Base (chemistry)

In a basic solution, phenolphthalein has a pink or red color.

In chemistry, a base is thought of as a substance which can accept protons or any chemical compound that yields hydroxide ions (OH-) in solution. It is also commonly referred to as any substance that can react with an acid to decrease or neutralize its acidic properties, change the color of indicators (e.g. turn red litmus paper blue), feel slippery to the touch when in solution, taste bitter, react with acids to form salts, and promote certain chemical reactions (e.g. base catalysis). Example of simple bases are sodium hydroxide and ammonia.Sodium Hydroxide (NaOH), also known as Caustic Soda or Lye dissociates in water to form hydroxide ions (OH^-) and sodium ions (Na⁺).

Bases have many practical uses. Several substances are commonly found in our homes. Household ammonia is a familiar cleaning agent. Sodium hydroxide (lye) is used for cleaning clogs and sink drains. Potassium hydroxide, also called caustic potash is used to make soft soap that dissolve in water with ease. Magnesium hydroxide in water (also called milk of magnesia) is used as an antacid or a laxative.

Acids and bases:

• Acid-base extraction
• Acid-base reaction
• Acid dissociation constant
• Acidity function
• Buffer solutions
• pH
• Proton affinity
• Self-ionization of water
• Acids:
  • Lewis acids
  • Mineral acids
  • Organic acids
  • Strong acids
  • Superacids
  • Weak acids
• Bases:
  • Lewis bases
  • Organic bases
  • Strong bases
  • Superbases
  • Non-nucleophilic bases
  • Weak bases

Alkali and Origin of the concept of base

Alkali (rather than base) was at first the antithesis of an acid. The word “alkali” is derived from the Arabic word al qalīy = "the calcined ashes." These plant ashes were regarded as having properties such as: the ability to reverse action of acids, and having detergent power. The formation of salts from the acid and alkali reaction led to the view that salts can be derived from two constituents of opposite natures. But not all non-acidic constituents possessed alkaline properties (example, oxides and hydroxides of heavy metals) hence; the concept of “base” was born. This concept was first introduced by the French chemist Guillaume François Rouelle in 1754. He noted that acids which in those days were mostly volatile liquids (like acetic acid) turned into solid salts only when combined with specific substances. These substances formed a concrete base for the salt and hence the name.

Definitions of acids and bases

Acids and bases form complementary pairs, so their definitions need to be considered together. There are three common groups of definitons: the Arrhenius, Brønsted-Lowry, and Lewis definitions, in order of increasing generality.

• Arrhenius: According to this definition, an acid is a substance that increases the concentration of hydronium ion (H₃O⁺) when dissolved in water, while bases are substances that increase the concentration of hydroxide ions (OH^-). This definition limits acids and bases to substances that can dissolve in water. Around 1800, many French chemists, including Antoine Lavoisier, incorrectly believed that all acids contained oxygen. Indeed the modern German word for Oxygen is Sauerstoff (lit. sour substance). English chemists, including Sir Humphry Davy at the same time believed all acids contained hydrogen. The Swedish chemist Svante Arrhenius used this belief to develop this definition of acid.
• Brønsted-Lowry: According to this definition, an acid is a proton (hydrogen nucleus) donor and a base is a proton (hydrogen nucleus) acceptor. The acid is said to be dissociated after the proton is donated. An acid and the corresponding base are referred to as conjugate acid-base pairs. Brønsted and Lowry formulated this definition, which includes water-insoluble substances not in the Arrhenius definition.
• Lewis: According to this definition, an acid is an electron-pair acceptor and a base is an electron-pair donor. (These are frequently referred to as "Lewis acids" and "Lewis bases," and are electrophiles and nucleophiles, respectively, in organic chemistry; Lewis bases are also ligands in coordination chemistry.) Lewis acids include substances with no transferable protons (ie H⁺ hydrogen ions), such as iron(III) chloride, and hence the Lewis definition of an acid has wider application than the Brønsted-Lowry definition. The Lewis definition can also be explained with molecular orbital theory. In general, an acid can receive an electron pair in its lowest unoccupied orbital (LUMO) from the highest occupied orbital (HOMO) of a base. That is, the HOMO from the base and the LUMO from the acid combine to a bonding molecular orbital. This definition was developed by Gilbert N. Lewis.

General properties

Some general properties of bases include:

• Taste: Bitter taste (opposed to sour taste of acids and sweetness of aldehydes and ketones)
• Touch: Slimy or soapy feel on fingers
• Reactivity:Caustic on organic matter, react violently with acidic or reducible substances
• Electric conductivity: Aqueous solutions or molten bases dissociate in ions and conduct electricity
• Litmus test: Bases turn red litmus paper blue.

Chemical Characteristics

Bases Ionization Constant and pH

A general equation can be written for the acceptance of H⁺ ions from water by a molecular base, B, to form its conjugate acid, BH⁺.

