Strong Acids

• Strong monoprotic acids dissociate completely in water.
• Use the initial concentration of the strong acid to calculate [H₃O⁺], which is used to determine pH.

pH for Strong Bases

• Alkali metal hydroxides (MOH) dissociate almost completely in water, producing metal cations (M⁺) and hydroxide anions (OH⁻) in a 1:1 molar ratio.
• The concentration of hydroxide ions ([OH⁻]) can be used to calculate [H₃O⁺] and then pH.

Strong Bases (continued)

• Some alkaline-earth metal hydroxides, while being strong bases, do not completely dissociate; this results in fewer hydroxide ions, impacting their pH to become lower.
• Metal oxides like CaO react with water to form alkaline solutions.