Formal charge

 Formal charge- defined as the difference between the no of valence electrons of that atom in an isolated of free state and the no of electron assigned to that atom in the lewis structure

= no. of valence electron in the free atom – no of non-bonding electrons (lone pair)- (no of bonding or shared electrons)/2

Charge on the molecule or ion = sum of all the formal charges

Significance-

·       It do not indicate real charge separation within the molecules. Indicating the charge on the atoms in the lewis structure helps in keeping track of the valence electrons in the molecule

·       Helpful in selection in lowest energy structure from a no of possible lewis structure for a given species

Lewis symbol and rule of Lewis dot structure -why does atom combine with other atom?

 Chemical bond- 

·       A force that acts between two or more atoms to hold them together as a stable molecule. 

·       This process accompanied by decrease in energy.  

·       Decrease in potential energy (P.E.) is proportional to Strength of the bond.  

·       Therefore, molecules are more stable than atoms. 

Cause of Chemical Combination  

(A) Tendency to acquire minimum energy  

(a) When two atoms approach to each other. Nucleus of one atom attracts the electron of another atom.  

(b) If net result is attraction, the total energy of the molecule decreases and a chemical bond form. 

 (d) So, Attraction is inversely proportional 1/energy and proportional to Stability.  

(e) Bond formation is an exothermic process during formation of a molecule by combining of molecules 

(B) Tendency to acquire noble gas configuration:  

(a) Atom combines to acquire noble gas configuration.  

(b) Only outermost electrons or valence electrons i.e. ns, np and (n-1) d shells electrons participate in bond formation or chemical reaction. 

(c) Inert gas elements do not participate in bond formation, since they have stable electronic configuration as 1s² or ns²np⁶) 

Lewis Symbols  

A Lewis symbol is a convenient way to represent the valence electrons, which are shown as dots placed on the sides, top, or bottom of the symbol for the element.  

Na

Lewis Structures  

·       The Lewis structure is a representation of a molecule that shows how the valence  electrons are arranged among the atoms in the molecule. 

·       These representations are named after G. N. Lewis, who conceived the idea while lecturing to a class of general chemistry students in 1902. 

·       Lewis structures are based on achieving a noble gas electron configuration by atoms 

·       Writing Lewis structures is a trial-and-error process. 

·       The following steps are followed in constructing dot formulae for molecules and polyatomic ions:  

      (i) Write a symmetrical 'skeleton' for the molecules and poly atomic ions. 

                   a. The least electronegative element is usually taken as the central                                                                   element except Hydrogen. 

                   b. Oxygen atoms do not bond to each other except in O₂, 0 ₃, peroxide and superoxides. 

    c.       Hydrogen actually bonds to an oxygen atom and not to the central atom in ternary acids (oxyacids) ex-Nitrous acid HN02, has the skeleton H-O-N=O 

(ii) Calculate the number of electrons available in the valency shell of all the atoms (sum of the valence electrons for a molecule.) this is expressed by A  

         For negatively charged ions add to the total number of electrons equal to the charge on the anion and for positively charged ions, subtract the number of electrons equal to the charge on the cations 

Example – 

1). Sum of valence electrons for PO₄^3-  = 1 x 5 (for P atom) + 4 x 6 (for 0 atoms) + 3 (for charge)   

                      = 5 + 24 + 3  

                      = 32 electrons.

  

  2). Sum of valence electrons for NH₄⁺ ion  = 1 x 5 (for nitrogen atom) + 4 x 1 (for H                                      atoms) 1 (for positive charge)   

                                                                      = 5+4-1 

                                                                      =  8 electrons. 

  iii). Calculate the total number of electrons needed by all atoms to achieve noble gas                                       configuration.     

                                      This number is represented by N. 

                                         For example: N for  H2S04 

                                          N =2x2+8xl+8x4 =4+8+32 = 44 electrons. 

  iv).  Calculate the total number of electrons shared. This is represented by S which is equal to

                 N - A. For example: S for H2SO4 .

