Bronsted Acids and Bases Quiz

The conjugate base of water (H2O) in the Brønsted-Lowry acid-base theory is hydroxide ion (OH–).

The conjugate base of hydrofluoric acid (HF) is fluoride ion (F–).

The conjugate acid of the oxide ion (O2−) is hydroxide ion (OH–).

The conjugate base of dimethylammonium ion ((CH3)2NH2+) is dimethylamine ((CH3)2NH).

In the reaction NH3 + H2O → NH4+ + OH–, water acts as a Bronsted acid.

The conjugate base of nitric acid (HNO3) is nitrate ion (NO3–).

A species that can act both as a Brønsted-Lowry acid and a base.

In the reaction NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH– (aq), H2O acts as an acid (donating a proton) while NH3 acts as a base (receiving a proton).

Reactions involving the formation of coordinate covalent bonds, where one reactant acts as the electron-pair acceptor (Lewis acid) and the other as the electron-pair donor (Lewis base).

Some chemical species, like HCl, cannot function simultaneously as both a Brønsted-Lowry acid and a base.

A species that is not likely to donate an electron pair; for example, BF₃.

A Brønsted acid that does not readily donate a proton.

An example of a law applicable in dilution of a weak electrolyte solution, typically for acids and bases, not neutral salts, like NaCl.

pH is the negative logarithm of the hydrogen ion concentration. To calculate for pH = -log([H+])

A measure of the concentration of hydrogen ions (H+) in a solution, expressed on a logarithmic scale.

Hydroxide ion concentration in a solution.

A measure of the concentration of hydroxide ions (OH⁻) in a solution, expressed on a logarithmic scale.

Determines the hydrogen ion concentration from pH, or the pH from H+ concentration.

The equilibrium constant for the autoionization of water.

pH greater than 7

pH less than 7

Negative logarithm of Kw (the ionic product of water).

A solution with a pH = 7 at 25°C.

Acids release H+ in water and bases release OH- in water.

A scale used to measure the acidity or alkalinity of a solution.

pH + pOH = pKw

The concentration of hydroxide ions in a 0.01M ammonia solution that is 5% ionized is 0.0005 M

The concentration of hydrogen ions in a 5  10⁻³ M H₂CO₃ solution with 10% dissociation is 10⁻³M

The pH of a 2.5 × 10⁻¹ molar HCN solution with a Ka of 4 × 10⁻¹⁰ is approximately 4.7.

The molarity of nitrous acid at which its pH reaches 2, with a Ka of 4.5 × 10⁻⁴ is approximately 0.3333 M.

The degree of dissociation of a weak acid (HA) with a 0.1 M solution and a pH of 3 is approximately 10%.

The pH of a 0.01 M acetic acid solution which is 1% ionised is approximately 3.

The degree of dissociation (α) for a 10⁻² M NH₄OH solution with a hydroxide ion concentration of 10⁻³ is 10%.

The pH of a 10⁻³ M CH₃COOH solution with a Ka of 10⁻⁵ is approximately 4.

The Ka value for a 10⁻² M HCN solution with a pOH of 10 is 10⁻⁴.

The Kb value for the fluoride ion is 2.9×10⁻¹¹ when the Ka value for hydrofluoric acid is 3.5×10⁻⁴

The ionization constant of X⁻ can be determined from the ionization constant of the acid (Ka), according to relationship between Ka and Kb with Kw.

A solution that resists changes in pH when small amounts of acid or base are added.

The buffer system in blood that maintains a nearly constant pH around 7.4.

A buffer solution typically consists of a weak acid and its conjugate base (or a weak base and its conjugate acid).

The pH of a buffer solution is determined using the Henderson-Hasselbalch equation.

The ratio of the concentration of the conjugate base to the concentration of the weak acid determines the pH of a buffer solution.

The pH of a buffer solution is approximately equal to the pKa of the weak acid component.

Strong bases react with the weak acid in the buffer, shifting the equilibrium and changing the buffer's pH.

Changing the ratio of weak acid to its salt alters the buffer's pH.

The amount of a weak acid that's ionized in a buffer solution

pOH can be calculated when a strong base is added to a weak acid buffer.

Adding small amounts of strong acids or bases (HCl or NaOH) to a buffer solution will cause little to no change in pH.

A mathematical equation that allows for the calculation of the pH of buffer solutions.

Solubility product constant for silver chloride, indicating its solubility in water.

The pH of pure water is 7.

Formation of solid AgCl when mixing solutions.

A salt derived from a strong base and a weak acid.

Solubility product constant for magnesium hydroxide.

The pH of a 1.0 M ammonium formate solution is approximately 6.5.

The pH at which Mg2+ ions begin to precipitate as Mg(OH)2 from a solution.

A salt that doesn't react with water (neutral salt).

pH = (1/2) * pKw + (1/2) * pKa - (1/2) * pKb

Solubility product constant for silver chromate.

Relative order of solubility for salts.

The pH of a 0.5M NaCN solution is approximately 11.5.

Calculating solubility of two salts in the same solution.

If salts M2X, QY2, and PZ2 have equal solubilities, Ksp(M2X) = Ksp(QY2) = Ksp(PZ2).

Determining Fluoride ion concentration in a solution saturated by two specific compounds.

The solubility product of silver chromate (Ag2CrO4) is 1.69 x 10^-12 using an Ag+ concentration of 1.5 x 10^-4 mol L-1.

The Ksp for AgCl, given a solubility of 1.43 x 10^-4g/100mL at 25°C., is approximately 1 x 10^-10

Finding equilibrium concentration of CO32- ion when two carbonate salts are present.

The constant for deprotonation of dimethylamine.

If salts M2X, QY2, and PZ3 have the same solubilities, Ksp (M2X) = Ksp(QY2) < Ksp (PZ3).

The percentage of dimethylamine molecules that are ionized in a solution.

Increasing the pH from 3 to 6 reduces the H+ ion concentration by 1000 times

The pOH of a beer is 10.0.

Constant for ionization of phenol.

Concentration of phenolate ion in a 0.05M phenol solution.

pH of a solution that resists drastic changes in pH.