In chemistry, neutralization or neutralisation (see spelling differences) is a chemical reaction in which acid and a base react with an equivalent quantity of each other. In a reaction in water, neutralization results in there being no excess of hydrogen or hydroxide ions present in the solution. The pH of the neutralized solution depends on the acid strength of the reactants.
In the context of a chemical reaction the term neutralization is used for a reaction between an acid and a base or alkali. Historically, this reaction was represented as
For example:
The statement is still valid as long as it is understood that in an aqueous solution the substances involved are subject to dissociation, which changes the ionization state of the substances. The arrow sign, →, is used because the reaction is complete, that is, neutralization is a quantitative reaction. A more general definition is based on Brønsted–Lowry acid–base theory.
Electrical charges are omitted from generic expressions such as this, as each species A, AH, B, or BH may or may not carry an electrical charge. Neutralization of sulfuric acid provides a specific example. Two partial neutralization reactions are possible in this instance.
After an acid AH has been neutralized there are no molecules of the acid (or hydrogen ions produced by dissociation of the molecule) left in solution.
When an acid is neutralized the amount of base added to it must be equal to the amount of acid present initially. This amount of base is said to be the equivalent amount. In a titration of an acid with a base, the point of neutralization can also be called the equivalence point. The quantitative nature of the neutralization reaction is most conveniently expressed in terms of the concentrations of acid and alkali. At the equivalence point:
In general, for an acid AHn at concentration c1 reacting with a base B(OH)m at concentration c2 the volumes are related by:
An example of a base being neutralized by an acid is as follows.
The same equation relating the concentrations of acid and base applies. The concept of neutralization is not limited to reactions in solution. For example, the reaction of limestone with acid such as sulfuric acid is also a neutralization reaction.
Such reactions are important in soil chemistry.
A strong acid is one that is fully dissociated in aqueous solution. For example, hydrochloric acid, HCl, is a strong acid.
A strong base is one that is fully dissociated in aqueous solution. For example, sodium hydroxide, NaOH, is a strong base.
Therefore, when a strong acid reacts with a strong base the neutralization reaction can be written as
For example, in the reaction between hydrochloric acid and sodium hydroxide the sodium and chloride ions, Na+ and Cl− take no part in the reaction. The reaction is consistent with the Brønsted–Lowry definition because in reality the hydrogen ion exists as the hydronium ion, so that the neutralization reaction may be written as
When a strong acid is neutralized by a strong base there are no excess hydrogen ions left in the solution. The solution is said to be neutral as it is neither acidic nor alkaline. The pH of such a solution is close to a value of 7; the exact pH value is dependent on the temperature of the solution.
Neutralization is an exothermic reaction. The standard enthalpy change for the reaction H+ + OH− → H2O is −57.30 kJ/mol.
The term fully dissociated is applied to a solute when the concentration of undissociated solute is below the detection limits, that is, when the undissociated solute's concentration is too low to measured. Quantitatively, this is expressed as log K < −2, or in some texts log K < −1.76. This means that the value of the dissociation constant cannot be obtained from experimental measurements. The value can, however, be estimated theoretically. For example the value of log K ≈ −6 has been estimated for hydrogen chloride in aqueous solution at room temperature. [1] A chemical compound may behave as a strong acid in solution when its concentration is low and as a weak acid when its concentration is very high. Sulfuric acid is an example of such a compound.
A weak acid HA is one that does not dissociate fully when it is dissolved in water. Instead an equilibrium mixture is formed:
Acetic acid is an example of a weak acid. The pH of the neutralized solution resulting from
is not close to 7, as with a strong acid, but depends on the acid dissociation constant, Ka, of the acid. The pH at the end-point or equivalence point in a titration may be calculated as follows. At the end-point the acid is completely neutralized so the analytical hydrogen ion concentration, TH, is zero and the concentration of the conjugate base, A−, is equal to the analytical or formal concentration TA of the acid: [A−] = TA. When a solution of an acid, HA, is at equilibrium, by definition the concentrations are related by the expression
The solvent (e.g. water) is omitted from the defining expression on the assumption that its concentration is very much greater than the concentration of dissolved acid, [H2O] ≫ TA. The equation for mass-balance in hydrogen ions can then be written as
where Kw represents the self-dissociation constant of water. Since Kw = [H+][OH−], the term Kw/[H+] is equal to [OH−], the concentration of hydroxide ions. At neutralization, TH is zero. After multiplying both sides of the equation by [H+], it becomes
and, after rearrangement and taking logarithms,
With a dilute solution of the weak acid, the term 1 + TA/Ka is equal to TA/Ka to a good approximation. If pKw = 14,
This equation explains the following facts:
In a titration of a weak acid with a strong base the pH rises more steeply as the end-point is approached. At the end-point, the slope of the curve of pH with respect to amount of titrant is a maximum. Since the end-point occurs at pH greater than 7, the most suitable indicator to use is one, like phenolphthalein, that changes color at high pH. [2]
The situation is analogous to that of weak acids and strong bases.
