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Lewis Structures: Visualizing Chemical Bonds and Electron Arrangements
Discover the power of Lewis structures in chemistry. Learn to draw diagrams for molecules like H2O and NaCl, predict molecular geometry, and understand bonding patterns. Perfect for students and aspiring chemists!

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  2. Examples0/8 watched
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Now Playing:Lewis structures – Example 0a
Intros
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  1. Introduction to dot and cross diagrams
  2. Introduction to dot and cross diagrams
    Lewis structures: introduction.
  3. Introduction to dot and cross diagrams
    How to draw covalent Lewis structures.
Examples
0/8 watched
  1. Draw Lewis structures for simple covalent compounds.
    Draw dot and cross structures for the following molecules:
    1. HCl

    2. N2_2

    3. NH3_3

    4. AlCl3_3

    5. NH4_{4} +

Introduction to bonding
Jump to:NotesConceptExampleFAQsPrerequisites
Notes
In this lesson, we will learn:
  • To construct Lewis (dot and cross) structures using knowledge of electrons shells and the octet rule.
  • To use Lewis diagrams to explain the difference between ionic and covalent bonding.
  • Some of the exceptions to the octet rule and the significance of this to theories of bonding.

Notes:
  • In C11.4.2: Ionic and covalent bonding we saw the two main types of chemical bond: ionic and covalent bonding. These are attractive forces that hold individual atoms together allowing them to form molecules and larger, more complex structures.
    We can obviously see that most atoms are just more stable existing in compounds or molecules, such as CH4 or O2, compared to being single isolated atoms, like H or O.
    In C11.4.1: Introduction to bonding we mentioned that theories of bonding exist to try to explain why this is true. One theory of bonding is valence bond theory, which uses Lewis structures to show the bonding in simple molecules.

  • In Lewis structures, dots and crosses are used to show the valence electrons, where dots or crosses are used for the central atom and the other is used for the atom(s) it is bonding to. They can be used to show covalent and ionic bonding between atoms in molecules and compounds, where electrons are shared or donated/accepted to achieve a full outer shell. Valence means outer shell or highest energy electrons; dont draw core electrons in Lewis structures!
    • When drawing these diagrams, look at our lesson Electronic structure: 288 rule for help when finding the number of outer shell electrons. Remember the octet rule, where electrons will try and fill their outer shell with an octet of (eight) electrons.

  • Lewis structures for covalent compounds can be drawn step-by-step as below, with CO2_2 and H2_2O): as examples:
    • Step 1: Write the chemical symbol for the central atom in the molecule and draw a ring around it (this will be the outer electron shell). If the molecule is diatomic, draw either one atom.
      • CO2:carbon atom diagram
      H2O: oxygen atom diagram
    • Step 2: Write the symbols of the other atoms equally spaced around it. Draw a ring around each of these so that their rings each have an overlap into the central ring. These rings are the outer electron shells.
      • CO2: carbon dioxide diagram


      H2O: water H2O diagram


    • Step 3: For the non-central atoms, fill in their outer shells of electrons using the 288 rule fromElectronic structure: 288 rule (288 rule). Remember to fill up to four electrons individually then start pairing them up. Mark all these electrons using either a dot or a cross. Be consistent with the dots and crosses; if you use crosses for one H atom, use it for all H atoms.

      • CO2:carbon dioxide with electrons diagram


      H2O: water H2O with electrons diagram

    • Step 4: LOOK CAREFULLY AT YOUR DIAGRAM. Now make sure any unpaired electrons are placed in the overlap with the central atom.
      This is a clue to show how many pairs of electrons (how many covalent bonds) there will be with the central atom. Each of these will be paired with an electron from the central atom too.

      • CO2:carbon dioxide with paired electrons diagram


      H2O: water H2O with proper electrons placement diagram

    • Step 5: Now fill the outer shell of electrons in the central atom. Use the other sign to the one you used with the non-central atom. Start by making pairs of dots and crosses with the electrons of the non-central atoms in the overlapping area. You can only pair dots and crosses in a 1:1 ratio. Now, any extra electrons can be placed outside the overlaps on the outer shell. Remember that if there are more than four electrons in the shell they must now be paired. Electrons in the overlapping area count for both overlapping atoms!

    • Step 6: Look carefully at your diagram,. If you have followed these steps, all of your atoms should have exactly 8 electrons (2 if it is hydrogen) in the outer shell. Remember that electrons in the overlaps count for both atoms.

      • CO2: carbon dioxide with all electrons diagram


      H2O: Water H2O with all electrons diagram

  • Dot and cross diagrams for ionic compounds can be drawn like this:
    • Step 1: Write the chemical symbols of the atoms side by side and draw a ring (the outer shell) around them but keep them separate. They must not overlap like in covalent compounds.

    • Step 2: Fill the outer shell with the correct number of electrons for each atom.

