Molecular/Ionic Compounds

On this page:

  • Types of bonding 
  • Properties of bonded compounds
  • Nomenclature
  • Lewis dot structures and formal charge
  • VSEPR 
  • Bond hybridization
  • Polyatomic ions
  • Worksheets

Types of Bonding

Covalent

Covalent bonds occur when two nonmetals share electrons. For example, HCl is a covalently bonded molecule because the atoms share both the single electron from the hydrogen and a single valence electron from the chlorine to fill the valence shells of both atoms. 


Covalent bonds can be polar or nonpolar. Polarity in bonds occurs due to differences in electronegativity between the atoms, so electrons can be pulled more towards one atom than the other, resulting in a slightly positive and slightly negative area. The most non-polar covalent bonds occur in the diatomic elements which are two of the same atom bonded together. This most often occurs in the elements Br, I, N, Cl, H, O, F. Many polar covalent bonds occur with fluorine as it is very electronegative. 

Ionic

Ionic bonds are the result of the exchange of electrons between a metal and a nonmetal. For example, NaCl is an ionic compound, because Na wants to lose an electron (low ionization energy) while Cl wants to gain an electron (high electron affinity). Once this occurs, the Na+ ion is attracted to the Cl- ion, and this strong electrostatic attraction creates the ionic bond. 


The lattice energy of an ionic compound is the magnitude of the electrostatic attraction within an ionic compound. If the lattice energy is higher, the compound will be harder to break apart, and if the lattice energy is lower, the compound will be easier to break apart. Based on Coulomb’s Law, the lattice energy will be higher in ionic compounds in which the ions are small and have higher magnitudes of charges. 

Metallic

Metallic bonds occur between two or more metal atoms. An alloy occurs when different kinds of metals are bonded together. As seen in the image below, a pure metal has all the same atom in it, an interstitial alloy has a smaller atom (metal or nonmetal) in between the large metal molecules, and a substitutional alloy switches out some pure metal atoms for atoms of a different metal. 

Single, Double, and Triple Bonds

A single covalent bond forms between nonmetal atoms when they share an electron pair. When more than one pair is shared, a double or triple bond forms. Double bonds are both stronger and shorter than single bonds, and triple bonds are both stronger and shorter than double bonds. 

Useful Bonding Images

Bond Polarity

Bond Polarity

Bond Polarity

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 This table can be used to predict bond polarity using an electronegativity chart to find the difference in electronegativities between elements

Ionic Bonding

Bond Polarity

Bond Polarity

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[Image courtesy of differencebetween.com]

Diatomic Elements

Diatomic Elements

Diatomic Elements

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[Image courtesy of scienceabc.com]

Types of Alloys

Diatomic Elements

Diatomic Elements

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[Image courtesy of Chem Libre Texts]

Properties of Bonded Compounds

Covalent

Covalent

Covalent

  • Low melting point
  • Poor thermal / electrical conductivity

Ionic

Covalent

Covalent

  • Crystalline structure
  • High melting point
  • Conductive when dissolved or liquid, but not as a solid


Metallic

Covalent

Metallic

  • Sea of electrons (electrons are free-flowing)
  • Conductive (both heat and electricity)
  • High melting point
  • Shiny/malleable/ductile


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Nomenclature

[Image courtesy of ChemTeam]

Lewis Dot Structures

What is a Lewis Dot Structure?

What is a Lewis Dot Structure?

What is a Lewis Dot Structure?

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A Lewis Dot Structure is a way of depicting the valence shell electrons present in a molecule. The octet rule, which states that most atoms want 8 electrons in their valence shell, forms the basis behind the Lewis Dot Structure theory

Step 1

What is a Lewis Dot Structure?

What is a Lewis Dot Structure?

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You will be given a molecule For example, let's choose ammonia, which is notated as NH3.

Step 2

What is a Lewis Dot Structure?

Step 2

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Determine the number of valence electrons in the compound using the periodic table. Add or subtract electrons if the molecule happens to be charged [Image courtesy of Newsweek.com]

Step 3

Formal Charge

Step 2

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In our example,  hydrogen has one valence electron per atom, and nitrogen has five valence electrons per atom. These electrons are represented by dots in the Lewis Dot Structure. Since the example molecule is NH3, there are three hydrogen atoms and one nitrogen atom.