Amines are examples of weak bases. The pH of the neutralized solution depends on the acid dissociation constant of the protonated base, pKa, or, equivalently, on the base association constant, pKb. The most suitable indicator to use for this type of titration is one, such as methyl orange, that changes color at low pH.
When a weak acid reacts with an equivalent amount of a weak base,
complete neutralization does not always occur. The concentrations of the species in equilibrium with each other will depend on the equilibrium constant, K, for the reaction, which is defined as follows:
The neutralization reaction can be considered as the difference of the following two acid dissociation reactions
with the dissociation constants Ka,A and Ka,B of the acids HA and BH+, respectively. Inspection of the reaction quotients shows that
K = Ka,A/Ka,B.
A weak acid cannot always be neutralized by a weak base, and vice versa. However, for the neutralization of benzoic acid (Ka,A = 6.5 × 10−5) with ammonia (Ka,B = 5.6 × 10−10 for ammonium), K = 1.2 × 105 >> 1, and more than 99% of the benzoic acid is converted to benzoate.
Chemical titration methods are used for analyzing acids or bases to determine the unknown concentration. Either a pH meter or a pH indicator which shows the point of neutralization by a distinct color change can be employed. Simple stoichiometric calculations with the known volume of the unknown and the known volume and molarity of the added chemical gives the molarity of the unknown.
In wastewater treatment, chemical neutralization methods are often applied to reduce the damage that an effluent may cause upon release to the environment. For pH control, popular chemicals include calcium carbonate, calcium oxide, magnesium hydroxide, and sodium bicarbonate. The selection of an appropriate neutralization chemical depends on the particular application.
There are many uses of neutralization reactions that are acid-alkali reactions. A very common use is antacid tablets. These are designed to neutralize excess gastric acid in the stomach (HCl) that may be causing discomfort in the stomach or lower esophagus. This can also be remedied by the ingestion of sodium bicarbonate (NaHCO3). Sodium bicarbonate is also commonly used to neutralise acid spills in laboratories, as well as acid burns.
In chemical synthesis of nanomaterials, the heat of neutralization reaction can be used to facilitate the chemical reduction of metal precursors. [3]
Also in the digestive tract, neutralization reactions are used when food is moved from the stomach to the intestines. In order for the nutrients to be absorbed through the intestinal wall, an alkaline environment is needed, so the pancreas produce an antacid bicarbonate to cause this transformation to occur.
Another common use, though perhaps not as widely known, is in fertilizers and control of soil pH. Slaked lime (calcium hydroxide) or limestone (calcium carbonate) may be worked into soil that is too acidic for plant growth. Fertilizers that improve plant growth are made by neutralizing sulfuric acid (H2SO4) or nitric acid (HNO3) with ammonia gas (NH3), making ammonium sulfate or ammonium nitrate. These are salts utilized in the fertilizer.
Industrially, a by-product of the burning of coal, sulfur dioxide gas, may combine with water vapor in the air to eventually produce sulfuric acid, which falls as acid rain. To prevent the sulfur dioxide from being released, a device known as a scrubber gleans the gas from smoke stacks. This device first blows calcium carbonate into the combustion chamber where it decomposes into calcium oxide (lime) and carbon dioxide. This lime then reacts with the sulfur dioxide produced forming calcium sulfite. A suspension of lime is then injected into the mixture to produce a slurry, which removes the calcium sulfite and any remaining unreacted sulfur dioxide.
An acid is a molecule or ion capable of either donating a proton (i.e. hydrogen ion, H+), known as a Brønsted–Lowry acid, or forming a covalent bond with an electron pair, known as a Lewis acid.
In chemistry, an acid–base reaction is a chemical reaction that occurs between an acid and a base. It can be used to determine pH via titration. Several theoretical frameworks provide alternative conceptions of the reaction mechanisms and their application in solving related problems; these are called the acid–base theories, for example, Brønsted–Lowry acid–base theory.
Hydroxide is a diatomic anion with chemical formula OH−. It consists of an oxygen and hydrogen atom held together by a single covalent bond, and carries a negative electric charge. It is an important but usually minor constituent of water. It functions as a base, a ligand, a nucleophile, and a catalyst. The hydroxide ion forms salts, some of which dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. The corresponding electrically neutral compound HO• is the hydroxyl radical. The corresponding covalently bound group –OH of atoms is the hydroxy group. Both the hydroxide ion and hydroxy group are nucleophiles and can act as catalysts in organic chemistry.
Hydrolysis is any chemical reaction in which a molecule of water breaks one or more chemical bonds. The term is used broadly for substitution, elimination, and solvation reactions in which water is the nucleophile.
In chemistry, pH, also referred to as acidity or basicity, historically denotes "potential of hydrogen". It is a scale used to specify the acidity or basicity of an aqueous solution. Acidic solutions are measured to have lower pH values than basic or alkaline solutions.
