CHEM 120B Acid / Base Equilibrium

Arrhenius acid

A compound that produces H+ ions in aqueous solutions.

Arrhenius base

A compound that produces OH– ions in aqueous solutions.

Hydronium ion

Formed when an H+ ion (proton) reacts with a water molecule, often represented as H3O+, it's a simplified representation of more complex ions.

Brønsted–Lowry acid

A H+ ion donor.

Brønsted–Lowry base

A H+ ion acceptor.

Acid ionization constant (Ka)

The equilibrium constant for the reaction of an acid with water.

Strength of an acid (Ka)

The larger the Ka value, the stronger the acid.

Behavior of strong acids in water

Strong acids are completely ionized in water.

Behavior of weak acids in water

Weak acids are partially ionized in water.

Characteristic functional group of carboxylic acids

Carboxylic acids have a COOH group.

Effect of oxygen atoms on oxoacids strength

The more oxygen atoms, the stronger the acid due to stabilization of the conjugate base.

Electronegativity and acid strength relation

Implied relationship, further details likely not covered in provided sources.

Examples of strong bases

Typically hydroxides of Group I and II metal ions.

Ionization of weak bases in water

Weak bases partially ionize in water.

Base ionization constant (Kb)

Measures the strength of a weak base in water.

Amines

Derivatives of ammonia (NH3) that contain a basic nitrogen atom.

Conjugate acid

Formed when a base accepts a proton (H+).

Conjugate base

Formed when an acid donates a proton (H+).

Conjugate acid-base pair example

HF(aq) (acid) and F–(aq) (conjugate base).

Leveling effect

In water, H3O+ is the strongest H+ donor, leading strong acids to share the same strength.

Autoionization of water

Water acts as both an acid and base undergoing self-ionization.

Kw value at 25°C

At 25°C, the ion product of water is Kw = [H3O+][OH–] = 1.0 × 10–14.

Definition of pH

A measure of the acidity or basicity of a solution defined as pH = −log[H3O+].

Neutral pH at 25°C

A neutral solution has pH = 7.

pH relation to acidity and basicity

pH < 7 is acidic, pH > 7 is basic.

Determining [H3O+] from pH

[H3O+] = 10−pH.

Percent ionization for a weak acid formula

Percent Ionization = ([H3O+]equilibrium / [HA]initial) × 100%.

Effect of acid concentration on percent ionization

Percent ionization decreases as [acid] increases.

[H3O+] in dilute weak acid solution

Larger [H3O+] in dilute solution compared to initial concentration due to larger percent ionization.

Polyprotic acids

Acids that have more than one ionizable H atom per molecule.

Ka values comparison in polyprotic acids

Generally Ka1 > Ka2 > Ka3.

First ionization step effect on pH in polyprotic acids

Only the first H+ significantly affects pH.

Anion as conjugate base concept

Every anion can be thought of as the conjugate base of an acid.

Strength of acid and conjugate base relation

Stronger acid has weaker conjugate base.

pH behavior of anions from strong acids

Conjugate bases of strong acids are pH neutral.

pH behavior of anions from weak acids

Conjugate bases of weak acids are basic.

Cations as acids and bases

Some cations can be thought of as the conjugate acid of a weak base.

Strength of base and conjugate acid relation

Stronger base has weaker conjugate acid.

pH behavior of cation counterions of strong bases

Cations from strong bases are pH neutral.

pH behavior of cations as conjugate acids

Conjugate acids of weak bases are acidic.

pH behavior of alkali and alkaline earth metal cations

These cations are pH neutral.

Ka and Kb relationship for conjugate pairs

Ka × Kb = Kw.

Ocean acidification definition

Decrease in ocean pH due to increased atmospheric CO2 absorption.

Effects of ocean acidification on marine organisms

Reduces CO3 2– needed for CaCO3 exoskeletons.

Buffer solutions definition

Solutions that resist changes in pH when acid or base is added.

Components of a buffer solution

Must contain weak acid and its conjugate base or weak base and its conjugate acid.

Acidic buffer neutralizing added base

Weak acid component neutralizes added strong base.

Acidic buffer neutralizing added acid

Conjugate base component neutralizes added strong acid.

Common ion effect

Addition of a common ion shifts equilibria, affecting pH.

Henderson-Hasselbalch equation for acidic buffer

pH = pKa + log ([base]/[acid]).

