MCAT General Chemistry - Acids and Bases

586

cocaine production

treat coca leaves with acid → cocaine: snorted or injected, water soluble

treat coca leaves with base → freebase/crack: smoked and inhaled, water insoluble

Arrhenius acid

dissociate to form an excess of H+ in solution; limited to aqueous acids and bases

contain H at the beginning of their formula

ex. HCl, HNO3, H2SO4

Arrhenius base

dissociate to form an excess of OH− in solution; limited to aqueous acids and bases

contain OH at the end of their formula

ex. NaOH, Ca(OH)2, Fe(OH)3

Brønsted–Lowry acid

species that donates hydrogen ions (H+); not limited to aqueous solutions; also always a Lewis acid

Brønsted–Lowry base

species that accepts hydrogen ions (H+); not limited to aqueous solutions; also always a Lewis base

Lewis acid

lone electron pair acceptor; not always Brønsted–Lowry acid

Lewis base

lone electron pair donor; not always Brønsted–Lowry base

amphoteric

reacts like an acid in a basic environment and like a base in an acidic environment

ex. water, conjugate base of a polyvalent acid, hydroxides of certain metals, amino acids that have a zwitterion intermediate

amphiprotic

can either gain or lose a proton

Acid naming

formed from anions with names that end in –ide = prefix hydro– + anion root + –ic

formed from oxyanions = prefixes and roots retained

• -ite = –ous acid.

• -ate = -ic acid

autoionization

water can react with itself, creating acid and base; reversible

hydronium ion (H3O+)

water with proton

hydroxide ion (OH−)

water sans proton

water dissociation constant, Kw

K_w = [H3O⁺][OH⁻] = 10⁻¹⁴ at 25°C (298 K)

concentrations of the hydrogen ions and hydroxide ions are always equal in pure water at equilibrium

pH scale

concentration scale for acidity; negative logarithm of concentration of protons/hydronium

pOH scale

concentration scale for basicity; negative logarithm concentration of hydroxide

pH-pOH scale

As pH increases, pOH decreases by the same amount

pH < 7 = pOH > 7 = relative excess of protons = acidic

pH > 7 = pOH < 7 = relative excess of hydroxide = basic

approximation of a p scale

if the nonlogarithmic value is written in proper scientific notation, it will be in the form n × 10−m, where n is a number between 1 and 10

p value ≈ m − 0.n

Strong acids and bases

completely dissociate into their component ions in aqueous solutions

HI, HBr, HCl, HClO₄, HNO₃, H₂SO₄

NaOH, KOH, etc.

weak acids and bases

only partially dissociate in aqueous solutions to achieve an equilibrium state

Ka/Kb < 1.0

acid dissociation constant (Ka)

smaller the Ka, the weaker the acid, and consequently, the less it will dissociate

inversely related to Kb

base dissociation constant (Kb)

smaller the Kb, the weaker the base, and consequently, the less it will dissociate

inversely related to Ka

conjugate acid

acid formed when a base gains a proton

conjugate base

base formed when an acid loses a proton

induction

Electronegative elements positioned near an acidic proton increase acid strength by pulling electron density out of the bond holding the acidic proton, weakening proton bonding and facilitating dissociation

neutralization reaction

Acids and bases may react with each other to form a salt and often water

HA (aq) + BOH (aq) → BA (s) + H2O (l)

hydrolysis

salt ions react with water to give back the acid or base

Strong acid + strong base

fully turn to salt and water; neutral solution

HCl + NaOH → NaCl + H2O

Strong acid + weak base

slightly acidic solution; usually does not make water

HCl + NH3 → NH4Cl

Weak acid + strong base

slightly basic solution; usually does not make water

HClO + NaOH → NaClO + H2O

Weak acid + weak base

final pH depends on the relative strengths of the reactants

HClO + NH3 → NH4ClO

Peptide Bond Formation

An acidic carboxyl group reacts with a basic amino group

acid equivalent

equal to one mole of H+/H3O+ ions

base equivalent

equal to one mole of OH− ions

polyvalent (polyprotic)

each mole of the (Brønsted-Lowry) acid or base liberates more than one acid or base equivalent

normality

quantity of acidic or basic capacity

ex. each mole of H3PO4 yields three equivalents of H3O+ → 2 M H3PO4 = 6 N

gram equivalent weight

the mass of a compound that produces one equivalent (one mole of charge)

Titration

procedure used to determine the concentration of a known reactant in a solution; adding small volumes of the titrant to a known volume of a solution of the titrand until completion of the reaction is achieved

titrant

solution of known concentration used to identify the pH of the titrand

titrand

a solution of unknown concentration to which the titrant is added

equivalence point

the number of acid equivalents present in the original solution equals the number of base equivalents added, or vice-versa

only 7 for strong acid-base due to full dissociation

NaVa = NbVb

where Na and Nb are the acid and base normalities and Va and Vb are the volumes of acid and base solutions

Speciation Plot

At any given pH, only two forms of the acid exist in solution; thus, each conjugate is titrated separately

pH meter

measures pH

indicator

weak organic acids or bases that have different colors in their protonated and deprotonated states; leads to a change in the absorption spectrum of the molecule; must always be a weaker acid or base than the acid or base being titrated

endpoint

The point at which the indicator changes to its final color; if done well, volume difference between it and the equivalence point is negligible

titration curve for a polyvalent acid or base

multiple equivalence points

ex. amino acids

half-equivalence point

occurs when half of a given species has been protonated (or deprotonated)

buffer solution

mixture of a weak acid/base and its salt that resist change in pH; generally maintained within 1 pH unit of the pKa value

bicarbonate buffer system

the H2CO3/HCO3− conjugate pair in the plasma component of the blood for maintaining the pH of the blood within a fairly narrow physiological range

Henderson–Hasselbalch equation

used to estimate the pH/pOH of a buffer solution

pH = pKa + log [A-]/[HA]

pOH = pKb + log [B+]/[BOH]

buffering capacity

the ability to which the system can resist changes in pH