Find the Shapes of molecule using the VSEPR Theory

BF₃:

                             Central atom is B. There are 3 fluorine atoms.

                    Valence electrons = 0

                            Negative Charge is = 0

                            Add all these, 3+3+0 = 6

                            Number of Bond (N/2) = 6/2 = 3

                            Table 1 tell us bond that it is a Triangular Planner.

BF₄^- :

                             Central atom is B. There are 4 fluorine atoms.

                  Valence electron = 3

                     Negative charge is = -1

                   Add all these, 3+4+1 = 8

                    Number of bond (N/2) = 8/2 = 4

Ammonia NH₃:

                                                 Central atom is N. There are 3 Hydrogen atoms.

                                                                         Valence electron = 5

                                           Add these, 5+3 = 8

                                           Number of bonds (N/2) = 8/2 = 2

                                  . _. .     Shape is Triangular pyramidal.

PCl₅:

                               Central atom is P. There are 5 Cl atoms.

                  Valence electron = 5 

                  Number of bond (N/2) = 10/2 = 5

                  Bonding electron = 5, Non Bonding electron = 0

           . _. .     Shape is Triangular bipyramidal.

CIF₃:

                       Central Atom is Cl.

                                       Valence electron = 7

   There three fluorine atoms, so the number of electron pairs shared is = 3

                       Sum is 10 No. of bond = 10/2 = 5

                         Shape is T-shaped.

                   

                                   

                                      

About Valence Shell electron-pair repulsion (VSEPR) Theory

VSEPR Theory

Valence –Shell electron-pair repulsion (VSEPR) Theory Predicts the shapes and geometries of molecules which may pr not obey the octet rule but have only single bonds. The model extensively uses the number of electron pairs surrounding the central atom to predict geometry of individual molecules and icons. It was being developed by Gillespie and Nyholm. VSEPR Theory is usually compared with valence bond theory. Which address molecular shape through orbitals that are energetically accessible for bonding.

Valence Bond theory concerns itself with the formation of s and p bond. Molecule orbital theory is another model for understanding how atoms and electrons are assembled into molecules and poly-atomic ions. VSEPR theory has been criticised for not being quantitative, and therefore, limited to the generation of crude though structurally accurate molecule geometries of contently-bonded molecule.

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The Geometric arrangement of atoms in molecules and icon may be predicted by means of the valence-shell electron-pair repulsion (VSEPR) theory. This theory predicts the shape of molecules which may or not obey the octet rule but have only single bond. VSEPR theory may be summarized as follows:

1.     The Shape of the molecule is determined by the repulsions between all of the electron pair present in the valence shell.

2.     A loan pair of electron takes up more space around the central atom than a bond-pair, since the loan-pair is one occupied by atom while the bond pair is shared by two nuclei. It follows that repulsion between two loan pair is greater than repulsion between a loan pair and a bond pair, which in turn is greater repulsion between two bond pairs. Thus the presence of loan pairs on the central atom causes slight distortion of the bond angle from the ideal shape. If the angle between a loan pair, the central atom and a bond pair in increased, it follow that the actual bond angle between the atom must be decreased.

3.     The magnitude of repulsion between bonded pair of electrons depends on the electro negativity differnce between the central atom and the other atoms.

4.     Triple bond cause more repulsion than double bond and double bond causes more repulsion thansingle bond. With very few exceptions, the prediction based on the VSEPR theory have been shown to be correct.

The AXE Method:

 The “AXE Method” of electron counting is commonly used when applying the VSEPR theory. The A is represents the central atom and always has an implied subscript one. X is represent the number of ligands atoms bonded to A. The E id represent the number of loan electron pair surrounding the central atoms.

Following steps are given to find out the shape of a molecule:

1.     Identify the central atom and count the number of valence electron.

2.     To this add number of other atoms.

3.     If is an icon, add negative charge and subtract positive charges. Represent the total number which we have calculate above as N.

4.     On dividing N by 2 and comparing the result we obtain the shapes of molecule as given  in table.

sTotal N/2	Number of atoms linked to central atom	Shape of moleculer       or ion	Example
2	2	Linear	HgCl2 , BeCl2
3	3	Trigonal Planar	BF3
5	2	Trigonal Bipyramidal	PCl5 , PF5
4	4	Tetrahedral	CH4 , BF4-
4	3	Trigonal Pyramidal	NH3 , PCl3
4	2	Angular	H2O
5	4	Distorted Tetrahedral	SF4 , IF4+

Types of Covalent Bonds-Sigma (s) and Pi (p) Bonds

Depending upon the type of overlapping, the covalent bonds are mainly of two types.

