Chem 2 chapter 17 and 18

GERD (gastroesophageal reflux disease)

• If the muscle between the stomach and esophagus does not close properly, acid leaks repeatedly

• It can be managed through lifestyle or dietary changes, acid-reducing medications, or surgical correction of the muscle in severe cases

Acids

• Sour taste

• Dissolve many metals

• Neutralize bases, forming salt and water

• Change blue litmus paper to red

• Strong acids are caustic and can cause severe burns.

• Can be binary (HCl), oxy (H2SO4, HNO3), or carboxylic (CH3COOH)

Bases

• Bitter taste

• Feel slippery

• Neutralize acids, forming salt and water

• Change red litmus paper to blue

• Strong bases are corrosive and can eat away/dissolve out body tissue.

• NaOH, KOH, Ca(OH)2, Mg(OH)2, and NH3 are common examples

Arrhenius Theory

• Acids are substances that produces H+ ions in aqueous solution. The released H+ ions combine with water (H2O) molecules, forming hydronium (H3O+) ions: H+ + H2O → H3O^+

• Bases are substances that produces OH− ions in aqueous solution

• During neutralization, the H+ from the acid combines with OH− from the base and form water (H2O) acid + base → salt + water

Brønsted-Lowry Theory

• An acid is a proton (H+) donor, where a base is a proton (H+) acceptor

• When HCl dissolves in water, it donates H+ to water - HCl is the acid (H+ donor) and water is the base (H+ acceptor) HCl 𝑎𝑞 acid + H2O 𝑙 base → Cl− 𝑎𝑞 + H3O+ 𝑎𝑞

• When NH3 dissolves in water, it accepts H+ from water – NH3 is the base (H+ acceptor) and water is the acid (H+ donor) NH3 𝑎𝑞 base + H2O 𝑙 acid ⇌ NH4+ 𝑎𝑞 + OH− 𝑎𝑞

• All Arrhenius acids and bases are also Brønsted–Lowry acids and bases, but not all Brønsted–Lowry substances fit the Arrhenius definition (e.g., ammonia (?), water).

Strong Acids

- Hydrochloric acid (HCl)

− Hydrobromic acid (HBr)

− Hydroiodic acid (HI)

− Sulfuric acid (H2SO4)

− Nitric acid (HNO3)

− Perchloric acid (HClO4)

Strong Bases

− Sodium hydroxide (NaOH)

− Potassium hydroxide (KOH)

− Barium hydroxide (Ba(OH)2)

− Lithium hydroxide (LiOH)

Weak Acids

− Hydrofluoric acid (HF)

− Acetic acid (HC2H3O2)

− Formic acid HCHO2

− Sulfurous acid (H2SO3)

− Carbonic acid (H2CO3)

− Phosphoric acid (H3PO4)

Weak Bases

− Ammonia (NH3)

− Aluminum hydroxide (Al(OH)3)

− Zinc hydroxide (Zn(OH)2)

Strong bases ionize completely in water. The water acts as an acid and accepts H+ from it.

In weak bases, only a small proportion them dissociates and accepts H+. The majority remains as molecules, undissociated.

Cations of strong bases and anions of strong acids do not ionize water; hence, they don’t affect solution pH. They result in neutral solution.

Cations of weak bases and anions of weak acids will ionize water and affect the pH of solution.

Every anion from an acid accepts a 𝐻+. It can be thought of as the conjugate base of an acid.

A− 𝑎𝑞 + H2O 𝑙 ⇌ HA 𝑎𝑞 + OH− 𝑎𝑞

The stronger the acid, the weaker the conjugate base (anion from salt).

Cations from weaker bases react with water, donating H+ ion. Hence, they are conjugate acids.

• Example: NH𝟒+(a conjugate acid of NH3). It is weakly acidic.