Step 4

Formal Charge

Formal Charge

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Next, you must combine these into the final Lewis Dot Structure. When two valence electrons are combined (shared between two atoms), they form a covalent bond, visually shown by a dash. 

Formal Charge

Formal Charge

Formal Charge

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Check your work using formal charge.  Read more about formal charge in the next section. In our example, the formal charge of nitrogen is given by 5-3-2=0, and the formal charge of all the hydrogens is given by 1-1=0. Since the  sum of formal charges corresponds with the charge of the molecule, our structure is correct. 

Resonance and Formal Charge

Resonance Structures

Resonance Structures

Resonance Structures

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Resonance structures occur when a double or triple bond can be placed anywhere within the Lewis Dot structure without changing the overall shape of the molecule. In the image above, the carbonate ion has a single double bond which can be placed with any of the oxygen molecules without changing the overall structure of the molecule, giving it a resonance structure. On a macroscopic level, resonance structures state that all the resonance bonds in a molecule are actually the same. In the example above, the resonance bond is stronger than a single bond but weaker than a double bond.  Common resonance structures include NO2, O3, and benzene (C6H6)

[Image courtesy of Chem Libre Texts]

Formal Charge

Resonance Structures

Resonance Structures

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As mentioned in the example above,  formal charge is a naming convention that can be used to check your work when creating a Lewis Dot structure. The formal charge of each atom is given by  (# of valence electrons)-(# of bonds)-(# of unshared electrons) and  should equal zero if the molecule has no charge. If the molecule is charged, the formal charges should add to the overall charge, and the most electronegative  atom should have the negative formal charge if there is one. 

VSEPR

What is VSEPR?

What is VSEPR?

What is VSEPR?

VSEPR (valence shell electron pair repulsion) is a theory for modeling 3-dimensional molecular structures based on the placement of a central atom, ligands (atoms that attach to the central atom), and lone electron pairs. The central idea behind the VSEPR theory is that electrons will move around in three dimensional space to become as far apart from each other as possible. 

First Steps

What is VSEPR?

What is VSEPR?

When identifying a VSEPR model, start with a Lewis Dot structure of the molecule.  Next, count the number of regions of electron density around the central atom in your Lewis Dot structure. Then count how many of these regions are lone electron pairs. 

Finding the Final Structure

Finding the Final Structure

Finding the Final Structure

Match the numbers you found out to the numbers in the following image. You may need to memorize this information. Then, follow the matched structure to obtain all of the information that goes with it.

Overall Molecule Polarity

Finding the Final Structure

Finding the Final Structure

The VSEPR structure of a molecule can be used to determine molecular polarity.  When there is symmetry in the VSEPR structure, such as something with a trigonal planar geometry, the net dipoles of the bonds cancel out, leaving a non-polar structure. When there is not symmetry, such as in a bent molecule, the net dipoles do not cancel out, leaving a polar molecule. The molecule is often polar when there is a lone electron pair present. 

Bond Hybridization

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The Theory

The theory behind bond hybridization is based on resonance structures and the equality between all the bonds in a given molecular structure. Electrons in different energy levels will form bonds with different energies, which is not consistent with the resonance structure theory. Instead, electrons in different energy levels that are shared between atoms form "degenerate" hybridized energy levels. 

What you need to know

To determine the bond hybridization in a molecule, count the regions of electron density as you would for a VSEPR structure. Remember, lone pairs count once, as do all bonds.


  • 2 regions: sp
  • 3 regions: sp2
  • 4 regions: sp3
  • 5 regions: sp3d
  • 6 regions: sp3d2


To remember this list, a tip is that the exponents add to the number of regions of electron density, and that the exponents for each letter represent pairs in that orbital, so s cannot exceed 1, p cannot exceed 3, and d cannot exceed 5.

Sigma and Pi Bonds

Sigma and Pi bonds are the basis of molecular orbital theory. Sigma bonds are all of the single bonds in a molecular structure as well as the first bond in any double or triple bond. Pi bonds are the second bond in a double bond and the second or third bonds in a triple bond.

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Polyatomic Ions

Memorizing the polyatomic ions and their charges from the list above is important in your ability to predict reaction, write formulas, and perform many other important tasks in chemistry. 

Worksheets

Molecular/ionic compounds nomenclature worksheet and answer key (pdf)

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Lewis Dot Structures and VSEPR worksheet and answer key (pdf)

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