Naming Binary Covalent Compounds

• Atoms bonded covalently yield molecules.
• Molecules are neutral in charge.
• Molecular formulas denote the types and numbers of atoms in a molecule.
• Valence electrons, located in the outermost energy level of an atom, facilitate bonding.
• The nucleus of one atom attracts the valence electrons of another, drawing atoms together.
• A covalent bond forms through the sharing of electrons between atoms.
• Covalent bonds require each atom to have space for an additional electron in its valence shell.
• Atoms bond to achieve a full valence shell, increasing stability.
• Covalently bonded atoms are more stable than separate atoms.
• Two non-metals can share electrons.
• Overlapping outer shells allows electron sharing between atoms.

Naming Binary Covalent Compounds

• Naming conventions for binary covalent compounds:
• 1 = mono
• 2 = di
• 3 = tri
• 4 = tetra
• 5 = penta
• 6 = hexa
• 7 = hepta
• 8 = octa
• 9 = nona
• 10 = deca
• Sulfur Trioxide example: SO3
• Covalent compounds use prefixes, unlike ionic compounds.
• Ionic compounds consist of a metal and a non-metal, while covalent compounds involve two non-metals.
• Unshared electron pairs are called lone pairs, are not involved in bonding, and belong exclusively to one atom.

Multiple Bonds

• Elements like carbon, nitrogen, and oxygen can share multiple electron pairs.
• Double bonds involve the sharing of four electrons between two atoms.

Covalent Bonds:

• Steps to represent covalent bonds for element combinations:
• Draw the dot diagram for each atom.
• Illustrate shared electrons with circles.
• Create the Lewis structure.
• Write the molecular formula.
• Name the compound.
• Example: H + Cl → H-Cl (Hydrogen Monochloride)

Exceptions to the Octet Rule

• Electron Deficient (hypovalent) molecules have less than 8 valence electrons.
• Boron (B) can be stable with 6 valence electrons.
• Beryllium (Be) can be stable with 4 valence electrons.
• Expanded Octet (hypervalency) molecules have more than 8 valence electrons.
• Phosphorus (P) can be stable with 10 valence electrons.
• Sulfur (S) can be stable with 12 valence electrons.

Polar vs Nonpolar Bonds

• Electronegativity difference of 0-0.3 indicates a nonpolar covalent bond where electrons are shared equally.
• Electronegativity difference of 0.3-1.7 indicates a polar covalent bond where electrons are shared unequally.
• Polar covalent bonds are indicated using lowercase delta (δ).
• The less electronegative atom gets a positive delta symbol (δ+).
• The more electronegative atom gets a negative delta symbol (δ-).
• An arrow in a bond diagram points toward the more electronegative element.

Properties of Molecules

• Small molecules are typically gases.
• Molecules are poor conductors of electricity.
• Intramolecular bonds (within molecules) are strong.
• Intermolecular bonds (between molecules) are weak, leading to low melting and boiling points.
• Larger molecules can be liquids or soft solids.
• Ionic compounds generally have higher melting and boiling points compared to molecules.
• Ionic compounds are hard, brittle, and good conductors of electricity.

Molecular Geometry

• VSEPR theory explains molecular geometry through Valence Shell Electron Pair Repulsion.
• Valence electron pairs around covalently bonded atoms repel each other.
• This repulsion leads valence electron pairs to be as far apart as possible, resulting in specific molecular shapes.
• Linear: 180° angle, two atoms bonded to a central atom with no lone pairs.
• Trigonal Planar: 120° angle, three atoms bonded to a central atom with no lone pairs.
• Bent: 109.5° (with 1 lone pair) or 120° (with 2 lone pairs), two atoms bonded to a central atom with lone pair(s).
• Tetrahedral: 109.5° angle, four atoms bonded to a central atom with no lone pairs.
• Trigonal Pyramidal: 109.5° angle, three atoms bonded to a central atom with one lone pair.

Molecular Polarity and Dipole-Dipole Forces

• The strongest intermolecular forces occur between polar molecules.
• Molecular polarity depends on both the type of covalent bonds and the shape of the molecule.
• More electronegative atoms in a polar molecule pull shared electrons more strongly, creating a dipole.
• In liquids or solids, the negative region of one polar molecule attracts the positive region of another.
• This attraction between polar molecules is termed dipole-dipole forces.

Hydrogen Bonding

• Hydrogen bonding is an intermolecular force resulting from polarity.
• It occurs when a hydrogen atom bonded to a highly electronegative atom becomes a positive dipole and is attracted to lone pair electrons on an electronegative atom in a nearby molecule.
• Hydrogen bonding is a special, strong, case of dipole-dipole forces.

London Dispersion Forces

• London dispersion forces are relatively weak intermolecular forces.
• They arise from the constant motion of electrons and the creation of instantaneous dipoles.
• Uneven electron distribution can create a brief dipole within an atom, inducing a dipole in an adjacent atom or molecule.
• All atoms and molecules experience London dispersion forces.
• This force is the only intermolecular force acting among noble gas atoms and nonpolar molecules.

Determining Molecular Polarity:

• The presence of polar bonds doesn't guarantee that a molecule is polar.
• Nonpolar molecules have either nonpolar bonds or polar bonds whose dipoles cancel out.
• Polar molecules have polar bonds with dipoles that do not cancel out.

Why is Water Polar?

• In water, H-O bonds are not balanced around the structure, so the pulls do not cancel.
• Lone pairs on oxygen cause asymmetry (imbalance).
• The electron cloud is shifted such that there is greater negative charge near oxygen and greater positive charge near hydrogen.

Adding Dipole Moments:

• Linear (nonpolar): Two identical polar bonds pointing in opposite directions will cancel each other out.
• Bent (polar): Two polar bonds with an angle of less than 180° between them will not cancel each other out.
• Trigonal Planar (nonpolar): Three identical polar bonds at 120° from each other will cancel.
• Tetrahedral (nonpolar): Four identical polar bonds in a tetrahedral arrangement (109.5°) will cancel.
• Trigonal Pyramidal (polar): Three polar bonds in a trigonal pyramidal arrangement (109.5°) will not cancel.
• In metallic compounds, electrons are shared between all atoms.
• A metallic bond is the force of attraction between the shared electrons and the positive nuclei of each metal atom.

Properties of Metals

• Malleable: Can be hammered into thin sheets.
• Ductile: Can be rolled into thin wire.
• Conduct heat and electricity
• Luster: Shininess.
• High melting point.

Property Comparison

• State at room temperature:
• Ionic compounds are solid.
• Molecular compounds are gas.
• Metallic compounds are solid.
• Boiling/melting points:
• Ionic and metallic compounds have high boiling/melting points.
• Molecular compounds have low boiling/melting points.
• Brittleness:
• Ionic compounds: Hard and brittle
• Molecular: N/A
• Metallic compounds: Not brittle
• Conduction of electricity:
• Ionic compounds: Conduct when liquid or melted.
• Molecular compounds: No
• Metallic compounds: Excellent conductors