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Chemical Bonding

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Title: Chemical Bonding


1
Chemical Bonding
2
Introduction
3
Introduction
4
Introduction
  • The Law of Definite Proportions (Joseph Louis
    Proust, 1799) and Daltons development of atomic
    theory (1803) lead to the recognition that atoms
    of an element had a characteristic combining
    ability with other atoms, which came to be called
    valence.
  • Existence of atoms suggested that compounds were
    composed of collections of atoms bound together
    by chemical bonds.

5
  • Atoms combine to form molecules.
  • The combining power of atoms to form molecules is
    called valency.
  • All atoms having unstable or incomplete outer
    shell have a tendency to gain or lose electrons ?
    tendency of atoms to complete and hence stabilize
    their outermost orbit of electrons which is
    mainly responsible for chemical combination
    between the atoms.

6
  • According to electronic theory of valency, a
    chemical bond is formed as a result of electronic
    interactions.
  • However, a molecule is formed only when electrons
    of the constituent atoms interact in such a way
    that the potential energy is lowered.

7
  • A bond as a force that holds groups of two or
    more atoms together and makes them function as a
    unit.
  • A chemical bond as an effect that causes the
    energy of two atoms close together to be markedly
    lower (by about 100 kJ per mole or more) than
    when they are far apart.
  • The forces that hold bonded atoms together are
    basically just the same kinds electrostatic
    attractions that bind the electrons of an atom to
    its positively-charged nucleus.

8
Lewis electron dot formulas
  • Lewis originated the idea of the electron pair
    bond.
  • In 1902, Lewis developed the concept of valence
    electrons and realized that all elements known to
    form simple ions by losing or gaining whatever
    number of electrons is needed to leave eight in
    the valence shell of each.
  • In 1916, Lewis published shared electron-pair
    theory.

9
  • Lewis Symbol/Structure --- representing atom
    singly or in combination
  • Only the valence electrons are shown.

10
  • Atoms having a tendency to lose electrons are
    called electropositive whereas the atoms which
    gain electrons are called electronegative.
  • The number of electrons gained or lost by an atom
    in order to acquire an inert gas configuration
    gives numerical value of the electrovalency of
    the atom.
  • A Lewis dot symbol consists of the symbol of an
    element and one dot for each valence electron in
    an atom of the element.

11
Type of Bonds
  • Electrovalent or ionic bond
  • Electropositive elements Electronegative
    elements
  • Covalent bond
  • Electronegative elements Electronegative
    elements
  • Coordinate bond
  • Electropositive elements Electropositive
    elements

12
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13
Ionic or Electrovalent Bonds
  • The ionic bond is formed due to the
    electrostatic attraction between stable ions
    formed by the complete transfer of electrons from
    one atom to another.
  • The atom which loses electrons acquires a
    positive charge, whereas the one which gains
    electrons becomes negatively charged.
  • Ionic bonds anions cations

14
Ionic Bonds
  • These charged atoms are called ions and held
    together by electrostatic attraction forces.
  • Such a mode of combination of atoms is called
    electrovalency and the bond formed between the
    atom is called electrovalent or ionic bond.
  • The compound thus formed is called an
    electrovalent or ionic compound.

15
Ionic Bonds
  • Ionic substances are formed when an atom that
    loses electrons relatively easily reacts with an
    atom that has a high affinity for electrons.
  • In other words, an ionic compound results when a
    metal reacts with a nonmetal.

16
Ionic Bonds
  • Example 1 NaCL
  • The electronic arrangement of Na and Cl atoms
    are
  • Na (11) ? (1s2, 2s2, 2p6, 3s1)
  • Cl (17) ? (1s2, 2s2, 2p6, 3s2, 3p5)
  • The electron from the outermost orbit of Na is
    completely transferred to the outermost orbit of
    Cl atom. As a result of this transfer, both atoms
    acquire inert gas structure.

