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Acids and Bases

Brønsted-Lowry Acids and Bases

Identifying Acids and Bases

In the Brønsted-Lowry model, acids are compounds that lose a proton in solution, and bases are compounds that accept a proton in solution.
Acids and bases were initially defined by their macroscopic properties: acids taste sour and can corrode metals, and bases feel slippery and are neutralized by adding acid. In 1887 Swedish chemist Svante Arrhenius defined an acid as a substance that releases an H+ ion (a single proton) when it is dissolved in water, and a base as a substance that releases an OH (hydroxide) ion. Although this definition holds true, many acid-base reactions do not involve hydroxide ions or do not occur in water. Reactions in which ammonia acts as a base, for example, do not release an OH ion.
HNO3(aq)+NH3(aq)NO3(aq)+NH4+(aq){\rm{HNO}_3}(aq)+{\rm{NH}_3}(aq)\rightarrow {{\rm{NO}_3}^-}(aq)+{{\rm{NH}_4}^+}(aq)
Proposed in 1923, the Brønsted-Lowry model offers a more precise description of what occurs on the molecular level in acid-base chemistry. A Brønsted-Lowry acid is a compound that can donate a proton to another compound in solution. A Brønsted-Lowry base is a compound that can accept a proton from another compound in solution. In other words, in an acid-base reaction, the acid always donates a proton to the base.

Ammonium Ionization

In ammonium ionization, nitric acid (HNO3) donates a proton (H+) to ammonia. The ammonia, acting as a base, accepts the proton, and there are no OH- ions involved. Protonated ammonia generates a new cation, ammonium (NH4+).
The NO3 anion resulting from the ammonium ionization reaction is the conjugate base of nitric acid (HNO3). A conjugate base is the molecule or ion remaining after a Brønsted-Lowry acid donates its proton to another molecule or ion (base). In this case NO3 is the conjugate base of HNO3. The NH4+ ion is the conjugate acid, the molecule or ion formed when a Brønsted-Lowry base has accepted a proton.
Acids, bases, and their conjugates are related by the transfer of protons. The Brønsted-Lowry acid donates a proton to the Brønsted-Lowry base, each of which has a conjugate in the products of the reaction.
Another important aspect of the Brønsted-Lowry model is its recognition that water both accepts and donates a proton, acting as both an acid and a base.

Ionization of Water

In the ionization of water, H+ ions can move from one water molecule to another. The molecule that becomes the positively charged hydronium ion acts as a base by accepting a proton. The molecule that becomes the negatively charged hydroxide ion acts as an acid by losing a proton.
Other compounds, such as HCO3, HSO4, and H2PO4, are amphiprotic, which means they are able to act as both a proton donor and a proton acceptor. They are all products of the first ionization of polyprotic acids, acids that can donate more than one proton or hydrogen atom per molecule to an aqueous solution.

Kw, the Ion Product Constant of Water

The ion product constant of water, Kw, is an equilibrium constant derived from the concentrations of H+ and OH present at equilibrium.
The equilibrium for water self-ionization is described by this equation:
Kw, the equilibrium constant for this reaction, known as the ionization product constant of water, can be defined as the product of the hydronium ion concentration, [H3O+], and the hydroxide ion concentration [OH].
At 25°C these concentrations in pure water are equal. Therefore knowing the value of Kw allows the calculation of [H3O+] and [OH].
Because the two concentrations are the same, the square of one of either concentration can be used to find the pH:
1.0×1014=[H3O+]21.0\times10^{-14}=\left[{\rm H_3}{\rm O^{+}}\right]^{2}
Take the square root of both sides of the equation to find the concentration.
[H3O+]=1.0×107\left[{\rm H_{3}O^{+}}\right]=1.0\times10^{-7}
The exponent indicates that the pH is 7.
As with many constants, the value of Kw varies with temperature. The ionization of water is an endothermic process, meaning it absorbs heat.
According to Le Chatelier's principle, a stress such as a change in the temperature, pressure, or concentration of a component will cause the equilibrium condition of a chemical system to change in a way that reduces the change in order to regain equilibrium. Heat added to the system will drive the reaction to the right, increasing the concentrations of H3O+ and OH and therefore increasing the value of Kw. The following chart shows this variation.

The Ion Product for Water at Various Temperatures

T (°C) Kw pH
0 0.114×10140.114\times10^{-14} 7.47
10 0.293×10140.293\times10^{-14} 7.27
20 0.681×10140.681\times10^{-14} 7.08
25 1.008×10141.008\times10^{-14} 7.00
30 1.471×10141.471\times10^{-14} 6.92
40 2.916×10142.916\times10^{-14} 6.77
50 5.476×10145.476\times10^{-14} 6.63

Temperature affects both Kw and pH.

An increase in temperature causes a decrease in pH. This means that not only has the concentration of H3O+ ions increased, so has the concentration of OH ions, because these two concentrations always remain equal. When there is an equal concentration of H+ and OH ions in a solution, the solution is neutral, even if pH has shifted because of the temperature.