2. Electronegativity. 3. Road Map. 4. Types Of Bonding. 5. Properties Controlled By Chemical Bond. 6. Polar Bonds. 7. Metallic Bonding. 8. Intermolecular Forces. Most important 1: chemical bonding occurs when one or more electrons are simultaneously attracted to two nuclei. Most important 2: A chemical bond between. Basic Concepts of Chemical Bonding. Cover to EXCEPT. 1. Omit Energetics of Ionic Bond Formation. Omit Born-Haber Cycle. 2. Omit Dipole Moments.
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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 . 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.
Log In Sign Up. Free Energy and the Equilibrium Constant ClO 4 If the C atom is excited. Multiple Bonds 9. MgCl2 and AlCl3 the polarization increases. Gases Exercises
The electrical conductivity suggests that it is easy to move electrons in any direction in these materials. The thermal conductivity also involves the motion of electrons. All of these properties suggest the nature of the metallic bonds between atoms.
Hydrogen Bonding Hydrogen bonding differs from other uses of the word "bond" since it is a force of attraction between a hydrogen atom in one molecule and a small atom of high electronegativity in another molecule. That is, it is an intermolecular force, not an intramolecular force as in the common use of the word bond.
When hydrogen atoms are joined in a polar covalent bond with a small atom of high electronegativity such as O, F or N, the partial positive charge on the hydrogen is highly concentrated because of its small size.
If the hydrogen is close to another oxygen, fluorine or nitrogen in another molecule, then there is a force of attraction termed a dipole-dipole interaction. Hydrogen bonding has a very important effect on the properties of water and ice. Hydrogen bonding is also very important in proteins and nucleic acids and therefore in life processes.
The atoms are held together because the electron pair is attracted by both of the nuclei. In the formation of 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 also called a dative covalent bond is a covalent bond a shared pair of electrons in which both electrons come from the same atom. The reaction between ammonia and hydrogen chloride If these colourless gases are allowed to mix, a thick white smoke of solid ammonium chloride is formed. 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.
Representing co-ordinate bonds In simple diagrams, a co-ordinate bond is shown by an arrow. The arrow points from the atom donating the lone pair to the atom accepting it. Molecular orbital theory Molecular orbital theory is more powerful than valence-bond theory because the orbitals reflect the geometry of the molecule to which they are applied.
We assume that the electrons would fill the molecular orbitals of molecules like electrons fill atomic orbitals in atoms.
We follow the Pauli exclusion principle. We follow Hund's Rule. One common approximation that allows us to generate molecular orbital diagrams for some small diatomic molecules is called the Linear Combination of Atomic Orbitals LCAO approach. The following assumptions lie at the core of this model. The major draw back is that we are limited to talking about diatomic molecules molecules that have only two atoms bonded together , or the theory gets very complex.
The MO theory treats molecular bonds as a sharing of electrons between nuclei. Unlike the V-B theory, which treats the electrons as localized balloons of electron density, the MO theory says that the electrons are delocalized.
That means that they are spread out over the entire molecule. Forming Molecular Orbitals Molecular orbitals are obtained by combining the atomic orbitals on the atoms in the molecule. Consider the H2 molecule, for example. One of the molecular orbitals in this molecule is constructed by adding the mathematical functions for the two 1s atomic orbitals that come together to form this molecule.
Another orbital is formed by subtracting one of these functions from the other, as shown in the figure below. One of these orbitals is called a bonding molecular orbital because electrons in this orbital spend most of their time in the region directly between the two nuclei.
Electrons placed in the other orbital spend most of their time away from the region between the two nuclei. The bonding molecular orbital concentrates electrons in the region directly between the two nuclei. Placing an electron in this orbital therefore stabilizes the H2 molecule.
Calculating Bond Order In molecular orbital theory, we calculate bond orders by assuming that two electrons in a bonding molecular orbital contribute one net bond and that two electrons in an antibonding molecular orbital cancel the effect of one bond. We can calculate the bond order in the O2 molecule by noting that there are eight valence electrons in bonding molecular orbitals and four valence electrons in antibonding molecular orbitals in the electron configuration of this molecule.
Thus, the bond order is two. The electrons in the Lewis structure are all paired, but there could be unpaired electrons in the molecular orbital description of a molecule. This [5 of 35]. Fluorine has a greater attraction for electrons or has higher electronegativity than hydrogen and the shared pair of electrons is nearer to the fluorine atom than hydrogen atom. For sake of clarity.
