Chapter+4+slides

4.1 Introduction to Covalent Bonding

• Covalent bonds result from the sharing of electrons between two atoms.

  • A covalent bond is a two-electron bond in which the bonding atoms share valence electrons.

  • A molecule is a discrete group of atoms held together by covalent bonds.

• Nonbonded Electron Pairs:

  • Unshared electron pairs are called nonbonded electron pairs or lone pairs.

  • Atoms share electrons to achieve the electronic configuration of the nearest noble gas.

    • Example: Hydrogen shares 2 electrons.

    • Other main group elements aim for an octet of electrons in their outer shell.

• Covalent Bonding and the Periodic Table:

  • Lewis structures represent electron-dot structures for molecules, indicating the location of all valence electrons.

4.2 Lewis Structures

• Covalent Compounds:

  • Covalent bonds are formed when two nonmetals combine or when a metalloid bonds with a nonmetal.

  • General bonding rules:

    • Hydrogen forms one bond.

    • Atoms with one to three valence electrons form one to three bonds.

    • Atoms with four or more valence electrons form enough bonds to achieve an octet, calculated by:

      • Predicted number of bonds = 8 - number of valence electrons.

• Lewis Structures:

  • A molecular formula indicates the number and identity of all atoms but does not clarify their bonding.

  • A Lewis structure shows connectivity between atoms and the location of bonding and nonbonding valence electrons.

• Drawing Lewis Structures:

  • Step 1: Arrange atoms that are likely bonded together (e.g., H and halogens on the periphery).

    • For example, in CH4, Carbon (C) typically forms four bonds.

  • Step 2: Count the number of valence electrons.

    • For main group elements, the valence electrons equal the group number.

  • Step 3: Arrange electrons around the atoms while ensuring:

    • Use one bond (two electrons) between every two atoms.

    • No more than 8 electrons for main group elements and 2 electrons for Hydrogen.

  • Step 4: If atoms still lack an octet after using all valence electrons, utilize multiple bonds to fill octets as necessary (one lone pair can convert into two electrons).

4.3 Exceptions to the Octet Rule

• Exceptions:

  • Hydrogen needs only 2 electrons.

  • Elements in group 3A may have fewer than 8 electrons (example: Boron can only have 6 electrons).

  • Elements in the third row, like Phosphorus (P) and Sulfur (S), can have more than 8 electrons due to available empty d orbitals.

4.5 Naming Covalent Compounds

• Naming Covalent Molecules:

  • Step 1: Name the first nonmetal by its element name and the second using the suffix “-ide”.

    • Examples: NO (nitrogen oxide), N2O4 (dinitrogen tetroxide).

  • Step 2: Use prefixes to indicate the number of atoms in each element, omitting “mono-” for the first element:

    • Common prefixes:

      • 1: Mono

      • 2: Di

      • 3: Tri

      • 4: Tetra

      • 5: Penta

      • 6: Hexa

      • 7: Hepta

      • 8: Octa

      • 9: Nona

      • 10: Deca.

4.6 Molecular Shape

• Determining Molecular Shape:

  • The Lewis structure provides information on atomic connectivity but does not indicate geometry.

  • Use VSEPR theory to determine molecular shape based on the arrangement of groups around an atom.

  • Common geometries:

    • Linear (2 groups): Bond angle of 180° (e.g., CO2).

    • Trigonal Planar (3 groups): Bond angles of 120° (e.g., BF3, H2CO).

    • Tetrahedral (4 groups): Bond angles of 109.5° (e.g., CH4).

    • Trigonal Pyramid (4 groups including 1 lone pair): Bond angles of 107° (e.g., NH3).

    • Bent (4 groups including 2 lone pairs): Bond angle of 105° (e.g., H2O).

4.7 Electronegativity and Bond Polarity

• Electronegativity:

  • Defined as an atom’s attraction for electrons in a bond.

  • Nonpolar bonds occur when electronegativities are equal or similar, leading to equal sharing of electrons.

  • Polar covalent bonds arise when electronegativities differ, resulting in a partial separation of charge (indicated by δ- and δ+).

• Electronegativity Differences:

  • Less than 0.5: Nonpolar

  • 0.5 to 1.9: Polar covalent

  • Greater than 1.9: Ionic (electrons transferred).

4.8 Polarity of Molecules

• Classifying Molecules:

  • The polarity of a molecule depends on the individual bond polarities and overall molecular shape.

  • Nonpolar molecules have no polar bonds or cancel out polarities.

  • Polar molecules feature one or more polar bonds with dipoles that do not cancel.

• Determining Molecules' Polarity:

  • Identify polar bonds based on electronegativity.

  • Count the groups around central atoms to assess shape.

  • Evaluate if dipoles reinforce or cancel.