Flash cards Introduction to Covalent Bonding and Lewis Structures

4.1 Introduction to Covalent Bonding

Overview of Covalent Bonds

Covalent bonds are formed through the sharing of electrons between two atoms. They represent a type of chemical bond in which the participating atoms share pairs of valence electrons. In this context, a molecule is defined as a distinct group of atoms that are bonded together by covalent bonds.

Electron Sharing and Valence Electrons

  1. Unshared Electron Pairs:

    • Unshared electron pairs are also referred to as nonbonded electron pairs or lone pairs.

  2. Noble Gas Configuration:

    • Atoms engage in electron sharing to achieve the electronic configuration that resembles the nearest noble gas from the periodic table.

    • For example:

      • Hydrogen shares 2 electrons to achieve stability.

      • Main group elements share electrons until they attain an octet (8 electrons) in their outermost shell.

Covalent Bond Formation

  1. Formation Conditions:

    • Covalent bonds primarily form when two nonmetals combine or when a metalloid bonds with a nonmetal.

  2. Predicted Number of Bonds:

    • An atom's bonding capacity can be determined:

      • Hydrogen (H): Forms one bond, since it possesses only one valence electron.

      • Valence Electron Counts:

      • Atoms with one, two, or three valence electrons will form 1, 2, or 3 bonds respectively.

      • Atoms with four or more valence electrons will form sufficient bonds to total 8 electrons in the outer shell, expressed through the formula:
        ext{Predicted Number of Bonds} = 8 - ext{Number of Valence Electrons}

General Bonding Rule

  • The general bonding rule (excluding hydrogen) states:
    ext{Number of Bonds} + ext{Number of Lone Pairs} = 4

Importance in Biological Chemistry

  • Covalent compounds play a critical role in biological systems, including various chemical compounds associated with the heart chemistry.

4.2 Lewis Structures

Molecular Formula vs. Lewis Structure

  • A molecular formula conveys the quantity and type of atoms in a compound but lacks information about bonding connectivity. Conversely, a Lewis structure depicts the connectivity between atoms and locates all bonding and nonbonding valence electrons.

General Rules for Drawing Lewis Structures

  1. Valence Electrons Only:

    • Only include valence electrons in the drawing.

  2. Octet for Main Group Elements:

    • All main group elements (except H) should achieve an octet (8 electrons).

  3. Hydrogen Electron Count:

    • Each hydrogen atom should have 2 electrons in bonding arrangements.

Steps to Draw a Lewis Structure

  1. Step [1]: Arranging Atoms:

    • Place bonded atoms adjacent to one another. Position hydrogen and halogens on the periphery, as they each can only form a single bond.

    • Example: Methane (CH4) must position hydrogen atoms appropriately to fulfill bonding rules.

  2. Step [2]: Counting Valence Electrons:

    • For main group elements, the number of valence electrons equals their group number in the periodic table.

    • Total the electrons from all atoms involved in the molecule to determine the total available electrons for the Lewis structure.

    • Example: For CH3Cl:

      • C has 4 electrons (1 × 4),

      • H has 3 electrons (3 × 1),

      • Cl has 7 electrons (1 × 7).

      • Total = 14 valence electrons.

  3. Step [3]: Arranging Electrons:

    • Establish one bond (2 electrons) between each pair of bonded atoms and ensure no atom (except H) has more than 8 electrons. H should have no more than 2 electrons.

    • Utilize remaining electrons to fill octets as lone pairs, starting with peripheral atoms.

  4. Step [4]: Completing Octets with Multiple Bonds:

    • If designated electrons lead to at least one atom lacking an octet, convert a lone pair into a bonding pair to create double bonds.

    • Example for CH3Cl:

      • With all 14 electrons utilized:

      • 2 electrons on each H: 8 on Cl; 8 with lone pairs yield a total utilized structure.

      • If an atom is still deficient, proceed to create multiple bonds as necessary.

Multiple Bonds Explained

  • A double bond consists of a pair of bonding electrons (2 electrons) being shared between two atoms, leading to 4 electrons in total. Conversely, a triple bond implicates the sharing of three pairs of electrons, equating to 6 electrons.

Example of Drawing Lewis Structures with Multiple Bonds

  • Draw the Lewis structure for ethylene (C2H4) using the following steps:

    • Step [1]: Arrange two carbon atoms and four hydrogen atoms.

    • Step [2]: Total valence electrons: 2 C (4 each) + 4 H (1 each) = 12 electrons.

    • Step [3]: Establish bonds, considering C needs an octet. Form a double bond between the carbons where necessary.

    • Step [4]: Confirm that each atom acquires an octet upon rearrangement.

4.3 Exceptions to the Octet Rule

Notable Exceptions

  1. Hydrogen:

    • Hydrogen requires only 2 electrons to achieve stability; hence it does not follow the octet rule.