Covalent compounds in solution react more slowly as compared with the ionic compounds which react instantaneously in solution. Covalent Bond By Mutual Sharing of Electrons The covalent bond is formed when two atoms achieve stability by the sharing of an electron pair. For example. The solubility of covalent compounds is..
The characteristic solubility of covalent compounds in non-polar solvents such as benzene and carbon tetrachloride can be described to the similar covalent nature of the molecules of solute and solvent i. There is no possibility of formation of double bonds in PH3. The polarity of bonds can lead to polarity of molecules and affect melting point.
Polarity of Bonds: A covalent bond is set up by sharing of electrons between two atoms. The polarity of a bond determines the kind of reaction that can take place at that bond and even affects the reactivity at nearby bonds.
The arrangement of electrons in a covalent molecule is often shown by a Lewis structure in which only valency shells outer shells are depicted. Bond polarities affect both physical and chemical properties of compounds containing polar bond. Thus covalent substances having giant molecules are insoluble in virtually all solvents due to the big size of the molecule unit.
SiCl 4 b. The magnitude of polarization depends upon following factors: As the charge on the cation increases. Polarisation of the anion increases as the size of the cation decreases i. MgCl2 and AlCl3 the polarization increases. Hence the electron cations behave as if they had a greater charge..
In SI units charge q is measured in coulombs C and the distance. This brings more and more covalent nature in the electrovalent compound. Dipole Moment: It is vector quantity and is defined as the product of the magnitude of charge on any of the atom and the distance between the atoms. It is represented by m. Whereas with the increasing charge of anion.
Cations with 18 electrons s2p6d10 in outermost shell polarize an anion more strongly than cations of 8 electrons s2p6 type. The d electrons of the 18 electron shell screen the nuclear charge of the cation less effectively than the s and p electrons of the electron shell. PbCl4 shows covalent nature. The larger the size of the anion. Copper I and Silver I halides are more covalent in nature compared with the corresponding sodium and potassium halides although charge on the ions is the same and the sizes of the corresponding ions are similar.
Similarly among NaCl.
Calculate the percentage ionic character of KCl. Greater the values of the dipole moment. We have seen that in a polar covalent bond between two atoms say A and B. Pauling gave a fairly accurate rule by which the nature of the bond can be predicted. According to this rule.
Dipole moment of methyl chloride is a vectorial addition of dipole moments of three C — H bonds and one C — Cl bond. Calculate the dipole moments of KCl. Dipole moment is a vector quantity and is often indicated by an arrow parallel to the line joining the point of charge and pointing towards the negative end e.
In general a polar bond is established between two atoms of different radii and different electronegativities while positive centres nuclei of different magnitudes combine to share an electron pair.
Greater the difference of electronegativity between A and B. The following points may be borne in mind regarding dipole moments: This bond is. When [7 of 35]. The solutions or fused mass do not allow the passage of electricity. The main properties are described below: Their melting and boiling points are higher than purely covalent compounds and lower than ionic compounds.
The compound consisting of the coordinate bond is termed coordinate compound. Like covalent compounds. Some examples of coordinate bond formation are given below: The properties of coordinate compounds are intermediate between the properties of electrovalent compounds and covalent compounds.
Such a bond is also called as dative bond.
A coordinate or a dative bond is established between two such atoms. The atom which contributes electron pair is called the donor while the atom which accepts it is called acceptor.
They combine to form two double bond and a coordinate bond as to achieve their octet completed. Carbon has four valency electrons and oxygen has six. These are sparingly soluble in polar solvents like water but readily soluble in nonpolar organic solvents. In NH3. Covalent bonds formed are of two types depending upon the way the orbitals overlap each other. Two electrons shared between two atoms constitute a bond.
If the C atom is excited. This is usually a full shell of electrons i. This occurs by excitation of the atom i. In H2O. The spins of the two electrons must be opposite antiparallel because of the Pauli exclusion principle that no two electrons in one atom can have all four quantum numbers the same.
In HF. There are now four x x unpaired electrons which overlap with singly occupied s orbitals on four H atoms. Hence they form tetrahedral structure. In this way the unpaired electrons are paired up.
This increases the number of unpaired electrons. Sigma bond s bond: The bond formed by the overlapping of two half filled atomic orbitals along their axis is known as sigma bond. In CH4. H has a singly occupied s-orbital that overlaps with a singly filled 2p orbital on F. The number of bonds formed by an atom is usually the same as the number of unpaired electrons in the ground state.
The hybrid orbitals always from s bond. A covalent bond results from the pairing of electrons one from each atom. Atoms with unpaired electrons tend to combine with other atoms which also have unpaired electrons. Due to the tetrahedral disposition of sp3 hybrid orbitals. It is proposed that from 2s orbital.