      • NaCl example: Sodium and chlorine atoms

    • Step 3: Move the electrons from the metal atom(s) to the non-metal atom(s). This should complete the outer shell of the non-metal atom(s) while leaving empty the outer shell of the metal atom(s). The atoms are now ions.

    • Step 4: Ions should have square brackets surrounding them with their charge on the top-right corner. This shows an ion has been formed. In order to show an electron was transferred and not shared, keep the original dot/cross signage you used. See below, where the Na outer electron is still a dot on Cl.

      • NaCl example: Sodium chloride

  • These Lewis diagrams should help to show the difference between bonding in covalent and ionic bonding in terms of electrons:
    • Covalent bonding involves a sharing of electrons to complete outer shells in both atoms.
    • Ionic bonding involves a discrete transfer of electrons usually from a metal atom to a non-metal atom, which also leads to complete outer shells.

  • Building Lewis structures using the octet rule is applying a theory of chemical bonding. These theories are made to try and explain what we observe and to make predictions that can be tested, but there are exceptions to the octet rule.

    • For example, the molecules NO (nitric oxide) and BF3 (boron trifluoride):
    • Nitric oxide (NO) would have a Lewis structure beginning with six electrons around oxygen, and five electrons around nitrogen. This means nitrogen needs three more electrons and oxygen two to satisfy the octet rule. This cannot be satisfied by:
      • A single bond (one pair of electrons) between the two atoms; it leaves only six electrons on nitrogen and seven on oxygen.
      • A double bond (or two pairs of electrons) between the two atoms; it leaves only seven electrons on nitrogen and eight on oxygen.
      • A triple bond between the two atoms; oxygen now has nine electrons in the outer shell, violating the octet rule.

    • Boron trifluoride (BF3) would have a Lewis structure beginning with three electrons around boron and seven around all three fluorine atoms each. The octet rule cannot be satisfied for boron. It has only three electrons, and with one covalent bond each to the three fluorine atoms, it is still has only six electrons when it should have eight.

    If the theory behind Lewis structures was complete and totally accurate, neither of these two molecules should exist. The fact that they do shows the theory is not totally accurate!
Concept

Introduction to Lewis Structures

Lewis structures are fundamental diagrams in chemistry that illustrate the bonding between atoms and the arrangement of valence electrons in molecules. Named after Gilbert N. Lewis, these structures provide a visual representation of chemical bonding and electron distribution. The introduction video serves as an essential starting point for understanding this crucial concept in chemistry. Lewis structures depict atoms as their elemental symbols surrounded by dots representing valence electrons. These diagrams are invaluable tools for chemists, helping to predict molecular geometry, reactivity, and bonding patterns. By showing how electrons are shared or transferred between atoms, Lewis structures offer insights into the formation of covalent and ionic bonds. They also aid in understanding concepts like resonance structures, formal charges, and octet rule exceptions. Mastering Lewis structures is vital for students and professionals alike, as they form the foundation for more advanced topics in organic and inorganic chemistry.

Example

Lewis structures: introduction.

Step 1: Understanding Lewis Structures

Lewis structures, also known as Lewis dot diagrams or electron dot structures, are a way to represent the valence electrons of atoms within a molecule. These diagrams help predict the bonding and structure of a substance, whether it is ionic or covalent. The primary goal is to visualize how atoms share or transfer electrons to achieve stable electron configurations.

Step 2: The Octet Rule

The octet rule is a fundamental concept in chemistry that states atoms tend to form bonds until they are surrounded by eight valence electrons. This rule is derived from the observation that atoms with eight electrons in their outer shell are particularly stable. In Lewis structures, this rule helps determine how atoms will bond and what kind of bonds (single, double, or triple) will form.

Step 3: Constructing Lewis Diagrams

To construct a Lewis diagram, follow these steps:

  • Identify the total number of valence electrons in the molecule.
  • Write the symbols for the atoms involved and connect them with single bonds initially.
  • Distribute the remaining electrons to satisfy the octet rule for each atom, starting with the outer atoms and moving inward.
  • If necessary, form double or triple bonds to ensure all atoms achieve an octet configuration.
These diagrams are sometimes called dot and cross structures because they use dots and crosses to represent electrons from different atoms.

Step 4: Differentiating Ionic and Covalent Bonds

Lewis structures can illustrate the difference between ionic and covalent bonds. In ionic bonding, electrons are transferred from one atom to another, resulting in the formation of positive and negative ions. In covalent bonding, electrons are shared between atoms, creating a bond through the overlap of electron clouds. The diagrams will show either the transfer or sharing of electrons, helping to visualize the type of bond formed.

Step 5: Exceptions to the Octet Rule

While the octet rule is a useful guideline, there are exceptions. Some molecules have atoms that do not follow the octet rule, either having fewer or more than eight electrons in their valence shell. These exceptions are significant because they highlight the limitations of the octet rule and the need for more advanced theories to explain bonding in certain compounds.