Sigma (s) and Pi (p) Bonds

When a bond is formed between two atoms by the overlap of their orbitals along the intern clear axis the resulting bond is called sigma (s) bond. Such type of overlap is also known as end or head on overlap. It is a strong bond and symmetrical. The overlapping along the inter nuclear axis can take place in any of the following ways;

1.     s-s Overlapping:

This type of overlap takes place between atoms having half-filled s-orbital in their outer most energy shell. For example, in the formation of hydrogen molecule, 1s orbital of one hydrogen atom overlap with 1s orbital of other hydrogen atom thus forming a sigma bond.

2.     s-p Overlapping:

In this case, half-filled s-orbital of one atom overlaps with the half-filled p-orbital of another atom. A simple example of this type is the formation of hydrogen fluoride. Here 1s orbital of hydrogen overlaps with 2p_z­ orbital of fluoride.

3.     p-p Overlapping:

This type of overlapping occurs when p-orbital of one atom overlaps with the p-orbital of the other as in case fluoride molecule. The molecule of fluorine is produced by the overlapping of 2p_z orbitals of the two fluorine atoms.

Pi (p) Bonds

 Pi-bond is formed by lateral or side wise overlapping of p-orbitals. Sideways overlap means overlapping of p-orbital in a direction perpendicular to the inter nuclear axis. A p-bond in not formed between two bonded atoms unless the two are held together with a s-bond. It is relative a weaker bond since the electrons are not strongly attracted by the nuclear nuclear of bonding atoms.

Bonding in Molecules Explained by Valence Bond Theory

H₂ Molecule:

The electron configuration of hydrogen atom in the ground state in 1s¹. In the formation of hydrogen molecule, two half filled 1s orbital of hydrogen atoms overlap along the inter-nuclear axis and thus by forming a s_s-s bond.

The formation of H₂

 In this diagram, H_A and H_B both have one electron the can be account for. It is known that the electron on H_A belongs to H_A and the electron on H_B belongs to H_B . However, When these two hydrogen atoms bond together, it is impossible to know which electron belongs to H_A or H_B.  Also, since these two molecules are the same, they have equal attraction on the electrons they share. This is because their orbitals overlap and they now share the electrons. The electron are allowed to spin in their respective orbitals.

Cl₂ Molecule:

 The electron configuration of Cl atom in the ground state is [Ne]3s² 3p_x² 3p_y² 3p_z¹ . The two half filled 3p_z atomic orbitals of two chlorine atoms overlap along the inter-nuclear axis and thus by forming a s_p-p bond. 

Valence Bond Theory

Valence Bond Theory

Valence Bond Theory is a qualitative method for predicting the behaviour of electrons in bonding. It focuses on the overlap of the outermost orbital where the value electrons reside. Electrons are thought to be concentrated along the intern clear axis, causing a density of negatively charged electrons between both atoms. Attracted by the nuclei, these shared electrons pull together their respective atoms to which the electrons belong, and the result is formation of a covalent bond. In the Following diagram, formation of a diatomic molecule is shown;

These hydrogen atoms come because of the electrostatic attractive between the nuclei and the electron density between them. Based on the principle of Valence Bond Theory, two electrons are required to make a single bond. Hydrogen has one valence electron, so two hydrogen atoms are needed to make a bond. This allows accurate prediction about the shaped of simple molecule. In case of H₂, the shape is simply linear. The valence bond theory was proposed by Heitler and London to explain of formation of covalent bond quantitatively using quantum mechanics. Later on, Linus Pauling improved this theory by introducing the concept of hydrogen of hybridisation.

The main postulates of this theory are following;

1.   A covalent bond is formed by the overlapping the two half filed valence atomic orbitals of two different atoms.

2.   The electrons in the overlapping orbitals get paired and confined between the nuclei of two atoms.

3.   The electron density between two bonded atoms increases due to overlapping. This confers stability to the molecule.

4.   Greater the extent of overlapping, stronger is the bond formed.

5.   The direction of the covalent bond is along is along the region of overlapping of the atomic orbitals so that it can be said that covalent is directional in nature.   