Applicability of Henderson-Hasselbalch equation

Applicable when initial concentrations are not very dilute.

Calculating pH after adding strong acid/base to buffer

Requires stoichiometry and equilibrium calculations.

Buffer range definition

pH range where buffer is effective.

Buffer capacity definition

Amount of acid/base a buffer can neutralize while maintaining pH.

Effect of buffer component concentration on pH resistance

Greater initial concentrations lead to less pH change.

pKa definition

pKa = -log[Ka].

pKb definition

pKb = -log[Kb].

Relation between pKa and pKb

pKa + pKb = 14.

pH at midpoint of weak acid-strong base titration

At midpoint, pH = pKa of weak acid.

Titration definition

Process where unknown concentration is added to known concentration.

Equivalence point in titration

Where moles of H3O+ = moles of OH–.

Indicator in titration definition

Chemical that changes color with pH for identifying equivalence point.

Titration endpoint definition

Point of pH change where indicator changes color.

Titration curve description

Plot of pH versus amount of titrant added.

pH at equivalence point of strong acid-strong base titration

Results in a neutral solution with pH = 7.

pH at equivalence in weak acid-strong base titration

Results in a basic solution with pH > 7.

Half-neutralization in weak acid-strong base titration

[HA] = [A–] and pH = pKa.

Equivalence points in diprotic acid titration

If Ka1 >> Ka2, there will be two distinguishable equivalence points.

Lewis acid definition

Substance that accepts a lone pair of electrons.

Lewis base definition

Substance that donates a lone pair of electrons.

Coordinate bond definition

Covalent bond formed when one atom donates a pair of electrons.

Ligand definition

A Lewis base bonded to a central metal ion of a complex ion.

Complex ion definition

Ionic species consisting of metal ion bonded to one or more Lewis bases.

Formation constant (Kf) definition

Equilibrium constant describing the formation of metal complexes.

Solubility definition

Amount of solute that dissolves, expressed in g/L.

Molar solubility (S) definition

Amount of dissolved solute expressed as mol/L.

Solubility product constant (Ksp) definition

Equilibrium constant for the formation of saturated solutions.

Calculating molar solubility (S) from Ksp

Using the equilibrium expression for dissolution of slightly soluble salt.

Common ion effect on solubility

Presence of a common ion shifts solubility equilibrium left, decreasing solubility.

pH effect on solubility of basic anion salts

Solubility increases as pH decreases in acidic solutions.

pH effect on solubility of neutral anion salts

Solubility is not significantly affected by pH.

Q in solubility context definition

Ion product using current concentrations of the ions.

Using Q and Ksp to predict precipitate formation

Q > Ksp indicates precipitate formation; Q < Ksp means no precipitate.

Buffer system changes related to ocean acidification

Increased CO2 forms H2CO3, which dissociates to HCO3– and H3O+, decreasing ocean pH.

Buffer component reacting with added H+ ions

Conjugate base reacts with H+ to form more weak acid.

Buffer component reacting with added OH– ions

Weak acid reacts with OH– to form more conjugate base.

Relationship among Kw, Ka, Kb

Kw = Ka × Kb for a conjugate pair.

pOH definition

pOH = −log[OH–].

relationship between pH and pOH at 25°C

pH + pOH = 14.

Calculating [OH–] from pOH

[OH–] = 10−pOH.

pH of equivalence point in titration

Depends on the pH of the formed salt solution.

Predominant species at equivalence point in weak acid-strong base titration

Conjugate base anion of the weak acid.

Blood buffer solution system

Consists of a mixture of H2CO3 and HCO3–.

Choosing an acid for buffer preparation

Choose acid with pKa closest to desired buffer pH.

Relationship between Ka and pKa inputs

pKa = -log(Ka) and Ka = 10-pKa.

Relationship between Kb and pKb

pKb = -log(Kb) and Kb = 10-pKb.

Ka for the first ionization of carbonic acid

Ka1 = 10^-6.37, from the pKa1 value of 6.37.

pKa and buffer range relationship

Effective buffer range is considered to be pKa ± 1.

Henderson-Hasselbalch equation components

[base] refers to conjugate base concentration, [acid] to weak acid concentration.

Significance of Ka1 >> Ka2 in polyprotic acid titration

Distinction of two equivalence points in titration curve.