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Understanding Covalent Bonding: VSEPR Theory and Molecular Structures, Lecture notes of Chemistry

Chemistry

A concise overview of covalent bonding, covering essential concepts such as bond length, bond energy, and the distinction between bond pair and lone pair electrons. It explains non-polar and polar bonds, coordinate/dative covalent bonds, and the valence shell electron pair repulsion (vsepr) theory. Examples of molecular structures like co2, bf3, h2o, nh3, and ch4, detailing their bond pairs, lone pairs, shapes, and angles. It also discusses electronegativity, polarity in larger molecules, intermolecular forces, and key notes on related topics, making it a valuable resource for understanding chemical bonding principles. Useful for high school and undergraduate students studying chemistry, providing a clear and structured explanation of covalent bonding and molecular structures. It covers essential concepts such as bond length, bond energy, and the vsepr theory, with detailed examples and notes.

Typology: Lecture notes

2021/2022

Available from 08/05/2025

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Chemical Bonding (Chapter 3)
Note: This is a final version of the notes for chapter 3. There will not be any further version of
notes for this, and this will be used for the final exam as preparation material
Covalent Bonding: It is the electrostatic forces of attraction between the nucleus of an atom and
the shared pair of electrons between the atoms in a covalent bond
How does Covalent Bonding occur?
1. When 2 atoms get closer together, the electron of an atom will repel the electron of
another atom, but it is compensated by the attraction the electron has towards the
nucleus
2. As atoms get closer, the electrons have more attraction towards the opposite nucleus,
but the electrons continue to repel
3. When the 2 nuclei get closer together, they start to repel each other
4. Since they have the same positive charge:
a. The most suitable situation will be when attractions of two electrons to two nuclei
are balanced by electron-electron and nucleus-nucleus repulsions
5. Being attracted to both atom nuclei, the 2 electrons stay in the region between
6. After bonding, each atoms attains the electronic configuration of noble gas
7. Due to attraction between opposite charges, covalent bond is a strong bond
Bond Length: distance between two nucleus
Bond Energy: energy needed to break a covalent bond
What are bond pair and lone pair electrons?
โ—Bond pair: electrons present in an atom which take part in bonding
โ— Lone pair: electrons present in an atom which doesnโ€™t take part in bonding
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Partial preview of the text

Download Understanding Covalent Bonding: VSEPR Theory and Molecular Structures and more Lecture notes Chemistry in PDF only on Docsity!

Chemical Bonding (Chapter 3)

Note: This is a final version of the notes for chapter 3. There will not be any further version of notes for this, and this will be used for the final exam as preparation material Covalent Bonding: It is the electrostatic forces of attraction between the nucleus of an atom and the shared pair of electrons between the atoms in a covalent bond How does Covalent Bonding occur?

  1. When 2 atoms get closer together, the electron of an atom will repel the electron of another atom, but it is compensated by the attraction the electron has towards the nucleus
  2. As atoms get closer, the electrons have more attraction towards the opposite nucleus, but the electrons continue to repel
  3. When the 2 nuclei get closer together, they start to repel each other
  4. Since they have the same positive charge: a. The most suitable situation will be when attractions of two electrons to two nuclei are balanced by electron-electron and nucleus-nucleus repulsions
  5. Being attracted to both atom nuclei, the 2 electrons stay in the region between
  6. After bonding, each atoms attains the electronic configuration of noble gas
  7. Due to attraction between opposite charges, covalent bond is a strong bond Bond Length: distance between two nucleus Bond Energy: energy needed to break a covalent bond What are bond pair and lone pair electrons? โ— Bond pair: electrons present in an atom which take part in bonding โ— Lone pair: electrons present in an atom which doesnโ€™t take part in bonding

Non-Polar and Polar Bond Non-Polar: โ— A covalent bond formed between two same atoms or two different atoms with small difference in electronegativity โ— Difference in electronegativity is 0, hence, both atoms attract the electron pair equally and no charge appears โ—‹ The whole molecule becomes neutral โ— Electronegativity difference is between 0 - <0. Polar Bond: โ— A covalent bond formed between two different atoms are polar covalent bond โ— Electrons are shared unequally since the more electronegative atoms pull the paired electron towards itself โ—‹ And away from the less electronegative atoms โ— Difference in charge comes in different areas of the atom โ—‹ Due to uneven spacing of electrons in atoms โ— One end of the molecule is partially + charge and the other end partially - charged โ— Electronegativity difference 0.6 - < Coordinate / Dative Covalent Bonds: A co-ordinate bond is formed when one atoms provides the electrons needed for a covalent bond For dative covalent bonding we need:

  1. One atom having a lone pair of electrons
  2. A second atom having a unfilled orbital, to accept the lone pair a. An electron-deficient compound
  3. In a co-ordinate bond, the head of the arrow points away from the lone pair towards the ion receiving the lone pair Note:
  • At temperatures < 800oC , aluminium chloride exists as AlCl 3
  • At temperatures > 800oC , aluminium chloride exists as Al 2 Cl 6
  • because at high temperatures, dative covalent bonds break

Things to consider when finding the shape of molecule:

  1. Find out the central atom of the molecule a. SO 2 CO 2 H 2 O CH 4
  2. Draw a dot-and-cross diagram of the molecule
  3. Find the total number of electrons around the central atom ( 6 for sulfur)
  4. Find the total number of electron pairs around the central atom (3 pairs)
  5. Electron Pairs โ†’ bond pairs โ†’ 2 โ†’ lone pairs โ†’ 1 CO 2 :
  • central atom: C
  • dot-and-cross diagram:
  • electrons around central atom: 4
  • electron pairs: 2
  • bond pairs: 2
  • lone pairs 0

BF 3 :

  • central atom: B
  • dot-and-cross diagram:
  • electrons around central atom: 6
  • electron pairs: 3
  • bond pairs: 3
  • lone pairs: 0

4 Electron Pairs H 2 O:

  • bond pairs = 2
  • lone pairs = 2
  • shape = bent
  • angle = 104.5o

NH 3 :

  • bond pairs = 3
  • lone pairs = 1
  • shape = triangular pyramidal
  • angle = 107o

CH 4 :

  • bond pairs = 4
  • lone pairs = 0
  • shape = tetrahedral
  • angle = 109.5o More Molecular Structures BF 3 :
  • bond pairs = 3
  • lone pairs = 0
  • shape = trigonal planar
  • angle = 120o

CO 2 :

  • bond pairs = 2
  • lone pairs = 0
  • shape = linear
  • angle = 180o

These are called sigma bonds (end-on overlap) Note: ฯ€bonds are only in double bonds Hybridization: Process of mixing atomic orbitals so each hybridized orbital has some characteristics of each of the orbital mixed Bond Polarity: It is a bond in which two atoms have different charges and the charges donโ€™t coincide

  • They have distinct + and - charge

Electronegativity H (2.1) Li (1.0) Be (1.5) B (2.0) C (2.5) N (3.0) O (3.5) F (4.0) Na (0.9) Mg (1.2) Al (1.5) Si (1.8) P (2.1) S (2.5) Cl (3.0)

  • Electronegativity increases because nuclear charge increases Li (1.0) Be (1.5) B (2.0) C (2.5) N (3.0) O (3.5) F (4.0)
  • charge density โ†‘ protons and electrons โ†‘ (F has the most charge density) Electronegativity: It is the ability of a particular atom, which is covalently bonded to another atom, to attract the bond pairs of electrons towards itself โ— The greater the power of electronegativity, the greater the power of an atom to attract the electrons in a covalent bond towards itself โ— Electronegativity increases across a period โ— Electronegativity increases up each group โ— Atoms with many electrons in nucleus have high electronegativity โ— The greater the difference in electronegativity of 2 atoms, the more polar the bond will be x < 0.5 0.5 - < 1.6 1.6 - 2.0 x > 2 Non-polar covalent Polar covalent If metal involved (ionic) If non-metal involved (covalent) Ionic
  • pd-pd is more stronger than id-id
  • CH 3 Cl had higher bp than CH 4 because it is a polar molecule CH 2 is asymmetrical due to presence of lone pairs Polarity of Molecules: Depending on angles between bonds individual bond dipoles can either reinforce or cancel each other
  1. If cancellation is complete, resulting molecule will have no dipole, hence non-polar
  2. If bond-dipoles reinforce each other, molecules with large dipoles are formed Overall dipole of a molecule depends on:
  • Polarity of each bonds
  • Arrangement of bonds in molecule
  • CHCl 3 is a polar molecule
  • The 3 Cl-Cl dipoles point in same direction
  • C - H bond is non-polar
  • Asymmetric electron distribution ASYMMETRIC IS POLAR
  • Polar bonds are arranged in a way, the dipoles cancel each other out
  • Non-Polar
  • Symmetric SYMMETRIC IS NON-POLAR H (2.2) Li (0.98) Be (1.57)

