Exam 3 Lewis Structures

Chapter Overview

Title: Chemistry: Atoms First 2e Chapter 4Author Contributions: Joe DePasquale, Northeastern University

Key Topics Covered:4.3 Naming Covalent Molecules, Acids, and Hydrated Compounds4.4 Lewis Symbols and Structures4.5 Formal Charges and Resonance

4.3 Naming Molecular Compounds

Definition: Compounds composed of two or more covalently bonded nonmetals.Naming Convention: Use prefixes to denote the number of atoms present:

• mono- (1)

• di- (2)

• tri- (3)

• tetra- (4)

• penta- (5)

• hexa- (6)

• hepta- (7)

• octa- (8)

• nona- (9)

• deca- (10)

• hemi- (1/2)

Naming Examples:

• SF6: sulfur hexafluoride

• Mono- is omitted if the first element is present in one atom: N2O4: dinitrogen tetroxide

• SO3: sulfur trioxide

• P4S10: tetraphosphorus decasulfide

Table 4.7: Names of Some Molecular Compounds

• SO₂: sulfur dioxide

• BCl₃: boron trichloride

• NO₂: nitrogen dioxide

• N₂O₄: dinitrogen tetroxide

• SF₆: sulfur hexafluoride

4.4 Lewis Symbols and Structures

Lewis Theory: Focuses on valence electrons in bonding. Dots represent valence electrons in elements.

Drawing Lewis Symbols:

• Place dots for valence electrons around the element's symbol; for instance, carbon (C) has 4 valence electrons, represented as

  • 
  C 
  • 

• Maximum of two dots per side should be drawn to avoid overcrowding.

Octet Rule

Atoms aim to achieve stable electron configurations (usually eight valence electrons). This stability is similar to that of noble gases, like neon, which naturally have full outer shells. For example, in the formation of NaCl, sodium (Na) loses one electron to achieve a stable configuration, while chlorine (Cl) gains that electron to also stabilize, following the octet rule.

Bonding Pairs vs Lone Pairs

• Bonding pairs: Electrons shared between atoms form covalent bonds. For instance in H2O, each hydrogen contributes one electron to create bonds with oxygen.

• Lone pairs: Nonbonding valence electrons belong to a specific atom, like in H2O where oxygen has two lone pairs that influence the molecule's shape.

Types of Bonds

• Single Bond: One pair of electrons shared (e.g., H–H bond in H2).

• Double Bond: Two pairs of electrons shared (e.g., in CO2, the carbon draws on two shared pairs of electrons from each oxygen).

• Triple Bond: Three pairs of electrons shared (strongest and shortest, as seen in N2).

Exercises

• Write names for given compounds. These exercises require writing chemical names and structures for specific compounds including hydrates and ions, such as determining the name of Cu(NO3)2, which is copper(II) nitrate.

Naming Acids

• Definition: Acids release H⁺ ions in water.

• Binary Acids: Use "hydro-" prefix, changing the nonmetal's name with "-ic", and adding "acid" (HCl becomes hydrochloric acid).

• Oxyacids: Omit "hydrogen"; change the anion name (-ate to -ic, -ite to -ous) (HNO3, as the anion (NO3) is a -ate becomes nitric acid).

Common Oxyacids

• HNO₃: nitric acid

• H₂SO₄: sulfuric acid

Naming Ionic Hydrates

Hydrate: Ionic compound with water molecules.Naming: Name the anhydrous compound and add "hydrate" with a Greek prefix for water molecules. For example, CuSO₄·5H₂O is named copper(II) sulfate pentahydrate because it contains five waters attached to the copper sulfate compound.

Lewis Structures for Compounds

Steps to write Lewis Structures:

1. Determine the skeletal structure (H is always terminal).

2. Calculate total valence electrons.

3. Distribute electrons to achieve octets or duets.

4. Form double/triple bonds if necessary to complete octets.For example, the Lewis structure for CO₂ shows a double bond between carbon and each oxygen.

Practice Questions

Questions on identifying bond types and drawing structures for common compounds like HF, CH₄, and N₂.