Covalent Compounds Notes

Covalent Compounds

• Covalent compounds involve non-metals combining with 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).

Ionic Compounds

• Ionic compounds involve a metal and a non-metal.
• Electrons are transferred from the metal to the non-metal.
• Atoms give or take electrons to achieve a full set of valence electrons (8 electrons, except for Hydrogen and Helium).

Properties of Covalent Compounds

• Covalent bonds are weaker than ionic bonds.
• Covalent compounds tend to have lower melting and boiling points.
• They are often 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 more than one way.

Covalent Bonds

• A covalent bond results from sharing valence electrons.
• A molecule is formed when two or more atoms bond covalently.
• The shared electrons are part of the outer energy levels of both atoms involved.

Diatomic Molecules

• $H<em>2$, $N</em>2$, $O<em>2$, $F</em>2$, $Cl<em>2$, $Br</em>2$, and $I_2$ are diatomic molecules.
• These 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

• Prefixes are used 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:
  • CO
  • $BF_3$
  • $C<em>2Cl</em>4$
  • NO
  • $NO_2$
  • $N<em>2O</em>4$
  • $P<em>4S</em>5$

Common Names

• Some compounds are known by their common names and not their formal names:
  • $H_2O$ - water
  • $NH_3$ - ammonia
  • $CH_4$ - methane

Acids

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

Practice (Naming)

• Formulas to Names:
  • NaBr
  • $N<em>2O</em>5$
  • $SiO_2$
  • HI
  • $CaCO_3$
  • MgO
  • $H_2S$
  • $SO_2$
  • $HNO_3$
• Names to Formulas:
  • 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.
  • Example: $F<em>2$, $H</em>2$
• 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: $O_2$
• 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.
  • Example: $N_2$, HCN

Lewis Dot Structure Practice

• 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

• Nitrogen monoxide
• Nitrogen dioxide

Less Than 8 Electrons

• Hydrogen (2 electrons)
• $SiH_4$
• Boron (6 electrons)
• $BF_3$

More Than 8 Electrons

• 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 Structure Activities

• Methane
• Ammonia
• $C<em>2H</em>2$
• $C<em>2H</em>4$
• Nitrogen tribromide
• Carbon tetrachloride
• Bromine
• Oxygen
• $NH_4^+$
• Sulfate
• $CH_2O$
• $NH_2Cl$
• 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 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 occurs 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.
• One end of the molecule is slightly negative, while the other is slightly positive.

Polar or Non-Polar?

• A molecule may have polar bonds and NOT have a dipole moment.
• This happens when the polar bonds cancel each other out.

Polarity Practice

• $NH_2Cl$
• Hydrosulfuric acid
• Boron trihydride