Unit 3 Study Guide and Review: Bonding

3a: Define chemical bonds.

Chemical bonding is fundamental to the study of chemistry since chemical bonds describe how molecules and compounds are formed.

    • Describe the nature of the chemical bond.

Chemical bonds describe an effect that occurs when one or more outer shell electrons are simultaneously attracted to two atomic nuclei. The forces holding a chemical bond together are electrostatic forces. A molecule is created when chemical bonds form. You can think about a molecule as an aggregate of atoms with distinct properties.

3b: Explain why most atoms form chemical bonds.

Most elements found on the periodic table readily form chemical bonds to create molecules.

    • Describe the energetics of bond formation.
    • Why is it often favorable to form a bond?

Atoms often bond together chemically to form molecules. Since heat is released during this process, chemical bonding is an example of an exothermic reaction. The more exothermic the reaction, the more stable the product. Since the formation of chemical bonds creates a more stable product, there is a natural tendency toward bond formation. We describe the bonds as energetically favorable since energy is released.

In general, the more stable structure that results (molecules) also has lower potential energy than the original individual atoms. Remember their energy was released during the chemical bonding process.

3c: Describe ionic, covalent, and metallic bonding.

Chemical bonds can take three forms depending on the atoms involved: ionic, covalent, and metallic.

    • Describe the ionic bond and the types of atoms that can form ionic bonds.
    • Describe the covalent bond and the types of atoms that can form covalent bonds.
    • Describe the metallic bond and the types of atoms that can form metallic bonds.

In an ionic bond, ions of opposite charge are attracted to each other via electrostatic forces: an electron is donated from the negative ions to the positive ion to form the bond.

Table salt, or sodium chloride (NaCl), is an example of a compound that has an ionic bond. In the case of NaCl, the chloride (Cl-) ion donates its extra electron to the sodium ion (Na+) to form the bond. Ionic solids, such as NaCl are crystalline in form.

Image of an ionic bond: table salt, or sodium chloride (NaCl).

In a covalent bond, the electrons in the bond are shared. Covalent bonds generally form among nonmetal atoms: each atom in the bond usually contributes one electron to form the bonding electron pair.

Image of a covalent bond.

We classify covalent bonds as nonpolar covalent bonds or polar covalent bonds. In a nonpolar covalent bond, the two atoms in the bond share the bonding electrons equally. In a polar covalent bond, the electrons are unevenly shared.

This means, there is a greater electron density near one of the atoms in the bond than the other. This is because one of the atoms has higher electronegativity than the other atom. Electronegativity is a measure of how much an atom pulls electrons toward itself in a covalent bond.

Metallic bonds form, not surprisingly, among metal atoms. Metal atoms have low electronegativities, and have empty, or near empty, outer electron shells. Consequently, metal atoms do not attract electrons, nor do they easily donate them.

A simple way to think about metal bonding is that positive metal ions are immersed in an electron fluid of free-flowing electrons. This leads to many properties of metals. For example, metals conduct electricity because the valence electrons are mobile throughout the material.

Image of atomic cores immersed in valence "electron fluid."

3d: Describe and give examples of Van der Waals forces as part of intermolecular forces.

Intermolecular forces are the forces that hold molecules of the same type together in the condensed phase.

    • Describe how each type of Van der Waals force works and the types of molecules that can be affected by the type of Van der Waals force.

Van der Waals forces, also known as intermolecular forces, describe the forces that exist between molecules that hold molecules together in the liquid or solid state.

In dipole-dipole force interactions, molecules with permanent dipole moments interact. A molecule has a permanent dipole moment if it has an uneven distribution of negative charge within the overall neutrally charged molecule. We call molecules with a permanent dipole moment polar molecules. A polar molecule has a partially positive area and a partially negative area within the molecule. 

Image of dipole-dipole force interactions.

When two polar molecules come together, the partially positive area of one molecule lines up with the partially negative area of the other molecule. The weak electrostatic attraction among the molecules leads to dipole-dipole interactions.

Image of two polar molecules.

In ion-induced dipole forces, an ion interacts with a nonpolar molecule. In this case, the electrostatic charge on the ion induces or forces the nonpolar molecule to momentarily develop a dipole. This is a weaker interaction because the nonpolar molecule does not have a permanent dipole, and is therefore not strongly attracted to the ion.

Image of ion-induced dipole forces.

In ion-dipole forces, an ion interacts with a polar molecule. The polar molecule lines up, so the opposite partial charge is attracted to the ion. This is a stronger interaction than ion-induced dipole because the polar molecule has a permanent partially-positive and partially-negative charge.

