Chemical Bonding: Lewis Structures & Polarity

Study Notes

Chemical Bonding

Lewis Structures

Definition: Diagrams that represent the valence electrons of atoms within a molecule.
Components:
- Dots: Represent valence electrons.
- Lines: Represent covalent bonds (shared pairs of electrons).
Steps to Draw:
1. Count total valence electrons.
2. Determine the central atom (usually the least electronegative).
3. Arrange surrounding atoms and connect them with single bonds.
4. Distribute remaining electrons to satisfy octet rule (8 electrons for main group elements).
5. Form multiple bonds if necessary to satisfy octet rule for all atoms.
Exceptions: Some elements can have expanded octets (e.g., P, S) or incomplete octets (e.g., H, Be, B).

Bond Polarity

Definition: Measure of how equally electrons are shared in a bond between two atoms.
Electronegativity:
- Ability of an atom to attract shared electrons.
- Scale (0 to 4): Fluorine (highest, ~4.0) to Cesium (lowest, ~0.7).
Types of Bonds:
- Nonpolar Covalent Bond: Equal sharing of electrons (electronegativity difference < 0.4).
- Polar Covalent Bond: Unequal sharing of electrons (electronegativity difference between 0.4 and 1.7).
- Ionic Bond: Complete transfer of electrons (electronegativity difference > 1.7).

Covalent Bonds

Definition: Chemical bonds formed by the sharing of electrons between atoms.
Types:
- Single Bonds: One pair of shared electrons (e.g., H-H).
- Double Bonds: Two pairs of shared electrons (e.g., O=O).
- Triple Bonds: Three pairs of shared electrons (e.g., N≡N).
Characteristics:
- Usually formed between nonmetals.
- Can be polar or nonpolar based on electronegativity differences.
- Lower melting and boiling points compared to ionic compounds.
Strength: Generally, triple bonds are stronger than double, and double bonds are stronger than single bonds.

Lewis Structures

Diagrams visually showing valence electrons of atoms in a molecule.
Components:
- Dots represent valence electrons.
- Lines indicate covalent bonds formed by shared electron pairs.
Steps to Draw:
- Calculate total valence electrons from all atoms.
- Identify the central atom, typically the least electronegative.
- Position surrounding atoms and connect with single bonds initially.
- Distribute remaining electrons to achieve the octet rule, which aims for 8 electrons around main group elements.
- Create multiple bonds if necessary to fulfill the octet for all atoms.
Exceptions:
- Some elements can expand their octets (e.g., phosphorus, sulfur).
- Elements like hydrogen, beryllium, and boron may have incomplete octets.

Bond Polarity

Indicates the distribution of electron sharing between two bonded atoms.
Electronegativity:
- Measures an atom's ability to attract shared electrons.
- Scale ranges from 0 to 4, with fluorine having the highest value (~4.0) and cesium the lowest (~0.7).
Types of Bonds:
- Nonpolar Covalent Bond: Electrons are shared equally, electronegativity difference is less than 0.4.
- Polar Covalent Bond: Electrons are shared unequally, electronegativity difference ranges from 0.4 to 1.7.
- Ionic Bond: Involves complete electron transfer, with an electronegativity difference greater than 1.7.

Covalent Bonds

Chemical bonds formed through the sharing of electron pairs between atoms.
Types of Covalent Bonds:
- Single Bonds: One pair of electrons shared (e.g., H-H).
- Double Bonds: Two pairs of electrons shared (e.g., O=O).
- Triple Bonds: Three pairs of electrons shared (e.g., N≡N).
Characteristics:
- Commonly occurs between nonmetals.
- Can exhibit polarity depending on the electronegativity difference between atoms.
- Generally, covalent compounds have lower melting and boiling points compared to ionic compounds.
Bond Strength: Triple bonds are stronger than double bonds, which in turn are stronger than single bonds.

Description

Explore the concepts of Lewis structures and bond polarity in chemical bonding. This quiz covers the steps to draw Lewis structures, exceptions to the octet rule, and the electronegativity scale. Test your knowledge on how electrons are shared in chemical bonds and what influences bond polarity.