Bonding, Molecular Structure, and Resonance: Study Guide for General Chemistry

Polyatomic Ions

Common Polyatomic Ions

Polyatomic ions are ions composed of two or more atoms covalently bonded, carrying a net charge. They are frequently encountered in chemical compounds and reactions.

• Nitrate, NO3-

• Carbonate, CO32-

• Bicarbonate, HCO3-

• Sulfate, SO42-

• Ammonium, NH4+

• Acetate, C2H3O2-

Example: Sodium sulfate contains the sulfate ion: Na2SO4.

Electronegativity and Bond Types

Electronegativity

Electronegativity is a measure of an atom's ability to attract shared electrons in a chemical bond. It is a key concept for predicting bond polarity and molecular properties.

• Definition: Electronegativity quantifies the tendency of an atom to attract electrons in a covalent bond.

• Trends: Electronegativity increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Highest Electronegativity: Fluorine (F) is the most electronegative element.

Example: Oxygen is more electronegative than sulfur.

Covalent and Polar Covalent Bonds

Covalent bonds involve the sharing of electron pairs between atoms. The nature of the bond depends on the difference in electronegativity between the atoms.

• Covalent Bond: Electrons are shared equally between atoms of similar electronegativity (e.g., H2, Cl2).

• Polar Covalent Bond: Electrons are shared unequally due to a difference in electronegativity (e.g., H2O, HF).

Example: In HCl, chlorine is more electronegative than hydrogen, resulting in a polar covalent bond.

Predicting Bond Type Using Electronegativity

The difference in electronegativity between two atoms can be used to predict the type of bond formed.

• Nonpolar Covalent: $ext{Difference} < 0.4$

• Polar Covalent: $0.4 ext{ to } 1.7$

• Ionic: $ext{Difference} > 1.7$

Example: The bond in NaCl is ionic, while the bond in CH4 is nonpolar covalent.

Lewis Structures and the Octet Rule

Writing Lewis Structures

Lewis structures represent the arrangement of valence electrons in molecules and ions. They are essential for understanding bonding and molecular geometry.

• Steps:

  1. Count total valence electrons.

  2. Arrange atoms and connect with single bonds.

  3. Distribute remaining electrons to complete octets (or duets for H).

  4. Use double/triple bonds if necessary.

• Octet Rule: Most atoms aim for eight electrons in their valence shell.

• Exceptions: Some molecules violate the octet rule (e.g., BF3, SF6).

Example: The Lewis structure for CO2 is O=C=O.

Resonance Structures

Resonance occurs when more than one valid Lewis structure can be drawn for a molecule or ion. The actual structure is a hybrid of all resonance forms.

• Definition: Resonance structures differ only in the placement of electrons, not atoms.

• Recognition: Molecules with delocalized electrons (e.g., O3, NO3-) exhibit resonance.

Example: The nitrate ion, NO3-, has three resonance structures.

Formal Charge

Formal charge helps determine the most stable Lewis structure for a molecule or ion.

• Calculation: $ext{Formal charge} = ext{Valence electrons} - ( ext{Nonbonding electrons} + rac{1}{2} imes ext{Bonding electrons})$

• Best Structure: The most stable structure has formal charges closest to zero and negative charges on the most electronegative atoms.

Example: In the carbonate ion, CO32-, formal charges help identify the best resonance form.

Bond Polarity and Dipoles

Bond Dipole Arrows

Bond dipoles indicate the direction of electron density shift in a polar bond. Arrows point toward the more electronegative atom.

• Arrow Notation: Draw an arrow from the less electronegative atom to the more electronegative atom, with a plus sign at the tail.

• Application: Used to show molecular polarity and predict interactions.

Example: In HCl, the dipole arrow points from H to Cl.

VSEPR Theory and Molecular Geometry

VSEPR Theory

Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the shapes of molecules and ions based on electron pair repulsion.

• Electron Domains: Regions of electron density (bonds and lone pairs) around a central atom.

• Basic Geometries: Linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral.

Example: CH4 is tetrahedral; CO2 is linear.

Bond Angles and Lone Pairs

Bond angles are determined by the number of electron domains and the presence of lone pairs, which compress bond angles.

• Linear: 180°

• Trigonal Planar: 120°

• Tetrahedral: 109.5°

• Trigonal Bipyramidal: 90°, 120°

• Octahedral: 90°

• Lone Pairs: Reduce bond angles due to increased repulsion.

Example: In NH3, the bond angle is about 107° due to one lone pair.

Summary Table: Electron Geometry and Bond Angles

Electron Geometry	Bond Angle	Example
Linear	180°	CO2
Trigonal Planar	120°	BF3
Tetrahedral	109.5°	CH4
Trigonal Bipyramidal	90°, 120°	PCl5
Octahedral	90°	SF6

Additional info: Lone pairs compress bond angles below ideal values; for example, H2O has a bond angle of about 104.5° due to two lone pairs.