Acids, bases, salts. Normal solutions. Titration.

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Acids. Acids are an important group of chemical compounds. Many are found in nature.

Examples of acids

● acetic acid (in vinegar)

● ascorbic acid (vitamin C)

● citric acid (in citrus fruit)

● tartaric acid (in grapes)

● tannic acid (in tea)

● lactic acid (in milk)

● sulfuric acid (H₂SO₄), nitric acid (HNO₃), and hydrochloric acid (HCl) — acids widely used in industry in producing metals, plastics, explosives, textiles, and dyes

Def. Acid. A compound containing hydrogen which, in water solution, forms no positive ions other than hydrogen, H⁺, ions.

Note that the H⁺ ion is a proton. Thus an acid can be viewed as a proton provider.

Def. Base. A compound containing the hydroxyl group which, when dissolved in water, forms no negative ions other than hydroxyl, OH^-, ions.

Note. A more general definition of a base is that it is any substance that can accept H⁺ ions i.e. protons.

Examples of bases

● ammonium hydroxide, NH₄OH (household ammonia)

● sodium hydroxide, NaOH (lye)

● calcium hydroxide, Ca(OH)₂ (present in limewater)

● magnesium hydroxide, Mg(OH)₂ (milk of magnesia)

Def. Neutralization. The union of hydrogen ions of an acid with hydroxyl ions of a base to form water.

Salts. The word “salt” has a specialized meaning in chemistry. Salts are the compounds other than water that are produced when an acid neutralizes a base. A salt can be defined as follows:

Def. Salt. A compound whose water solution contains a positive ion other than the hydrogen ion and a negative ion other than the hydroxyl ion.

Examples of salts

● sodium chloride, NaCl (table salt)

● calcium carbonate, CaCO₃

● barium chloride, BaCl₂

● sodium sulfate, Na₂SO₄

Def. Indicator. A substance used to show, by means of a change in color, whether a solution is an acid or a base.

Properties of acids

1. All acids contain ionizable hydrogen and a nonmetallic ion or radical.

Examples.

1) Hydrochloric acid (HCl) in water solution yields

HCl → H⁺ + Cl^-

2) Nitric acid (HNO₃) in water solution yields

HNO₃ → H⁺ + NO₃^-

3) Sulfuric acid (H₂SO₄) ionizes in two steps, depending on the amount of its dilution:

H₂SO₄ → H⁺ + HSO₄^-

HSO₄^- → H⁺ + SO₄⁼

The first step represents the ionization in a rather concentrated solution. In a dilute solution both reactions occur. In a concentrated solution acid salts containing the HSO₄^- may be produced whereas in a dilute solution regular salts containing the SO₄⁼ radical are produced.

2. All acids have a sour taste.

Examples. Lemons, limes and grapefruit are sour. Vinegar is sour.

3. Acids cause some substances to change color, allowing them to be used as acid indicators.

Example. An acid will turn blue litmus paper red.

4. Acids neutralize bases. If we add an acid to a base in the proper proportions each will destroy the properties of the other.

5. Acids act on some metals, setting the hydrogen free and producing a salt.

Example. Zn + H₂SO₄ → ZnSO₄ + H₂↑

6. Acids react with the oxides of metals forming salts and water.

Example. CuO + H₂SO₄ → CuSO₄ + H₂O

7. Acids act on carbonates liberating carbon dioxide and producing a salt and water.

Example. CaCO₃ + 2HCl →CaCl₂ + CO₂↑ + H₂O

Sulfuric acid (H₂SO₄). A dense, oily liquid with a high boiling point. Concentrated sulfuric acid contains about 95% sulfuric acid and 5% water. Ordinary dilute sulfuric acid is made by pouring 1 part of concentrated sulfuric acid into 6 parts of water. The specific gravity of concentrated sulfuric acid is 1.84.

Nitric acid (HNO₃). 100% nitric acid is a volatile liquid that is too unstable to be put on the market. Commercial concentrated nitric acid is fairly stable and contains 68% nitric acid dissolved in water. Ordinary dilute nitric acid is usually made by adding 1 part of nitric acid to 5 parts of water. It contains about 10% nitric acid. The specific gravity of concentrated nitric acid is 1.42.

