CHEMICAL BONDS PDF

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approach to chemical bonding;. • explain the octet rule and its limitations, draw Lewis structures of simple molecules;. • explain the formation of different types of . Covalent Bonds. Covalent bond: bond in which one or more pairs of electrons are shared by two atoms. Covalent bonding occurs when two (or more) elements . Chemical Bonds – atoms combine, Electron distribution changes Ionic Bonds is the transfer of electrons (ex. salts, crystals) and is composed of a positive.


Chemical Bonds Pdf

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Claim: chemical bonds are what hold atoms together to form the more complicated aggregates that we know as molecules and extended solids. Most important. PDF | Chemical bonding is one of the key and basic concepts in chemistry. The learning of many of the concepts taught in chemistry, in both secondary schools. 1,onding in Chemistry 4 0 al Golum1)in in the sprirrg of , and is mairll) iilter~ ded for Llle ur~dcrgraduate scuden t in chemistry. \\rho desires an iiltroductiorl.

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May be insoluble in H2O, may dissolve in organic solvents Giant molecular Solids of high melting point and boiling point Quartz Properties: Insoluble in H2O and other solvents In covalent bonds, electron pairs are shared between atoms. Documents Similar To Chemical bonding and structure. Aston Hamilton. Connor Marshman. James Wong. Mini Pg. Sachin Rane. Satria Adi Nugroho.

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Ishtiaque Hossain. Popular in Physics. Mico Andres. Mayarani Praharaj. Ananthesh Rao. Youhan Karemina. Gavie Marquos. Arjun Singh Arora. The low density of ice also can be explained in terms of Notes hydrogen bonding. Due to the difference in the electronegativity between hydrogen and the other electronegative atom, the bond connecting them becomes polar.

The hydrogen atom acquires a positive charge while the electronegative atom bears the negative charge. Hydrogen bonding results from the electrostatic interaction between the positively charged hydrogen atom and the negatively charged electronegative atom. The second electronegative atom may be a part of the same molecule or it may belong to a different molecule.

Accordingly, there are two types of hydrogen bonds. If the hydrogen bond is formed between two different molecules it is called intermolecular hydrogen bond. When the hydrogen bond exists within the same molecule, it is called intramolecular hydrogen bonding.

Salicyldehyde ad o-nitrophenol are two common examples of the molecules showing intramolecular hydrogen bonding whereas in water, intermolecular hydrogen bonding exists. The theories of bonding that we have discussed so far do not say anything about the shape of the molecules. It was Chemical Bonding later refined and extended by Nyholm and Gillespie in This theory focuses on the electron pairs present in the valence shell of the central atom of the molecule and can be stated in terms of two postulates: POSTULATE 1 The electron pairs both bonding and non-bonding around the central atom in a Notes molecule arrange themselves in space in such a way that they minimize their mutual repulsion.

In other words, the chemical bonds in the molecule will be energetically most stable when they are as far apart from each other as possible. Let us take up some examples. BeCl2 is one of the simple triatomic molecules. In this molecule, the central atom, beryllium has an electronic configuration of 1s2 2s2.

That is it has two electrons in its valence shell. In the process of covalent bond formation with two chlorine atoms two more electrons are contributed one by each chlorine atom to the valence shell.

Thus there are a total of 4 valence electrons or two pairs of valence electrons. According to the postulate given above, these electron pairs would try to keep as far away as possible. It makes the two electron pairs to be at an angle of o which gives the molecule a linear shape.

In boron trifluoride, the central atom, boron has an electronic configuration of 1s 2 2s2 2p1. That is, it has three electrons in its valence shell. In the process of covalent bond formation with three fluorine atoms three more electrons are contributed one by each fluorine atom to the valence shell. Thus there are a total of 6 valence electrons or three pairs of valence electrons.

It makes the three electron pairs to be located at an angle of o which gives the molecule a planar trigonal shape. F B F F Planar trigonal shape Thus different molecules would have different shapes depending on the number of valence shell electrons involved. The geometric shapes associated with various numbers of electron pairs surrounding the central atom are given in Table 5. Geometric arrangements of electron pairs around central atom. The shapes of molecules containing a combination of lone pairs and bond pairs would be distorted from the above mentioned shapes.

