Free Hydrogen Ions do not Exist in WaterNeutralization

There are three major classifications of substances known as acids or bases. The Arrhenius definition states that an acid produces H+ in solution and a base produces OH-. This theory was developed by Svante Arrhenius in 1883. Later, two more sophisticated and general theories were proposed. These are the Brønsted-Lowry and the Lewis definitions of acids and bases. The Lewis theory is discussed elsewhere.

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The Arrhenius Theory of Acids and Bases

In 1884, the Swedish les-grizzlys-catalans.orgist Svante Arrhenius proposed two specific classifications of compounds; acids and bases. When dissolved in an aqueous solution, certain ions were released into the solution. An Arrhenius acid is a compound that increases the concentration of H+ ions that are present when added to water. These H+ ions form the hydronium ion (H3O+) when they combine with water molecules. This process is represented in a les-grizzlys-catalans.orgical equation by adding H2O to the reactants side.

\< HCl_(aq) \rightarrow H^+_(aq) + Cl^-_(aq) \>

In this reaction, hydrochloric acid (\(HCl\)) dissociates completely into hydrogen (H+) and chlorine (Cl-) ions when dissolved in water, thereby releasing H+ ions into solution. Formation of the hydronium ion equation:

\< HCl_(aq) + H_2O_(l) \rightarrow H_3O^+_(aq) + Cl^-_(aq) \>

The Arrhenius theory, which is the simplest and least general description of acids and bases, includes acids such as HClO4 and HBr and bases such as \(NaOH\) or \(Mg(OH)_2\). For example the complete dissociation of \(HBr\) gas into water results generates free \(H_3O^+\) ions.

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This theory successfully describes how acids and bases react with each other to make water and salts. However, it does not explain why some substances that do not contain hydroxide ions, such as \(F^-\) and \(NO_2^-\), can make basic solutions in water. The Brønsted-Lowry definition of acids and bases addresses this problem.

An Arrhenius base is a compound that increases the concentration of OH- ions that are present when added to water. The dissociation is represented by the following equation:

\< NaOH \; (aq) \rightarrow Na^+ \; (aq) + OH^- \; (aq) \>

In this reaction, sodium hydroxide (NaOH) disassociates into sodium (Na+) and hydroxide (OH-) ions when dissolved in water, thereby releasing OH- ions into solution.


Note

Arrhenius acids are substances which produce hydrogen ions in solution. Arrhenius bases are substances which produce hydroxide ions in solution.

Free Hydrogen Ions do not Exist in Water

Owing to the overwhelming excess of \(H_2O\) molecules in aqueous solutions, a bare hydrogen ion has no chance of surviving in water. The hydrogen ion in aqueous solution is no more than a proton, a bare nucleus. Although it carries only a single unit of positive charge, this charge is concentrated into a volume of space that is only about a hundred-millionth as large as the volume occupied by the smallest atom. (Think of a pebble sitting in the middle of a sports stadium!) The resulting extraordinarily high charge density of the proton strongly attracts it to any part of a nearby atom or molecule in which there is an excess of negative charge. In the case of water, this will be the lone pair (unshared) electrons of the oxygen atom; the tiny proton will be buried within the lone pair and will form a shared-electron (coordinate) bond with it, creating a hydronium ion, \(H_3O^+\). In a sense, \(H_2O\) is acting as a base here, and the product \(H_3O^+\) is the conjugate acid of water:

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Although other kinds of dissolved ions have water molecules bound to them more or less tightly, the interaction between H+ and \(H_2O\) is so strong that writing “H+(aq)” hardly does it justice, although it is formally correct. The formula \(H_3O^+\) more adequately conveys the sense that it is both a molecule in its own right, and is also the conjugate acid of water.

The equation "HA → H+ + A–" is so much easier to write that les-grizzlys-catalans.orgists still use it to represent acid-base reactions in contexts in which the proton donor-acceptor mechanism does not need to be emphasized. Thus, it is permissible to talk about “hydrogen ions” and use the formula H+ in writing les-grizzlys-catalans.orgical equations as long as you remember that they are not to be taken literally in the context of aqueous solutions.