17
Ionic Bonds
  • The Na becomes Na (2, 8) and Cl become Cl (2,
    8, 8).

18
  • Example 2
  • The formation of lithium flouride (LiF)

Li

(or LiF)
1s2 2s2 2p6
1s2
Li
Li
e -

e -

19
Ionic Bonds
  • Electrovalent compounds exhibit following
    properties
  • Electrovalent compounds are generally hard
    solids.
  • They have high melting and boiling point.
  • Electrovalent compounds are generally sparingly
    soluble in organic solvents.
  • Electrovalent compounds in solid state are poor
    conductors of electricity. But when dissolved in
    solvents of relatively high dielectric constant,
    they exhibit a strong electrical conductivity.
    They also conduct electricity in the molten
    state.

20
Lattice energy of ionic compounds
  • The stability of solid ionic compounds depends on
    the interactions of all cations and anions.
  • A quantitative measure of the stability of any
    ionic solid is its lattice energy --- the energy
    required to completely separate one mole of a
    solid ionic compound into gaseous state.

21
Covalent Bond
  • Formation of molecules by the sharing of
    electrons between combining atoms is called
    covalency and the bond formed is called covalent
    bond or covalent linkage.
  • Compounds containing this type of linkage are
    called covalent compounds.
  • In covalent bond formation the inert gas
    configuration of the two concerned atoms is
    achieved by sharing equal number of electrons.

22
Covalent Bond
  • The sharing of electrons to form chemical bond
    between two atoms is described by showing pairs
    of electrons between the bonded atoms.
  • If one pair of electrons is shared, the bond
    formed is called single bond, whereas sharing of
    two or three pairs of electrons leads to the
    formation of double or triple bonds respectively.
  • Due to sharing of electrons, both the bonding
    atoms acquire inert gas configuration.

23
Covalent Bond
  • Example 1 the formation of hydrogen molecule
    from 2 hydrogen atoms.
  • The covalent bond is represented by ().

24
Covalent Bond
  • Example 2 the formation of CCl4
  • Example 3 the formation of CH4

25
Some other example of the formation of covalent
bonds
  • Single bond
  • Double bonds

26
  • Triple bonds
  • In covalent compounds the numerical value of
    covalency of any element in the molecule is the
    number of electron pairs shared between the
    atoms. Thus the valency of hydrogen in H2 is one,
    oxygen in O2 is two, nitrogen in N2 is three and
    carbon in CH4 is four.

27
  • Two types of covalent bonds
  • Polar covalent bonds
  • When a bond is formed between unlike atoms, the
    bonding electrons will not be equally shared.
    Example HCl
  • The shared electrons will be shifted more
    towards the atom having higher electronegativity
    and this will result in the accumulation of a
    negative charge on it. The other atom will carry
    an equivalent positive charge.

28
  • Example HCl
  • The chlorine atom acquires small amount of ve
    charge because of its higher electronegativity
    and hydrogen has an equivalent ve charge. The d
    and d represents respectively the small ve and
    ve charge.
  • Electronegativity
  • The ability of an atom to attract toward itself
    the electron in a chemical bond.

29
  • Non-polar covalent bonds.
  • When a bond is formed between atoms of the same
    element, the bonding electrons are equally shared
    on account of equal electronegativity of the
    atoms.
  • In case of such a bond, the centre of ve charge
    coincides with the centre of ve charge in the
    molecule.
  • For example, bonds involved in the formation of
    H2, Cl2, O2, N2 etc. are non-polar bonds.

30
  • Covalent compounds possess following general
    characteristics
  • Covalent compounds possess definite geometrical
    shapes. They exhibit isomerism because covalent
    bonds are rigid and possess directional
    characteristics.
  • Covalent compounds are mostly liquids and gases.
    The solid compounds are generally volatile.
  • They are highly soluble in organic solvents but
    slightly soluble in water. Some compounds like
    HCl and NH3 readily dissolve in water because
    they react with water.

31
  • The melting and boiling points of covalent
    compounds are relatively low because the forces
    involved in covalent compounds are less strong
    than those involved in electrovalent or ionic
    compounds.
  • These compounds do not contain ions. Therefore,
    when dissolved, they do not conduct electricity.
    They even do not conduct electricity in the
    molten state.

