Chapter 12 Study Guide Notes

Chemical Bonds

• Chemical bonds arise from strong electrostatic forces of attraction between atoms, involving the valence electrons.
• Types of chemical bonds:
  • Ionic: Exchange of electrons.
  • Covalent: Sharing of electrons between two atoms.
  • Metallic: Sharing of electrons among many metal atoms (discussed later).

Valence Electrons and Lewis Dot Symbols

• Most substances are compounds resulting from bonded atoms.
• Valence electrons are the highest energy level s and p electrons in an atom and determine bonding behavior.
• Lewis dot symbols represent valence electrons: atomic symbol surrounded by dots for each valence electron.
  • Example:
    • Lithium ([He]2s1): Li.
    • Magnesium ([Ne]3s2): .Mg.
    • Fluorine ([He]2s22p5): F
• Elements in the same group have the same number of valence electrons, hence the same number of dots.
• Noble gases have eight valence electrons (except He, which has two), forming a complete octet.

Formation of Ions and the Octet Rule

• Atoms gain or lose electrons to achieve a noble gas electron configuration (octet rule).
• The octet rule signifies the desire of atoms to have a complete set of eight valence electrons, similar to noble gases.
• Metal atoms (s-block) lose valence electrons to attain the electron configuration of the preceding noble gas, forming cations (positively charged).
  • Example:
    • Na {.} \rightarrow Na^+ + e^- ([Ne]3s1 [Ne])
    • Mg \rightarrow Mg^{2+} + 2e^- ([Ne]3s2 [Ne])
• Non-metals (p-block) gain electrons to achieve the electron configuration of the noble gas in the same period, forming anions (negatively charged).
  • Example:
    • F + e^- \rightarrow F^- ([He]2s22p5 [Ne])
    • P + 3e^- \rightarrow P^{3-} ([Ne]3s23p3 [Ar])
• Group 14 elements do not readily form ions due to the difficulty of gaining or losing four electrons.

Ionic Compounds

• Ionic compounds are formed through the electrostatic attraction between anions and cations (ionic bonds).
• Compounds are neutral, with equal numbers of positive and negative charges.
• Metal cations and non-metal anions combine to form ionic compounds.
• Chemical formulas represent the group of atoms necessary for neutrality, with subscripts indicating the number of atoms of each kind.
• The cation is always listed first in a chemical formula.

Covalent Bonds

• Covalent bonds involve the sharing of one or more pairs of electrons between atoms.
• Atoms held together by covalent bonds form a molecule.
• Molecules can be represented by molecular formulas (showing type and number of atoms) or structural formulas (showing how atoms are connected).

Octet Rule in Covalent Bonds

• Covalent bonds form between non-metal atoms.
• Atoms share electrons to achieve an octet of valence electrons (or two for hydrogen), resembling a noble gas configuration.
• A shared pair of electrons can be represented by a line segment connecting the atoms.
• Electrons not involved in bonding are called unshared pairs (lone pairs).
• Atoms may share more than one pair of electrons, forming double bonds (two pairs) or triple bonds (three pairs).

Diatomic and Polyatomic Molecules

• Diatomic molecules consist of two atoms (e.g., O2, N2, F2, Cl2, Br2, I2, H2).
• Seven elements exist as diatomic molecules to achieve an octet of valence electrons.
• Polyatomic molecules contain more than two covalently bonded atoms.
• Complex molecules can contain hundreds of atoms.
• Simple polyatomic molecules have more than one atom bonded to a central atom, all achieving an octet (or two for hydrogen).

Polyatomic Ions

• A group of atoms may form a polyatomic ion by gaining or losing electrons.