Hydrochloric acid (HCl). Hydrochloric acid is made by dissolving hydrogen chloride gas in water (hydrogen chloride gas is extremely soluble in water). Concentrated hydrochloric acid contains about 38% by weight of hydrogen chloride. Ordinary dilute hydrochloric acid is made by adding 1 part of concentrated hydrochloric acid to 4 parts of water. Such a solution contains from 6% to 8% of hydrogen chloride. The specific gravity of concentrated hydrochloric acid is around 1.20.

Def. Acid anhydride. The oxide of a nonmetal which combines with water to form an acid.

Examples.

1) From the reaction

CO₂ + H₂O ⇄ H₂CO₃

we see that CO₂ is an acid anhydride of H₂CO₃.

2) From the reaction

SO₃ + H₂O ⇄ H₂SO₄

we see that SO₃ is an acid anhydride of H₂SO₄.

Two general methods for preparing acids

1. Through the reaction of water with the oxide of a metal (i.e. with an acid anhydride).

Examples.

1) We can prepare sulfurous acid by combining water with SO₂

SO₂ + H₂O ⇄ H₂SO₃

2) We can prepare sulfuric acid by combining water with SO₃

SO₃ + H₂O ⇄ H₂SO₄

2. Through the reaction of sulfuric acid with a salt. In this general method we heat a salt of the acid we wish to prepare with sulfuric acid.

Examples.

1) To prepare hydrochloric acid we heat sodium chloride and sulfuric acid

NaCl + H₂SO₄ →NaHSO₄ + HCl↑

2) To prepare nitric acid we heat sodium nitrate and sulfuric acid

NaNO₃ + H₂SO₄ →NaHSO₄ + HNO₃↑

Properties of bases

1. A base contains a metal and one or more hydroxyl (OH) groups.

Examples. NaOH, KOH, Ca(OH)²

The number of OH groups is equal to the valence of the metal. When a base dissociates, the metal of the base becomes the positive ion and the hydroxyl (OH) radical forms the negative ion.

NaOH → Na⁺ + OH^-

KOH → K⁺ + OH^-

Ca(OH)² → Ca⁺⁺ + 2 OH^-

2. Soluble bases have a bitter taste.

3. A solution of a soluble base is caustic, attacks the skin, and feels slippery.

4. Soluble bases affect indicators, such as litmus paper, causing a color change (insoluble bases generally do not).

5. Bases neutralize acids.

6. Bases react with the oxides of nonmetals forming salts and water.

Example. CO₂ + 2 NaOH → NaCO₃ + H₂O

Naming of bases. The name of a base is formed by adding to the name of the metallic ion the word hydroxide.

Examples. NaOH is sodium hydroxide. Zn(OH)₂ is zinc hydroxide. Bi(OH)₃ is bismuth hydroxide.

Def. Basic anhydride. The oxide of a metal which combines with water to form an base.

Examples. In the reactions

CaO + H₂O → Ca(OH)₂

MgO + H₂O → Mg(OH)₂

CaO and MgO seen to be basic anhydrides.

Three common methods for preparing bases

1. Some active metals react with water to form bases. Such active metals as sodium and potassium react with water with great vigor, liberating hydrogen and forming strong, active bases:

2Na + 2 HOH →2 NaOH + H₂↑

2K + 2 HOH →2 KOH + H₂↑

2. The oxides of some metals unite with water to form bases (i.e. basic anhydrides unite with water to form bases).

Example. CaO + H₂O → Ca(OH)₂

3. Insoluble bases may be formed from soluble bases by treating a soluble salt of some metal with a soluble base.

Example. If we add the soluble salt ferric chloride to the soluble base sodium hydroxide we obtain as a precipitate the insoluble base ferric hydroxide.

FeCl₃ + 3 NaOH →Fe(OH)₃↓+ 3 NaCl

Nearly all dense metals form insoluble bases.

pH scale. The pH of a solution is a measure of the acidity (or alkalinity) of the solution. Neutral solutions have a pH of exactly 7. Numbers less than 7 indicate the presence of an acid and numbers greater than 7 indicate the presence of a base. pH can be measured electrically with a pH meter or with various color change indicators.

The pH of an aqueous solution is the negative log of the activity of the hydrogen ion in the solution.

Normal solutions

Def. Normal solution. A solution containing 1 gram-equivalent weight of solute per liter of solution i.e. a solution with a normality of 1. Also termed a 1 normal or1-N solution.

Examples.