32
Ionic vs Covalent Compounds
33
Coordinate Bonds
  • When, an atom having a complete octet donates a
    pair of free valence electrons to another atom
    which is short of two electrons, the resulting
    bond is known as coordinate bond. Thus both the
    atoms acquire inert gas configurations.
  • The atom which donates a pair of electrons is
    called donor and the other atom which accepts
    the electrons is called acceptor. The
    coordinate bond is similar to covalent bond
    except that both the shared electrons are donated
    by one atom.

34
  • The formation of covalent and coordinate bond is
    illustrated below
  • When one atom furnishes both electrons for the
    formation of a covalent bond as described above,
    the process is called coordination.

35
  • Since one atom donates an electron pair and the
    other accepts, the molecule acquires polarity.
    These bonds, therefore, are also known as semi
    polar bonds or dative bond.
  • Coordinate bond is represented by the symbol (?).
  • Example

36
  • Following are the main characteristics of
    coordination compounds
  • Coordinate compounds, like covalent compounds,
    exhibit space isomerism. This is due to
    directional characteristics possessed by
    coordinate linkage.
  • The B.P. and M.P. of these compounds have
    intermediate value between electrovalent and
    covalent compounds.
  • They are only slightly soluble in water and most
    of them are soluble in organic solvents.

37
Lewis Structure
  • The Lewis structure is a representation of a
    molecule that shows how the valence electrons are
    arranged among the atoms in the molecule.

38
Writing Lewis Structure
  • Write the skeletal structure of the compound
    showing what atoms are bonded to what other
    atoms.
  • Count the total number of valence electrons
    present, referring, if necessary to Lewis dot
    symbols.
  • Draw a single covalent bond between the central
    atom and each of the surrounding atoms. Complete
    the octets of the atoms bonded to the central
    atom.
  • If the octet rule is not met for the central
    atom, try double or triple bonds between the
    surrounding atoms and the central atom, using the
    lone pairs from the surrounding atoms.

39
  • Example
  • Write the Lewis structure of nitric acid, HNO3,
    in which the three O atoms are bonded to the
    central N atom and the ionizable H atom is bonded
    to one of the O atoms.
  • Answer
  • Step 1 The skeletal structure of HNO3 is
  • O N O H
  • O

40
  • Step 2 The outer-shell electron configurations
    of N, O, and H are
  • N 2s2 2p3
  • O 2s2 2p4
  • H 1s1
  • Thus, there are 5 (3 x 6) 1 24, valence
    electrons to account for in HNO3.

41
  • Step 3 Draw a single covalent bond between N and
    each of the three O atoms and between one O atom
    and the H atom. Then we fill in electrons to
    comply with the octet rule for the O atoms.
  • O N O H
  • O

42
  • Step 4 We see that this structure satisfies the
    octet rule for all the O atoms but not for the N
    atom. Therefore we move a ion pair from one of
    the end O atoms to form another bond with N.
  • O N O H
  • O

43
Formal Charge Lewis Structure
  • Electrons are shared in a bond, we must divide
    the electrons in a bonding pair equally between
    the atoms forming the bond. The difference
    between the valence electrons in an isolated
    atom and the number of electrons assigned to that
    atom in a Lewis structure is called that atoms
    formal charge.
  • The equation for calculating the formal charge on
    an atom in a molecule is given by
  • FC total number of valence electrons in the
    free atom total number of nonbonding electrons
    - ½ (total number of bonding electrons)

44
  • Example
  • Write formal charges of the carbonate ion!
  • Answer The Lewis structure for carbonate ion
  • O
  • O C O

2-
45
  • The FC on the atoms can be calculated as follows
  • The C atom FC 4 0 ½ (8) 0
  • The O atom in CO FC 6 4 ½ (4) 0
  • The O atom in C-0 FC 6 6 ½ (2) -1
  • Thus the Lewis formula for CO32- with FC is
  • O
  • O C O
  • Note that the sum of the FC is -2, the same as
    the charge on carbonate ion
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