Covalent Compounds Notes

Covalent Compounds

Ionic Compounds

• Formed from a metal and a non-metal.

• Electrons are transferred from the metal to the non-metal.

• Atoms transfer electrons to achieve a full set of valence electrons (8 electrons, except for Hydrogen and Helium).

Covalent Compounds

• Formed from two non-metals.

• Electrons are shared between the atoms.

• Atoms share electrons to achieve a full set of valence electrons (8 electrons, except for Hydrogen and Helium).

Properties of Covalent Compounds

• Covalent bonds are generally weaker than ionic bonds, resulting in lower melting and boiling points.

• Covalent compounds tend to be liquids or gases at room temperature.

• Covalent compounds do not conduct electricity because they do not have ions.

• Also known as molecular compounds.

• Nonmetals can combine in multiple ways.

Covalent Bonds

• 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: $H<em>2$, $N</em>2$, $O<em>2$, $F</em>2$, $Cl<em>2$, $Br</em>2$, and $I_2$

• Form when two atoms of each element share electrons.

• Two-atom molecule is more stable than 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

• $BF_3$

• $C<em>2Cl</em>4$

• NO

• $NO_2$

• $N<em>2O</em>4$

• $P<em>4S</em>5$

Trivial Names

• Some compounds are known by their common names:

  • $H_2O$ - water

  • $NH_3$ - ammonia

  • $CH_4$ - methane

Acids

• Water solutions of some molecules are acidic.

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

Naming Acids

• Acids are divided into two categories:

  • Acids WITHOUT oxygen (binary acids)

  • Acids WITH oxygen (oxyacids)

Naming Binary Acids (Acids WITHOUT Oxygen)

• Prefix: Hydro-

• Suffix: -ic

• Examples:

  • HCl - Hydrochloric acid

  • HF - Hydrofluoric acid

  • HBr - Hydrobromic acid

  • HI - Hydroiodic acid

  • $H_2S$ - Hydrosulfuric acid

Naming Oxyacids (Acids WITH Oxygen)

• No prefix

• Suffix: -ic

• Examples:

  • $HNO_3$ - Nitric acid

  • $H<em>2CO</em>3$ - Carbonic acid

  • $H<em>2SO</em>4$ - Sulfuric acid

  • $H<em>3PO</em>4$ - Phosphoric acid

  • $HC<em>2H</em>3O_2$ - Acetic acid

Examples

• NaBr

• $N<em>2O</em>5$

• $SiO_2$

• HI

• $CaCO_3$

• MgO

• $H_2S$

• $SO_2$

• $HNO_3$

Examples (Naming)

• Ammonium nitrate

• Hydrosulfuric acid

• Potassium iodide

• Phosphorus trioxide

• Carbonic acid

• Carbon tetrachloride

Lewis Dot Structure

• Single Covalent Bond: Atoms share one pair of electrons.

  • $F_2$

  • $H_2$

• Double Covalent Bond: Atoms share two pairs of electrons.

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

  • $O_2$

• Triple Covalent Bond: Atoms share three pairs of electrons.

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

  • $N_2$

  • HCN

Examples (Lewis Dot Structure)

• Carbon monoxide

• Carbon dioxide

• Water

• $H<em>2O</em>2$

• 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 (Exceptions)

• Nitrogen monoxide

• Nitrogen dioxide

Less Than 8 Electrons (Exceptions)

• Hydrogen (2 electrons)

• $SiH_4$

• Boron (6 electrons)

• $BF_3$

More Than 8 Electrons (Exceptions)

• Sulfur (8, 10, or 12 electrons)

  • $SF_6$

• Phosphorus (8, 10, or 12 electrons)

  • $PCl_5$

• Xenon (8, 10, or 12 electrons)

  • $XeF_4$

Lewis Dot Activity

• Methane

• Ammonia

• $C<em>2H</em>2$

• $C<em>2H</em>4$

• Nitrogen tribromide

• Carbon tetrachloride

Lewis Dot Activity (Continued)

• Bromine

• Oxygen

• $NH_4^+$

• Sulfate

• $CH_2O$

• $NH_2Cl$

Lewis Dot Activity (Continued Again)

• Phosphorus pentachloride

• Hydrosulfuric acid

• Boron trichloride

• HCN

• Nitrogen tribromide

• Nitrite

• Nitrate

Bond Polarity

• Covalent bonds involve sharing electrons between atoms, resembling a tug-of-war.

Nonpolar Covalent Bonds

• Electrons are shared equally.

• Bonds between diatomic molecules are nonpolar (same electronegativity).

• Bonds between carbon and hydrogen atoms are also nonpolar (similar electronegativities).

Polar Covalent Bonds

• Electrons are shared unequally.

• More electronegative atom has a stronger pull on electrons.

Dipole Moment

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

• One end is slightly negative, the other is slightly positive.

Polar vs. Non-Polar Molecules

• A molecule may have polar bonds but no dipole moment if the polar bonds cancel each other out.

Practice (Polarity)

• $NH_2Cl$

• Hydrosulfuric acid

• Boron trihydride