Chemical Bonding Concepts

Chemical Bonding

• A compound is a substance made of two or more different elements combined chemically.
• Attractive forces, called chemical bonds, hold the atoms together in a compound.
• The Octet Rule describes how atoms tend to gain, lose, or share electrons to achieve an electron arrangement with eight electrons in their outermost energy level.
• Noble gases have stable electron configurations with eight or two electrons in their outermost level, so they do not react.
• Transition metals often do not follow the octet rule
• Hydrogen, lithium, and beryllium often achieve a full outer shell with just two electrons.
• An ion is a charged atom or group of atoms that has gained or lost electrons.
• Ionic bonding occurs when electrons are transferred completely from one atom to another, forming ions that attract each other.
• A cation is a positively charged ion.
• An anion is a negatively charged ion.
• An ionic bond is the force of attraction between oppositely charged ions.

Ionic Compounds: Formulas

• Chemical formulas represent compounds using element symbols and numbers to show how many atoms of each element are present.
• Ionic compounds are usually formed from metal groups I and II combining with non-metal groups VI and VII.
• Ionic compounds are neutral, so the positive and negative charges in the compound must be equal.
• Common Polyatomic ions (include their charges and formulas)
  • Hydroxide (OH⁻)
  • Nitrate (NO₃⁻)
  • Hydrogencarbonate (HCO₃⁻)
  • Permanganate (MnO₄⁻)
  • Carbonate (CO₃²⁻)
  • Chromate (CrO₄²⁻)
  • Dichromate (Cr₂O₇²⁻)
  • Sulfate (SO₄²⁻)
  • Sulfite (SO₃²⁻)
  • Thiosulfate (S₂O₃²⁻)
  • Phosphate (PO₄³⁻)
  • Ammonium (NH₄⁺)

Writing Ionic Compound Formulas

• Determine the charges of the ions involved.
• Find the lowest common multiple of the absolute values of the charges.
• Use the lowest common multiple to determine the number of each ion needed to balance the charges.
• Write the formula using the subscripts to indicate the number of each ion.

Covalent Bonding

• Covalent bonding involves the sharing of electrons between atoms.
• A molecule is a group of atoms joined together, the smallest particle of an element or compound that can exist independently.
• Single, double, or triple bonds represent the number of electron pairs shared between atoms.
  • A single bond = one shared pair
  • A double bond = two shared pairs
  • A triple bond = three shared pairs
• Chlorine gas (Cl2) shares a pair of electrons, a single covalent bond, so Cl2 is a molecule.
• Oxygen needs two pairs of electrons (two covalent bonds) to reach a full outer shell of 8 electrons, forming a molecule (O2).
• Hydrogen gains a pair, and hydrogen atoms gain a full outer shell of 2 electrons so H2 is a molecule

Shapes of Covalent Molecules

• Use VSEPR (valence shell electron pair repulsion) theory to predict the shape of a molecule. Count the bond pairs and lone pairs around the central atom.

Electronegativity

• Electronegativity is the tendency of an atom to attract shared electrons in a covalent bond
• A large electronegativity difference indicates polar covalent bonds (unequal sharing).
• A small difference indicates nonpolar covalent bonds (equal sharing)
• Polar bonds result in a slightly positive end (δ+) and a slightly negative end (δ-)

Types of Covalent Bonds

• nonpolar covalent: bonds with equal sharing of electrons
• polar covalent: bonds with unequal sharing of electrons resulting from a difference in electronegativity between the atoms

Intermolecular Forces

• Intramolecular bonds occur within a molecule.
• Intermolecular forces occur between molecules. Examples include Van der Waals forces, dipole-dipole forces, and hydrogen bonding
• Van der Waals forces: weak attractive forces between non-polar molecules due to temporary dipoles
• Dipole-dipole forces: attractive forces between the positive end of one polar molecule and the negative end of another polar molecule.
• Hydrogen bonding: a strong type of dipole-dipole force that occurs when hydrogen is bonded to a highly electronegative atom.
• These interactions greatly affect boiling/melting point. Larger molecules generally (but not always) have higher boiling points due to stronger intermolecular forces.