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Unit 2

2.0 Unit 2 Overview: Molecular and Ionic Bonding

4 min readjune 29, 2021

dalia

Dalia Savy


AP Chemistry 🧪

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2.0: Unit Overview

Welcome to Unit 2 of AP Chemistry! Now that we've learned about atomic structure and the basics in unit 1, we can build upon this new knowledge by discussing Molecular and Ionic Bonding. While the last unit went in-depth about the structure of the atom and electrons, this unit will discuss at large the ins and outs of molecular and ionic compounds - compounds with more than one atom or ion bonded together. 
Throughout the unit, we will cover:
  • the different types of bonds
  • how electron orbitals relate to covalent bonding
  • how chemists represent different kinds of bonds using:
    • Lewis Dot Diagrams
    • VSEPR
According to the CollegeBoard, "In Unit 2, students apply their knowledge of atomic structure at the particulate level and connect it to the macroscopic properties of a substance. The structure and arrangement of atoms, ions, or molecules, as well as the forces between them, can determine both the chemical and physical properties of materials. These forces, called chemical bonds, are distinct from typical intermolecular interactions. Let's get started!

The Big Picture of Unit 2

The overarching question of unit 2 is "How are molecular compounds arranged?". Bonding is all about potential energy. When atoms and ions bond to form compounds, they decrease the potential energy between them and create different compounds. For example, two hydrogen atoms and an oxygen atom may react to form H2O, otherwise known as water. Water is known as a covalent compound in which electrons are shared within a molecule. Other compounds such as NaCl are ionic compounds. In ionic bonds, electrons are directly transferred from a metal such as sodium to a nonmetal like chlorine to form a positive cation and a negative anion.
https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2FScreen%20Shot%202021-06-29%20at%201.34-rWv9vP7LHP9F.png?alt=media&token=f0155431-d451-49a9-a9c6-f059eb0d214a
Along with the types of bonds, you'll also learn how chemists theorize about why bonds happen and how they model molecules. Molecules are displayed using Lewis Dot Diagrams, which show elements, bonds, and valence electrons. However, Lewis Structures aren't always perfect, which leads to chemists using the concepts of formal charge and resonance to adjust. Chemists have also developed theories like VSEPR and the localized model of covalent bonds for explaining molecules and molecular geometry. These models aren't the only models that chemists use. You may also learn about the Molecular Orbital Model, but the AP exam likely won't test it outside of basic conceptual knowledge of the model.

The Details of Unit 2

2.1-2.4: The Three Types of Bonds

There are three major types of bonds: ionic bonds, metallic bonds, and covalent bonds. Compounds utilizing both a metal and a nonmetal will use ionic bonds.  The nonmental will often "steal" an electron from the metal, usually an alkali or alkaline earth metal. The metal then becomes a positively charged ion, and the nonmetal becomes a negatively charged ion. They are attracted to each other by the electromagnetic force. When atoms "share" electrons among themselves, they are connected via covalent bonds. We'll see that this sharing occurs via overlapping orbitals. Metallic bonds form within metals and alloys (mixtures of metals) in which electrons are delocalized in an "electron sea."
You'll also learn the connection between bonding and intramolecular forces, that is, forces that occur between atoms within a molecule. Bonds occur to minimize the potential energy between atoms. When atoms are too close, there is repulsion; when they are far apart, there's no minimal potential energy between them, which you can see in the following graph:
https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2FScreen%20Shot%202021-06-29%20at%201.41-jbrHO9HNGcQW.png?alt=media&token=9f2ea08a-e67c-469c-a84e-5aa52f44f8b7

2.5-2.7: Bonding Models

After learning the different kinds of bonds, we'll discuss some other models chemists use to represent bonds and molecules. To visually display molecular structures, chemists use Lewis Dot Diagrams. These structures help simplify what a molecule is and what atoms connect to which. While it may be obvious in something like SO (it's pretty to see that this will involve S bonded directly to O), it isn't as evident in a molecule like C2H5OH, so Lewis Structures make it easier.
https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2FScreen%20Shot%202021-06-29%20at%201.42-veUtSOTqp8Ft.png?alt=media&token=33638931-52de-417d-9444-47f113ce286c
Lewis Dot Diagrams can also represent aspects of molecules such as resonance and formal charge. Resonance is a phenomenon in which a molecule has multiple optimal Lewis Structures. This often occurs when there is a double bond in a molecule that can be swapped around without creating more formal charge. For example, the nitrate ion exhibits resonance:
https://firebasestorage.googleapis.com/v0/b/fiveable-92889.appspot.com/o/images%2FScreen%20Shot%202021-06-29%20at%201.43-SJcUH98pBBTQ.png?alt=media&token=189fc49b-afa2-4c5c-b20c-46a9d0a40cc6
In reality, the double bond is delocalized and shared among the three single bonds. Along with resonance, chemists also observe formal charge within Lewis structures. Formal charge describes the charge felt by an atom within a molecule. Finally, we'll discuss bonding theory and how sigma and pi bonds can describe orbitals within bonds. VSEPR theory and hybridization will also be applied to molecular geometry to round off the unit.

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