Image of ion-dipole forces.

In dipole-induced dipole interactions, a polar molecule interacts with a nonpolar molecule. The dipole in the polar molecule induces or causes a momentary dipole in the nonpolar molecule. The resulting interaction is weak because the nonpolar molecule only has momentary and not permanent dipole.

Image of dipole-induced dipole interactions.

Finally, London forces describe the forces between two nonpolar molecules. A momentary dipole that randomly forms in one molecule induces an induced dipole in the other nonpolar molecule. This is a very weak intermolecular force.

Image of London Forces.

3e: Explain VSEPR theory.

First, review how to write electron configurations for atoms in Section 3: The Aufbau Rules, in Electrons in Atoms. This will help you understand how to draw a Lewis Dot structure to apply VSEPR theory.

Next, read Section 2: Lewis Dot Structures to understand how Lewis dot structures work and review how to use VSEPR theory.

Section 3: How to Draw Lewis Dot Structures, provides a step-by-step description of how to draw Lewis dot structures in The Shared-Electron Covalent Bond.

To write a Lewis dot structure, first write the symbols for the elements in a simple diagram to show how the elements will be connected.

For example, for ammonia, NH3, or hydroxylamine, you would write the following:

Image of Ammonia and Hydroxylamine.

Then, using the Aufbau principle, draw electron dot structures for each of the elements. These drawings tell us how many valence shell electrons are in each atom, based on their group on the periodic table. The group number on the periodic table equals the number of valence electrons.

For example, hydrogen is in group 1A, so it has one valence electron. Nitrogen is in group 5A so it has five valence electrons, and oxygen is in group 6A so it has six valence electrons.

Image of Position of Valence Electrons.

Finally, bring the atoms together in a way that places eight electrons around each atom wherever possible. Note that hydrogen is an exception to this rule and will only ever have two electrons.

Image of Position of Valence Electrons.

After studying the basics of Lewis electron dot structures, we can begin to understand VSEPR Theory.

VSEPR Theory is an acronym for Valence Shell Electron Pair Repulsion Theory. It describes the three-dimensional shapes of molecules.

    • Explain why chemists use VSEPR to describe the shapes of molecules.
    • What forces determine the shape of molecules?

VSEPR Theory focuses on the valence electron pairs in the outermost electron shells of the atoms involved in bonding. These pairs are the electrons that can form chemical bonds.

We assume that electrons involved in bonding exist in between the two atoms being bonded. We also know that similar or “like” charges repel each other via electrostatic forces. Therefore, the lone pairs of electrons, or nonbonding valence electrons in the molecule are repelled by the bonding pair and by all of the other lone pairs of electrons in the molecule.

Because they are repelled, all of the valence electron pairs will adopt shapes that make them as far apart from each other as possible in a three-dimensional space. This description provides the basis of VSEPR Theory and predicts the shapes different molecules will take based on the number of bonding and nonbonding valence electron pairs.

Image that shows minimum repulsion between electron clouds.

3f: Predict the shape of molecules or polyatomic ions using VSEPR theory.

Remember that we use VSEPR Theory to predict the shape molecules and polyatomic ions will take on based on their Lewis electron dot structure.

Again, it is important to have a strong foundation in Lewis electron dot structures, so be sure to review Sections 2 and Section 3, in The Shared-Electron Covalent Bond.

    • Use VSEPR to predict the shape of a given molecule or polyatomic ion.

Molecules with the formula AX2 have two bonding electron pairs that are 180o apart to maximize the distance between them. We call them linear molecules.

Image of linear molecule showing 180 degree separation.Image of a linear molecule.


Molecules with the formula AX3 have three bonding electron pairs that are 120o apart to maximize distance between them. We call them trigonal planar molecules.

Image of trigonal planar molecules showing 120 degree separation.

The most common configuration we see in chemistry is the AX 4 molecule. These molecules have bonding electron pairs 109.5 o apart in three dimensional space. These molecules are called tetrahedral molecules.

Methane, CH4, is an example of a tetrahedral molecule.

Image of methane, a tetrahedral molecule, showing 109.5 degree bond angles.

Sometimes, one or more of the valence electron pairs in a molecule are lone pairs of electrons. This alters the geometry of the molecule since the lone pairs take up a bit more space than bonding electrons do.

A molecule of the form AX3E is based in tetrahedral geometry, but has one lone pair of electrons (E). Ammonia, NH3, is an example of this type of molecule. Here, the molecule takes a trigonal pyramidal shape and the bonding angles are approximately 107o.

Image of ammonia, a trigonal pyramidal shape, showing 107 degree angle.

