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Lecture 8

Chemical Bonding

Shaun Williams, PhD

Types of Bonds

  • Chemical bond
    • A force that holds atoms together in a molecule or compound
  • Two types of chemical bonds
    • Ionic Bonds
    • Covalent Bonds
graphic of calcium carbonate, an example of an ionic bonding compound, and carbon dioxide, an example of a covalent bonding compound.

Ionic Bond

  • A bond created by electrostatic attraction between oppositely charged ions
  • Occurs between a metal and a nonmetal
  • Electrons transferred between the cation (positively charged ion) and the anion (negatively charged ion)
  • Extremely strong bonds
graphic of the attraction and repulsion pairs of ions in a ionic crystal.

Covalent Bond

graphic of calcium carbonate, an example of an ionic bonding compound, and carbon dioxide, an example of a covalent bonding compound.
  • A bond created by the sharing of electrons between atoms
  • Occurs between two nonmetals (resulting in a neutral overall charge)
  • Electrons not transferred in this case
  • Electrons shared in pairs typically
  • Weaker bonds than ionic bonds

Polar vs Nonpolar

Electronegativity

Electronegativity Values

The periodic table showing the elements electronegativity values. They are increasing from bottom to top and from left to right.

Ionic Bonding

Structures of Ionic Crystals

A graphic showing the cleaving of an ionic crystal along the plane between layers of ions.

A More Complicated Ionic Crystal

A graphic showing the ion structure of the CaF2 compound.

Covalent Bonding

  • Single covalent bond
    • A covalent bond that consists of a pair of electrons shared by two atoms
    • Each atom contributes one electron to the bond
      • The orbitals overlap to allow the electron pair to be located around both atoms
    • Lewis formula
      • The atoms are shown separately and the valence electrons are represented by dots
A graphic of the sharing of electrons in CF4 and HF. A pair of electrons, or two dots, are shared between each pair of atoms.

More on Covalent Bonding

  • Multiple covalent bonds
    • Covalent bonds that consist of more than one pair of electrons shared by two atoms
    • Double bond
      • Sharing of two pairs of electrons (4 electrons total)
      • In Lewis Dot structures, a double bond is represented by 4 dots or 2 parallel lines.
    • Triple bond
      • Sharing of three pairs of electrons (6 electrons total)
      • In Lewis Dot structures, a triple bond is represented by 6 dots or 3 parallel lines.
A graphic showing the atoms in O2 sharing four electrons (double bond) and the atoms in N2 sharing six electrons (triple bond).

Steps for Writing Lewis Dot Structures

  1. Write an atomic skeleton.
    • The arrangement of atoms is usually symmetrical.
    • In a molecule of two different elements, the one with the greater number of atoms usually surrounds the one with the lesser number of atoms.
    • The central atom, the one surrounded by the other atoms, tends to be the one that is less electronegative and is present in the least quantity. This atom usually forms the greater number of bonds and is found further toward the bottom left side of the periodic table.
    • Hydrogen atoms are generally on the outside of the molecule.
    • The chemical formula may give clues about the arrangement of atoms.

Steps for Writing Lewis Dot Structures (cont.)

  1. Sum the valence electrons from each atom to get the total number of valence electrons.
  2. Place two electrons, a single bond, between each pair of bonded atoms.
  3. If you have not placed all the valence electrons in the formula, add any remaining electrons as unshared electron pairs, consistent with the octet rule.
    • Add pairs of electrons first to complete the octet of atoms surrounding the central atom. Then add any remaining electrons in pairs to the central atom.
  4. If necessary to satisfy the octet rule, shift unshared electrons from non-bonded position on atoms with completed octets to positions between atoms to make double or triple bonds.

Writing Lewis Dot Structures

  1. Write an atomic skeleton for \(\chem{CH_2O}\):

    A carbon atom with a hydrogen atom on the left and right and an oxygen atom above.