Double bond has one s and one p bond. In ground state. Pi bond p bond: The bond formed by the lateral overlapping of half filled atomic orbitals is known as pi bond. In the excited atom. A p bond is formed when a s bond already exists between the combining atoms. In A — B molecule the bond formed is s bond. It is a process of intermixing of atomic orbitals with small difference in energy and belonging to the same atom. Triple bond has one s and two p bonds. The sidewise overlapping takes place to less extent.
Each of these four sp3 orbital possesses one electron and overlaps with 1s orbitals of four H atoms thus forming four equivalent bonds in methane molecule. At this stage the carbon atom undoubtedly has four half-filled orbitals and can form four bonds.
Types of hybridization and spatial orientation of hybrid orbitals: The geometry and shapes of various species on the basis of VSEPR theory along with hybrid state of central atom is given below in tabular form. The remaining two sp2 orbitals of each carbon form s bonds with H atoms.
When three out of the four valence obritals of carbon atom in excited state hybridize. If 2s and 2p. The molecule is a planar one. The unhybridized 2p. Thus the carbon to carbon double bond in ethene is made of one s bond and one p bond. Since the energy of a p bond is less than that of a s bond.
Two such carbon atoms are involved in the formation of alkenes compounds having double bonds. L Examples BeF2. Method of predicting the Hybrid state of the central atom in covalent molecules of polyatomic ions: The hybrid state of the central atom in similar covalent molecule or polyatomic ion can be predicted by using the generalized formula as described below: ClO 4 1.
CCl 4. BeCl 2. PCl 5 SF HgCl 2 2BF3. CO 3 2CH 4. The magnitude of repulsions between bonding pairs of electrons depends on the electronegativity difference between the central atom and the other atoms. In NH3 and N atom has four electron pairs in the outer shell. But in ether. What angle would you expect for them. Thus the presence of lone pairs on the central atom causes slight distortion of the bond angles from the ideal shape. This may be summarized as: Double bonds cause more repulsion than single bonds.
Effect of Lone Pairs: Molecules with four electron pairs in their outer shell are based on a tetrahedron. A lone pair of electrons takes up more space round the central atom than a bond pair.
In a similar way. The shape of the molecule is determined by repulsions between all of the electron pairs present in the valence shell. In CH4 there are four bonding pairs of electrons in the outer shell of the C atom. If the angle between a lone pair. Because of the lone pair. It follows that repulsion between two lone pairs is greater than repulsion between a lone pair and a bond pair.
The shape of the H2O molecule is based on a tetrahedron with two corners occupied by bond pairs and the other two corners occupied by lone pairs. Whilst it might be expected that two lone pairs would distort the bond angles in an octahedral as in XeF4 but it is [13 of 35].
The order of repulsion between lone pairs and bond pairs of electrons follows the order as: Lone pair. In BrF5. In H2O the O atom has four electron pairs in the outer shell. Thus in ClF3. There are no lone pairs. Thus PCl5 is highly reactive. Three electrons form bonds to F. Gaseous PCl5 is covalent. The lone pairs always occupy the equatorial positions in an triangle. Thus in I3— ion. The high electronegativity of F push the bonding electrons further away from N than in NH3.
The electronic structure P is 1s22s22p63s23p3. In the PCl5 molecule the valence shell of the P atom contains five electron pairs: Effect of Electronegativity: NF3 and NH3 both have structures based on a tetrahedron with one corner occupied by a lone pair.
The electronic configuration of Cl is 1s22s22p63s23p5. The lone pairs occupy all three equatorial positions and the three atoms occupy the top.
Lone pairs distort the structures as before. The chlorine atom is at the centre of the molecule and determines its shape. All five outer electrons are used to form bonds to the five Cl atoms. Symmetrical structures are usually more stable than asymmetrical ones. Molecules with five pairs of electrons are all based on a trigonal bipyramid.
Lone pair bond pair repulsions are next strongest. Mulliken put forward a theory [15 of 35]. For example the nitrogen in trimethyl amine and trisilyl amine has a lone pair electron at nitrogen but nitrogen in trimethyl amine has pyramidal shape while in N trisilyl amine nitrogen has planer shape because in trimethyl amine there is repul.
The most stable structure will be the one of lowest energy. In F. This confirms that the correct structure is III. F Sulphur hexafluoride SF6: The electronic structure of S is S 1s22s22p63s23p6. It was noted previously that a trigonal bipyramid is not a regular shape since the bond angles are not all the same.
As a general rule. Three different arrangements are theoretically possible. Back Bonding: The interaction between an empty orbital and lone pair electron known as back bonding.
