Category: 2. Bases Properties and Examples

In the chemical properties of acids and bases, we now focus on bases.

• Bases change the colour of litmus from red to blue.
• They are bitter in taste.
• Bases lose their basicity when mixed with acids.
• Bases react with acids to form salt and water. This process is called Neutralisation Reaction(Read).
• They can conduct electricity.
• Bases feel slippery or soapy.
• Some bases are great conductors of electricity.
• Bases like sodium hydroxide, potassium hydroxide, etc are used as electrolytes.
• Alkalis are bases that produce hydroxyl ions (OH-) when mixed with water.
• Strong alkalis are highly corrosive in nature whereas other alkalis are mildly corrosive.
• The pH value of bases ranges from 8-14.
• Alkalis and ammonium salts produce ammonia.
• Hydrogen gas is evolved when metals react with a base.
• Bases are classified on the basis of strength, concentration and acidity.
• The different kinds of acids are strong base acid, weak base acid, concentrated base, dilute base, monoacidic base, diacidic base and triacidic base.

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)           K_a(NH₄⁺) = 5.6 x 10-10F^–(aq) + H₂O(l) ↔ HF(aq) + OH^–(aq)                   K_b(F^–) = 1.4 x 10-11

Since K_a(NH₄⁺) > K_b (F^–), the reaction of ammonia with water is more favorable. Therefore, the resulting solution is slightly acidic.

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.

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

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         wateracid             base

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

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)

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)     Then, 

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.

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.

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 (i.e. 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.