Salt

            Salt are generally prepared by action between acid and base. Salt is a neutral substance which possesses neither acidic nor basic property nor it shows no effect on indicators such as litmus paper. Salts need high heat for melting and boiling. Electrolysis takes place in salty solution which means decomposition of an electrolyte by the passage of electric current.

            Copper oxide is a metallic oxide (a salt) which, when treated with sulphuric acid produces copper sulphate or complete replacement of hydrogen of an acid by a metal or electropositive radical.

      H₂SO₄ + NaOH = NaHSO₄ + H₂O (partial replacement)

     H₂SO₄ + 2NaOH = NaHSO₄ + 2H₂O (complete replacement)

            When salts are dissolved in water, these ionize giving positively and negatively charged ions other than hydrogen ion (H⁺) or, hydroxyl (OH^-) in solution. In some cases hydrogen and hydroxyl ions may also be present in addition to other positive or negative ions, the presence of which is essential characteristics of salt.

General methods of Preparation of salts:

1. All nitrates are soluble.

2. All salts of sodium, potassium and ammonium are soluble.

3. All carbonates are insoluble except sodium, potassium and ammonium carbonates.

4. All sulphates are soluble in water except for barium and lead sulphates which are insoluble and calcium sulphate which is only very soluble.

5. All chlorides are soluble in water except silver and lead chloride.

Some soluble salts are given below:

Chemical name	Molecular formula	Common name
Sodium chloride Sodium carbonate Sodium sulphate Magnesium sulphate	NaCl Na2CO3.10H2O Na2SO4.10H2O MgSO4.7H2O	Common salt Washing soda Glauber’s salt Epsom salt

           

Some salts are insoluble in water. They have industrial importance. Insoluble salts are used to produce different colors. Some are listed below:

Common name of color	Molecular formula	Color produced
Lead sulphate Copper carbonate Lead carbonate Zinc chromate Titanium dioxide	PbSO4 CuCO3 PbCO3 ZnCrO3 TiO2	White Green White Yellow white

Uses of Salts:

Salt	Common name	Uses
Sodium carbonate Potassium nitrate Magnesium sulphate Sodium chloride Ammonium chloride	Washing soda Saltpeter Epson salt Common salt	Softening water, making glass Fertilizer, gunpowder Medicine Cooking fertilizer

Base

Any such substance which gives hydroxyl ion (OH^-) when dissolves in wateror, give salt and water when treated with acid or neutralizes the acid is called base.

            KOH = K⁺ + OH^-

            NaOH = Na⁺ + OH^-

            Ba(OH)₂ = Ba⁺⁺ + 2OH^-

Combination of oxygen with metal results in the formation of metallic oxides. Metallic oxides are basic. E.g. copper oxide (CuO), sodium oxide (Na₂O). Some metallic oxides are soluble in water and some are insoluble in water. Metallic oxides which dissolve in water produce corresponding hydroxide which is base. 

            Na₂O + H₂O = 2NaOH

A highly soluble hydroxide is an alkali. All alkalis are bases but all bases are not alkalis.

            Some common alkalis are given below:

           

Name of Alkali	Molecular Formula
Potassium hydroxide Sodium hydroxide Calcium hydroxide Ammonium hydroxide	KOH NaOH Ca(OH)2 NH4OH

Ammonium hydroxide gives strong smell of ammonia and it is injurious to health.

Characteristics of Base:

1. They produce hydroxyl ions when dissolved in water and the strength of base depends upon degree of ionization. A base which gives more hydroxyl ions is a strong base.

2. Strong bases posses a bitter taste and soapy touch.

3. They turn red litmus into blue, turmeric brown and phenolphthalein into pink.

4. Base reacts with acids to form salt.

2KOH + H₂SO₄ = K₂SO₄ + H₂O

Preparation of Base:

            Bases can be prepared by following methods:

1. By the action of water on certain metals.

Ca + 2H₂O = 2NaOH + H₂

2. By the action of water on basic oxides.

Na₂O + H₂O = 2NaOH

3. By double decomposition.

 Na₂CO₃ + H₂O = CaCO₃ + 2NaOH

Acid

folic acid is very good and beneficial for physical

Introduction
    We use colored solution to detect acidity or basicity of substance in the lab. These substances can be prepared by mixing colored parts of plant body in ethanol (C₂H₅OH). The color in plant leave is due to the presence of pigments such as chlorophyll and chromoplast. Colored solution is used widely in chemistry to detect acidity and basicity.

