General Chemistry/Acids and bases neutralize each other

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An acid (from Latin acidus, meaning sour or tart) is a w:chemical compound generally defined by its reactions with complementary chemicals, designated bases (see w:Acid-base reaction theories). An acid tends to give a proton and can be represented by the generic formula AH. In water, there is the following reaction:

$\text{AH}+{\text{H}}_2\text{O}\iff {\text{A}}^{-}+{\text{H}}_3{\text{O}}^{+}$

There is a distinction between w:weak acids and w:strong acids. For a strong acid, no AH remains in solution:

Failed to parse (syntax error): {\displaystyle \mbox{AH} +\mbox{H}_2\mbox{O} \rightarrow \mbox{A}^- + \mbox{H}_3\mbox  :<math>\mbox{Ka} = {\mbox{A}^- \mbox{H}_3\mbox{O}^+ \over AH}}

Some of the stronger acids include the hydrohalic acids - HCl, HBr, and HI - and the oxyacids, which tend to contain central atoms in high oxidation states surrounded by oxygen - including HNO₃, H₂SO₄, HClO₄.

Acidity is typically measured using the w:pH scale.

Acidic (chemistry), the opposite to basic, reacting with basics to form salts. Acidic (geology), of rock: containing more than 65% of silica.

Acids are generally:

• Taste: sour when dissolved in water
• Touch: strong acids have a stinging feeling
• Reactivity: acids react violently with many w:metals
• Electrical Conductivity: acids are w:electrolytes

Acids in Food

• w:citric acid: found in w:citrus fruits
• w:lactic acid: found in w:dairy products such as w:yoghurt and sour w:milk
• w:carbonic acid: found in w:carbonated beverages
• w:acetic acid: found in w:vinegar

Different Definitions of Acid/Base

The word acid comes from the w:Latin acidus meaning sour. Chemically though the term acid has a more specific meaning.

The Swedish chemist w:Svante Arrhenius defined an acid to be a substance that gave up hydrogen ions (H+) when dissolved in water, while bases are substances that give up hydroxide ions (OH-). Notice that this definition limits acids and bases to substances that can dissolve in water. Later on, w:Bronsted and w:Lowry defined an acid to be a w:proton donor and a base to be a proton acceptor. In this definition, even substances that are insoluble in water can be acids and bases. The most general definition of acids and bases is the Lewis definition. A w:Lewis acid is an w:electron acceptor, while a w:Lewis base is an electron donor. Acid/base systems are different from redox reactions in that there is no change in oxidation state.

Acid number

This is used to quantify oxidation. It is the quantity of base, expressed in milligrams of w:potassium hydroxide, that is required to neutralize the acidic constituents in 1 g of sample.

AN = (V_eq-b_eq)×N×56.1/W_oil ).

V_eq is the amount of titrant (ml) consumed by crude oil sample and 1ml w:spiking solution at the equivalent point, and b_eqbeq is the amount of titrant (ml) consumed by 1ml spiking solution at the equivalent point.

The molarity concentration of titrant (N) is calculated as such: N = 1000×W_KHP/(204.23×V_eq).

In which, W_KHP is the amount (g) of KHP in 50ml of KHP standard solution, and Veq is the amount of titrant (ml) consumed by 50ml KHP standard solution at the equivalent point.

Acid number (mgw:KOH/g oil) for w:biodiesel is preferred to be lower than 3.

Bases

In w:chemistry, a base is a compound that is the opposite of an w:acid in the sense that it will neutralize an acid. Common bases include compounds such as some metal oxides and hydroxides, and w:ammonia.

An acid "donates" H+ ions to the solution, while a base "accepts" H+ ions _or_ donates OH- ions. Both of those actions will decrease the hydrogen ion (H+) concentration, and thus increase pH (-log[H+])

Soluble bases (w:alkalis) produce w:hydroxyl w:ion (OH^-) in aqueous solution and have a w:pH above 7.

Example:

The amino group (NH2) acts as a base by accepting a H+ ions from the solution. It does this by forming a w:coordiate covalent bond with the unshared pair of electrons belonging to the nitrogen atom. This decreases the hydrogen ion concentration.

Sodium hydroxide (NaOH) decomposes into Na+ and OH-, lowering the hydrogen ion concentration because the hydroxide ion will accept hydrogen ions to form water.

