Covalent Compounds Study Notes

Covalent Compounds

4.1 Introduction to Covalent Bonding

Covalent bonds result from the sharing of electrons between two atoms.

  • Covalent Bond: A two-electron bond in which the bonding atoms share valence electrons.

  • Molecule: A discrete group of atoms held together by covalent bonds.

4.1 Introduction to Covalent Bonding (Continued)

Unshared electron pairs are called nonbonded electron pairs or lone pairs. Atoms share electrons to achieve the electronic configuration of the noble gas closest to them in the periodic table:

  • Hydrogen shares 2 electrons.

  • Other main group elements share electrons until they reach an octet of electrons in their outer shell.

4.1A Covalent Bonding and the Periodic Table

Lewis Structures

  • Lewis Structures: Electron-dot structures for molecules that show the location of all valence electrons.

Formation of Covalent Bonds

Covalent bonds are formed when:

  • Two nonmetals combine.

  • A metalloid bonds to a nonmetal.

Number of Covalent Bonds

To determine how many covalent bonds a particular atom will form:

  • Hydrogen forms one bond, with its one valence electron.

  • Atoms with one, two, or three valence electrons form one, two, or three bonds, respectively.

  • Atoms with four or more valence electrons form enough bonds to give an octet. The following formula is used:
    Predicted number of bonds=8−number of valence electrons\text{Predicted number of bonds} = 8 - \text{number of valence electrons}

General Bonding Rule

  • General Rule for Bonding (except for H): Number of bonds + Number of lone pairs = 4

4.1B Focus on the Human Body

  • There are many covalent compounds related to the chemistry of the heart.

4.2 Lewis Structures

  • A molecular formula shows the number and identity of all of the atoms in a compound, but does not indicate which atoms are bonded to each other.

  • A Lewis structure shows the connectivity between atoms, as well as the location of all bonding and nonbonding valence electrons.

4.2A Drawing Lewis Structures

General Rules for Drawing Lewis Structures

  1. Draw only valence electrons.

  2. Give every main group element (except hydrogen) an octet of electrons.

  3. Give each hydrogen 2 electrons.

How to Draw a Lewis Structure
Step [1]: Arrange the Atoms

  • Place H and halogens on the periphery, since they can only form one bond.
    Example for CH4 included in resource.

Step [2]: Count the Valence Electrons

  • For main group elements, the number of valence electrons is equal to the group number. This gives a total number of electrons used in the Lewis structure.
    Example for CH3Cl: 14 total valence electrons

Step [3]: Arrange Electrons Around Atoms

  • Place one bond (two electrons) between every two atoms.

  • For main group elements, give no more than 8 electrons.

  • For H, give no more than 2 electrons.

  • Use all remaining electrons to fill octets with lone pairs, starting with peripheral atoms.

Step [4]: Use Multiple Bonds, If Needed

  • If all valence electrons are used and an atom still does not have an octet, use multiple bonds to complete the structure.
    Example: CH3Cl

4.2B Multiple Bonds

Concept of Multiple Bonds

  • One lone pair of electrons can be converted into one bonding pair of electrons for each 2 electrons needed to complete an octet in a Lewis Structure.

  • Double Bond: Contains four electrons in two 2-electron bonds.

  • Triple Bond: Contains six electrons in three 2-electron bonds.

Example: Drawing Lewis Structure for C2H4
Step 1: Arrange the Atoms
Step 2: Count the Valence Electrons

  • Total of 12 electrons available.

Step 3: Add the Bonds and Lone Pairs

  • Total of 12 valence electrons are used up in constructing the Lewis Structure.

Step 4: Adjust Electrons to Form a Double Bond

  • Change one lone pair into one bonding pair of electrons, forming a double bond to ensure each C attains an octet.

4.3 Exceptions to the Octet Rule

  • Most common elements generally follow the octet rule.

  • Hydrogen is an exception; it needs only 2 electrons in bonding.

Elements in Group 3A

  • Elements in group 3A do not have enough valence electrons to form an octet in a neutral molecule.

Elements in the Third Row

  • Elements in the third row have empty d orbitals available to accept electrons, thus elements like P and S may have more than 8 electrons surrounding them.

4.4 Resonance

Drawing Lewis Structures for Polyatomic Ions

  • To draw, add one electron for each negative charge and subtract one electron for each positive charge.
    Example: For CN−: 10 electrons total.

