Question

Rank the following $0.10\mathrm{M}$ solutions in order of increasing $\mathrm{pH}$. a. HI, HF, NaF, NaI b. ${\mathrm{NH}}_4\mathrm{Br},\mathrm{HBr},\mathrm{KBr},{\mathrm{NH}}_3$ c. ${\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{NH}}_3{\mathrm{NO}}_3,{\mathrm{NaNO}}_3,\mathrm{NaOH},{\mathrm{HOC}}_6{\mathrm{H}}_5,{\mathrm{KOC}}_6{\mathrm{H}}_5,{\mathrm{C}}_6{\mathrm{H}}_5{\mathrm{NH}}_2$, ${\mathrm{HNO}}_3$

Short answer

In summary, the solutions can be ranked based on their increasing pH as follows: a. HI, HF, NaF, NaI b. HBr, NH4Br, KBr, NH3 c. HNO3, HOC6H5, C6H5NH3NO3, NaNO3, C6H5NH2, NaOH, KOC6H5

Step-by-step solution

Set a: Identify the nature of each solution

HI: Hydroiodic acid, a strong acid HF: Hydrofluoric acid, a weak acid NaF: Sodium fluoride, forms a weak base (F-) in solution NaI: Sodium iodide, forms a neutral solution (I- is not a base)

Set a: Rank the solutions in increasing pH

Most acidic: HI (strong acid) < HF (weak acid) < NaF (weak base) < NaI (neutral) Order: HI, HF, NaF, NaI

Set b: Identify the nature of each solution

NH4Br: Ammonium bromide, forms a weak acid (NH4+) in solution HBr: Hydrobromic acid, a strong acid KBr: Potassium bromide, forms a neutral solution (Br- is not a base) NH3: Ammonia, a weak base

Set b: Rank the solutions in increasing pH

Most acidic: HBr (strong acid) < NH4Br (weak acid) < KBr (neutral) < NH3 (weak base) Order: HBr, NH4Br, KBr, NH3

Set c: Identify the nature of each solution

C6H5NH3NO3: Benzylammonium nitrate, forms a weak acid (C6H5NH3+) in solution NaNO3: Sodium nitrate, forms a neutral solution (NO3- is not a base) NaOH: Sodium hydroxide, a strong base HOC6H5: Phenol, a weak acid KOC6H5: Potassium phenoxide, a strong base C6H5NH2: Aniline, a weak base HNO3: Nitric acid, a strong acid

Set c: Rank the solutions in increasing pH

Most acidic: HNO3 (strong acid) < HOC6H5 (weak acid) < C6H5NH3NO3 (weak acid) < NaNO3 (neutral) < C6H5NH2 (weak base) < NaOH (strong base) < KOC6H5 (strong base) Order: HNO3, HOC6H5, C6H5NH3NO3, NaNO3, C6H5NH2, NaOH, KOC6H5

Key concepts

Acid-Base Equilibrium

Understanding the concept of acid-base equilibrium is crucial for ranking the pH of solutions, as it involves the balance between the concentrations of acids and bases in a solution. An acid donates protons (H+) in an aqueous solution, while a base accepts protons. The strength of an acid or base is determined by its tendency to donate or accept protons, respectively.

For instance, when we see strong acids like hydroiodic acid (HI) and hydrobromic acid (HBr), they completely disassociate in water, releasing a large number of hydrogen ions which significantly lowers the pH. Conversely, weak acids like hydrofluoric acid (HF) and phenol (HOC₆H₅) do not dissociate completely; hence, they contribute less to the hydrogen ion concentration, resulting in a higher pH than strong acids.

On the base side of equilibrium, strong bases like sodium hydroxide (NaOH) disassociate completely to produce a large concentration of hydroxide ions (OH-), increasing the pH. Weak bases such as ammonia (NH₃) only partially accept protons, leading to less increase in pH. The principles of acid-base equilibrium allow us to predict the behavior of substances in solution and to rank their pH accordingly.

Strong and Weak Acids and Bases

The distinction between strong and weak acids and bases is paramount for predicting the pH levels of different solutions. Strong acids, such as nitric acid (HNO₃), and strong bases like potassium phenoxide (KOC₆H₅) ionize completely in water, which means they donate all their available hydrogen ions or accept hydrogen ions, respectively, without reserve. This complete ionization has a prominent effect on the pH, making solutions of strong acids or bases have very low or very high pH values.

In contrast, weak acids and bases only partially ionize in solution. For example, hydrofluoric acid (HF) and aniline (C₆H₅NH₂) do not release or accept all their potential hydrogen ions in solution. This incomplete reaction means that the concentration of hydrogen or hydroxide ions is less than it would be in a strong acid or base, leading to more moderate changes in pH. Consequently, in a mixture of various solutions, strong acids and bases will always be at the extremes of pH ranks, with weak acids and bases falling somewhere in the middle.

Ionic Compounds in Solution

Ionic compounds in solution play a significant role in determining the pH. When these compounds dissolve in water, they disassociate into ions. The nature of the resulting ions—whether they are acidic, basic, or neutral—will affect the pH of the solution. Neutral ionic compounds like sodium iodide (NaI) and sodium nitrate (NaNO₃) do not affect the pH significantly since their constituent ions (I- and NO₃-, respectively) do not react with water to form hydrogen or hydroxide ions.

However, ionic compounds such as sodium fluoride (NaF) and ammonium bromide (NH₄Br) can affect the pH. In the case of NaF, fluoride ions (F-) act as weak bases and can attract protons from water, leading to an increase in pH. Similarly, the ammonium ion (NH₄+) from NH₄Br can donate a proton to water, functioning as a weak acid and thus lowering the pH. Understanding the behavior of ions in solution allows us to predict how an ionic compound will influence the acidity or basicity of a solution, and this knowledge is crucial when ranking the pH of different solutions.