Indicator             An indicator is a chemical substance which has capacity to show the termination of a chemical reaction by changing its color. The indicator does not take part in chemical reaction and does not effects the reaction showing change in color. The indicators are widely used to identify whether a substance is acid or base. Colored solution prepared from red-rose and beet-root is examples of indicator. Litmus paper is the most common in lab ot identify whether a substance is base or acid. Litmus is prepared from lichen plant.
            Blue litmus changes into red in acidic medium and red litmus changes into blue in basic medium. Phenolphthalein, methyl orange, turmeric, etc are examples of indicator.

Litmus can detect whether a solution is acid or base but it cannot show the strength of acidity or basicity. Universal indicators show the strength of an acid or base. Universal indicator is a mixture of ordinary indicators. Universal indicator is a mixture of ordinary indicators. Universal indicator changes the color itself according to the strength of an acid or base.

pH value of some substances:

Substance	pH
Water Salt solution Sugar Ethanol Lemon juice Toothpaste Ammonia solution Apple Butter Soda Milk of magnesium Washing soda Kerosene Vinegar	7 7 7 7 2.5 9 10 3 6 8.5 9 11.5 7 3

From above table we come to know that in food materials also acidic and basic property exists. Basicity is present in sodium hydrogen carbonate (baking soda) and lemon juice is acidic. Universal indicator is used to measure the pH scale of solution. This scale is called pH meter.

Acid
            An acid is a substance which gives hydrogen ion when dissolved in water. An acid neutralizes the base.
            HCl = H⁺ + Cl^-
            H₂SO₄ = 2H⁺ + SO^{- -}₄
This hydrogen of an acid can be replaced partially or completely b y metals like zinc and magnesium.
Example: Zinc + Sulphuric acid = Zinc sulphate + Hydrogen
                Zn + H₂SO₄ = ZnSO₄ + H₂

Some common acid:

Name of the acid	Molecular Formula
Hydrogenchloric acid Sulphiric acid Nitric acid Phosphoric acid Carbon acid	HCl H2SO4 HNO3 H3PO4 H2CO3

Characteristic of Acid:

1. Acids posses sour taste. 
2. Acids turn blue litmus into red.
3. Acids neutralize base.                                                                    HCl + NaOH = NaCl + H₂O  
4. Hydrogen of an acid is displaced by metals.                                    Zn + H₂SO₄ = ZnSO₄ + H₂
5.  Acid produces salt and water when treated with basic oxide        CaO + 2HCl = CaCl₂ + H₂O  
6. Many carbonates evolve carbon dioxide when treated with acids.                                                                                               E.g. CaO + 2HCl = CaCl₂ + H₂O

  Preparation of Acids

         Acids can be prepared by following methods.

1. By direct combination of elements.

H₂ + Br₂ = 2HBR

2. By action of water on acid anhydride. When acid anhydride is treated with water, corresponding acids are produced.

CO₂ + H₂O = H₂CO₃

3. By oxidation of non-metals. When nitric acid acts on phosphorous, and sulphur, phosphoric and sulphuric acids are formed.

P + 5HNO₃ = H₃PO₄ + 5NO₂ + H₂O

Balancing of Chemical Equation by Oxidation Number Method

The principal underlying this universally applicable of balancing equations is that electrical charge must be conserved in the course of chemical reaction like any increase in oxidation number must be compensated by a decrease.

               In brief, the method consists in selection of such coefficients for the oxidizing and reducing agents as will ensure to make the total decrease in oxidation number for the former equal to the total increase in oxidation number for the latter.

               Following rules can be adapted for it:

a) Assign oxidation numbers to the atoms that change.

b) Choose the proper ratio of oxidizing agent to the reducing agent so that the oxidation number change balanced,

c) Make appropriate change in the coefficient of the products corresponding to change in the coefficient of reactions in step (b).

d) Balance the oxygen atoms on both sides by adding H₂O to the side that is deficient in oxygen.

e) Balance the hydrogen atoms on both sides by adding H⁺ to the side that is deficient in oxygen.

f) The equation is balanced if the reaction is taking place in acidic solution. If the reaction proceeds in basic solution, and sufficient number of OH^- to get rid of H⁺ but add equal number of OH^- on both sides.

g) Cancel any duplication that might have crept in on both sides of the equation.