                                                S = N-A = 44- 32 =12 electrons. 

   v) Place the shared pair of electrons into the skeleton, using double and triple bonds only when                   necessary. 

  vi) Place the additional unshared (lone) pairs of electrons to fill the octet to every atom except                        hydrogen which can have only 2 electrons as the total comes equal to A 

 

States of matter intermolecular forces Part -1

 

States of matter

There are two types of properties shown by individual or the collection of a large number of particles

A)        1)   Individual/atomic properties- are shown by single particle or atom of a matter

Examples- atomic size, atomic/ionic radii, ionisation enthalpy, electron gain enthalpy, electronegativity, electro positivity

B)      2). Bulk properties-

                           those properties which are associated with the collection of a large number of individual particles are called bulk properties.

Example - Melting, boiling, solubility, water having wetting properties

Water exist in three states of matter as

                                                                   Solid -     Ice

                                                                  liquid – water

                                                                  Gas - water vapour

Triple point- the state at which all three phases exists simultaneously is called triple point

i.e. there are three states of matter as         a) solid   b) liquid     c) gas

on the basis of shape, volume, rigidity, fluidity, interparticle spaces, interparticle force, K.E. and Diffusion  

Mainly, the states of matter depend upon –

                                                       1). Energies of molecules

                                                       2). On the manner in which the individual molecules bind/aggregate

Note -Chemical properties of a substance do not change with the change of its physical state but the rate of chemical reaction do depend upon the physical state.

          Therefore  ,it become necessary for a chemist to know the knowledge of state of matter and the physical laws which govern the behaviour of matter in different states

Intermolecular forces-

·       The forces of attraction and repulsion between interacting particles (atoms or molecules) are called intermolecular forces.

·       These exist all matters ( solid liquid gas)  and responsible for many structural feature and physical properties of matter.

·       this term does not involved the electrostatic/ionic force and covalent bonds

·       intermolecular forces are weaker forces.

·       The attrative intermolecular forces are known as Vander Waals forces, in honour of Dutch scientists Johannes who studied the effect of these molecular forces on the real gases

Intermolecular forces  arise due to the following type of interactions:-

1). Dispersion forces or London force or instantaneous dipole – induced dipole forces

      ·            This force of attraction was firstly proposed by the German physicist Fritz London.

      ·           Dispersion forces are weaker intermolecular forces

      ·          These forces exert among the atoms(noble gases) or molecules(H₂ O₂ N₂ CO2  which are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed.

       ·         Temporary dipoles can occur in non- polar molecules when the electrons that constantly orbit the nucleus occupy a similar location or unsymmetrical that means the charge cloud become more on one side than the other  by chance(momentarily).it is also called instantaneous dipole.

                     This instantaneous or transient dipole distorts the electron density of the other atom or molecule.

    i.e. Temporary dipoles can induce a dipole in neighbouring molecules, initiating an attraction called London force.

·       These forces are always attractive and interaction energy is proportional to 1

2. Dipole- induced dipole interaction·       This force of attraction was firstly proposed by the German physicist Fritz London.

·       Dispersion forces are weaker intermolecular forces

·       These forces exert among the atoms(noble gases) or molecules(H₂ O₂ N₂ CO2  which are electrically symmetrical and have no dipole moment because their electronic charge cloud is symmetrically distributed.

·       Temporary dipoles can occur in non- polar molecules when the electrons that constantly orbit the nucleus occupy a similar location or unsymmetrical that means the charge cloud become more on one side than the other  by chance(momentarily).it is also called instantaneous dipole.

·       This instantaneous or transient dipole distorts the electron density of the other atom or molecule.

i.e. Temporary dipoles can induce a dipole in neighbouring molecules, initiating an attraction called London force.

·       These forces are always attractive and interaction energy is proportional to 1/r⁶

2. Dipole- diploe (Keesom) force-

     ·       This interaction was first studied by Kessom in 1912.

     ·       These forces occur in molecules which have permanent electric dipole such as HCl, NH3, H2O etc.

