Chemistry 2.3 Lewis Structures Flashcards

Study Notes

Molecular Structures

Lewis Dot Structures represent molecular structures.
Sodium (Na) has 1 valence electron; Chlorine (Cl) has 7 valence electrons.
Ionic bond forms between Na and Cl: Na donates its electron to Cl, resulting in Na⁺ and Cl⁻.

Lewis Dot Structures and Covalent Bonding

Lewis Dot Structures are primarily used for covalent bonding.
Two chlorine atoms share unpaired electrons to achieve filled octets, indicating a covalent bond.

Octet Rule and Exceptions

Most atoms prefer a filled octet, with a total of 8 electrons being optimal due to shell structure.
Key exceptions include:
- Hydrogen (H) seeks 2 electrons.
- Beryllium (Be) typically seeks 4 electrons.
- Boron (B) and Aluminum (Al) typically seek 6 electrons.
Elements in the 3rd row and below can exceed the octet by utilizing d orbitals.
Molecules with an odd number of electrons (e.g., NO) cannot satisfy the octet rule for all atoms.

Lewis Structure Example: Methane (CH₄)

Total of 8 valence electrons: Carbon has 4, each hydrogen has 1.
Carbon is central because it forms 4 bonds; Hydrogen requires 2 electrons each.
All valence electrons are utilized, making the structure stable.

Lewis Structure Example: Ammonia (NH₃)

Total of 8 valence electrons, with nitrogen making 3 bonds to hydrogen atoms.
2 electrons remain after bonding that are assigned to the central nitrogen atom, resulting in a stable structure.

Lewis Structure Example: Nitrogen Trichloride (NCl₃)

Total of 26 valence electrons, with nitrogen as the central atom bonded to 3 chlorine atoms.
Each chlorine atom gets 6 electrons to complete its octet, while nitrogen is happy with the configuration.

Lewis Structure Example: Hydrogen Cyanide (HCN)

Total of 10 valence electrons; Carbon as the central atom bonded to H and N.
6 electrons are given to N, requiring sharing from nitrogen to achieve a triple bond and stability for carbon.

Lewis Structure Example: Carbon Dioxide (CO₂)

Total of 16 valence electrons, with carbon bonded to two oxygen atoms.
Each oxygen is filled first, leading to the formation of double bonds to satisfy carbon's octet.

Lewis Structures and Formal Charge

Formal charge should ideally be 0 for the best structural representation.
For CO₂:
- Two double bond structure gives zero formal charges for all atoms.
- A triple bond scenario leads to a +1 and a -1 formal charge, indicating it is less favorable.

Lewis Structure Example: Sulfur Tetrafluoride (SF₄)

Total of 28 valence electrons; Sulfur as central atom bonded to 4 fluorine atoms.
Sulfur can violate the octet rule due to its ability to utilize d orbitals.

Lewis Structure Example: Xenon Difluoride (XeF₂)

Total of 22 valence electrons; Xenon bonded to 2 fluorines with 3 lone pairs added to fulfill valence needs.
Again, xenon can violate the octet rule thanks to its position in the periodic table.

Polar Molecules

Polar bonds occur when electronegativity differences exist.
Examples include:
- N-H bond in NH₃ (polar due to nitrogen's higher electronegativity).
- C-O bond in CO₂ (polar due to oxygen's higher electronegativity).
A difference in electronegativity >0.4 indicates a likely polar bond, except for H-C bonds and bonds between identical elements.

Description

This quiz focuses on the Lewis Dot Structures, specifically molecular structures and how elements like Sodium and Chlorine interact. Learn the significance of valence electrons and the formation of ionic bonds through illustrative examples. Test your understanding of how to represent these molecular arrangements accurately.