B

C

N

O

F

Na (0.93) Mg (1.31) Al (1.61) Si (1.90)

P

S

Cl (3.16) K (0.82) Ca (1.00) Ga (1.81) Ge (2.01) As (2.18) Se (2.55) Br (2.96) Rb (0.82) Sr (0.95) In (1.78) Sn (1.96) Sb (2.05) Te (2.1)

I

Cs (0.79) Ba (0.89) Tl (1.62) Pb (1.87) Bi (2.02) Pa (2.0) At (2.2) Fr (0.7) Ra (0.9)

N - H โ†’ 3.04 - 2.2 = 0.84 (Polar)

  • Pyramidal Structure
  • No longer symmetrical
  • Hence it is a polar molecules Note: Polarity results from an equal sharing of electrons
  • Polar Molecules
  • Asymmetrical โ— If there is presence of lone pairs , then the molecule is asymmetrical , hence a polar molecule โ— If there is no presence of lone pairs , the molecule is symmetrical , hence, a non-polar molecule Intermolecular Forces: Forces of attraction between simple covalent molecules
  1. Van Der Waals Forces (id-id forces)
  2. Pd-Pd Forces (permanent dipole)
  3. Hydrogen Bonding H 2 O --- H 2 O (hydrogen bonding) Cl 2 --- Cl 2 (Id-Id forces) CO 2 --- CO 2 (Id-Id forces) HCl --- HCl (Pd-Pd forces)

If non-polar, id-id forces If polar: H with a N, O, F?

  • If no: pd-pd
  • If yes: H-bonds Id-Id โ†’ Pd-Pd โ†’ H-bond (weakest to strongest) If particles have same type of IMF present, the larger the number of electrons, the stronger the IMFโ€™s
  • Symmetrical (Non-Polar)
  • Asymmetrical (Polar) Intermolecular Forces:
  1. Instantaneous Dipole / Temporary Dipole โ†’ Non-Polar Molecules (I 2 / Cl 2 )
  2. Permanent Dipole โ†’ Polar Molecules โ†’ Shape of the Molecule / Symmetry โ†’ Polarity of each bond
  3. Hydrogen bonding a. HCl โ†’ Polar Molecule and Permanent Dipole b. H 2 O โ†’ Polar Molecule and Hydrogen Bonding Hydrogen Bonding โ†’ Intermolecular Bonding (between the molecules) Criteria of Hydrogen Bonding:
  4. One of the molecule has to be hydrogen bonded to either F, O or N
  5. Another molecule has to have F, O or N with a lone pair of electrons which it can donate

Tiny Notes

  • CH is completely non-polar (no electronegativity)
  • above 800, AlCl 3 - below 800, Al 2 Cl 6
  • polar molecules soluble in water
  • bond polarity:
    • the two elements have different electronegativities
    • bonding pair is unequally shared
  • a molecule is polar because the dipoles don't cancel out
  • sp3 hybridization angle = 109.5o
  • sp2 hybridization angle = 120o
  • sp hybridization angle = 180o
  • write permanent dipole and NOT pd-pd
  • bonding polarity: bonding pairs unequally shared, two elements have different electronegativities
  • if the bond angle happens to be 120 degrees, it is trigonal planar
  • when there are three atoms bonded to a central atom, which has a lone pair, the shape is pyramidal
  • metallic bonds exists between metal atoms
  • a higher relative mass means lower bp
  • if ions can't be formed, it isn't a molecule
    • single bonds have only 1 sigma bond
      • double bonds have 1 sigma bond and 1 pi bond
      • triple bonds have 1 sigma bond and 2 pi bonds
  • a dative covalent bond is formed in CO
  • aluminium oxide has COVALENT bonding
  • planar molecules have no lone pairs