Covalent and Ionic Compounds

Covalent Compounds

  • Formed between non-metal and non-metal atoms.

  • Involve the sharing of electrons between atoms to achieve a full set of valence electrons (8 electrons, except for Hydrogen and Helium).

Electrons Sharing

  • Atoms share electrons to achieve a full set of valence electrons, similar to ionic compounds seeking a full valence shell.

Properties of Covalent Compounds

  • Lower melting and boiling points: Covalent bonds are generally weaker than ionic bonds; therefore, covalent compounds tend to have lower melting and boiling points.

    • Often exist as liquids or gases at room temperature.

  • Non-conductive: Covalent compounds do not conduct electricity because they do not have ions.

  • Also known as molecular compounds.

  • Nonmetals can combine in multiple ways.

Covalent Bond

  • A covalent bond results from the sharing of valence electrons.

  • A molecule is formed when two or more atoms bond covalently.

  • Shared electrons are considered part of the outer energy levels of both atoms involved.

Diatomic Molecules

  • Examples: H<em>2H<em>2, N</em>2N</em>2, O<em>2O<em>2, F</em>2F</em>2, Cl<em>2Cl<em>2, Br</em>2Br</em>2, and I2I_2.

  • Form when two atoms of each element share electrons.

  • They exist this way because the two-atom molecule is more stable than the individual atoms.

Naming Covalent Compounds

  • Use prefixes to distinguish between different compounds.

  • First nonmetal: prefix (except "mono") + name of element.

  • Second nonmetal: prefix + name of element + "-ide" ending.

Prefixes

  • 1 - Mono

  • 2 - Di

  • 3 - Tri

  • 4 - Tetra

  • 5 - Penta

  • 6 - Hexa

  • 7 - Hepta

  • 8 - Octa

  • 9 - Nona

  • 10 - Deca

Examples of Covalent Compounds

  • CO

  • BF3BF_3

  • C<em>2Cl</em>4C<em>2Cl</em>4

  • NO

  • NO2NO_2

  • N<em>2O</em>4N<em>2O</em>4

  • P<em>4S</em>5P<em>4S</em>5

Common Names

  • Some compounds are known by their common names instead of their formal names.

    • H2OH_2O – water

    • NH3NH_3 – ammonia

    • CH4CH_4 – methane

Acids

  • Water solutions of some molecules are acidic and are named as acids.

  • If a compound produces hydrogen ions (H+H^+) in solution, it is an acid.

Categories of Acids

  • Acids WITHOUT oxygen (binary acids).

  • Acids WITH oxygen (oxyacids).

Naming Acids WITHOUT Oxygen (Binary Acids)

  • Prefix: Hydro-

  • Suffix: -ic

  • Examples:

    • HCl: Hydrochloric acid

    • HF: Hydrofluoric acid

    • HBr: Hydrobromic acid

    • HI: Hydroiodic acid

    • H2SH_2S: Hydrosulfuric acid

Naming Acids WITH Oxygen (Oxyacids)

  • No prefix.

  • Suffix: -ic

  • Examples:

    • HNO3HNO_3: Nitric acid

    • H<em>2CO</em>3H<em>2CO</em>3: Carbonic acid

    • H<em>2SO</em>4H<em>2SO</em>4: Sulfuric acid

    • H<em>3PO</em>4H<em>3PO</em>4: Phosphoric acid

    • HC<em>2H</em>3O2HC<em>2H</em>3O_2: Acetic acid

Practice Examples

  • NaBrNaBr

  • N<em>2O</em>5N<em>2O</em>5

  • SiO2SiO_2

  • HI

  • CaCO3CaCO_3

  • MgO

  • H2SH_2S

  • SO2SO_2

  • HNO3HNO_3

More Practice Examples

  • Ammonium nitrate

  • Hydrosulfuric acid

  • Potassium iodide

  • Phosphorus trioxide

  • Carbonic acid

  • Carbon tetrachloride

Lewis Dot Structure

  • Single Covalent Bond: A bond in which atoms share one pair of electrons.

    • Examples: F<em>2F<em>2, H</em>2H</em>2

  • Double Covalent Bond: A bond in which atoms share two pairs of electrons.

    • A double bond is shorter and stronger than a single bond.

    • Example: O2O_2

  • Triple Covalent Bond: A bond in which atoms share three pairs of electrons.

    • A triple bond is shorter and stronger than a single and a double bond.

    • Examples: N2N_2, HCN

Lewis Structure Practice

  • Carbon monoxide

  • Carbon dioxide

  • Water

  • H<em>2O</em>2H<em>2O</em>2

  • Sulfur dioxide

  • Sulfur trioxide

  • Ammonia

  • Dinitrogen monoxide

  • Carbon tetrachloride

Exceptions to the Octet Rule

  • Odd number of electrons.

  • Stable with less than 8 electrons.

  • Stable with 8, 10, or 12 electrons.

Odd Number of Electrons

  • Nitrogen monoxide

  • Nitrogen dioxide

Less than 8 Electrons

  • Hydrogen (2 electrons)

  • SiH4SiH_4

  • Boron (6 electrons)

  • BF3BF_3

More than 8 Electrons

  • Sulfur (8, 10, or 12 electrons): SF6SF_6

  • Phosphorus (8, 10, or 12 electrons): PCl5PCl_5

  • Xenon (8, 10, or 12 electrons): XeF4XeF_4

Lewis Dot Activity

  • Methane

  • Ammonia

  • C<em>2H</em>2C<em>2H</em>2

  • C<em>2H</em>4C<em>2H</em>4

  • Nitrogen tribromide

  • Carbon tetrachloride

Lewis Dot Activity continued

  • Bromine

  • Oxygen

  • NH4+NH_4^+

  • Sulfate

  • CH2OCH_2O

  • NH2ClNH_2Cl

Lewis Dot Activity Continued

  • Phosphorus pentachloride

  • Hydrosulfuric acid

  • Boron trichloride

  • HCN

  • Nitrogen tribromide

  • Nitrite

  • Nitrate

Bond Polarity

  • Covalent bonds involve the sharing of electrons between atoms.

  • However, this sharing is more like a tug-of-war between the atoms.

Nonpolar Covalent Bonds

  • Electrons are shared equally.

  • Bonds between diatomic molecules are nonpolar because the atoms have the SAME electronegativity.

  • Bonds between carbon atoms and hydrogen atoms are also nonpolar because they have very similar electronegativities.

Polar Covalent Bonds

  • Electrons are shared unequally.

  • Unequal sharing takes place because the more electronegative atom has a stronger electron attraction and will have a stronger pull on the electrons.

Dipole Moment

  • A molecule with a dipole moment is a polar molecule.

  • This means one end of the molecule is slightly negative while the other is slightly positive.

Polar vs Non-Polar Molecules

  • A molecule may have polar bonds and NOT have a dipole moment.

  • This happens when the polar bonds cancel each other out.

Polarity Practice

  • NH2ClNH_2Cl

  • Hydrosulfuric acid

  • Boron trihydride