B(aq) + H₂O(l) ⇌ BH⁺(aq) + OH^-(aq)

$K_{b}={[ConjugateAcid]\cdot [OH^{-}] \over [ConjugateBase]}$ Then, $K_{b}={[BH^{+}]\cdot [OH^{-}] \over [BH]}$

The equilibrium constant Kb is also called the Base Ionization Constant. It refers to the reaction in which a base forms its conjugate acid by removing an H⁺ ion from water.

The pH of (impure) water is a measure of its acidity. In pure water, about one in ten million molecules dissociate into hydronium ions (H₃O⁺) and hydroxide ions (OH⁻), according to the following equation:

2H₂O(l) ⇌ H₃O⁺(aq) + OH^-(aq)

A base accepts (removes) hydronium ions (H₃O⁺) from the solution, or donates hydroxide ions (OH^-) to the solution. Both actions will lower the concentration of hydronium ions, and thus raise pH. By contrast, an acid donates H₃O⁺ ions to the solution or accepts OH⁻, thus lowering pH.

For example, if 1 mole of sodium hydroxide (40 g) is dissolved in 1 litre of water, the concentration of hydroxide ions becomes [OH⁻] = 1 mol/L. Therefore [H⁺] = 10⁻¹⁴ mol/L, and pH = −log 10⁻¹⁴ = 14.

The basicity constant or pK_b is a measure of basicity and related to the pKa by the simple relationship pK_a + pK_b = 14.

Base Strenght

A "Strong Base" is one which hydrolyzes completely, deprotonating acids in an acid-base reaction, hence, raising the pH of the solution towards 14. Compounds with a pH of more than about 13 are called strong bases. Strong bases, like strong acids, attack living tissue and cause serious burns. They react differently to skin than acids do so while strong acids are corrosive, we say that strong bases are caustic. Common examples of strong bases are the hydroxides of alkali metals and alkaline earth metals like NaOH and Ca(OH)₂. Very strong bases are even able to deprotonate very weakly acidic C-H groups in the absence of water. Superbases are a class of especially basic compounds and harpoon bases are a special class of strong bases with poor nucleophilicity.

Examples of Strong Bases (Hydroxide compounds) in descending strenght:

• Potassium hydroxide (KOH)
• Barium hydroxide (Ba(OH)₂)
• Cesium hydroxide (CsOH)
• Sodium hydroxide (NaOH)
• Strontium hydroxide (Sr(OH)₂)
• Calcium hydroxide (Ca(OH)₂)
• Lithium hydroxide (LiOH)
• Rubidium hydroxide (RbOH)

The cations of these strong bases appear in groups 1 and 2 of the periodic table (alkali and alkaline earth metals).

Even stronger bases are:

• Sodium hydride (NaH)
• Lithium diisopropylamide (LDA) (C₆H₁₄LiN)
• Sodium amide (NaNH₂)

A "Weak Base" is one that does not fully ionize in solution. When a base ionizes, it takes up a hydrogen ion from the water around it, leaving an OH- ion behind. Weak bases have a higher H⁺ concentration than strong bases. Weak bases exist in chemical equilibrium in the same way weak acids do. The Base Ionization Constant K_b indicates the strength of the base. Large K_bs belong to stronger bases. The pH of a base is greater than 7 (where 7 is the neutral number; below 7 is an acid), normally up to 14. Common example of a weak base is ammonia, which is used for cleaning.

Examples of Weak Bases:

• Alanine (C₃H₅O₂NH₂)
• Ammonia (water) (NH₃ (NH₄OH))
• Dimethylamine ((CH₃)₂NH)
• Ethylamine (C₂H₅NH₂)
• Glycine (C₂H₃O₂NH₂)
• Hydrazine (N₂H₄)
• Methylamine (CH₃NH₂)
• Trimethylamine ((CH₃)₃N)

Acid - Base Neutralization

Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called neutralization. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium ion (H₃O⁺) concentration in water, where as bases reduce this concentration. Bases react with acids to produce salts and water.

A salts positive ion comes from the base and its negative ion comes from the acid. Considering a metal hydroxide as a base the general reaction is:

HX(aq) + MOH(aq) → MX(aq) + HOH(l)

acid base salt water

Salts of strong bases and strong acids

A strong acid HCl (hydrochloric acid) reacts with a strong base NaOH (sodium hydroxide) to form NaCl (salt = sodium chloride) and water. If the amounts of the acid and the base are in the correct stoichiometric ratio, then the reaction will undergo complete neutralization where the acid and the base both will lose their respective properties.

HCL(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)

strong strong salt water

acid base

Salts of strong bases and weak acids

A strong base NaOH (sodium hydroxide) added to a weak acid CH₃COOH (acetic acid) in 1L of solution, forming NaCH₃COO (sodium acetate) and water.