The following examples will cleared the method:

Example: Oxidation of ammonia by copper oxide to give copper, nitrogen and water

a) Writing skeleton equation with oxidation number of copper and nitrogen

                          Oxidation numbers      _{+2         -3           0    0}

       Skeleton                  CuO + NH₃ =   Cu + N₂ + H₂O

b)  Oxidation number of Cu change from +2 to 0 and that of N changes from -3 to 0. The Cu atom is going from +2 to 0 have to gain two units of negative charge. The nitrogen atom is going from -3 to 0 has to get rid 3 units of negative charge.

c) To balance the oxidation number change multiplies CuO by 3 and NH₃ by 2. Then provide same number of Cu and N atoms on the right. The equation will be

          3CuO + 2NH₃ = 3Cu + N₂ + H₂O.

d) Balancing oxygen by multiplying H₂O by 3, we get the required balanced equation.

   3CuO + 2NH₃ = 3Cu + N₂ + 3H₂O.

Covalency, Oxidation State and Oxidation Number

   

Covalency of an element is defined as a number indicating its combining capacity. For example, it is represents the number of hydrogen atoms which can combine with a given atom. It also represents the number of single bonds which an atom can form. It is also defined as the number of electrons its atom is able to share. In any case covalency is a pure number and has no plus or minus sign associated with it.

            Oxidation number is defined as the charge which an atom appears to have when electrons are counted. It is positive or negative. For example, in ammonia the covalency of nitrogen is three but its oxidation number is -3.

            In ionic compounds the oxidation state of an element is the same as the charge on the ion formed from an atom of the element. For example, in potassium bromide potassium is said to be in the +1 oxidation state and bromine in -1 oxidation state. Is ionizes as

                                    KBr = K⁺ + Br

In other words, oxidation numbers of potassium and bromide are +1 and -1 respectively.

Oxidation state of Aluminium in Al2O₃ is +3 and the total oxidation number of two aluminium atoms is +6. Thus oxidation number of two aluminium atoms is +6. Thus oxidation state of an element is its oxidation number pre atom

 There may actually be a difference between the magnitude of covalency and the oxidation number. Example :

    CH₄ (Methane)                            CH₃Cl (Methylchloride)        

   CH₂Cl₂ (Methylene chloride)      CHCl₃ (Chloroform)         

    CCl₄ (Carbon tetrachloride)

In each case one atom of carbon shares a total of 4 pairs of electrons with other atoms. Carbon atom is, therefore, tetracovalent in each case.

            Oxidation number for carbon in CH₄, CH₃Cl, CH₂Cl_2, CHCl₃ and CCl₄   is -4, -2, 0, +2 and +4 respectively.

            Thus, while covalency of carbon remains constant (=4) in each case, its oxidation number varies from -4 and +4.

Oxidation and Reduction in term of Oxidation Numbers.

            The term oxidation refers to any chemical change involving increase in oxidation number where as the term reduction applies to any chemical change involving decrease in oxidation number.

            Consider the following chemical changes:

                        2H₂ + O₂ = 2H₂O

Here, oxidation number of hydrogen changes from 0 (H₂) to +1 (+H₂O). It is therefore, a case of oxidation of hydrogen.

The oxidation number of oxygen decreases from 0 (O₂) to -2 (H₂O). It is therefore, a case of reduction of oxygen.

In the same reaction, oxidation number of hydrogen increases and that of oxygen decreases, i.e. hydrogen undergoes oxidation while oxygen undergoes reduction. Thus oxidation and reduction occur together.

Oxidation-Reduction Reaction

Oxidation Number

 Oxidation is any reaction in which an atom or ion loses electrons. On the other hand, reduction is defined as any reaction in which an atom or ion gains electrons. As oxidation involves removal of electrons, it has also been termed de-electronation. Similarly reduction may be referred to as electronation.

Change of ferrous (Fe²⁺) to ferric (Fe³⁺), stannous (Sn²⁺) to stannic (Sn⁴⁺), manganate (MnO^2-₄) to permanganate (MnO^-₄) are all cases of oxidation as each one of them involves loss of electrons.

