Acids and Bases

How does the Brønsted-Lowry definition of acids and bases differ fundamentally from earlier definitions?

Ammonia (NH3) is considered a base even though it does not contain OH- ions in its formula. What mechanism allows it to act as a base?

What property defines a substance as amphiprotic?

A scientist discovers a new compound and finds that its aqueous solution has a pH of 9. Based on this information, how would you classify the compound?

Given the equation for pH, $pH = -log[H_3O^+]$, how does a tenfold increase in the hydronium ion concentration, $[H_3O^+]$, affect the pH of a solution?

At room temperature, the self-ionization of water is described by the equation $Kw = [H_3O^+][OH^-] = 1.0 imes 10^{-14}$. If a solution has a hydronium ion concentration of $1.0 imes 10^{-5}$ M, what is the hydroxide ion concentration, and what is the pH?

How does the behavior of strong acids and bases in aqueous solutions differ from that of weak acids and bases?

For a weak acid, the acid dissociation constant, $K_a$, is defined as $K_a = rac{([H_3O^+][A^-])}{[HA]}$. What does a smaller $K_a$ value indicate about the strength of the acid?

A buffer solution is prepared to maintain a stable pH. What is the primary mechanism by which a buffer resists changes in pH upon the addition of small amounts of acid or base?

The buffer capacity is defined as the amount of acid or base that can be added before the pH begins to change significantly. How is buffer capacity quantitatively defined?

How would the Arrhenius definition classify $CH_3COOH$ in an aqueous solution?

Which statement correctly describes the roles of Johannes Bronsted and Thomas Lowry in defining acids and bases?

In a chemical reaction, water can act as both a proton donor and a proton acceptor. What term describes substances with this dual capability?

What approach did Soren Sorensen introduce in 1909 to quantify acidity and basicity in aqueous solutions more conveniently?

Given $pKw = pH + pOH = 14$ at room temperature, if a solution has a $pH$ of 6.0, what is its $pOH$ and how does this relate to its acidity or basicity?

How does calculating the hydronium ion concentration in a strong acid solution differ from that of a weak acid solution?

For a weak base, the base dissociation constant, $K_b$, is represented as $Kb = rac{([OH^-][B^+])}{[B]}$. What does this constant indicate about the strength of the weak base?

How does adding a strong acid to a buffer solution affect the equilibrium and what components of the buffer system counteract this addition?

A laboratory technician prepares a buffer solution for an experiment but finds that adding only a small amount of acid causes a significant pH change. What is the most likely reason for this?

Why is the concept of buffer capacity important in biological systems, such as blood, and what happens if the buffer capacity is exceeded?