Classifying salt solutions

• strong acid and strong base = neutral solution

• cations that are strong bases and anions that are weak acids = basic solutions

• cations that are weak bases and anions that are strong acids = acidic solutions

• cations from weak bases and anions from weak acids may be acidic or basic

Acids consisting of two more ionizable hydrogens are called polyprotic acids.

Second and third ionizations in polyprotic acids are usually too weak to significantly affect pH. Thus, for most pH problems, only the first ionization is considered.

Binary Acids (H–X)

Strength increases with weaker H–X bonds and higher electronegativity of X (bond strength is more influential than EN)

• Trends: Increases across a period (↑ EN), increases down a group (↓ bond strength)

Oxyacids (H–O–Y)

• Strength increases with more electronegative Y and more attached oxygens

• Trends: Increases across a period (↑ EN, ↑ oxidation state), decreases down a group.

Lewis Acids and Bases

Lewis theory defines acids and bases in terms of transfer of an electron pair.

• Acids are electron pair acceptors; bases are electron pair donors.

Lewis theory does not require the release or transfer of H+ or OH−

Lewis Acids

• Are compounds or ions with empty orbitals that can accept electron pairs.

• Examples: H⁺ (empty 1s orbital), BF₃ (boron has empty 2p orbital), Al³⁺ (small, highly charged metal cation).

• They accept electron pairs and accommodate on the empty orbitals

Lewis Bases

• Are species with lone electron pairs available for donation.

• Examples: OH⁻, NH₂⁻, H2O, O2−, and N3− have lone pairs that they can donate.

Lewis Acid–Base Reactions

• The base donates a pair of electrons to the acid. The product that forms is called an adduct, to mean a combined product.

• For example, the boron in boron trichloride (Lewis acid) accepts a lone pair from the oxygen atom in dimethyl ether (Lewis base).

• Arrhenius and Brønsted-Lowry acid–base reactions are also Lewis reactions.

Acidosis

Acidosis compromises the ability of hemoglobin to carry O2, causing it to release O2 where it is not needed and bind excess H+ instead.

Buffer Solutions

• A buffer is a solution that maintains a nearly constant pH of a given system by neutralizing added acids or bases.

• It contains either a weak acid and its conjugate base, or a weak base and its conjugate acid.

• The weak acid neutralizes added bases, while its conjugate base neutralizes added acids.

Acidic Buffers

• Acidic buffers are made by adding certain amount of a weak acid and its conjugate base.

• If a strong base is added, the acetic acid neutralizes it.

• If a strong acid is added, the acetate (conjugate base) neutralizes it.

common ion effect

The degree of ionization is very small. In the case of acidic buffers, it becomes even smaller due to the presence of the conjugate base of the acid, such as the acetate ion.

This is known as the common ion effect because the solution contains two substances (𝐶𝐻3𝐶𝑂𝑂𝐻 and 𝐶𝐻3𝐶𝑂𝑂𝑁𝑎) that share a common ion (𝐶𝐻3𝐶𝑂𝑂−).

Using ICE Analysis or HHE?

HHE is the easiest and shortest way, but it is only good enough when the “x is small” approximation is applicable.

To use HHE, the following conditions should be met.

– The initial concentrations of acid and salt are not very dilute.

– The 𝐾a is fairly small.

For most problems, the initial acid and salt concentrations should be over 100 to 1000 times larger than the value of 𝑲𝒂.

Buffer Effectiveness

The effectiveness of a buffer depends on two factors: (1) the relative amounts of buffer acid and base, and (2) the absolute concentrations of buffer acid and base. A buffer is most effective:

(1) When the [acid] and [conjugate base] are equal. The difference in the [base] [acid] ratio should at least be between 0.1 and 10.

(2) When the [acid] and [conjugate base] are large.

Buffer Capacity

Buffer capacity is the amount of acid or base that can be added to a buffer without causing a large change in pH.

• It increases:

– with increasing absolute concentrations of buffer components

– as the relative concentrations of the components become more nearly equal