Water, H 20, is an example of a molecule with AX 2>E 2 geometry. In this molecule, there are two bonding pairs of valence electrons and two lone pairs. This molecule has a bent geometry with a bond angle around 104.5 o.

Image of a molecule has a bent geometry, showing 104.5 degree angle.

Occasionally, we see molecules with five or six valence electron pairs that adopt other, more complex geometric shapes.

To review how VSEPR theory describes different molecular geometries, see the descriptions in Molecular Geometry.

3g: Explain how the shapes of molecules are accounted for by hybridization theory.

Hybridization theory accounts for molecular shapes that have been seen in experiments when combining atomic orbitals into molecular orbitals.

    • Explain why hybridization theory is needed.
    • Use a hybridization diagram to construct molecular orbitals for a tetrahedral compound.
    • Describe why hybridization theory gives us molecular shapes.

Hybridization theory presents the limitations of VSEPR theory. In VSEPR theory, we assumed the electrons were in their atomic orbitals. If this were the case, many molecules we know to exist should never form.

For example, beryllium hydride, BeH2, is a known compound, but beryllium (Be) has a set of paired electrons in its valence shell (the 2s atomic orbital). The paired electrons in the atomic orbitals are stable and should not combine with other atomic orbitals to form bonds.

However, we know beryllium hydride, BeH2, exists. The only way for this to happen is for an electron from the beryllium valence shell to move to the next energy level, 2p, so there are unpaired electrons. We know this does not happen because it would require too much energy. Therefore, a new model of bonding is needed to explain molecules such as BeH2 which cannot be explained by VSEPR.

In hybridization theory, the electrons from the atoms involved in the bond are hybridized or combined into new molecular orbitals which are combinations of the input atomic orbitals. This is a mathematical construct that gives us molecular orbitals that possess properties that are consistent with what we observe about molecules.

Hybrid orbitals are constructed by combining the wavefunctions, ψ, of the atomic orbitals involved in the bond. Because the wavefunction describes a wave, the wavefunctions interfere constructively and destructively, to create a new shape for the hybrid molecular orbital.

Images of wavefunctions.

To use a hybrid orbital diagram, we first fill the electron configurations for the atomic orbitals.

Review how to write electron configurations for atoms in Section 3: The Aufbau Rules, in Electrons in Atoms.

Then, we combine the electrons from both atoms together to fill in the hybrid molecular orbitals created by the atomic orbitals. We always fill the molecular orbitals starting at the lowest energy level. For each energy level, we fill in one electron per molecular orbital at a time and then fill the second electrons per molecular orbital in the energy level before moving on to the next energy level.

In linear molecules, s and p atomic orbitals combine to form hybrid sp orbitals.

Image of the origin of an sp hybrid orbital.

In our example of beryllium hydride, BeH 2, we can use hybrid orbitals to explain why this molecule exists and has a linear shape. Two sp hybrid orbitals are formed in this molecule and overlap with the 1s orbitals of each hydrogen atom to form covalent bonds.

Image of beryllium hydride, and its two sp hybrid orbitals.

We can use hybridization theory to explain larger molecules as well. Here we will look at a trigonal planar molecule, BF 3. In this case, the atomic orbitals of B hybridize to become three sp2 hybrid orbitals. These sp 2 hybrid orbitals have unpaired electrons that can combine with the unpaired electrons of fluorine to create covalent bonds.

Image of sp2 hybrid orbitals in BF3.

When the sp 2 hybrid orbitals are formed, they create trigonal planar geometry:

Image of the origin of sp2 hybrid orbitals, and trigonal planar molecular geometry.

Finally, we will describe the most important geometry in chemistry, tetrahedral, and how hybridization theory predicts the molecular geometry. We will use methane, CH 4, as our example.

Here, the electrons in carbon hybridize to form four sp3 hybrid orbitals. Each of the four sp3 hybrid orbitals contains an unpaired electron that can combine with a hydrogen electron to form a bond.

Image of sp3 hybrid orbitals

The sp 3 orbitals form the expected tetrahedral molecular geometry.

Image of origin of sp3 hybrid orbitals, and tetrahedral molecular geometry.


3h: Explain what determines molecular polarity.

Molecular polarity occurs in molecules where the electrons are not evenly distributed.

    • Define molecular polarity.
    • Define electronegativity.
    • Define dipole moment.
    • Determine if a molecule is polar or nonpolar.

Electronegativity measures how well an atom attracts electrons toward itself in a bond.

An atom that is highly-electronegative attracts electrons toward itself more strongly. On the other hand, an atom that has low electronegativity and does not attract electrons toward itself.