Let us take an example of three molecules namely, methane, ammonia and water.

Covalent bond

All the three contain a total of 4 electron pairs around their central atom. But the nature of these is different in the three cases. In methane molecule the central carbon atom has 4 valence electrons and it shares 4 electrons with four hydrogen atoms. So there are a total of 4 bond pairs and according to Table 5.

In case of ammonia also there are four pairs of electrons but their nature is different. Three of these are bond pairs while one is a lone pair.

Similarly, in case of water again there are four pairs of electrons ; two are bond pairs while two are lone pairs. Due to the differences in the mutual repulsion between bond pair - bond pair and lone pair - bond pair the molecular shape would be slightly distorted from the expected tetrahedral shape. The number and nature of electron pairs and the geometries of these three molecules are given in Table 5. Molecular geometries of molecules with 4 electron pairs with different combinations of lone pairs and bond pairs.

H H2O 2 2 angular or bent O.. We have so far learnt that a chemical bond formation between two atoms can occur by transfer ionic bonding or sharing covalent bonding of electrons. The processes of bond formation and the bonding in simple molecules can be conveniently represented in terms of electron — dot structures. This is in contradiction with the probabilistic orbital representation of the electron that you have learnt in lesson 3.

Let us learn how do we explain the process of bond formation in terms of modern theories that incorporate the wave mechanical representation of atom. Intext Question 5. These do not take into account the wave mechanical or quantum mechanical principles. After the development of quantum mechanical description of atomic structure two more theories were proposed to explain the bonding between atoms.

These are called modern theories of chemical bonding. Let us discuss these theories in brief. Linus Pauling and others further developed it. In this approach the process of chemical bond formation can be visualised as the overlapping of atomic orbitals of the two atoms as they approach each other.

The strength of the bond depends on the effectiveness or extent of the overlapping. Greater the overlapping of the orbitals, stronger is the bond formed.

Let us take the example of bonding in hydrogen molecule to understand the VB approach. Notes Suppose that the two hydrogen atoms are at infinite distance from each other. Their electrons are in their respective 1s orbitals and are under the influence of the corresponding nuclei. As the two atoms approach each other their 1s orbitals begin to overlap which lead to decrease in energy, Fig.

At a distance equal to the bond length the overlapping is maximum and the energy is minimum.

The overlapping can be equated to the sharing of electrons between the atoms. The electrons occupying the shared region of orbitals are under the influence of both the nuclei. Formation of hydrogen molecule from overlapping of two hydrogen atoms This simple approach can be used to explain the bonding in simple diatomic molecules like HF,F2 etc.. However, to explain bonding in molecules containing more than two atoms some additional concepts like excitation and hybridisation need to be used.

The atomic number of beryllium is 4. Its electronic configuration is 1s2 2s2. In order to form bonds with two hydrogen atoms the valence electrons 2s2 of beryllium atom must overlap with the 1s electrons of the two hydrogen atoms.

Since the valence shell of beryllium atom contains both the electrons in the same orbital i. Pauling got over this problem by suggesting that in the process of bond formation an electron from the 2s orbital of beryllium atom gets momentarily excited to the empty 2p orbital as shown below.

The problem is still not over. The two bonds formed by these overlaps would be of different nature. One of these would involve overlapping of 2s orbital of beryllium with 1s orbital of hydrogen while the other would involve overlapping of 2p orbital of beryllium with 1s orbital of hydrogen. However, experimentally the two bonds are found to be equivalent.

This problem is solved with the help of a concept called hybridisation of orbitals. Notes According to this two or more than two non equivalent orbitals having different energies and shapes of comparable energies mix or hybridize and give rise to an equal number of equivalent same energies and shapes hybrid orbitals. In case of BeCl2 the two singly occupied orbitals 2s and 2p hybridize to give two sp- hybrid orbitals.