Titration is a common laboratory method of quantitative chemical analysis to determine the concentration of an identified analyte. A reagent, termed the titrant or titrator, is prepared as a standard solution of known concentration and volume. The titrant reacts with a solution of analyte to determine the analyte's concentration. The volume of titrant that reacted with the analyte is termed the titration volume.
A buffer solution is a solution where the pH does not change significantly even on dilution or even if an acid or base is added at constant temperature. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many living systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood, and bicarbonate also acts as a buffer in the ocean.
In chemistry, hydronium (hydroxonium in traditional British English) is the common name for the cation [H3O]+, also written as H3O+, the type of oxonium ion produced by protonation of water. It is often viewed as the positive ion present when an Arrhenius acid is dissolved in water, as Arrhenius acid molecules in solution give up a proton (a positive hydrogen ion, H+) to the surrounding water molecules (H2O). In fact, acids must be surrounded by more than a single water molecule in order to ionize, yielding aqueous H+ and conjugate base. Three main structures for the aqueous proton have garnered experimental support: the Eigen cation, which is a tetrahydrate, H3O+(H2O)3, the Zundel cation, which is a symmetric dihydrate, H+(H2O)2, and the Stoyanov cation, an expanded Zundel cation, which is a hexahydrate: H+(H2O)2(H2O)4. Spectroscopic evidence from well-defined IR spectra overwhelmingly supports the Stoyanov cation as the predominant form. For this reason, it has been suggested that wherever possible, the symbol H+(aq) should be used instead of the hydronium ion.
In chemistry, an acid dissociation constant is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction
Solubility equilibrium is a type of dynamic equilibrium that exists when a chemical compound in the solid state is in chemical equilibrium with a solution of that compound. The solid may dissolve unchanged, with dissociation, or with chemical reaction with another constituent of the solution, such as acid or alkali. Each solubility equilibrium is characterized by a temperature-dependent solubility product which functions like an equilibrium constant. Solubility equilibria are important in pharmaceutical, environmental and many other scenarios.
A pH indicator is a halochromic chemical compound added in small amounts to a solution so the pH (acidity or basicity) of the solution can be determined visually or spectroscopically by changes in absorption and/or emission properties. Hence, a pH indicator is a chemical detector for hydronium ions (H3O+) or hydrogen ions (H+) in the Arrhenius model.
In chemistry, there are three definitions in common use of the word "base": Arrhenius bases, Brønsted bases, and Lewis bases. All definitions agree that bases are substances that react with acids, as originally proposed by G.-F. Rouelle in the mid-18th century.
The self-ionization of water (also autoionization of water, and autodissociation of water, or simply dissociation of water) is an ionization reaction in pure water or in an aqueous solution, in which a water molecule, H2O, deprotonates (loses the nucleus of one of its hydrogen atoms) to become a hydroxide ion, OH−. The hydrogen nucleus, H+, immediately protonates another water molecule to form a hydronium cation, H3O+. It is an example of autoprotolysis, and exemplifies the amphoteric nature of water.
A weak base is a base that, upon dissolution in water, does not dissociate completely, so that the resulting aqueous solution contains only a small proportion of hydroxide ions and the concerned basic radical, and a large proportion of undissociated molecules of the base.
An acid–base titration is a method of quantitative analysis for determining the concentration of Brønsted-Lowry acid or base (titrate) by neutralizing it using a solution of known concentration (titrant). A pH indicator is used to monitor the progress of the acid–base reaction and a titration curve can be constructed.
The Brønsted–Lowry theory (also called proton theory of acids and bases) is an acid–base reaction theory which was first developed by Johannes Nicolaus Brønsted and Thomas Martin Lowry independently in 1923. The basic concept of this theory is that when an acid and a base react with each other, the acid forms its conjugate base, and the base forms its conjugate acid by exchange of a proton (the hydrogen cation, or H+). This theory generalises the Arrhenius theory.
Acid salts are a class of salts that produce an acidic solution after being dissolved in a solvent. Its formation as a substance has a greater electrical conductivity than that of the pure solvent. An acidic solution formed by acid salt is made during partial neutralization of diprotic or polyprotic acids. A half-neutralization occurs due to the remaining of replaceable hydrogen atoms from the partial dissociation of weak acids that have not been reacted with hydroxide ions to create water molecules.
In chemistry, a strong electrolyte is a solute that completely, or almost completely, ionizes or dissociates in a solution. These ions are good conductors of electric current in the solution.
The Charlot equation, named after Gaston Charlot, is used in analytical chemistry to relate the hydrogen ion concentration, and therefore the pH, with the formal analytical concentration of an acid and its conjugate base. It can be used for computing the pH of buffer solutions when the approximations of the Henderson–Hasselbalch equation break down. The Henderson–Hasselbalch equation assumes that the autoionization of water is negligible and that the dissociation or hydrolysis of the acid and the base in solution are negligible.
Neutralization is covered in most general chemistry textbooks. Detailed treatments may be found in textbooks on analytical chemistry such as
Applications