Rules for Writing Lewis Structures

• Diatomic Molecules
  1. Write the symbol of each atom and determine the number of valence electrons.
  2. Add the valence electrons from each atom to obtain the total number of valence electrons.
  3. Write the atomic symbols of the two atoms with the atom that had the fewest valence electrons on the left.
  4. Place an octet of electrons around the atom on the left (two electrons if it is hydrogen).
  5. Place the remaining electrons around the other atom in pairs.
  6. If the atom on the right does not have an octet of electrons, move one or more pairs of electrons from the other atom into the bond until it does.
  7. Replace the dots in the bonds with line segments, one for each pair of electrons.
• Polyatomic Molecules
  1. Write the symbol of each atom and the number of valence electrons.
  2. Add the valence electrons from each atom to obtain the total number of valence electrons.
  3. Arrange the Lewis symbols of the atoms so that the atom with the fewest valence electrons is in the center and the others are arranged around it. (Hydrogen can never be the central atom)
  4. Place a pair of electrons between the central atom and each of the other atoms.
  5. Place the remaining electrons around the other atoms in pairs so each has an octet (two for hydrogen). If any remain place them around the central atom.
  6. If the central atom does not have an octet of electrons, move one or more pairs of electrons from one of the other atoms to form multiple bonds until all the atoms have a complete octet.
  7. Replace the dots in the bonds with line segments, one for each pair of electrons.
• Polyatomic Ions
  1. Write the symbol of each atom and the number of valence electrons.
  2. Add the valence electrons from each atom to obtain the total number of valence electrons.
  3. If the ion is negative add additional electrons equal to the charge. If the ion is positive subtract electrons equal to the charge.
  4. Arrange the symbols of the atoms so that the atom with the fewest valence electrons is in the center and the others are arranged around it. (Hydrogen can never be the central atom)
  5. Place a pair of electrons between the central atom and each of the other atoms.
  6. Place the remaining electrons around the other atoms so each has an octet. If any remain place them around the central atom.
  7. If the central atom does not have an octet of electrons, move one or more pairs of electrons from the other atoms into the bond until it does.
  8. Replace the dots in the bonds with line segments, one for each pair of electrons. Enclose the ion in brackets and indicate the charge outside the brackets.

Shapes of Molecules (VSEPR Theory)

• Molecules have three-dimensional shapes.
• Valence Shell Electron Pair Repulsion (VSEPR) theory: electron pairs around the central atom (shared and unshared) repel each other and take positions as far apart as possible.
• To determine the shape of a simple molecule:
  1. Count the number of atoms bonded to the central atom.
  2. Count the number of unshared electron pairs around the central atom.
  3. The sum of these is the total number of electron groups.

Bond Polarity

• Electrons in covalent bonds are not always shared equally.
• Electronegativity determines the pull on shared electrons.
• Polar bond: unequal sharing of electrons, leading to slight negative and positive charges on either side (overall neutral).
• Non-polar bond: equal sharing of electrons, neutral overall.
• Electronegativity differences:
  • Polar: 0.4 - 1.9
  • Non-polar: 0 - 0.4
  • Ionic: greater than 1.9 (involving a metal and non-metal)

Examples of Bond Polarity

• I-Br: |2.5 - 2.8| = 0.3 Non-polar (electrons shared equally).
• S-O: |2.5 - 3.5| = 1.0 Polar (oxygen side is negative, sulfur side is positive).
• Na-Cl: |0.9 - 3.0| = 2.1 Ionic (electrons are exchanged).

Polar Molecules

• Molecules with polar bonds may be polar themselves, depending on their shape.
• Asymmetrical charge distribution: polar molecule.
• Symmetrical charge distribution: non-polar molecule.

VSEPR Shapes

Intermolecular Forces

• Covalent molecules are held together in solids and liquids by intermolecular forces.
• Ionic compounds form large crystals held by strong electrostatic attraction, while molecules are discrete groups.
• Dipole-dipole forces: attraction between the negative side of one polar molecule and the positive side of another.
• Stronger polarity results in stronger attraction.

Hydrogen Bonding

• Hydrogen bonding: a strong intermolecular force involving a hydrogen atom bonded to an oxygen atom attracted to an oxygen atom on another molecule.
• It is not a true bond, but a strong force that holds water molecules together in liquid water and ice.

Comparison of Bonds

• Ionic bonds are very strong due to the attraction of ions, resulting in high-melting solids.
• Covalent bonds are also strong, forming discrete molecules that can be solids, liquids, or gases, depending on intermolecular forces. More often a gas than an ionic bond. Resulting materials depends on the strength of the intermolecular forces.