Chemical Bonds and Molecular Structure

• Covalent bonding involves the sharing of electrons between non-metal atoms.

Covalent vs. Ionic Bonds

• Covalent bonds form between non-metal atoms, while ionic bonds form between metal and non-metal atoms.
• Covalent compounds are called molecular compounds.
• Electrons are shared in covalent bonds but transferred in ionic bonds.
• Ionic compounds are formed through the removal of electrons from metals and the addition of electrons to non-metals.
• Ionic compounds are generally hard and brittle with high melting points, while covalent compounds are soft or hard with lower melting points.
• Ionic compounds are generally soluble in water, while covalent compounds are less soluble.
• Ionic compounds conduct electricity in a molten or aqueous state but covalent compounds do not.
• Examples of ionic compounds: NaCl, MgO (solids at room temperature).
• Examples of covalent compounds: H2O, CO2, CH4 (solid, liquid, or gas).

Types of Covalent Bonds

• Single covalent bonds: One pair of electrons is shared between two non-metal atoms.
• Double covalent bonds: Two pairs of electrons are shared between two non-metal atoms.
• Triple covalent bonds: Three pairs of electrons are shared between two non-metal atoms.

Single Covalent Bonds

• A single covalent bond is formed when atoms share one pair of electrons.
• A dash represents a bond consisting of two electrons.
• Unshared pairs of electrons, also called lone pairs, are represented as dots.

Drawing Lewis Electron-Dot Structures

• Lewis structures show valence electrons around the atomic symbol.
• Electron positions are symbolic, not literal.
• Steps to draw Lewis structures include:
• Write the electron configuration.
• Find the valence electrons.
• Draw the Lewis dot structure for each atom.
• Draw the Lewis dot structure for the molecule.
• Molecule drawing steps include:
• Gather information, drawing Lewis structures, and determining the total number of valence electrons.
• Arrange atoms to show bonding, with halogens and hydrogen usually on the end and carbon in the center.
• Distribute dots to satisfy the octet rule (except for H, Be, B).
• Draw the bonds as long dashes.
• Verify the structure by counting the number of valence electrons.

Multiple Bonds

• Double bonds occur when two atoms share two pairs of electrons, represented by two dashes (four electrons).
• Triple bonds occur when two atoms share three pairs of electrons, represented by three dashes = 6 electrons.

Bond Length and Bond Energy

• Bond length is the distance between the nuclei of two bonded atoms.
• Bond energy is the energy required to break a bond.
• Shorter bonds have higher bond energy and are stronger.

Bond Length and Energy Values:

• C-C: 154 pm, 346 kJ/mol
• C=C: 134 pm, 612 kJ/mol
• C≡C: 120 pm, 835 kJ/mol
• C-N: 147 pm, 305 kJ/mol
• C=N: 132 pm, 615 kJ/mol
• C≡N: 116 pm, 887 kJ/mol
• C-O: 143 pm, 358 kJ/mol
• C=O: 120 pm, 799 kJ/mol
• C≡O: 113 pm, 1072 kJ/mol
• N-N: 145 pm, 180 kJ/mol
• N=N: 125 pm, 418 kJ/mol
• N≡N: 110 pm, 942 kJ/mol

Binary Molecular Compounds

• Definition: Compounds composed of two non-metal atoms.
• Naming rules: first element name + second element root + suffix -ide
• Use prefixes to indicate the number of atoms; examples: CO (carbon monoxide), N2O (dinitrogen monoxide).

Naming Molecular Compounds

• Nonmetals are to the right of the stair-step line on the periodic table.
• Name the first element and add -ide to the second element.
• Use prefixes to specify the number of each atom.
• Omit the prefix "mono-" if the first element has only one atom.

Prefixes to Memorize

• 1: Mono
• 2: Di
• 3: Tri
• 4: Tetra
• 5: Penta
• 6: Hexa
• 7: Hepta
• 8: Octa
• 9: Nona
• 10: Deca

Writing Formulas with Nonmetals

• Prefixes indicate the number of atoms so charges do not need balancing.

Examples:

• Nitrogen Trichloride = NCl3
• Sulphur Dibromide = SBr2
• Dihydrogen Monoxide = H2O

Electronegativity

• Atoms share electrons equally or unequally.
• The difference in electronegativity can be used to predict the type of bond:
• Nonpolar covalent bonds: bonding electrons shared equally.
• Polar covalent bonds: shared electrons more likely to be found around the more electronegative atom.

Polarity

• An atom's strength is measured by electronegativity.
• The larger the electronegativity, the stronger the atom attracts electrons.
• The more electronegative atom gets a δ- (partial negative) charge, and the less electronegative atom gets a δ+ (partial positive) charge.

Dipole Moment

• Occurs in polar molecules with partially positive and partially negative sides.
• Represented by an arrow pointing towards the more electronegative element with a plus sign on the other end.

Polar Molecules

• When placed between oppositely charged plates, align themselves.
• The partially negative side attracts the positive plate, and vice versa.

Determining Polarity

• Determined by the difference in electronegativity between two bonded atoms: nonpolar (<0.5), polar (0.5-2), ionic (>2).

Electronegativity Notes

• Polar molecules have positive and negative ends (dipoles).
• δ ("delta") means "partial" in math and science.
• Positive end: δ+
• Negative end: δ-

Examples

• EN of H = 2.2 and EN of Cl = 3.2
• DEN of HCl = 3.2 - 2.2 = 1 (Polar molecule)

Electronegativity Difference for Hydrogen Halides

Naming Covalent Compounds

• First Name:
• • Name the first element in the formula (Usually least electronegative)
• • Requires a prefix if more than one of them
• Second Name:
• • ends in -ide
• • Requires a prefix (if one or more )

Molecular Shape

• Is a three-dimensional shape that helps determine physical and chemical properties

Determining Molecular Shapes

• Use the VSEPR Theory:
• Valence
• Shell
• Electron
• Pair
• Repulsion (between electron pairs.)
• Molecule adopts the shape that minimizes electron pair repulsions.

Applying VSEPR Theory

1. Draw the Lewis structure of the molecule and identify the central atom.
2. Count the number of electron charge clouds (lone and bonding pairs) surrounding the central atom.
3. Predict molecular shape by assuming that clouds orient so they are as far away from one another as possible.

Molecular Shapes and Polarity

• Molecular shape affects overall polarity.
• Polarity affects properties like boiling point due to attractions between molecules.
• Carbon dioxide (CO2) is linear has no molecular dipole, while water (H2O) is bent and has an overall molecular dipole.