Chemical bonds form when atoms share or transfer electrons. There are several main types of bonds: - Ionic bonds form when metals transfer electrons to nonmetals to form positive and negative ions that are attracted to each other. Ionic compounds are crystalline and dissolve in water. - Covalent bonds form when atoms share two or more valence electrons to achieve stability. Covalent bond strength depends on the number of electron pairs shared. Covalent compounds exist as discrete molecules. - Metallic bonds result from the attraction between positively charged metal ions and delocalized electrons in the "sea of electrons" in the solid metal. Metallic bonding explains the properties of metals like conductivity.

2. MZGIN M. AYOOB
Types of chemical bond

3. › Force that holds groups of two or more atoms together and makes
the atoms function as a unit.
› Atoms or ions are held together in molecules or compounds by
chemical bonds.
› The type and number of electrons in the outer electronic shells of
atoms or ions are instrumental in how atoms react with each
other to form stable chemical bonds.
› Over the last 150 years scientists developed several theories to
explain why and how elements combine with each other.
Whats the chemical bond

4. › named after Gilbert N. Lewis, who introduced
› it in his 1916
› also known as Lewis dot diagrams, electron dot
diagrams, "Lewis Dot formula" Lewis dot structures,
and electron dot structures)
› are diagrams that show
the bonding between atoms of a molecule and the
lone pairs of electrons that may exist in the molecule.

6. Symbols of atoms with dots to represent the valence-
shell electrons
1 2 13 14 15 16 17 18
H He:
          
Li Be  B   C   N   O  : F  :Ne :
       
          
Na Mg  Al  Si  P S :Cl  :Ar :
       

7. › Strong, STABLE bonds require lots of
energy to be formed or broken
› weak bonds require little E

8. – Occurs between like atoms of a metal in the free
state

9. › electrons are transferred between
valence shells of atoms
› ionic compounds are
made of ions
• ionic compounds are called Salts or Crystals
NOT MOLECULES

10. › Always formed between metals and non-metals
[METALS ]
+ [NON-METALS ]
-
Lost e-
Gained e-

13. Ionic Bonds
› Metal to nonmetal.
› Metal loses electrons to form cation.
› Nonmetal gains electrons to form
anion.
› The electronegativity between the
metal and the nonmetal must be >
than 2.
› Ionic bond results from + to −
attraction.
› Larger charge = stronger attraction.
› Smaller ion = stronger attraction.
› Lewis theory allows us to predict the
correct formulas of ionic compounds.

14. The formations of ionic bond governed by the following
factors:
1. Ionization energy:
• Formation of ionic bond metal atom loses electron to form cation
• Energy required for this equal to ionization energy
• Alkali metals have lowest ionization energy, thus have more
tendency to form cation
2. Electron gain enthalpy:
• Electron released in the formation of cation are to be accepted
by the other atom taking part in the ionic bond formation
• Electron accepting tendencies depend on upon the electron
gain enthalpy
• Defined as energy released when isolated gaseous atom takes
up an electron to form anion.
• Greater the negative enthalpy, easier the formation of anion

15. 3. Lattice energy:
• Combination of oppositely charged ions to form ionic crystal, with
release of energy is referred as lattice energy
• Higher value of lattice energy, greater will be the stability of
compound
• Magnitude of lattice energy gives idea about the strength of interionic
forces
• Size of ions:
• In case of similar ions inter-nuclear distance is lesser due to which
inter-ionic attraction is greater and hence the magnitude of lattice
energy will be larger
• Charge on the ions:
• Ions have higher charge exerts stronger forces of attraction and hence
larger amount of energy is released. Thus value of lattice energy is
higher

16. • Ionic compound exist in solid state
• The network of ions have a definite geometric pattern which
depends on the size and charge of ions
• Posses high melting and boiling points due to strong
electrostatic force of attraction between the ions
• Good conductor of electricity in molten or dissolved state
• Does not conduct electricity in solid state as ions are not free
to move
• Are soluble in polar solvent like water as solvent interacts
with the ions of ionic solid
•The chemical reactions between ionic compounds in aqueous
solution involves the combination between their ions, such
reactions are called ionic reactions.

17. • Between nonmetallic elements of similar
electronegativity.
• Formed by sharing electron pairs
• Stable non-ionizing particles, they are not
conductors at any state
• Examples; O2, CO2, C2H6, H2O, SiC

18. The Convalent Bond
• Shared electrons are attracted to the nuclei
of both atoms.
• They move back and forth between the
outer energy levels of each atom in the
covalent bond.
• So, each atom has a stable outer energy
level some of the time.

20. Chemical Bonds
Covalent bonds form when atoms share 2 or more
valence electrons.
Covalent bond strength depends on the number of
electron pairs shared by the atoms.
single
bond
double
bond
triple
bond
< <

21. Chemical bond

22. Sigma bonds
Sigma bonds are covalent bonds formed by direct overlapping between two
adjacent atom's outer most orbitals. The single electrons from each atom's orbital
combine to form an electron pair creating the sigma bond
π bonds
pi bonds are covalent chemical bonds where two lobes of one involved atomic
orbital overlap two lobes of the other involved atomic orbital

23. The hydrogen bond
A hydrogen bond is formed between H atom attached to an
electronegative atom, and an electronegative atom that possesses a
lone pair of electrons

24. when electrons are
shared equally
H2 or Cl2

25. Oxygen Atom Oxygen Atom
Oxygen Molecule (O2)

26. when electrons are shared
but shared unequally
H2O

27. - water is a polar molecule because oxygen is more
electronegative than hydrogen, and therefore electrons
are pulled closer to oxygen.

28. • Compounds formed exist as discrete molecules
•Weak intermolecular force due to small molecular size
•Mainly exist in liquid or gaseous state
•Sugar, urea, starch etc. exist in solid state
•Low melting and Boiling points due to weak attractive forces
•Poor conductor of electricity in fused or dissolved state
•Less soluble in water
•Gives molecular reaction

29. ›A covalent bond is formed by two atoms
sharing a pair of electrons. The atoms are
held together because the electron pair is
attracted by both of the nuclei.
›In a simple covalent bond, each atom
supplies one electron to the bond - but that
doesn't have to be the case.
›A co-ordinate bond is a covalent bond (a
shared pair of electrons) in which both
electrons come from the same atom.

30. Example: NH4
+

31. Properties of Coordinate bond :
1. Are generally soluble in water and organic solvents
2. Boiling and melting points of these compounds are less than
electrovalent compounds but are higher than covalent
compounds
3. Compounds ionize in aqueous solution giving simple and complex
ions
4. These bonds are also directional and stereoisomerism is also found
5. Molecules possess definite shape and definite bond angles, thus
have definite geometry

32. Metallic bonding is the electrostatic attraction between
the positively charged atomic nuclei of metal atoms and
the delocalised electrons in the metal. In the solid state,
both metallic and ionic compounds possess ordered
arrays of atoms or ions and form crystalline materials
with lattice structures.
Metallic bond

33. Metallic Bonding
› The model of metallic bonding can
be used to explain the properties of
metals.
› The luster, malleability, ductility,
and electrical and thermal
conductivity are all related to the
mobility of the electrons in the
solid.
› The strength of the metallic bond
varies, depending on the charge
and size of the cations, so the
melting points of metals vary as
well.

34. Sea of Electrons
• Electrons are free to move through
the solid.
• Metals conduct electricity.