Acids and bases
60 flashcards covering Acids and bases for the MCAT Chem / Phys / Psych / Soc section.
Acids and bases are fundamental concepts in chemistry that describe substances based on their behavior in water. Acids are compounds that donate hydrogen ions (protons), leading to solutions with a low pH, like vinegar or stomach acid, which can corrode materials and play roles in digestion. Bases, on the other hand, accept hydrogen ions or donate hydroxide ions, resulting in higher pH solutions, such as soap or baking soda, which often neutralize acids and are essential in cleaning and industrial processes.
On the MCAT, acids and bases frequently appear in the Chem/Phys section through multiple-choice questions testing equilibrium, pH calculations, titrations, and buffer systems. Common traps include confusing strong acids with weak ones or misapplying concepts in biological contexts, like blood pH regulation. Focus on mastering definitions (Arrhenius and Bronsted-Lowry), acid-base reactions, and their applications in medicine to avoid errors on passage-based or standalone questions.
Practice sketching titration curves for quick mastery.
Terms (60)
- 01
Arrhenius acid
An Arrhenius acid is a substance that increases the concentration of hydrogen ions (H+) when dissolved in water.
- 02
Arrhenius base
An Arrhenius base is a substance that increases the concentration of hydroxide ions (OH-) when dissolved in water.
- 03
Bronsted-Lowry acid
A Bronsted-Lowry acid is a proton (H+ ion) donor in a chemical reaction.
- 04
Bronsted-Lowry base
A Bronsted-Lowry base is a proton (H+ ion) acceptor in a chemical reaction.
- 05
Lewis acid
A Lewis acid is an electron pair acceptor in a chemical reaction.
- 06
Lewis base
A Lewis base is an electron pair donor in a chemical reaction.
- 07
pH scale
The pH scale measures the acidity or basicity of a solution, defined as the negative logarithm of the hydrogen ion concentration, ranging from 0 to 14, with 7 being neutral.
- 08
pOH
pOH is the negative logarithm of the hydroxide ion concentration in a solution, and it relates to pH by the equation pH + pOH = 14 at 25°C.
- 09
Strong acid
A strong acid completely dissociates into its ions in aqueous solution, such as hydrochloric acid (HCl).
- 10
Weak acid
A weak acid partially dissociates in aqueous solution, establishing an equilibrium, like acetic acid (CH3COOH).
- 11
Strong base
A strong base completely dissociates in aqueous solution to produce hydroxide ions, such as sodium hydroxide (NaOH).
- 12
Weak base
A weak base partially reacts with water to produce hydroxide ions, like ammonia (NH3).
- 13
Ka
Ka is the acid dissociation constant, a measure of the strength of a weak acid, calculated as the equilibrium constant for the dissociation reaction.
- 14
Kb
Kb is the base dissociation constant, indicating the strength of a weak base by measuring the extent of its reaction with water.
- 15
pKa
pKa is the negative logarithm of the acid dissociation constant (Ka), used to compare the strengths of acids, with lower values indicating stronger acids.
- 16
pKb
pKb is the negative logarithm of the base dissociation constant (Kb), where a lower pKb value signifies a stronger base.
- 17
Kw
Kw is the ion product constant for water, equal to 1.0 × 10^-14 at 25°C, representing the product of hydrogen ion and hydroxide ion concentrations in pure water.
- 18
Conjugate acid
A conjugate acid is the species formed when a base accepts a proton, differing from the base by one hydrogen ion.
- 19
Conjugate base
A conjugate base is the species formed when an acid donates a proton, differing from the acid by one hydrogen ion.
- 20
Acid-base reaction
An acid-base reaction involves the transfer of a proton from an acid to a base, often resulting in the formation of a conjugate acid and conjugate base.
- 21
Neutralization reaction
A neutralization reaction occurs when an acid and a base react to form water and a salt, typically reaching a pH of 7 for strong acid-strong base pairs.
- 22
Buffer solution
A buffer solution resists pH changes when small amounts of acid or base are added, consisting of a weak acid and its conjugate base or a weak base and its conjugate acid.
- 23
Henderson-Hasselbalch equation
The Henderson-Hasselbalch equation calculates the pH of a buffer solution as pH = pKa + log([conjugate base]/[weak acid]).
- 24
Titration
Titration is a technique to determine the concentration of an unknown solution by adding a solution of known concentration until the reaction reaches the equivalence point.
- 25
Equivalence point
The equivalence point in a titration is when the moles of acid equal the moles of base, marked by a sharp change in pH.
- 26
pH at equivalence point
For a strong acid-strong base titration, the pH at the equivalence point is 7, but for weak acid-strong base, it is greater than 7.
- 27
Acid-base indicator
An acid-base indicator is a weak organic acid that changes color over a specific pH range, used to detect the endpoint of a titration.
