Chemical bonding - Inorganic Presentation

1. INORGANIC Pharmacy-II
Presentation On-Chemical Bonding
SUBMITTED TO
MD.ASSADUJAMAAN
Lecturer, Dept. Of Pharmacy, NUB
SUBMITTED BY
1. Saidur Rahman , ID-938
2. Mubassera Anee , ID-940
3. Zakia Sultana , ID-951
4. Malvin Sajib , ID-955
5. Abdul Hannan , ID-959
6. Swarnali Sultana , ID-960

2. CHEMICAL BONDING
 A chemical bond is defined as
the attractive force that hold
two or more atoms together in a
molecule or an ion.

3. WHY DO ATOMS COMBINE
Atoms combine because they are not stable
unless their valence orbital is full of
electrons. In order to fill this outer orbital they
will share, give, or take electrons from other
atoms.

4. HOW DO ATOMS COMBINE
Atoms can atoms combine in the following two
ways :
1. Transfer of one or more electrons from the
valence shell of one atom to the valence shell
of another atom.
2. One , two or three electron pairs of the valence
shell of both the combining atoms are shared
between them.

5. TYPES OF CHEMICAL BONDS
 Two types:
• Strong Bonds
• Weak Bonds
 Strong Bonds:
I. Ionic Bond
II. Covalent Bond
III. Coordinate or Dative Covalent Bond
IV. Metallic Bond
 Weak Bond
I. Hydrogen Bond
II. Vander Waals Interaction

6. IONIC BOND
An ionic bond is a type of chemical bond that involves
a metal and a non-metal ion (or polyatomic ions such
as ammonium) through electrostatic attraction. In short, it is a
bond formed by the attraction between two oppositely charged
ions.

7. How are ionic BONDS FORMED
The metal donates one or more electrons, forming a positively charged
ion or cation with a stable electron configuration. These electrons then
enter the non metal, causing it to form a negatively charged ion or
anion which also has a stable electron configuration. The electrostatic
attraction between the oppositely charged ions causes them to come
together and form a bond.
 For example, common table salt is sodium chloride.
When sodium (Na) and chlorine (Cl) are combined, the sodium atoms
each lose an electron, forming cations (Na+), and the chlorine atoms
each gain an electron to form anions (Cl−). These ions are then
attracted to each other in a 1:1 ratio to form sodium chloride (NaCl).
 The removal of electrons from the atoms is endothermic and causes
the ions to have a higher energy. There may also be energy changes
associated with breaking of existing bonds or the addition of more
than one electron to form anions. However, the attraction of the ions
to each other lowers their energy.

9. Factors or conditions for favoring ionic compounds
I. Ionization Enthalpy (Energy) : Lesser the value of ionization enthalpy,
greater the tendency of the atom to form cation. For example, alkali
metals form cations easily because of the low value of ionization
energies.
II. Electron gain enthalpy : Greater the value of electron gain enthalpy,
more the tendency of the atom to form anion .For example ,halogens
have high electron gain enthalpies within the respective periods and form
ionic compounds easily .
III. Lattice enthalpy : It is the energy released when the close packing of
the gaseous ions of the opposite charge forms one mole of ionic solid.
Magnitude of lattice energy gives an idea about the inter-ionic forces and
it also gives the measure of the stability of the ionic compound which
depends upon the following factors.
IV. Size of the ion : Smaller the size of the ion s, lesser the inter nuclear
distance and greater the inter ionic interaction, hence, larger the
magnitude of lattice energy
V. Charge on the ions : Larger the magnitude of the charge on the ions
greater will be the attractive forces and higher the negative value of

10. covalent BOND
 A chemical bond in which two atoms share some
of their valence electrons, thereby creating
a force that holds the atoms together.
 Usually each atom contributes one electron to
form a pair of electrons that are shared by both
atoms.

11. A clear example is the molecule of chlorine, chlorine occurs in nature
as a molecule composed of 2 atoms of chlorine, chlorine atoms are
linked by a covalent bond produced by the sharing of 2 electrons.
During this process two atoms are join together to form one molecule,
ignoring the molecular orbital theory, bonding / anti bonding and in
order to explain it simply, we can say that 2 atomic orbital (Cl + Cl)
combine to form a new molecular orbital (Cl2).
The orbital are defined as regions of the atoms or molecules where
the electrons are localized.