Physical Properties of Ionic and Covalent Compound

Melting Point

Ionic compounding are typically solid and usually have high melting and boiling points. In contrast covalent compound are typically gases, liquids or low melting solids. These differences occur because of differences in bonding and structure. Ionic compounds are made up of positive and negative ions arranged in a regular way in a lattice. The attraction between ions is electrostatic and is non-directional, extending equally in all direction. Melting of the compound involves breaking of the lattice. This requires considerable energy and hence melting and boiling points are usually high and the compounds are very hard. Compounds with covalent bonds are usually made up of discrete molecule. The bonds are directional and strong covalent bonding forces hold the atom together to make a molecule. In the solid molecules are held together by weak Vander Waals forces. To melt or boil the compound we simply have to supply amount of energy to break Vander Waal forces. Hence compounds are often gases, liquid or soft solids with low melting points.

§  Melting- Change from solid to liquid

§  Melting Point- Specefic tempreture whem melting occurs

§  Each pure substance has a specific melting point.

Example;

               Melting point for water-0^oC (32^oF)

               Melting point for Nitrogen- (-209.9^oc)

               Melting point for Silver- 961^.93 ^oC

              Melting point for Corban- 3500.0^oC

Conductivity

Ionic compounds conduct electricity when the compound is melted or in aqueous solution. Conduction is achieved by the ions migrating towards the electrodes under the influence of electric potentials. Covalent compounds contain neither ions nor mobile electrons so they are unable to conduct electricity in either solid or gaseous state.

Solubility

Ionic compounds are usually soluble in polar solvents. There are solvents if high dielectric constant such as water or mineral acids. Covalent compounds are not normally soluble in polar solvents but they are soluble in non-polar solvents of low dielectric constant such as benzene or CCl₄.

Discus of coordinate Bond

Coordinate Bond

A covalent bond results from sharing of a pair of electrons between two atoms, where atom contributes one electron to the bond. It is also possible to have an electron pair bond where both the electrons come from one of the two bonding atoms and there is no contribution from the other atom. Such bonds are called coordinate bonds or dative bonds, Coordinate bond is a special type of covalent bond in which both the bonded electron come from one of the two binding atoms. One common example is formation of ammonium ion. Even though the ammonia molecule has electronic configuration it can react with a hydrogen ion (H⁺) by donating a loan pair of electron from N atom to H⁺  ion forming the ammonium ion NH⁺₄ .

Covalent bonds are usually shown as a straight line joining the two atoms, and coordinate bonds as arrows indicating which atom is donating the electron. Similarly, ammonium donates its lone pair to boron trifluoride and by this means the boron atom attains noble gas configuration.

In a similar way a molecule of BF₃ can form a coordinate bond by accepting a loan pair from a F^- ion.

Double and Triple Bonds

Sometimes more than two electrons are shared between a pair of atoms. If four electrons are shared then there are two bonds this arrangement is called a double bond. If six electrons are shared then there are three bonds and this is called a triple bond.

Discus of metallic Bond

Metallic Bond

A less mentioned type of bonding is the metallic bond. In this type of bonding, all metallic atoms lodes their valence electron from a pool of electrons, which is mobile. Leaving the valence electrons, the remainder portion of the metal atom is a positive ion called “kernel”.

For example, in lithium each atom contributes are valence electron to the pool leaving behind Li⁺ ions; in case of Mg, each atom contributes two valence electron to the pool leaving behind Mg²⁺ ions. These positive ions or kernels are held in the three dimensional space in a definite pattern in the sea of mobile electrons. This model is called electron gas model because the electrons are free to move in all direction like the molecules of a gas.

   Structure of metals (electron sea model)

The simultaneous attraction between the kernels and the mobile electrons which holds the kernels together is called metallic bond. Metallic bonding may be seen as an extreme example of delocalization of electrons over a large system of covalent bonds, in which every atom participates. This type of bonding is often very strong resulting in the tensile strength of metals. However, metallic bonds are more collective in nature than other types, and so they allow metal crystals to deform more easily, because they are composed of atom attracted to each other, but not in any particularly-oriented ways. This results in the malleability of metals. The sea of electrons in metallic bonds causes the characteristically good electrical and thermal conductivity of metals, and also their “shiny” reflection of most frequencies of white light.

All types of chemical bonds can be explained by quantum theory, but, in practice, simplification rules allow chemists to predict the strength, irrationality, and polarity of bonds. The octet rule and VSEPR theory are two examples, More sophisticated theories are valence theory which includes orbital weatherization and resonance, and the linear combination  of atomic orbitals which is the basis for the extraordinary molecular orbital method.