(LM22) Lewis Dot Structure Applications

Lewis Dot Structure Applications

Overview of Lewis Structures

  • It is essential to understand the different varieties of Lewis dot structures and how to approach them effectively.

  • Drawing Lewis structures using the outlined five-step approach may present challenges, reflected in various test questions.

Types of Lewis Structures

  • The following represent various types of Lewis structures:

    • Multiple Bonds: Structures that contain double or triple bonds (e.g., C=C, C≡C).

    • Resonance Structures: Structures where multiple equivalent forms can be drawn for a compound (e.g., Ozone, O₃).

    • Multiple Central Atoms: Compounds with more than one central atom.

    • Hypovalent Compounds: Molecules where a central atom has fewer than 8 electrons (e.g., boron compounds).

    • Hypervalent Compounds: Molecules where a central atom holds more than 8 electrons (e.g., sulfur compounds, phosphorus compounds).

    • Radicals: Molecules that contain an unpaired electron.

Drawing Correct Lewis Structures

  • The same general procedures apply for obtaining a Lewis structure across various types.

  • Understanding these variations helps in confidently deriving correct structures.

Explanation of Resonance

  • Resonance: A phenomenon where a molecule can be represented by two or more valid Lewis structures. This arises due to symmetry allowing the positioning of double bonds in various locations.

  • Commonly observed in certain compounds where bonds must be evenly distributed across multiple sites.

Examples of Resonance:

  • Compounds such as formaldehyde (H₂C=O) exhibit resonance with structures including:

    • Bond sharing among several atoms, such as three bonds across two locations or four bonds across three locations.

    • Benzene (C₆H₆): exhibits resonance having 9 bonds shared among six locations.

Formal Charge and Lewis Structures

  • Formal Charge: Used to determine the most stable Lewis structure among multiple valid forms.

    • Definition:

      \text{Formal Charge} = \text{(Valence Electrons)} - ( \text{Non-bonding Electrons}) - \frac{1}{2}( \text{Bonding Electrons})

  • Ideally, all atoms in a stable structure should have a formal charge of 0, or the smallest possible formal charge difference.

Octet Rule Exceptions

  • The concept of the octet rule is fundamental but has exceptions:

Types of Exceptions:

  1. Hypervalent Compounds:

    • Examples: sulfur difluoride (SF₄), phosphorus pentafluoride (PF₅).

    • Central atoms can accommodate more than 8 electrons due to available d orbitals.

  2. Hypovalent Compounds:

    • Examples include BF₃, AlCl₃ where central atoms such as Boron and Aluminum hold fewer than 8 electrons.

  3. Radicals:

    • Molecules such as Nitric Oxide (NO) with an odd number of electrons, hence containing an unpaired electron.

Drawing Structures with Multiple Central Atoms

  • Approach remains the same as single central atom structures:

    1. Lay out all atoms involved.

    2. Count total valence electrons and assign to peripheral atoms.

    3. Distribute remaining electrons around central atoms and form bonds as required.

  • Notable examples that utilize multiple central atoms include:

    1. Ethylene (C₂H₄)

    2. Glucose (C₆H₁₂O₆)

    3. Amino Acids.

Radicals and Stability

  • The presence of a radical signifies an unpaired electron, complicating molecule stability.

  • For instance, in NO, there exist 11 valence electrons leading to one unpaired electron.

Conclusion on Lewis Structures

  • Use of formal charge helps elucidate the best structures to draw when multiple Lewis structures exist.

  • Understanding and applying formal charge can greatly influence the clarity and accuracy of drawing Lewis dot structures, including considerations like charge distribution among atoms and stability.

Personal notes:

  • B is an exception to the octet rule (can form stable w/ 6 valence electrons)

  • Resonance: a phenomenon where a molecule can be represented by two or more valid Lewis structures

  • The hypervalent case – more than 8 electrons surround the central atom

  • The hypovalent case– fewer than 8 electrons are around the central atom

  • The radical case– an odd number of electrons are in the valence

  • No electrons: Lewis acid

    One electron: radical or magnet

    Two unpaired electrons: Lewis base

  • for formal charge:

  • Overall:

  • formal charge is (valence electrons of atom) - (bonds - dots)

    • The one with the lowest energy — or lowest number — shows the correct Lewis structure

      • Also the most stable

  • Lone pairs is (valence electrons - 8*number of atoms that are not central atoms (look at subscript))/2

    • Ex: ClF_3

      • 1. Find valence electrons 7 + 7(3) = 28

      • 2. use da formula (28 - 8×3)/2 = 2

  • valence electrons is (valence electrons of central atom) + (valence electrons of non-central atom)(subscript of the non-central atom)

    • Ex: PCl_5 = 5 + 7(5) = 40