$H^{+}+O{H}^{-}\leftrightarrow H_{2}O$

Acid-base reaction theories

Originally, w:acids and bases were defined only by properties, such as the fact that acids tasted sour and turned certain plant w:dyes from one color to another, while bases tasted bitter, changed the colors of the same plant dyes in the opposite direction, and neutralized the activity of acids. A scientific definition was first proposed by the French w:chemist w:Antoine Lavoisier.

Lavoisier's definition

Since Lavoisier's knowledge of strong acids was mainly restricted to the oxyacids, which tend to contain central atoms in high oxidation states surrounded by oxygen, such as HNO₃ and H₂SO₄, and since he was not aware of the true composition of the hydrohalic acids, HCl, HBr, and HI, he defined acids in terms of their containing w:oxygen, which in fact he named from Greek words meaning "acid-former". When the elements w:chlorine, w:bromine, and w:iodine were identified and the absence of oxygen in the hydrohalic acids was established, this definition had to be rejected.

The Arrhenius definition

w:Svante Arrhenius provided the first modern definition of acids and bases in w:1884. In w:water, a dissociation takes place:

H₂O → H⁺ + OH^-

A compound causing an increase in H⁺ and a decrease in OH^- is an acid and one causing the reverse is a base.

An Arrhenius acid, when dissociated in water, typically yields positively-charged w:hydrogen ion and a complementary negative en:ion.

An Arrhenius base, when dissociated in water, typically yields negatively-charged w:hydroxide ion and a complementary positive w:ion.

The positive ion from a base can form a salt from the negative ion of an acid. For example, two moles of the base w:sodium w:hydroxide (NaOH) can combine with one mole of sulfuric acid (H₂SO₄) to form two moles of w:water and one mole of sodium w:sulfate.

2NaOH + H₂SO₄ → 2H₂O + Na₂SO₄

The protonic (Brønsted-Lowry) definition

The Brønsted-Lowry definition, formulated independently by its two proponents in w:1923, revolves around an w:acid's ability to donate w:protons (H⁺) to another compound, called a base, in a w:chemical reaction.

A base is a w:proton acceptor. In Brønsted-Lowry acid-base reactions, there is a "competition" between two bases for a proton. so that if X and Y are two species, the equilibrium

HX + Y^- ↔ HY + X^-:

occurs. Both HX and HY are Brønsted-Lowry acids; both X^- and Y^- are Brønsted-Lowry bases. If the reaction runs mostly to the left, then HY is the stronger acid and X^- the stronger base; if the reaction runs mostly to the right, then HX is the stronger acid and Y^- the stronger base.

Acids and bases in the Brønsted-Lowry system occur in conjugate pairs; in the reaction

HX → H⁺ + X^-

HX is denoted the w:conjugate acid of the base X^-, and X^- is denoted the w:conjugate base of the acid HX.

Some compounds, like w:water, can act either as an acid or a base, and are called amphoteric compounds. Stronger acids also typically oxidize metals, forming salts and releasing hydrogen.

See w:pH for a measure of proton concentration frequently used for measuring acidity and alkalinity using this definition.

The solvent-system definition

This definition is based on a generalization of the earlier Arrhenius definition. If we consider a solvent which can be dissociated into a positive species X and a negative species Y:

XY ↔ X⁺ + Y^-

or

2XY ↔ X₂Y⁺ + Y^-

or

2XY ↔ X⁺ + XY₂^-

a compound causing an increase in X⁺ (or X₂Y⁺) and a decrease in Y^- (or XY₂^-) is an acid and one causing the reverse is a base. For example in liquid w:sulfur dioxide (SO₂), w:thionyl compounds (formally supplying SO⁺²) behave as acids, and w:sulfites (supplying SO₃^-2) behave as bases.

In this more general sense, aprotic compounds (those which do not donate protons), can still react with bases, and the terms "acid" and "base" can still be used for reactions in aprotic or non-aqueous environments.

The w:electronic (Lewis) definition

The more general definition offered by Lewis in w:1923 (the same year as the Brønsted-Lowry definition) describes the reactivity of an acid in terms of its ability to accept a pair of electrons from a base, defined as an electron-pair donor. In general, an acid reacts with a base by forming a new w:covalent bond utilizing an empty orbital of the acid to share the extra electron pair of the base.

The Usanovich definition

The most general definition is that of the Russian chemist Usanovich, and can basically be summarized as defining an acid as anything that accepts negative species or donates positive ones, and a base as the reverse. This tends to overlap the concept of w:redox (w:oxidation-w:reduction), and so is not highly favored by chemists.

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