Hund and R. Lone pair occupy two of the corners. These questions cannot be explained by valence bond theory. Thus in SF6. It therefore follows that all the corners are not equivalent. There are three bond pairs and two lone pairs. These factors indicate that structure III is the most probable. The great repulsion occurs between two lone pairs. All six of the outer electrons are used to form bonds with the F atoms.
The molecular orbital configuration of O2. When two atomic orbitals combine. Magnetic Behaviour: If all the molecular orbitals in species are spin paired. Order of energy of various molecular orbitals is as follows: For O2 and higher molecules s1s. Bond order B. According to this theory. Molecular orbitals are formed by the combination of atomic orbitals of comparable energy and proportional symmetry.
While an electron in atomic orbital is influenced by one nucleus. The number of molecular orbitals formed is equal to the number of combining atomic orbitals. First BMO are filled. It may be defined as the half the difference between the number of electrons present in the bonding orbitals and the anti-bonding orbitals i.
This type of bonding results between the positive and negative ends of different molecules of the same or different substances. N by a covalent bond.
The attractive force that binds hydrogen atom of one molecule with electronegative atom of the other molecule of the same or different substance is known as hydrogen bond. Bond length is inversely proportional to bond order. In a hydrogen compound.
Examples are o-nitro-phenol. PH3 VA group elements hydrides. From these plots it may be seen that although in case of SbH3. O and F. This type of bonding results between hydrogen and an electronegative element both present in the same molecule.
The solubility of the compound also decreases. The increase in boiling point is due to association of several molecules of the compound. Properties Explained by Hydrogen Bonding a Strength of certain acids and bases can be explained on the basis of hydrogen bonding. H2O and HF hydrides suddenly increase with a further decrease of their molecular weights. This type of bonding is generally present in organic compounds.
Hence compound becomes more volatile. Having no power to form H-bonds. An organic substance is said to be insoluble in water if it does not form hydrogen bonding with water.
If the melting points and boiling points of the hydrides of the elements of IVA. The existence of H-bonding in these molecules gives polymerized molecules NH3 n. The organic compound like alkanes. The density of water increases. Ice has an open cage like structure with a large empty space due to the existence of H-bonds.
In o-nitrophenol. Explain which of the two has higher boiling point? Both have hydrogen bonding. But due to larger distance between —NO2 and —OH group in p-nitrophenol. Two Hwater molecule dotted lines are hydrogen bonds. Circles indicate oxygen atoms. In the crystal structure of ice.
The explanation of this fact is as 2. Thus in ice every water molecule is associated with four other water molecules by H-bonding in a tetrahedral fashion. Melting point of ionic compound is more than covalent compounds Sol. B Solutions: B both ortho and paranitrophenol have intramolecular hydrogen bonding..
A HgCl2 Which one of the following compounds has sp2 hydridization? Which of the following hydrogen bonds is the strongest? Orthonitrophenol is steam volatile but paranitrophenol is not because A orthonitrophenol has intramolecular hydrogen bonding while paranitrophenol has intermolecular hydrogen bonding.
D Van der Waals forces are dominant in orthonitrophenol. A lone pair of electrons in an atom implies [Kurukshetra CET] A a pair of valence electrons B a pair of electrons C a pair of electrons involved in bonding D a pair of valence electrons not involved in bonding Which species has the maximum number of lone pair of electrons on the central atom?
C orthonitrophenol has intermolecular hydrogen bonding and paranitrophenol has intramolecular hydrogen bonding. O The maximum possible number of hydrogen bonds a water molecule can form is  A 2 B 4 C 3 D 1 [21 of 35]. A CO2 A SiH3 3N: B CH3 3N: Which one or more among the following involve s pp.
ClO 4 D CN- NO3 2 C CN. In OF 2. Which of the following structures is linear? I are 4. Select the correct statements: A The heat of hydration of the dipositive alkaline earth metals ions decrease with an increase in their ionic size.
C Model Questions Practice Questions Statement — 2 is NOT a correct explanation for Statement — 1. D Statement — 1 is False. Statement — 2 is True. Statement — 2 is False. XeF2 Which one of the following sequences represents the increasing order of the polarizing power of the cationic species. Each question has 4 choices A. CO2 B BF3. C Statement — 1 is True. IF5 D CF4. B Statement — 1 is True. PCl3 C PF5. XeF4 and XeF6. SF4 [24 of 35]. I3 Which are the species in which central atom undergoes sp3 hybridization?
Statement — 2 is a correct explanation for Statement — 1. Code A Statement — 1 is True.