  2. Group 3A Elements:

    • These atoms (e.g., boron) may not fulfill the octet requirement in neutral configurations.

  3. Third-row Elements (P, S):

    • They possess empty d orbitals, permitting them to accommodate over 8 electrons (an expanded octet).

4.4 Resonance

Utilizing Resonance in Lewis Structures

  • When drawing Lewis structures for polyatomic ions, add one electron for each negative charge and subtract for each positive charge.

    • Example for CN−:

    • C (4 electrons) + N (5 electrons) - 1 (negative charge) = 10 valence electrons.

    • Confirm that all are utilized, ensuring each atom achieves octet fulfillment as necessary.

Drawing Resonance Structures

  • Two Resonance Structures:

    • Resonance structures portray identical atomic arrangements but differ in electron configurations. The true structure is a hybrid, providing stability through the distribution of lone pairs and bonded electrons.

Significance of Resonance in Environmental Chemistry

  • Ozone (O3):

    • Ozone demonstrates resonance through two forms. Formed in the upper atmosphere from O2 and oxygen atoms, ozone shields the Earth's surface from harmful ultraviolet radiation, critical for ecological preservation.

4.5 Naming Covalent Compounds

Process of Naming

Steps for Naming Examples:

  1. Step [1]: Identify the first nonmetal and use its elemental name.

    • For example, in NO2: Nitrogen oxide.

    • In N2O4: Also nitrogen oxide.

  2. Step [2]: Use prefixes to indicate the quantity of each element. Certain conventions apply:

    • Prefix “mono-” is omitted for the first element if it has one atom.

    • Vowel changes may occur to avoid awkwardness in phonetics (ex. monoxide instead of monooxide).

Common Prefixes Table

  • Table 4.1 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:

    • NO2 is nitrogen dioxide.

    • N2O4 is dinitrogen tetroxide.

4.6 Molecular Shape

Lewis Structure Limitations

  • Lewis structures reveal bonding connections but do not provide geometric or shape information for molecules. To ascertain the three-dimensional arrangement:

  1. Count Surrounding Groups:

    • A group may represent either an atom or a lone pair of electrons.

  2. VSEPR Theory Application:

    • Evaluates the optimal arrangement of groups around an atom to maximize separation.

Geometric Shapes Based on Group Count

  1. Two Groups:

    • Geometrically linear with a bond angle of approximately 180° (e.g., CO2).

  2. Three Groups:

    • Trigonal planar with a bond angle of approximately 120° (e.g., H2CO).

  3. Four Groups:

    • Tetrahedral with a bond angle of approximately 109.5° (e.g., CH4).

  4. Trigonal Pyramidal and Bent Shapes:

    • If one or more lone pairs exist, shapes adjust to a trigonal pyramid (NH3) or bent (H2O) geometry, with angles around 107.5° or close.

Common Molecular Shape Summary Table

  • Table 4.2 Total Groups vs. Shape:

    • 2 Groups: Linear, 180°

    • 3 Groups: Trigonal planar, ~120°

    • 4 Groups: Tetrahedral, ~109.5°.\

4.7 Electronegativity and Bond Polarity

Electronegativity Defined

  • Electronegativity quantifies an atom's capability to attract electrons in a bond, reflecting an atom's propensity for gaining electrons.

Bond Polarity Explained

  1. Nonpolar Bonds:

    • If two bonded atoms have equal or similar electronegativities, the bond is considered nonpolar, as electrons are shared equally.

  2. Polar Covalent Bonds:

    • When electrodes differ, leading to uneven sharing of electrons, resulting in a partial separation of charge or a dipole. The more electronegative atom attracts electrons more strongly, symbolized by δ− and δ+ showcasing electron density skew towards the more electronegative atom.

Bond Type Classification Table

  • Table 4.3 Electronegativity Difference vs. Bond Type:

    • Differentiates bond types based on electronegativity differences:

      • Less than 0.5: Nonpolar

      • 0.5 to 1.9: Polar covalent

      • Greater than 1.9: Ionic bond.

4.8 Polarity of Molecules

Determining Polarity

  • A molecule's classification as polar or nonpolar derives from:

  1. Individual Bond Polarities

  2. Overall Molecular Shape.

  • Characteristics of Nonpolar Molecules:

    • Typically, they exhibit no polar bonds or do have polar bonds that cancel one another out through molecular geometry.

  • Polar Molecules:

    • Typically possess one or more polar bonds which do not cancel, resulting in an overall dipole.

Method for Assessing Molecular Polarity**

  • For structures having multiple polar bonds, evaluate:

  1. Identify polar bonds via electronegativity differences.

  2. Determine the atomic arrangement through group counting for geometry.

  3. Assess if individual dipoles work to cancel or reinforce the polarization in the molecule.