Handbooks Chemical Bonding 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 . 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.

Chemical Bonding Pdf

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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.

Acid-Base Reactions 4. Oxidation-Reduction Reactions 4. Concentration of Solutions 4. Solution Stoichiometry and Chemical Analysis 4. Reactions in Aqueous Solution Exercises 4. Thermochemistry 5. The Nature of Energy 5. The First Law of Thermodynamics 5. Enthalpy 5. Enthalpy of Reaction 5. Calorimetry 5. Hess's Law 5. Enthalpies of Formation 5. Foods and Fuels 5. Thermochemistry Exercises 5. Electronic Structure of Atoms 6. The Wave Nature of Light 6.

Quantized Energy and Photons 6. Line Spectra and the Bohr Model 6.

8.S: Basic Concepts of Chemical Bonding (Summary)

The Wave Behavior of Matter 6. Quantum Mechanics and Atomic Orbitals 6. Many-Electron Atoms 6. Electron Configurations 6. Electron Configurations and the Periodic Table 6. Electronic Structure of Atoms Exercises 6. Periodic Properties of the Elements 7. Development of the Periodic Table 7. Effective Nuclear Charge 7. Sizes of Atoms and Ions 7.

Ionization Energy 7. Electron Affinities 7.

Metals, Nonmetals, and Metalloids 7. Group Trends for the Active Metals 7. Group Trends for Selected Nonmetals 7. Periodic Properties of the Elements Exercises 7. Basic Concepts of Chemical Bonding 8.

Ionic Bonding 8. Covalent Bonding 8. Bond Polarity and Electronegativity 8. Drawing Lewis Structures 8. Resonance Structures 8. Exceptions to the Octet Rule 8. Strength of Covalent Bonds 8. Basic Concepts of Chemical Bonding Exercises 8.

Molecular Geometry and Bonding Theories 9. Molecular Shapes 9. Molecular Shape and Molecular Polarity 9. Covalent Bonding and Orbital Overlap 9. Hybrid Orbitals 9.

Multiple Bonds 9. Molecular Orbitals 9. Second-Row Diatomic Molecules 9. Exercises 9. Gases Characteristics of Gases Pressure The Gas Laws The Ideal Gas Equation Further Applications of the Ideal-Gas Equations Gas Mixtures and Partial Pressures Kinetic-Molecular Theory Molecular Effusion and Diffusion Real Gases - Deviations from Ideal Behavior Exercises Liquids and Intermolecular Forces Intermolecular Forces Some Properties of Liquids Phase Changes Vapor Pressure Phase Diagrams Structure of Solids Bonding in Solids Liquids and Intermolecular Forces Exercises Solids and Modern Materials Classes of Materials Materials for Structure Materials for Medicine Materials for Electronics Materials for Optics Materials for Nanotechnology Properties of Solutions The Solution Process Saturated Solutions and Solubility Factors Affecting Solubility Ways of Expressing Concentration Colligative Properties Colloids Properties of Solutions Exercises Chemical Kinetics Factors that Affect Reaction Rates Reaction Rates Concentration and Rates Differential Rate Laws Temperature and Rate Reaction Mechanisms Catalysis Chemical Equilibrium The Concept of Equilibrium The Equilibrium Constant Heterogeneous Equilibria Calculating Equilibrium Constants Applications of Equilibrium Constants Acid—Base Equilibria Acids and Bases: A Brief Review The Autoionization of Water The pH Scale Strong Acids and Bases Weak Acids Weak Bases Acid-Base Properties of Salt Solutions Acid-Base Behavior and Chemical Structure Lewis Acids and Bases Acid—Base Equilibria Exercises Additional Aspects of Aqueous Equilibria Choose Sodium Na.

What type of element is it? How many valence electrons does it have? Choose Fluorine F. Circle: Ionic or Covalent? Choose the appropriate number of atoms to make the bond. Record the number of each atom below: 4. Watch the final animation closely it will play continuously.

Describe the change in the number of valence electrons in the atoms as the bond is successfully formed: b. What does the negative - charge indicate mention specific subatomic particles in your answer? American Association of Chemistry Teachers 3 d. What is the final overall charge? Record the name and molecular formula for the compound below: Reset the selected data using the reset symbol.

Ionic bond is Illustration 2: Electrical Conductivity: When an electrovalent compound is molten or dissolved in a solvent of high dielectric constant e.. Non-polar solvents like benzene and carbon tetrachloride do not solvate the ions as their dielectric constants are low.

As a result. Ionic compounds like sulphates and phosphates of barium and strontium are insoluble in water because lattice energy is greater than hydration energy.

Hence the lone pair repels the bond pairs of NF3 more than it does in NH3. Dissolution is also favoured by the high dielectric constant of the solvents such as water. The hydrogen end of the molecule. In covalent compounds. It is further classified as polar or non-polar depending upon the fact whether the electron pair is shared unequally between the atoms or shared equally.

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..

Simulation: Ionic & Covalent Bonding

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.

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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.

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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.Proteins For nonmetals, the number of valence electrons is the same as the group number.

PCl3 C PF5. It is also a major factor governing the properties of ions in solution. Carbohydrates It has zero dipole moment. Kim Joshua Bolo.