- 28
Common ion effect
The common ion effect suppresses the dissociation of a weak acid or base when a salt containing a common ion is added, shifting the equilibrium.
- 29
Le Chatelier's principle in acids
Le Chatelier's principle predicts that adding acid or base to an equilibrium system will shift the reaction to counteract the change, such as in buffer solutions.
- 30
Hydrolysis of salts
Hydrolysis of salts occurs when the ions of a salt react with water, potentially making the solution acidic, basic, or neutral depending on the parent acid and base.
- 31
Acidic salt
An acidic salt is formed from a strong acid and a weak base, resulting in a solution with pH less than 7 due to hydrolysis.
- 32
Basic salt
A basic salt is formed from a weak acid and a strong base, leading to a solution with pH greater than 7 because of ion hydrolysis.
- 33
Amphoteric substance
An amphoteric substance can act as both an acid and a base, such as water or amino acids, depending on the reaction conditions.
- 34
Autoionization of water
Autoionization of water is the process where water molecules produce equal amounts of H+ and OH- ions, governed by the equilibrium constant Kw.
- 35
pH of salt solution
The pH of a salt solution depends on the nature of the ions; for example, salts of strong acids and strong bases are neutral, while others may be acidic or basic.
- 36
Calculating pKa
pKa can be calculated from the pH and the ratio of conjugate base to acid concentrations using the Henderson-Hasselbalch equation.
- 37
Acidity of carboxylic acids
Carboxylic acids are relatively strong organic acids due to the resonance stabilization of their conjugate bases.
- 38
Resonance effects on acidity
Resonance effects increase acidity by stabilizing the negative charge on the conjugate base, as seen in phenols compared to alcohols.
- 39
Amino acid as acid
An amino acid can act as an acid by donating a proton from its carboxylic acid group.
- 40
Isoelectric point
The isoelectric point is the pH at which an amino acid or protein has no net charge, calculated from the average of relevant pKa values.
- 41
Zwitterion
A zwitterion is a molecule with both positive and negative charges, such as an amino acid at its isoelectric point.
- 42
Buffer range
The buffer range is the pH interval over which a buffer is effective, typically within one pH unit of its pKa.
- 43
Polyprotic acid
A polyprotic acid can donate more than one proton per molecule, like sulfuric acid, which has two dissociation steps.
- 44
Diprotic acid
A diprotic acid, such as carbonic acid, donates two protons and has two Ka values corresponding to each dissociation.
- 45
Successive Ka values
Successive Ka values for polyprotic acids decrease, as each proton is removed from a increasingly negatively charged species.
- 46
Solubility and pH
The solubility of some salts, like metal hydroxides, increases in acidic conditions due to the reaction of the anion with H+ ions.
- 47
HCl as a strong acid
Hydrochloric acid (HCl) is a strong acid that fully dissociates in water, making it a common laboratory reagent.
- 48
NaOH as a strong base
Sodium hydroxide (NaOH) is a strong base that completely dissociates in water to provide hydroxide ions.
- 49
Weak acid equilibrium
For a weak acid, the equilibrium expression is Ka = [H+][A-]/[HA], where the dissociation is incomplete.
- 50
Base hydrolysis
Base hydrolysis is the reaction of a weak base with water to produce OH- ions, as in the case of ammonia forming NH4+ and OH-.
- 51
pH of weak base
The pH of a weak base solution is calculated using Kb and the base hydrolysis equilibrium, resulting in a pH greater than 7 but less than that of a strong base.
- 52
Dilution of acids
Dilution of acids decreases their concentration, but for strong acids, it does not affect their degree of dissociation.
- 53
Heat of neutralization
The heat of neutralization is the energy released when an acid and base react, approximately constant for strong acid-strong base reactions at 57 kJ/mol.
- 54
Entropy in acid-base reactions
Acid-base reactions often increase entropy due to the formation of more disordered products, favoring spontaneity.
- 55
Gibbs free energy for acids
For acid dissociation, Gibbs free energy relates to Ka via ΔG = -RT ln(Ka), indicating spontaneity for strong acids.
- 56
Lewis acid in complexes
In coordination compounds, a Lewis acid is a metal ion that accepts electron pairs from ligands, which act as Lewis bases.
- 57
Bronsted-Lowry vs. Lewis acids
Bronsted-Lowry acids donate protons, while Lewis acids accept electron pairs, so all Bronsted-Lowry acids are Lewis acids but not vice versa.
- 58
pKa trends in organic acids
pKa trends show that acidity increases with electron-withdrawing groups and decreases with electron-donating groups on the molecule.
- 59
Acid strength and electronegativity
Acid strength increases with the electronegativity of the atom bonded to the acidic hydrogen, as it stabilizes the conjugate base.
- 60
Base strength and atom size
Base strength decreases with increasing size of the atom bearing the negative charge, due to poorer stabilization of the charge.