CH₃COOH (aq) + NaOH(aq) → NaCH₃COO (aq) + H₂O(l)

weak weak salt water

acid base

Salts of weak bases and strong acids

Weak bases react with strong acids to form acidic salt solutions. The conjugate acid of the weak base determines its pH. For example, NH₃ (ammonia) is added to HCl (hydrochloric acid) to form NH₄Cl (ammonium chloride).

NH₃(aq) + HCl(aq) → NH₄Cl(aq)

weak strong salt

base acid

As soon as the salt is formed it reacts with water, resulting in a slightly acidic solution.

Salts of weak bases and weak acids

Salt solutions containing acidic cations and basic anions such as NH4F (ammonium fluoride) have two possible reactions:

NH₄⁺(aq) + H₂O(l) ↔ H₃O⁺(aq) + NH₃(aq) Ka(NH₄⁺) = 5.6 x 10-10 F^-(aq) + H₂O(l) ↔ HF(aq) + OH^-(aq) Kb(F^-) = 1.4 x 10-11

Since Ka(NH₄⁺) > Kb (F,sup>-), the reaction of ammonia with water is more favourable. Therefore, the resulting solution is slightly acidic.

Alkalis

Confusion between base and alkali

The terms "base" and "alkali" are often used interchangeably, since most common bases are alkalis. It is common to speak of "measuring the alkalinity of soil" when what is actually meant is the measurement of the pH (base property). Similarly, bases which are not alkalis, such as ammonia, are sometimes erroneously referred to as alkaline.

Note that not all or even most salts formed by alkali metals are alkaline; this designation applies only to those salts which are basic.

While most electropositive metal oxides are basic, only the soluble alkali metal and alkali earth metal oxides can be correctly called alkalis.

This definition of an alkali as a basic salt of an alkali metal or alkali earth metal does appear to be the most common, based on dictionary definitions [1][2], however conflicting definitions of the term alkali do exist. These include:

• Any base that is water soluble [3][4]. This is more accurately called an Arrhenius base.
• The solution of a base in water [5].

Alkali salts

Most basic salts are alkali salts, of which common examples are:

• sodium hydroxide (often called "caustic soda")
• potassium hydroxide (commonly called "potash")
• lye (generic term, for either of the previous two, or even for a mixture)
• calcium carbonate (sometimes called "free lime")
• magnesium hydroxide is an example of an atypical alkali: it is a weak base (cannot be detected by phenolphthalein) and it has low solubility in water.

Alkaline soil

Soil with a pH value higher than 7.4 is normally referred to as alkaline. This soil property can occur naturally, due to the presence of alkali salts. Although some plants do prefer slightly basic soil (including vegetables like cabbage and fodder like buffalograss), most plants prefer a mildly acidic soil (pH between 6.0 and 6.8), and alkaline soils can cause problems.

Alkali lakes

In alkali lakes (a type of salt lake), evaporation concentrates the naturally occurring alkali salts, often forming a crust of mildly basic salt across a large area.

Examples of alkali lakes:

• Redberry Lake, Saskatchewan, Canada.
• Tramping Lake, Saskatchewan, Canada.

Neutralization of acids

When dissolved in water, the base sodium hydroxide decomposes into hydroxide and sodium ions:

NaOH → Na⁺ + OH^-

and similarly, in water hydrogen chloride forms hydronium and chloride ions:

HCl + H₂O → H₃O⁺ + Cl^-

When the two solutions are mixed, the H₃O⁺ and OH⁻ ions combine to form water molecules:

H₃O⁺ + OH^- → 2 H₂O

If equal quantities of NaOH and HCl are dissolved, the base and the acid exactly neutralize, leaving only NaCl, effectively table salt, in solution. Acids have pH levels below 7

Weak bases, such as soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as lye or ammonia, can cause a violent exothermic reaction, and the base itself can cause just as much damage as the original acid spill.

Alkalinity of non-hydroxides

Both sodium carbonate and ammonia are bases, although neither of these substances contains OH⁻ groups. That is because both compounds accept H⁺ when dissolved in water:

Na₂CO₃ + H₂O → 2 Na⁺ + HCO₃^- + OH^-

NH₃ + H₂O → NH₄⁺ + OH^-

Bases as heterogeneous catalysts

Basic substances can be used as insoluble heterogeneous catalysts for chemical reactions. Examples are metal oxides such as magnesium oxide, calcium oxide, and barium oxide as well as potassium fluoride on alumina and some zeolites. A great deal of transition metals make good catalysts, many of which form basic substances. Basic catalysts have been used for hydrogenations, the migration of double bonds, in the Meerwein-Ponndorf-Verlay reduction, the Michael reaction, and many other reactions.

See also

• Acid-base reaction theories
• Acid

References

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External links

• Acid-Base equilibrium diagrams, pH calculation and titration curves simulation and analysis - freeware Link

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