1. A normal solution of HCl would consist of 36.5 g of HCl in a liter of water since the gram-equivalent weight of HCl is 36.5 g.

2. A normal solution of K₂SO₄ would consist of 87 g of K₂SO₄ in a liter of water since the gram-equivalent weight of K₂SO₄ is 87 g (K₂SO₄ has a formula weight of 174 and a valence of 2, so 174/2 = 87).

● A 2-N solution is one with a normality of 2, a 0.1-N solution is one with a normality of 0.1, and a x-N solution is a solution with a normality of x.

● Normal acid solutions neutralize normal acid solutions milliliter for milliliter. Equal volumes of solutions of equal normality are equivalent.

Titration. Suppose a chemist would like to know the percentage of acetic acid in a particular vinegar. To find out he would titrate the vinegar against an acid of known normality using a pair of burettes in a process called titration. See Fig. 1. He proceeds as follows:

1. He fills one of the burettes with the vinegar and the other with 0.1-N base.

2. He draws off 10 ml of vinegar into an Erlenmeyer flask and dilutes it with about 150 ml of water.

3. He adds two drops of a solution of phenolphthalein which is an indicator that is colorless in an acid solution but red in a base solution.

4. He now draws off from the base burette just enough of the 0.1-N base to neutralize the acid in the vinegar, adding it drop by drop until the solution turns a pale pink.

Suppose it takes 80 ml of the 0.1-N base to neutralize the 10 ml of the vinegar. That would mean that 80 ml of base was equivalent to 10 ml of acid and so the acid must be eight times as concentrated as the base. Thus the vinegar must be 0.8-N acetic acid. If 1-N acetic acid contains 60 g of acetic acid per liter, then 0.8-N acetic acid contains 48 g (0.8 × 60) per liter. A liter of 0.8-N acetic acid weighs approximately 1000 g. In 100 g of the sample there is then 4.8 g of acetic acid and the strength is thus 4.8%.

Properties of salts. There are hundreds of salts and their properties vary widely. One of the most important properties of a salt is its solubility.

Solubility of salts in water

1. Common sodium, potassium, and ammonium compounds are soluble in water.

2. Common nitrates, acetates, and chlorates are soluble.

3. Common chlorides are soluble except for silver, mercurous, and lead.

4. Common sulfates are soluble except for calcium, barium, strontium, and lead. (Lead chloride is soluble in hot water.)

5. Common carbonates, phosphates, and silicates are insoluble except sodium, potassium, and ammonium.

6. Common sulfides are insoluble except for calcium, barium, strontium, magnesium, sodium, potassium, and ammonium.

Methods of preparing salts

1. Direct union of the elements.

Example. 2 Na + Cl₂ → 2 NaCl

2. Replacement of the hydrogen of an acid by a metal.

Example. 2 Na + 2 HCl →2 NaCl + H₂↑

3. Reaction of the oxide of a metal with an acid.

Example. Na₂O + 2HCl → 2 NaCl + H₂O

4. Reaction of the oxide of a nonmetal with an base.

Example. CO₂ + Ca(OH)₂ → CaCO₃↓ + H₂O

5. Neutralize a base with an acid to form a salt.

Example. Sodium hydroxide is added to hydrochloric acid until the solution is neutral.

NaOH + HCl → NaCl + H₂O

The solution is evaporated and sodium chloride remains.

Many different salts can be prepared by neutralization.

6. Two salts can be prepared at one time by double replacement.

Example. Add a solution of sodium sulfate to a solution of barium chloride

and two salts are formed — sodium chloride and barium sulfate.

BaCl₂ + Na₂SO₄ →2 NaCl + BaSO₄↓

Barium sulfate is insoluble and precipitates out. The water can then be evaporated to obtain crystals of sodium chloride.

Note. It is impossible to separate two salts unless one is soluble and the other is insoluble.

7. Action of an acid on a carbonate to form a salt.

Example. 2HCl + Na₂CO₃ → 2 NaCl + H₂O + CO₂↑

Naming of salts. Salt names are formed from the names of their constituent ions.

Examples. Ba(NO₃)₂ is barium nitrate. FeCl₃ is ferric chloride. FeCl₂ is ferrous chloride. BaSO₄ is barium sulfate.

Table 1 gives the names of various acids and the names of the negative ions of the salts derived from the acids.

References

Dull, Brooks, Metcalfe. Modern Chemistry.