Acetic Acid and Protogenic Solvents Quiz

A substance capable of accepting a proton (H+ ion). Examples include acetone, ether, and liquid ammonia.

A substance capable of donating a proton (H+ ion). Examples include formic acid, acetic acid, sulfuric acid, liquid HCl, and liquid HF.

A substance which accepts a proton (H+ ion) to form a new species.

A substance which donates a proton (H+ ion) to form a new species.

The ability to donate or accept a proton (H+ ion).

A scale used to measure the acidity or alkalinity of a solution, with values ranging from 0 to 14.

A solution with a pH value between 0 and 7 is considered acidic.

A solution with a pH value between 7 and 14 is considered alkaline or basic.

The pH of a solution can be considered in terms of a numeric scale having values from 0 to 14, which expresses in a quantitative way the degree of acidity (7 to 0) and alkalinity (7-14).

The H ion concentration of a solution varies from approximately 1 in a 1 M solution of a strong acid to about 1 × 10-14 in a 1 M solution of a strong base.

The equilibrium constant for the dissociation of a weak acid in water. It represents the ratio of the product of the concentrations of the dissociated ions (H+ and conjugate base) to the concentration of the undissociated acid.

A substance capable of accepting a proton (H+ ion) from another molecule. This leads to the formation of a new species.

A substance capable of donating a proton (H+ ion) to another molecule. This leads to the formation of a new species.

The process of transferring a proton (H+ ion) from one molecule to another. This process is key in acid-base reactions.

A solvent that can both donate and accept protons (H+ ions). Examples include water and alcohols.

Acids and bases that donate or accept only one proton (H+) in a reaction.

One proton (H+) is released during the reaction.

One proton (H+) is accepted during the reaction.

They form a new species after donating or accepting a proton.

A single proton (H+) is transferred between them.

In an equilibrium reaction, if the value of 'x' in the equilibrium constant expression (Kc) is significantly smaller than the initial concentration of reactants (c), then 'x' can be disregarded in the calculation. This simplified approach is valid when the change in concentration due to the reaction is negligible compared to the original concentration. By assuming that the change in concentration is negligible, we can simplify the calculation and obtain an approximation.

The equilibrium constant (Kc) is a numerical value that describes the relative amounts of reactants and products at equilibrium in a reversible reaction. It is a measure of the extent to which a reaction proceeds to completion. Kc is calculated by dividing the product of the concentrations of products raised to their stoichiometric coefficients by the product of the concentrations of reactants raised to their stoichiometric coefficients.

A solution's molarity is a measure of its concentration, defined as the number of moles of solute present in one liter of solution. A 0.1 M solution contains 0.1 moles of solute in one liter of solution.

The degree to which a weak acid dissociates in solution is represented by its ionization constant (Ka). It reflects the equilibrium constant for the acid's dissociation reaction.

Acetate ions (CH3COO-) and hydronium ions (H3O+) are formed when acetic acid (CH3COOH) dissolves in water. Acetic acid is a weak acid that only partially ionizes in solution.

In solutions containing two weak acids, the concentrations of the conjugate bases (Cb1 and Cb2) are assumed to be zero. Additionally, the concentrations of the acids (Ca1 and Ca2) are significantly greater than their respective dissociation constants (K1 and K2). This simplifies the calculation of pH and allows for the direct use of the simplified equation.

Solutions containing a salt of a weak acid and a weak base, like ammonium acetate, dissociate almost completely in water. The cation (e.g., NH4+) acts as an acid (HB1), and the anion (e.g., Ac-) acts as a base (B2-) in the equilibrium.

The pH of a solution containing a salt of a weak acid and a weak base is determined by the relative strengths of the acid and base components of the salt.

The equilibrium constant for the hydrolysis of a salt of a weak acid and a weak base is given by the product of the acid dissociation constant (Ka) and the base dissociation constant (Kb).

The pH of a solution containing a salt of a weak acid and a weak base can be calculated using the Henderson-Hasselbalch equation, which relates the pH to the pKa of the acid component and the ratio of the concentrations of the conjugate acid and base.