This is called sp- hybridisation.

These hybrid orbitals lie along the z- direction and point in opposite directions. Let us Clpz Be Clpz take up some more cases involving hybridisation of s and p orbitals. Formation of BeCl2 ; sp hybridisation Boron trichloride sp2 hybridisation: In boron there are five electrons and the electronic configuration is 1s2, 2s2,2p1.

There are three electrons in the valence shell of boron atom. In order to form bonds with three chlorine atoms one of the electrons from the 2s orbital of boron atom is excited to its 2p orbital. Boron Atom Excitation Ground state 1s 2s 2p Boron Atom Excited state 1s 2s 2p One 2s orbital and two 2p orbitals hybridise to give three sp2 hybridized orbitals.

This is called sp2 - hybridisation. T h e s e shown below, fig. Formation of BCl3 ; sp2 hybridisation. Bonding in Methane sp3 hybridisation: In case of methane the central atom, carbon, has an electronic configuration of 1s2, 2s2 ,2p2. In order to form bonds with four hydrogen atoms one of the electrons from the 2s orbital of carbon atom is excited to the 2p orbital.

Carbon Atom Excitation Carbon Atom Ground state Excited state 1s 2s 2p 1s 2s 2p One 2s orbital and three 2p orbitals of the carbon atom then hybridize to give four sp 3 hybridized orbitals. This is called sp3- hybridisation.

Carbon Atom Excitation Carbon Atom Ground state Excited state 1s 2s 2p 1s Sp3 Hybridised These four sp3 hybrid orbitals are directed towards the corners of a regular tetrahedron.

These hybrid orbitals then form bonds with the 1s orbitals of hydrogen atoms to give a methane molecule as shown below, Fig. Formation of CH4 ; sp hybridisation. Phosphorus pentachloride sp3d hybridisation: P ground state P excited state sp3d hybridisation Five sp3d hybrid orbitals are formed which are directed towards the corners of a trigonal bipyramidal Fig.

Thus PCl5 molecule has a trigonal bipyramidal geometry. SF6 sp3d2 hybridisation: Trigonal bipyramidal geometry Fig. Octahedral geometry of SF6 of PCl5 molecule.

Let us see how does the concept of hybridisation help us in understanding bonding between pairs of such atoms. In the case of bonding in ethane C2H6 , two carbon atoms are bonded to each other and each carbon atom is bonded to three hydrogen atoms.

Chemical bonding and structure.pdf

You would recall that in the case of methane the valence orbitals of carbon atom undergo sp3 hybridisation. In ethane each carbon atom undergoes sp 3 hybridisation to give four sp3 hybridized orbitals. Becke and K. Edgecombe in The Journal of Chemical Physics 92, This function and its use in chemistry and physics is reviewed in Andreas Savin and coworkers, Angewandte Chemie Int.

Savin in Chemie in unserer Zeit 31, — What makes ELF useful in the present case is that it allows us to have a local view of chemical bonding even in metallic systems. PDF high quality version about KByte Small lithium clusters of up to 40 atoms are found to be dominated by multi-center bonds that involve three, four or even as much as six centers, see the middle figure above. Within this picture, the metallic phase of lithium is characterized by multi-center bonds and very gentle variations of ELF accross space.Such covalent substances are low-boiling-temperature liquids such as ethanol , and low-melting-temperature solids such as iodine and solid CO2.

In the process of suggesting the process of chemical bonding Lewis provided a very convenient way of representing bonding in simple molecules. Minh Tieu. Formation of BCl3 ; sp2 hybridisation. In other words, the chemical bonds in the molecule will be energetically most stable when they are as far apart from each other as possible. In boron there are five electrons and the electronic configuration is 1s2, 2s2,2p1.

Write down the molecular orbital electronic configuration of the following species Chemical Bonding and compute their bond orders. Such a bond is called covalent bond and the compounds so obtained are called covalent compounds. The overlapping region responsible for bonding is situated between the two atoms i. This leaves two mutually perpendicular unhybridised p orbitals each on both the carbon atoms.

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