     ·       A polar molecule has separate centre of positive and negative charges and the ends of dipoles possess “partial charge” (δ=delta=partial)

    ·       The partial charges are always less than the unit electronics charge (1.6 X 10^-19C).

    ·       These forces arise due to interaction between oppositely charged ends of the polar molecules.

    ·       The positive end of one molecule attract the negative end of the other molecule and vice versa

     ·       Greater the dipole moment of the molecules, greater is the forces of attraction.

     ·       It is stronger than London force

      ·       Interaction energy is proportional to 1/r³ (between stationary polar molecules and is proportional to 1/r⁶

   Besides dipole-dipole interaction, polar molecule can interact by London forces also.as a result the net intermolecular forces in polar molecules increase.

4. Hydrogen Bond

Salt analysis chemistry practical part -1

SYSTEMATIC PROCEDURE FOR INORGANIC QUALITATIVE ANALYSIS  

I. PRELIMINARY TESTS 

experiment	observation	interference
1. Colour of the substance is noted.           2. A little of the salt is heated in a dry test tube                   3. FLAME TEST: -a little of the salt is made into a paste with one or two drops of con.HCl in a watch glass. A little of the paste is shown in Nonluminous flame by platinum wire	a) Blue  b) Pink   c)Green   d)Colourless       a) A colourless gas which turned lime water milky   b) A colourless gas with burnt smell of sulphur   c) Colourless gas with pungent smell. Turns Red litmus paper blue.  d) Reddish brown vapours turning acidified ferrous sulphate paper brown   a) Blue colour     b) Pale green     c)Brick Red     d)Crimson Red	a) May be due to presence of Copper.  b) May be due to presence of Manganese.   c) May be due to presence of Nickel.   d) Absence of Copper, Nickel, Manganese.  a) Presence of Carbonate.     b) Presence of Sulphide.     c) Presence of Ammonium salt.       d) Presence of Nitrate.       a) Presence of Copper     b) Presence of Barium.     c) Presence of Calcium.     d) Presence of Strontium.

 
ANALYSIS OF ANIONS  

The scheme of analysis of anions consists of identification and confirmatory tests. Once the presence of radical is detected by identification tests, proceed to do the confirmatory tests of that ion. 

IDENTIFICATION TESTS
1.ACTION OF Dil.HCl -To a little of the substance about 1ml of dil.HCl is added slowly and heated.           2. ACTION OF Dil.H2SO4 a little of the substance is rubbed with Dil. H2SO4 in a watch glass     3. Action of Con. H2SO4-   To a small amount of given salt taken in a dry test tube,  add 2-3 ml of Con. H2SO4 and gently warmed.	a) Vigorous or brisk effervescence with evolution of colourless, odorless gas   b) Colourless gas with smell of rotten eggs.   c)No characteristic observation       a) smell of vinegar     b) No characteristic observation   a) Reddish brown vapours turning moist fluorescent paper red.   b) Colourless gas with pungent smell giving dense white fumes with a glass rod dipped in NH4OH solution is shown in the mouth of test tube.   c) Violet coloured vapours turning starch paper blue or violet.   d) No characteristic observation	a) Presence of Carbonate ion is identified.   b) Presence of Sulphide ion is identified.   c) Absence of carbonate & sulphide.       a) Presence of acetate ion is identified.   b) Absence of acetate     a) Anion may be bromide ion is identified.    b) Anion may be chloride ion is identified.         c)Anion may be Iodide ion is identified.     d) Absence of bromide, chloride, iodide
4. To a small amount of given salt taken in a dry test tube, add 2-3 ml of Con. H2SO4 and heated strongly. To boiling solution add paper ball or Copper tunings is added.     5. Action of Barium chloride -To the little of salt solution a little of barium chloride solution is added.	a) Reddish brown vapours turning acidified ferrous sulphate paper brown.     b) No characteristic observation     a) A white precipitate insoluble in dil.HCl is formed.     b) No characteristic observation	a) Presence of Nitrate is identified.       b) Absence of Nitrate.       a) Presence of sulphate is identified     b) Absence of Sulphate.