            Fe²⁺ = Fe³⁺ + 1 electron

            Sn²⁺ = Sn⁴⁺ + 2 electron

            MnO^2-₄ = MnO^-_{4 +} 1 electron

 Similarly formation of mercurous ions (Hg₂²⁺) from mercuric ions (Hg₂²⁺) and formation of chlorine ions (Cl^-) from chlorine atoms are cases of reduction as both involve a gain of electrons.

            Hg₂²⁺ + 2 electrons = Hg₂²⁺

                    Cl + 1 electron = Cl^-

 In order to keep track of electron shifts in oxidation-reduction reaction, it is convenient to use the concept of oxidation state of various atoms involved in these reactions. The oxidation number is defined as the formal charge which an atom appears to have when electrons are counted in accordance with the following rules:

1.  Electrons shared between two unlike atoms are counted with more electronegative atom. For example, the electron pair shared between H and Cl in H⁺¹:Cl^-1 is counted with more electronegative Cl. As a result of it hydrogen having lost share in the electron pair appears to have +1 charge and chlorine appears to have -1 charge. Hence oxidation numbers of H and Cl are +1 and -1 respectively. 
2. Electrons shared between two like atoms are divided equally between the two sharing atoms. For example, in hydrogen molecule, H:H, the electron pair is equally shared between the two atoms. Thus both the atoms appears to have no charge i.e. oxidation number of hydrogen is zero in hydrogen molecule.

In the molecule of water given in the margin, the two electron pairs shared between oxygen and the two hydrogen atoms are counted with the more electronegative atom. Hence in water, oxidation number of each H is +1 and that of the O atom is -2.

H⁺¹: O^-2: H⁺¹

Counting of electron like this is very hard. The following operation rules derived from the above will be very convenient:

1) In the elementary or uncombined state, the atoms are assigned an oxidation number zero.

2) In compounds, the oxidation number of fluorine is always -1.

3)     In compounds, the group IA elements (Li, Na, K, Rb, Cs and Fr) have an oxidation number +1 and the group IIA elements (Be, Mg, Ca, Sr, Ba, and Ra) have an oxidation number +2.

4) Oxidation number of hydrogen in compounds is generally +1 except in metallic hydrides wherein its oxidation number is -1.

5) In compounds, the oxidation number of oxygen is generally -2 except in FeO wherein oxidation number of fluorine is -1 and that of oxygen is +2. In hydrogen peroxide molecule the electron pair shared between O and H is counted with O but the other electron pair shared between two O atoms is equally shared. The number of electrons counted with each O is, therefore, seven (i.e. one more than its own electrons). The oxygen atom, therefore, appears to have -1 charge or its oxidation number in H₂O₂ is -1.

6) In neutral molecules, the sum of oxidation numbers of all the atoms is zero.

7) For complex ions (charged species), the sum of the oxidation numbers of all the atoms is equal to the net charge on the ion.

With the help of above rules, we can calculate the oxidation number of an atom present in a molecule or complex ion.

Example: What is the oxidation number of S in (a) H₂SO₄  (b) H₂S₂O₇ and (c) Na₂S₂O₃?

a) Let the oxidation number of S be x.

Sum of oxidation numbers of various atoms in H2SO4

= 2× (+1) + x + 4× (-2)

= 2+x-8 = x-6

This sum must be zero (rule 6). Hence

X - 6 = 0

x = 6

Or Oxidation number of S in H₂SO₄ = +6

b) Sum of oxidation numbers of various atoms in H₂S₂O₇

₌ 2 × (+1) + 2x + 7 × (-2)

= 2 + 2x – 14 = 2x – 2

Putting 2x – 12 = 0 as above, we have

 X = +6.

 Oxidation number of S in H₂S₂O₇ = +6

Ionic Equation

Molecular and Ionic equations

  On adding silver nitrate solution to a sodium chloride solution we get a white precipitate of silver chloride. Sodium chloride (a salt) is produced on mixing sodium hydroxide and hydrochloric acid in equivalent proportion. These reactions are usually represented by the following chemical equations:

NaCl +AgNO₃ = AgCl +NaNO₃

NaOH + HCL = NaCl + H₂O

These equations are termed molecular equation. If all strong bases and acids and most salts are completely ionized, it follows that molecular equations given above are not strictly appropriate when no molecules are actually present. Chemists differ somewhat in their views on writing such equations. Some argue that since it is simpler to write molecular equations, there is no harm in writing such equations in the molecular form as long as the chemist is fully aware of the ionic character of the reacting substances. Others insist on expressing highly ionized compounds only in the form of ionic symbols which is obviously more exact. It is of course essential to know the ionic or non-ionic nature of the compounds in advance.