  2. Sum the valence electrons from each atom to get the total number of valence electrons.
    • Carbon is in Group IVA (14), so it has 4 valence electrons.
    • Each hydrogen contributes 1 valence electron (H is in Group IA (1)).
    • Oxygen contributes 6 valence electrons because it is in Group VIA (16).
    • Total number of valence electrons: \( 4 + \left( 1 \times 2 \right) + 6 = 12 \)

Writing Lewis Dot Structures (cont.)

  1. Next, bond the electrons around each atom in a single bond first, then use double bonds as necessary.

    Adding two dots (electrons) between each neighboring atom pair and three pair isolated to the oxygen. One of the pair on the oxygen is then moved to form a double bond with the carbon so that both carbon and oxygen have complete octets.

Exceptions to the Octet Rule

Bonding on Carbon Compounds

A graphic showing carbon molecules of various lengths.

Functional Groups

  • Functional group
    • A group that is introduced into or substituted in a hydrocarbon chain
    • Gives the hydrocarbon its characteristic properties
    • The group has a heteroatom, an atom other C and H
      • Typically O, S, and N
    • Alcohol
      • A hydroxyl group (-OH) replaces a hydrogen atom in the formula for a hydrocarbon
A carbon molecule with alcohol groups added at four locations and two carbon atoms bridged by an oxygen atom.

Shapes of Molecules

How do we determine the shape of molecules?

VSEPR Derivative Structures

General Formula Number of Bonded Atoms Number of Lone Pairs Molecular Shape Bond Angle Examples
\(\chem{AX_2}\) 2 0 Linear
Three atoms in a straight line.
\(180^\circ\) \(\chem{BeCl_2}\), \(\chem{CO_2}\), \(\chem{HCN}\)
\(\chem{AX_3}\) 3 0 Trigonal Planar
Three atoms equally spaced around a central atom.
\(120^\circ\) \(\chem{BF_3}\), \(\chem{BH_3}\), \(\chem{SO_3}\), \(\chem{NO_3^-}\)
\(\chem{AX_2}\) 2 1 Bent
Three spaced equally spaced around a central atom. Two of the spaces are occupied by an atom while the final space is occupied by a lone pair of electrons
\(120^\circ\) \(\chem{SO_2}\), \(\chem{NO_2^-}\)

More VSEPR Derivative Structures

General Formula Number of Bonded Atoms Number of Lone Pairs Molecular Shape Bond Angle Examples
\(\chem{AX_4}\) 4 0 Teetrahedral
Four atoms equally spaced around a central atom.
\(109.5^\circ\) \(\chem{CH_4}\), \(\chem{CH_2Cl_2}\), \(\chem{SiCl_4}\), \(\chem{POCl_3}\), \(\chem{BrO_4^-}\)
\(\chem{AX_3}\) 3 1 Trigonal Pyramidal
Four spaced equally spaced around a central atom. Three of the spaces are occupied by atoms while the fourth is occupied by a lone pair of electrons
\(109.5^\circ\) \(\chem{NH_3}\), \(\chem{PF_3}\), \(\chem{NH_2Cl}\)
\(\chem{AX_2}\) 2 2 Bent
Four spaced equally spaced around a central atom. Two of the spaces are occupied by atoms while the other two are occupied by lone pairs of electrons
\(109.5^\circ\) \(\chem{H_2O}\), \(\chem{F_2O}\), \(\chem{BrO_2}\), \(\chem{SO_2}\), \(\chem{SCl_2}\)

Steps for VSEPR Structures

  1. Draw a Lewis formula.
  2. Count the number of atoms bonded to the central atom, and count unshared pairs on the central atom.
  3. Add the number of atoms and the number of unshared electron pairs around the central atom. The total indicates the parent structure.
  4. The molecular shape is derived from the parent shape by considering only the positions in the structure occupied by bonded atoms.

Polarity of Molecules

"Like" Dissolves "Like"

  • Ionic salts and polar liquids dissolve better in polar liquids than in nonpolar liquids
  • Nonpolar liquids dissolve better in other nonpolar liquids than in polar liquids
A picture of a container holding two liquids. The yellow liquid is seperated and floating on top of a clear layer.

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