\< HCl \; (aq) + NH_3 \; (aq) \rightarrow NH_4^+ \; (aq) + Cl^- \; (aq) \>



Strong and Weak Acids and Bases

Strong acids are molecular compounds that essentially ionize to completion in aqueous solution, disassociating into H+ ions and the additional anion; there are very few common strong acids. All other acids are "weak acids" that incompletely ionized in aqueous solution. Acids and bases that dissociate completely are said to be strong acids, e.g.:

\(HClO_4(aq) \rightarrow H^+_(aq) + ClO^-_4(aq)\) \(HBr_(aq) \rightarrow H^+_(aq) + Br^-_(aq)\) \(CH_3O^-_(aq) + H_2O_(l) \rightarrow CH_3OH_(aq) + OH^-_(aq)\) \(NH^-_2(aq) + H_2O_(l) \rightarrow NH_3(aq) + OH^-_(aq)\)

Here the right-handed arrow (\(\rightarrow\)) implies that the reaction goes to completion. That is, a 1.0 M solution of HClO4 in water actually contains 1.0 M H+(aq) and 1.0 M ClO4-(aq), and no undissociated HClO4.

Conversely, weak acids such as acetic acid (CH3COOH) and weak bases such as ammonia (NH3) dissociate only slightly in water - typically a few percent, depending on their concentration and exist mostly as the undissociated molecules.

STRONG ACIDS: HCl, HNO3, H2SO4, HBr, HI, HClO4 WEAK ACIDS: All other acids, such as HCN, HF, H2S, HCOOH

Strong acids such as \(HCl\) dissociate to produce spectator ions such as \(Cl^-\) as conjugate bases, whereas weak acids produce weak conjugate bases. This is illustrated below for acetic acid and its conjugate base, the acetate anion. Acetic acid is a weak acid (Ka = 1.8 x 10-5) and acetate is a weak base (Kb = Kw/Ka = 5.6 x 10-10)

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Like acids, strong and weak bases are classified by the extent of their ionization. Strong bases disassociate almost or entirely to completion in aqueous solution. Similar to strong acids, there are very few common strong bases. Weak bases are molecular compounds where the ionization is not complete.

WEAK BASES: All other bases, such as NH3, CH3NH2, C5H5N

Note

The strength of a conjugate acid/base varies inversely with the strength or weakness of its parent acid or base. Any acid or base is technically a conjugate acid or conjugate base also; these terms are simply used to identify species in solution (i.e acetic acid is the conjugate acid of the acetate anion, a base, while acetate is the conjugate base of acetic acid, an acid).



pH Scale

Since acids increase the amount of H+ ions present and bases increase the amount of OH- ions, under the pH scale, the strength of acidity and basicity can be measured by its concentration of H+ ions. This scale is shown by the following formula:

pH = -log

with being the concentration of H+ ions.

To see how these calculations are done, refer to Calculating the pH of the solution of a Polyprotic Base/Acid

The pH scale is often measured on a 1 to 14 range, but this is incorrect (see pH for more details). Something with a pH less than 7 indicates acidic properties and greater than 7 indicates basic properties. A pH at exactly 7 is neutral. The higher the , the lower the pH.

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Figure 4. The pH scale shows that substances with a pH greater than 7 are basic and a pH less than 7 are acidic.

Lewis Theory

The Lewis theory of acids and bases states that acids act as electron pair acceptors and bases act as electron pair doners. This definition doesn"t mention anything about the hydrogen atom at all, unlike the other definitions. It only talks about the transfer of electron pairs. To demonstrate this theory, consider the following example.

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This is a reaction between ammonia (NH3) and boron trifluoride (BF3). Since there is no transfer of hydrogen atoms here, it is clear that this is a Lewis acid-base reaction. In this reaction, NH3 has a lone pair of electrons and BF3 has an incomplete octet, since boron doesn"t have enough electrons around it to form an octet.

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Figure 2. The Lewis structures of ammonia and boron trifluoride.

Because boron only has 6 electrons around it, it can hold 2 more. BF3 can act as a Lewis acid and accept the pair of electrons from the nitrogen in NH3, which will then form a bond between the nitrogen and the boron.

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Figure 3. The Lewis structure of \(H_3NBF_3\), which resulted from the coordinate covalent bond between nitrogen and boron.

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This is considered an acid-base reaction where NH3 (base) is donating the pair of electrons to BF3. BF3 (acid) is accepting those electrons to form a new compound, H3NBF3.