Acids, Bases, and pH

• pH solutions from 0-14 determine the strength of acids and bases.
• Lower pH solutions from 0-6 are strong acids.
• Higher pH solutions from 8-14 are bases.
• Solutions measuring pH 7 are neutral.

Acids

• An acid tastes sour
• Acids will turn litmus paper red
• Acids feel irritating/corrosive
• Acids have a pH less than 7
• Acids are a chemical substance that has a pH lower than seven
• Acids are hydrogen-containing substances capable of donating a proton to another substance.
• Lemon juice, soda, and coffee are examples of acid substances

Bases

• A base tastes bitter or soapy
• Bases turns litmus paper blue
• Bases feel slippery
• Bases have a pH greater than 7
• A base is a chemical substance that has a pH higher than seven
• A base is a molecule or ion able to accept a hydrogen ion from an acid
• Baking soda, ammonia and soap, and bleach are examples of base substances

Neutralization

• A strong acid and a strong base will neutralize to make salt and water.
• HCl + NaOH → NaCl + H2O is an example of this reaction.

Acid and Bases Theory

Arrhenius Theory

• Acids ionize to form hydrogen (+) ions in aqueous solution
• HCL + H2O -> H3O+ + C1- is an example of the reaction
• Bases ionize to form hydroxide ions OH (-) in aqueous solution
• NaOH + H2O -> Na+ + OH- is an example of the reaction

Bronsted Lowry Theory

• Acids: proton (hydrogen ion) DONOR, acids turn to as conjugate base once the proton is lost. A weak acid turns into a strong conjugate base
• H2SO4 +H2O <-> HSO4- + H3O+ is an example of this theory
• Bases: proton (hydroxide ion) DONOR, bases turn to as conjugate acids with the addition of the proton to the base. A weak base becomes a strong conjugate acid

Lewis Theory

• Applies to acid base reactions that don't involve H+ transfer
• Acids: electron ACCEPTOR (electrophiles), metal cations, group 3A atoms, transition metalsk
• NH3 + H* -> NH4 is an example of this reaction
• Bases: electron DONOR (nucleophiles), compounds with lone pairs

Acid strength

• Acid strength refers to the ability to donate hydrogen ions (H+) or protons in a solution
• It depends on the degree of ionization or dissociation in water.

Strong Acids

• Completely dissociate in water, releasing all of their H+ ions.
• Hydrochloric acid (HCl), Sulfuric acid (H2SO4), and Nitric acid (HNO3) are examples of strong acids

Weak Acids

• Partially dissociate in water, releasing only some H+ ions.
• Acetic acid (CH3COOH), Formic acid (HCOOH), and Carbonic acid (H2CO3) are examples of weak acids

Acid Dissociation Equation

• Shows the process in which an acid (HA) donates a proton (H +) to water, forming hydronium ions (H3O+) and its conjugate base (A−).
• The general formula is: HA (aq) + H2O (l) → H3O+ (aq.) + A- (aq.)

Ionization Constant (Ka)

• Indicates the relative strength of the acid or base
• An acid with Ka of 1 x 10-1 is stronger (more ionized) than one with a Ka of 1 x 10-3
• Stronger Acids: Ka is larger and pKa is smaller where a base with a Ka of 1 x 10-7 is weaker (less ionized) than one with a Ka of 1 x 10-9.
• Stronger Base: Ka is smaller and pKa is larger

Acid dissociation constant expression, (Ka)

• Quantitative measure of the strength of an acid in solution.
• Represents the extent to which an acid dissociates into its ions in water.
• Ka = [H+][A-]/[HA]

pKa

• pKa is The negative log of the ionization constant
• An acid with a pKa of 5 (Ka = 1 x 10-5) is weaker (less ionized) than one with a pKa of 3 (Ka = 1 x 10-3)
• A base with a pKa of 9 (Ka = 1 x 10-9 ) is stronger (more ionized) than one with a pKa of 7 (Ka = 1 x 10-7).

Ionization for acids and bases

• Weakly acidic drugs are less ionized in acid media than in alkaline media.
• When the pka of an acidic drug is greater than the pH of the medium in which it exists, it will be more than 50% in its nonionized (molecular) form
• pH is strength of an aqueous solution
• Ka is Strength of acid molecule
• AUTOIONIZATION OF WATER: When proton is transferred from one molecule to one another to produce Hydronium ion[H3O+] and a Hydroxide ion [OH-]

Acid and Base Strength

• (Ka)(Kb) = Kw - Constant
• Kw = 1.0 x 10-14
• Because in pure water [H3O+] = [OH-]
• [H3O+] = (1.0 × 10-14)1/2 = 1.0 × 10-7
• pKw = pKa + pKb
• pKw = - log Kw
• pKw the negative logarithm of water ion constant, Kw
• pka = -log Ka
• pKa: negative logarithm of acid dissociation constant, Ka
• pkb = -log Kb
• pKb → negative logarithm of base dissociation constant

pH Equation

• Measures the strength of an aqueous acid solution
• pH = -log[H+] - ACID
• pOH = -log[H+] - BASIC
• pH + pOH = pKw
• pH + pOH = 14

Calculating pH

• pH = – log [H3O+ ]
• pOH = -log [OH-]
• pH + pOH = 14

Henderson-Hasselbalch Equation

• The relationship between pH, pKa, and the concentrations of the acid and its conjugate base in a buffer solution can be described by the Henderson-Hasselbalch equation.
• pH = pKa + log ([A-]/[HA])
• pH is The measure of acidity of the solution.
• pKa is The negative logarithm of the acid dissociation constant Ka of the weak acid.
• [A-] is the concentration of the conjugate base
• [HA] is the concentration of the weak acid.