12. There are three types of covalent bond depending upon the number of shared electron
pairs.
I. SINGLE COVALENT BOND
II. DOUBLE COVALENT BOND
III. TRIPLE COVALENT BOND
SINGLE COVALENT BOND
A covalent bond formed by the mutual sharing of one electron pair between
two atoms is called a “Single Covalent bond”
Examples:

13. DOUBLE COVALENT BOND
A covalent bond formed between two atoms by the mutual sharing of two
electron pairs is called a "double covalent bond".
Examples:
TRIPLE COVALENT BOND
A covalent bond formed by the mutual sharing of three electron pairs is
called a "Triple covalent bond".
Examples:

14. POLAR COVALENT BOND
A covalent bond formed between two different atoms is known as Polar covalent bond.
For example when a Covalent bond is formed between H and Cl , it is polar in nature
because Cl is more electronegative than H atom . Therefore, electron cloud is shifted
towards Cl atom. Due to this reason a partial -ve charge appeared on Cl atom and an
equal +ve charge on H atom.
Examples:
NON-POLAR BOND
A covalent bond formed between two like atoms is known as Non-polar bond.
Since difference of electro negativity is zero therefore, both atoms attract electron pair
equally and no charge appears on any atom and the whole molecule becomes neutral.
Examples:

15.  Coordinate or Dative Covalent Bond
A dipolar bond, also known as coordinate
link, coordinate covalent bond, dative bond,
or semi polar bond, is a description
of covalent bonding between two atoms in which
both electrons shared in the bond come from the
same atom.

16. The reaction between ammonia and hydrogen chloride
• If these colorless gases are allowed to mix, a thick white
smoke of solid ammonium chloride is formed.
• Ammonium ions, NH4
+, are formed by the transfer of a
hydrogen ion from the hydrogen chloride to the lone pair of
electrons on the ammonia molecule.

17. When the ammonium ion, NH4
+, is formed, the fourth hydrogen is
attached by a dative covalent bond, because only the hydrogen's
nucleus is transferred from the chlorine to the nitrogen. The hydrogen's
electron is left behind on the chlorine to form a negative chloride ion.
Once the ammonium ion has been formed it is impossible to tell any
difference between the dative covalent and the ordinary covalent
bonds. Although the electrons are shown differently in the diagram,
there is no difference between them in reality.

19. metallic bond
The chemical bond characteristic of metals, in
which mobile valence electrons are shared
among atoms in a usually stable crystalline
structure.

20. The structure of a metallic bond is quite different from
covalent and ionic bonds. In a metal bond, the valence
electrons are delocalized, meaning that an atom's electrons
do not stay around that one nucleus. In a metallic bond, the
positive atomic nuclei (sometimes called the 'atomic
kernels') are surrounded by a sea of delocalized electrons
which are attracted to the nuclei
Positive atomic nuclei (+) surrounded by delocalized electrons (∙)

21. PROPERTIES OF METALIC BOND
Formed between atoms of metallic
elements
Electron cloud around atoms
Good conductors at all states, lustrous,
very high melting points
Examples; Na, Fe, Al, Au, Co

22.  hydrogen bond
An electrostatic attraction between a hydrogen
atom in one polar molecule (as of water) and a
small electronegative atom (as of oxygen,
nitrogen, or fluorine) in usually another molecule
of the same or a different polar substance

23. Hydrogen bonds only form between hydrogen and oxygen (O), nitrogen (N) or fluorine (F).
Hydrogen bonds are very specific and lead to certain molecules having special properties due
to these types of bonds. Hydrogen bonding sometimes results in the element that is not
hydrogen (oxygen, for example) having a lone pair of electrons on the atom, making it polar.
Lone pairs of electrons are non-bonding electrons that sit in twos (pairs) on the central atom
of the compound. Water, for example, exhibits hydrogen bonding and polarity as a result of
the bonding. This is shown in the diagram below. Because of this polarity, the oxygen end of
the molecule would repel negative atoms like itself, while attracting positive atoms, like
hydrogen. Hydrogen, which becomes slightly positive, would repel positive atoms (like other
hydrogen atoms) and attract negative atoms (such as oxygen atoms). This positive and
negative attraction system helps water molecules stick together, which is what makes the
boiling point of water high (as it takes more energy to break these bonds between water
molecules).
In addition to the four types of chemical bonds, there are also three categories bonds fit into:
single, double, and triple. Single bonds involve one pair of shared electrons between two
atoms. Double bonds involve two pairs of shared electrons between two atoms, and triple
bonds involve three pairs of shared electrons between two atoms. These bonds take on
different natures due to the differing amounts of electrons needed and able to be given up.

24.  Vander Waals Interaction
Van der Waals forces include attractions between atoms,
molecules, and surfaces. They differ from covalent and
ionic bonding in that they are caused by correlations in the
fluctuating polarizations of nearby particles (a
consequence of quantum dynamics).