  • All valence electrons are used, but C lacks an octet.

Resonance Structures

  • Resonance Structures: Two or more Lewis structures with the same arrangement of atoms but different arrangements of electrons.

  • The true structure is a hybrid of the resonance structures, stabilizing a molecule by spreading out lone pairs and electron pairs over a larger area.

Focus on the Environment: Ozone (O3)

  • Ozone can be drawn with two resonance structures and is formed in the upper atmosphere by the reaction of O2 and oxygen atoms, shielding the Earth’s surface from destructive UV radiation.

4.5 Naming Covalent Compounds

Steps in Naming a Covalent Molecule

  1. Name the first nonmetal by its element name and the second using the suffix “-ide.”
    Example:

  • a) NO2: nitrogen oxide

  • b) N2O4: nitrogen oxide.

  1. Add prefixes to show the number of atoms of each element using Table 4.1 for common prefixes.

  • The prefix “mono-” is omitted when only one atom of the first element is present but is retained for the second element if applicable.

  • If the combination would place two vowels next to one another, omit the first vowel.
    Examples:

  • mono + oxide = monoxide (not monooxide).




  2. Common Prefixes in Nomenclature:


    Number of Atoms

    Prefix



    1

    Mono



    2

    Di



    3

    Tri



    4

    Tetra



    5

    Penta



    6

    Hexa



    7

    Hepta



    8

    Octa



    9

    Nona



    10

    Deca





    Examples Continued:



    • a) NO2: nitrogen dioxide

    • b) N2O4: dinitrogen tetroxide

    4.6 Molecular Shape

    Understanding Molecular Geometry

    The Lewis structure only indicates which atoms are connected without showing geometry. To determine the shape around an atom, first ascertain how many groups (either atoms or lone pairs) surround the atom. Use Valence Shell Electron Pair Repulsion (VSEPR) theory to determine the geometry, wherein the most stable arrangement keeps groups as far apart as possible.

    Common Molecular Shapes

    • Linear: An atom surrounded by two groups is linear with a bond angle of 180°.

    • Trigonal Planar: An atom surrounded by three groups has bond angles of 120°.

    • Tetrahedral: An atom surrounded by four groups has bond angles of 109.5°.

    Group Configurations and Shapes

    • CO2 example: Linear — 180°

    • BF3 example: Trigonal Planar — 120°

    • CH4 example: Tetrahedral — 109.5°

    Four Groups Around an Atom
    1. Tetrahedral Geometry

    • If four groups around the atom include one lone pair, the geometry becomes trigonal pyramidal with bond angles of 107°.

    • NH3 (ammonia) serves as an example.

    2. Bent Geometry

    • If four groups include two lone pairs, the geometry is bent with bond angles of 105°.

    • H2O serves as an example.

    4.7 Electronegativity and Bond Polarity

    Definition of Electronegativity

    • Electronegativity: A measure of an atom’s attraction for electrons in a bond. It signifies how much a specific atom “wants” electrons.

    Bond Types Based on Electronegativity Difference

    1. Nonpolar Bonds: If the electronegativities of two bonded atoms are equal or similar, the bond is nonpolar with equal sharing of electrons between the atoms.

    2. Polar Covalent Bonds: When atoms have different electronegativities, it results in a polar covalent bond or dipole, characterized by a partial separation of charge where electrons are pulled towards the more electronegative atom.

    Table 4.3 Electronegativity Difference and Bond Type

    Electronegativity Difference

    Bond Type

    Electron Sharing

    Less than 0.5 units

    Nonpolar

    Electrons are equally shared.

    0.5 to 2.0 units

    Polar Covalent

    Electrons are unequally shared towards the more electronegative element.

    Greater than 2.0 units

    Ionic

    Electrons are transferred from the less electronegative element to the more electronegative element.

    4.8 Polarity of Molecules

    Dependent Factors for Molecular Polarity

    Classification of a molecule as polar or nonpolar relies on:

    • The polarity of the individual bonds.

    • The overall shape of the molecule.

    Determining Molecular Polarity with Multiple Bonds

    1. Identify all polar bonds based on electronegativity differences.

    2. Determine the shape around individual atoms by counting groups.

    3. Assess if individual dipoles reinforce or cancel.