Bonds are connections between atoms. A solid grasp of valence shell electron pair repulsion (VSEPR) theory will help you understand how elements that differ by one or two atomic numbers behave. According to VSEPR theory, the number of electrons an element has corresponds with its chemical properties. For example, sodium differs from neon and potassium by one atomic number, but it resembles potassium, not neon. Sodium and potassium both have one valence electron, which explains their similar properties, while neon is a stable element with eight valence electrons. We use VSEPR to predict the three-dimensional structure, or geometry, of molecules.
Completing this unit should take you approximately 5 hours.
Now that we understand atomic structure and electron configurations, we are ready to learn how valence electrons combine to form chemical bonds between atoms.
To begin our exploration of bonding, we need to define the two main types of bonds: covalent, and ionic. Covalent bonds occur mostly between nonmetal atoms. In a covalent bond, the electrons are shared between atoms. Ionic bonds occur between metal ions and nonmetal ions, or polyatomic ions. In an ionic bond, the positively charged metal ion or ions are attracted to the negatively charged nonmetal ion or ions. It is an electrostatic interaction.
Watch these videos for an overview of the differences between the two types of bonding. Note the types of elements that are involved in each type of bonding.
Watch this video, which reviews ionic and covalent bonds, and introduces another common type of bonding called metallic bonds. Metallic bonds occur between metal atoms.
When discussing covalent bonds, we need to make one further definition. Nonpolar covalent bonds are covalent bonds where the electrons are equally shared in the bond. Polar covalent bonds are covalent bonds where the electrons are not equally shared in the bond. This occurs when the bonding atoms have different electronegativities. Watch this video to learn how to identify polar covalent bonds.
Now that we have an overview of covalent and ionic bonding, we can see how quantum mechanics informs our understanding of bonding.
As you read this text, note that the section "Classical Models of the Chemical Bond" describes ionic bonding, which is largely based on the attractive forces among charged particles. This section also discusses models of covalent bonding that describe bonds as areas of shared electrons between atoms. While this is simplistic, it does accurately describe the three-dimensional shape of molecules.
The section "Quantum Models of Chemical Bonding" briefly outlines hybrid orbital theory, which is also often used to describe the three-dimensional shape of molecules. We will explore these models further in later sections of this unit.
Now that you understand the types of chemical bonds, we explore Lewis Dot Diagrams, a common way to draw covalently-bonded molecules. In Lewis dot diagrams, we use small dots to show the valence (outer shell) electrons of an atom. We can then combine the valence electrons of different atoms together to form covalent bonds, using a set of rules.
Watch the following videos in order. These four videos guide you through the steps for making Lewis dot diagrams for small molecules, and show worked examples.
Most Lewis dot structures are guided by the Octet Rule, which states that bonded atoms must have eight valence electrons. However, there are exceptions to the octet rule. Some atoms, such as hydrogen, helium, and boron, have less than an octet. Many atoms, such as those past sulfur on the periodic table, can have an expanded octet, meaning more than eight valence electrons. Watch this video, which outlines these cases.
In some cases, we can draw more than one equivalent Lewis structure for a given molecule or polyatomic ion. These equivalent structures are called resonance structures. The true structure of these types of molecules or polyatomic ions are actually a hybrid, or mixture, of the resonance structures. This is called a resonance hybrid.
Watch these videos, which introduce resonance structures with the example of the nitrate, Na3- ion, which has three resonance structures. All three structures are equivalent – they only differ by the placement of the double bond.
Sometimes we can draw more than one Lewis structure for a given molecule or polyatomic ion that are not equivalent. In this case, we use a concept called formal charge to determine, which Lewis structure is best. Formal charge is not actually a charge; rather, it is just a system to keep track of electrons in a given Lewis structure.
Watch these two videos, which teach you the rules for assigning formal charge to atoms in a Lewis structure, and show examples of using formal charge to determine the best possible Lewis structure for a given molecule.
Read this text to review the material that was discussed in the videos you have just watched.
Now that we have studied bonding and how to draw chemical structures, we can investigate how molecules interact with each other. Molecules interact with each other through intermolecular forces; forces that hold molecules together, but are not covalent or ionic bonds. As we will learn, there are different types of intermolecular forces that occur between different types of molecules.
Read this text. The first two sections describe the differences between bonding interactions and intermolecular force interactions in terms of electrostatics, or charge interactions. You should focus on section 3, which outlines the different types of intermolecular forces and the types of molecules that experience these different forces.
Now that we can draw Lewis structures, we can determine the shape, or molecular geometry, of molecules. This is important because the shape of a molecule often determines its reactivity, its intermolecular forces, and other properties.
For example, carbon dioxide, CO2, and water, H2O, are both small molecules with three atoms total. However, CO2 has a linear shape. This makes it nonpolar, and it is a gas at room temperature. Water has a bent, or v shape. This makes water a polar molecule, and, as we know, it is a liquid at room temperature.
The main way we understand molecular geometry is through Valence Shell Electron Pair Repulsion Theory, or VSEPR. The idea of VSEPR is that pairs of electrons in the valence, or outer shell, in a molecule will be repelled and get as far apart as possible in three-dimensional space. VSEPR is based on Lewis dot structures.
Read this text, which outlines the different geometries predicted by VSEPR theory. Pay attention to the pictures that show the shapes of these different geometries.
Next complete this VSEPR practice assessment, which gives you the opportunity to practice applying your understanding of Lewis structures and VSEPR.
While VSEPR works well to describe molecular geometry, it does not take into account what we know about atomic structure from quantum mechanics. How do electrons in atomic valence orbitals form bonds? The hybrid orbital model helps to explain this.
Read these sections, which describe how atoic orbitals combine to form hybrid orbitals in molecules. The molecular shapes and the bond angles of hybrid orbitals match the molecular shapes and bond angles predicted by VSEPR theory.
Read this page to further explore dipole moments, ionic character, and ionic solids.
Take this assessment to see how well you understood this unit.