Step 6: Practical Applications of Lewis Structures

Lewis structures are part of valence bond theory, which helps predict the existence, stability, and type of bonds in molecules. For example, they can explain why water is H2O and not H3O, or why nitrogen forms N2 molecules. By constructing these diagrams, chemists can predict molecular shapes, bond angles, and reactivity, making them a valuable tool in both theoretical and practical chemistry.

Step 7: Drawing Lewis Structures for Covalent and Ionic Compounds

To draw Lewis structures for covalent compounds, start by identifying the central atom (usually the least electronegative) and arrange the other atoms around it. For ionic compounds, show the transfer of electrons from the metal to the non-metal, resulting in the formation of ions. Use brackets to indicate the charge on each ion. These visual representations help in understanding the nature of the bonds and the overall structure of the compound.

Step 8: Predicting Molecular Stability

By using Lewis structures, chemists can predict whether a molecule will be stable or not. If a proposed structure does not satisfy the octet rule or results in an unstable arrangement of electrons, it is likely that the molecule will not exist in nature. This predictive power is one of the key benefits of mastering Lewis structures.

FAQs

Here are some frequently asked questions about Lewis structures:

1. How do you draw Lewis structures step by step?

To draw Lewis structures, follow these steps: 1) Count the total valence electrons. 2) Identify the central atom. 3) Connect atoms with single bonds. 4) Place remaining electrons as lone pairs. 5) Check for octets and adjust bonds if necessary.

2. What is the Lewis structure for H2O?

The Lewis structure for H2O (water) shows the oxygen atom in the center with two single bonds to hydrogen atoms. The oxygen also has two lone pairs, giving it a total of 8 electrons (octet). The structure resembles a bent shape.

3. How do you determine formal charges in Lewis structures?

Calculate formal charge using the formula: FC = (# valence e-) - (# non-bonding e-) - ½(# bonding e-). Assign formal charges to atoms and aim for structures with the lowest possible charges, ideally zero.

4. What are resonance structures?

Resonance structures are multiple valid Lewis structures for a molecule where only the electron positions differ. The actual structure is a hybrid of these forms. They're common in molecules with delocalized electrons, like benzene or carbonate ion.

5. How do Lewis structures help predict molecular geometry?

Lewis structures show the arrangement of electron pairs around atoms, which directly relates to molecular geometry through VSEPR theory. The number of electron domains (bonding and non-bonding) determines the overall shape of the molecule.

Prerequisites

Before diving into the intricacies of Lewis structures, it's crucial to have a solid foundation in related chemical concepts. Two key prerequisite topics that significantly enhance your understanding of Lewis structures are molecular geometry and VSEPR theory, as well as conjugation and resonance structures.

Understanding molecular geometry prediction is essential when working with Lewis structures. This concept helps you visualize the three-dimensional arrangement of atoms in a molecule, which is directly related to how electrons are distributed in Lewis structures. The VSEPR (Valence Shell Electron Pair Repulsion) theory, a key component of molecular geometry, explains how electron pairs around a central atom influence the shape of a molecule. This knowledge is invaluable when drawing Lewis structures, as it helps you predict and represent the correct arrangement of atoms and bonds.

Similarly, familiarity with resonance structures is crucial for a comprehensive understanding of Lewis structures. Resonance occurs when a single Lewis structure cannot adequately represent the electron distribution in a molecule. By studying conjugation and resonance, you'll learn how to draw multiple Lewis structures for a single molecule and understand how these structures contribute to the overall stability and properties of the compound.

The interplay between Lewis structures and these prerequisite topics is significant. When you draw a Lewis structure, you're essentially creating a two-dimensional representation of a molecule's electron distribution. Your knowledge of molecular geometry allows you to mentally translate this 2D structure into a 3D model, giving you a more accurate picture of the molecule's true form. Additionally, understanding resonance helps you recognize when a single Lewis structure is insufficient and how to represent more complex electron distributions.

Moreover, these prerequisite topics enhance your ability to predict molecular properties based on Lewis structures. For instance, your understanding of molecular geometry can help you deduce polarity and intermolecular forces, while knowledge of resonance structures aids in predicting reactivity and stability. This interconnectedness demonstrates why a strong grasp of these foundational concepts is vital for mastering Lewis structures and, by extension, more advanced topics in chemistry.

In conclusion, investing time in understanding molecular geometry and resonance structures will significantly enhance your ability to work with Lewis structures. These prerequisite topics provide the necessary context and tools to not only draw accurate Lewis structures but also to interpret and apply them in various chemical scenarios. By building this strong foundation, you'll be better equipped to tackle more complex chemical concepts and problem-solving in your studies and beyond.