 Some rules that should follow while writing ionic equations are given below:

a) All strong electrolytes are expressed in ionic symbols if they are soluble in water and all weak electrolytes and covalent substances are expressed in the molecular form.

b)     Any electrolyte which is highly insoluble in water is generally written in the molecular form to indicate its insolubility. Actually, this is incorrect since many such solids are ionic in crystal form as well.

c)  In addition to the atoms which must balance on both sides of the equation, the ionic charges must also balance.

Example: Potassium chloride and silver nitrate react to produce potassium nitrate (soluble) and silver chloride (insoluble).

Molecular: KCl + AgNO₃ = AgCl + KNO₃

Ionic:  K⁺ + Cl^- + NO^-₃ = AgCl + K⁺ + NO^-₃

The Ion-Electron method of balancing Equation for Oxidation-Reduction Method(by use of Half reactions)

            In this method the reaction is split up into two half-reactions. In one half-reaction the oxidizing agent picks up electrons and gets reduced, and in the other reducing agent is oxidized by supplying electrons. The two half-reaction are balanced separately and added in such a way that the electrons on the left of one and on the right of the other cancel out.

            This produced seems to imply that electrons are produced in redox reaction and travel the solution from the reducing agent to the other species which is reduced. This is not true.

Following steps should be followed while balancing the equation by this method:

a)  Separate the oxidizing and reducing agents.

b)  Write down one half-equation for the oxidizing agent changing into its reduced form.

c)  Write down the other half-equation showing the reducing agent changing into its oxidized form.

d)  Balance the atoms other than H and O for each half-reaction by adjusting coefficients, if necessary.

e)  Balance the oxygen atoms on the two sides by adding H₂O to the side deficient in oxygen.

f)  Balance the hydrogen atoms on the two sides by adding H⁺ to the side deficient in hydrogen.

g) Equalize the charge on both sides by adding electrons (e^-) to the side deficient in negative charge.

h)  If the reaction proceeds in basic solution, add enough OH^- on both side of the half-reaction to get rid of H⁺ appearing there. Combine H⁺ and OH^- to give H₂O and remove H₂O duplication.

i)  Add these two balanced half-reactions in such a way that the electrons appearing on the right of one half-reaction and on the left of the other cancel. For this each half-reaction will be multiply by some number before addition.

Example: write balanced equation for the oxidation of ferrous to ferric ion by dichromate ion in acid solution. The dichromate ion by dichromate yields cr³⁺.

For one half-Reaction

a) Writing down the reactant and product of half-reaction for the oxidizing agent changing into its reduced form:

       Cr₂O₇^2- = Cr³⁺

b) Balancing the atoms other than oxygen:

       Cr₂O₇^2- = 2Cr³⁺

c) Balancing oxygen atoms by adding 7H₂O on the right:

       Cr₂O₇^2- = 2Cr³⁺ + 7H₂O

d) Balancing hydrogen atoms by adding 14H⁺ on the left:

       Cr₂O₇^2- + 14H⁺ = 2Cr³⁺ + 7H₂O

e) Equalizing the charge by adding 6 electrons on the left:

             Cr₂O₇^2- + 14H⁺ + 6e^- = 2Cr³⁺ + 7H₂O……. (1)

For the other Half-Reaction

a) Writing down the reactant and product of the half-reaction for the reducing agent changing into its oxidized form.

b) Equalizing the charge by adding one electron on the right:

          Fe²⁺ = Fe³⁺ + e^-…………… (2)

Adding two half-reactions

    Multiplying the half-reaction (2) by 6 and adding to the half-reaction (1)

                     Cr₂O₇^2- + 14H⁺ + 6e^- = 2Cr³⁺ + 7H₂O

                           6Fe²⁺ = 6Fe³⁺ +6 e^-

               Cr₂O₇^2- + 6Fe²⁺+ 14H⁺ = 2Cr³⁺+6Fe³⁺ + 7H₂O