When determining polarity, chemists examine the relative electronegativities of an atom in a bond. The location of each element in the periodic table indicates this trend.

Fluorine is the most electronegative atom. The elements that appear closest to fluorine in the periodic table (top right) are the most electronegative. Elements that are located far away from fluorine in the periodic table (bottom left) are the least electronegative and said to be electropositive.

Image of a graph that shows electronegativities of elements in the periodic table.

A polar covalent bond occurs among atoms with different electronegativities. The electrons in the bond are more attracted to the more electronegative atom and therefore spend more time closer to that atom. So, in these bonds, the electrons are not evenly shared.

We say polar covalent bonds have a dipole moment, which is a vector that points from the less electronegative atom to the more electronegative atom.

Vectors describe a mathematical quantity that also have a direction associated with them. Because dipole moments are vectors, we can sum the vectors to determine if a molecule is polar. If the vectors do not cancel out, the molecule is polar because one area of the molecule has a higher electron density than the rest. If the vectors do cancel out, the molecule is nonpolar because the electron distribution in the molecule is even.

Review a mathematical treatment of dipole moments in Section 2: Molecular Dipole Moments, in Polar Covalence.

In the figure below, we can see how dipole moments determine molecular polarity. The oxygen molecule, O2, is homonuclear, or made up of only one element type. Therefore, there cannot be a difference in electronegativity and it is a nonpolar molecule.

Carbon monoxide, CO, consists of two different types of atoms (carbon and oxygen). Oxygen is more electronegative and therefore the dipole moment goes from the carbon to the oxygen (NOTE: The figure from your text contains an error here). Because there is a dipole moment, carbon monoxide is a polar molecule.

Carbon dioxide, CO2, on the other hand, also contains carbon-oxygen bonds. However, the two dipole moments are 180 degrees opposite each other and cancel each other out. Consequently, CO2 is a nonpolar molecule that contains polar bonds.

Image of a homonuclear molecule, carbon monoxide, and carbon dioxide.

3i: Draw resonance structures.

We create resonance structures for molecules in which some electrons are delocalized. In general, resonance structures describe a molecule or polyatomic ion in which we could write more than one equivalent Lewis electron dot structure.

    • Why are resonance structures necessary?
    • Draw resonance structures for a given molecule or polyatomic ion.

First, review the rules for writing Lewis electron dot structures in Section 2: Lewis Dot Structures, and, Section 3: How to Draw a Lewis Dot Structure, in The Shared-Electron Covalent Bond.

Resonance structures are necessary when there are multiple equivalent Lewis electron dot structures that could be written for a given molecule or polyatomic ion. Generally, resonance structures involve double bonds.

Let’s look at the example of the nitrate ion, NO3. In the nitrate ion Lewis electron dot structure, we make two single N-O bonds and one double N=O bond. When we draw this, it does not matter which nitrogen-oxygen bonds we make single or double. We say that all three possibilities are resonance structures and we denote these with double arrows between them:

Image of the Nitrate Ion, NO3.

In reality, the true structure of the nitrate ion is a superposition, or combination of all three of these resonance structures. Each bond is really about 1 ⅓ of a bond, or bond order.

Image of the Nitrate Ion as a superposition.

We write resonance structures because we cannot accurately draw the true structure of compounds with resonance.

Review the section for more examples of molecules and polyatomic ions with resonance structures, Multiple Equivalent Structures: Resonance, in The Shared-Electron Covalent Bond.

Unit 3 Vocabulary

      • Bent geometry
      • Bond order
      • Chemical bond
      • Covalent bond
      • Dipole moment
      • Dipole-dipole force
      • Dipole-induced dipole force
      • Electron fluid
      • Electronegativity
      • Exothermic
      • Homonuclear
      • Hybrid (or molecular) orbital
      • Hybridization theory
      • Ion
      • Ion-dipole force
      • Ion-induced dipole force
      • Ionic bond
      • Lewis Dot structure
      • London force
      • Metallic bond
      • Molecular orbital
      • Molecular polarity
      • Molecule
      • Nonpolar molecule
      • Permanent dipole/polar molecule
      • Polar covalent bond
      • Polar molecule
      • Resonance structures
      • Sp hybrid orbital
      • Sp2 hybrid orbital
      • Sp3 hybrid orbital
      • Superposition
      • Tetrahedral molecules
      • Trigonal planar molecule
      • Valence
      • Valence electron pair
      • Valence shell electron
      • Van der Waals force
      • VSEPR theory
      • Wavefunction
Last modified: Tuesday, June 11, 2019, 8:41 AM