Draw Lewis Dot Structures

Drawing Lewis dot structures (also known as Lewis structures or Lewis diagrams) can be confusing, particularly for a beginning chemistry student. However, these structures are helpful in understanding the bonding and valence electron configurations of different atoms and molecules. The complexity of the drawing will vary depending on whether you are creating a Lewis dot structure for a diatomic (two atom) covalent molecule, a larger covalent molecule, or ionically bonded molecules.

Steps

Drawing Diatomic Covalent Structures

  1. Write the atomic symbol for each atom. Write the two atomic symbols side by side. These symbols will represent the atoms present in the covalent bond. Be sure to leave enough space between the atoms to draw your electrons and bonds.[1]
    • Covalent bonds generally occur between two nonmetals.
  2. Determine the degree of the bond between the two atoms. Atoms can be held together by a single, double, or triple bond. Generally, this will be dictated by each atom’s desire to reach a full Study the Chemical Theory of Valence Bonds with eight electrons (or in the case of hydrogen, two electrons). To find out how many electrons each atom will have, multiply the bond degree by two (each bond involves two electrons) and then add the number of unshared electrons.[2]
  3. Add your bonds to the drawing. Each bond is represented with a line between the two atoms. For a single bond, you would simply draw one line from the first atom to the second. For a double or triple bond, you would draw two or three lines respectively.[3]
    • For example, N2 (nitrogen gas) has a triple bond connecting the two nitrogen atoms. So, its bond will be notated in a Lewis diagram as three parallel lines connecting the two N atoms.
  4. Draw unbound electrons. Some of the valence electrons in one or both of the atoms may not be involved in a bond. When this happens, you should represent each remaining electron with a dot around it’s respective atom. In most cases, neither atom should have more than eight electrons bound to it. You can check your work by counting each dot as one electron and each line as two electrons.[4]
    • For example, O2 (oxygen gas) has two parallel lines connecting the atoms, with two pairs of dots on each atom.

Creating Lewis Structures for Larger Covalent Molecules

  1. Determine which atom is your central atom. This atom is usually least Study the Elements of the Periodic Table. As such, it is most capable of forming bonds with many other Study the Chemical and Physical Properties of Atoms in the Periodic Table. The term ‘central atom’ is used because all the other atoms in the molecule are bonded to this particular atom (but not necessarily to each other).[2]
    • Atoms like phosphorus and carbon are often central atoms.
    • In some more complex molecules, you may have multiple central atoms.
  2. Consider the valence electrons of the central atom. As a general (but not all-exclusive) rule, atoms like to be surrounded by eight valence electrons (this is known as the octet rule). When the central atoms bonds to the other atoms, the most stable configuration is one that will satisfy the octet rule (in most cases). This can help you determine the number of bonds that will be between the central atom and the other atoms because each bond represents two electrons.[1]
    • Some large atoms such as phosphorus can break the octet rule.
    • For example, carbon dioxide (CO2) has two oxygens covalently double-bonded to the central atom, carbon. This allows the octet rule to be satisfied for all three atoms.
    • Phosphorus pentachloride (PCl5) breaks the octet rule by having five bonding pairs around the central atom. This molecule has five chlorine atoms covalently single-bonded to the central atom, phosphorus. The octet rule is satisfied for each of the five chlorine atoms, but it is exceeded for the phosphorus atom.
  3. Write the symbol of your central atom. With larger covalent molecules, it is best to start the drawing with the central atom. Resist the urge to write all of the atomic symbols at the same time. Leave plenty of room around the central atom to place your other symbols after you have determined their place.[4]
  4. Show the electron geometry of the central atom. For each unshared electron pair, draw two small dots right next to each other around the central atom. For each single bond, draw a line going away from the atom. For double and triple bonds, instead of one line, draw two or three, respectively. This maps out where the other molecules can bond to the central atom.[2]
  5. Add remaining atoms. Each remaining atom in the molecule will attach to the one of the bonds coming from the central atom. Write the symbol for each of these atoms at the end of one of the bonds you placed around the central atom. This indicates that electrons are being shared between that atom and the central atom.[1]
  6. Fill in remaining electrons. Count each bond as two electrons (double and triple bonds as 4 and 6 electrons, respectively). Then add electron pairs around each atom until the octet rule is satisfied for that atom. You can check your work on each atom by counting each dot as one electron and each bond as two electrons. The sum should be eight.[4]
    • Of course, exceptions include atoms that exceed the octet rule and hydrogen, which only has zero or two valence electrons at any given time.
    • When a hydrogen molecule is covalently bonded to another atom, it will have no other unshared electrons surrounding it.

Making Lewis Structures for Ions

  1. Write the atomic symbol. The atomic symbol for an ion will the be the same as the atomic symbol for the atom that formed it. Leave enough space on the paper around the symbol to be able to add electrons and brackets later. In some cases, ions are polyatomic (more than one atom) molecules and are designated by writing the atomic symbol for all atoms in the molecule.[1]
    • To create the symbol for polyatomic ions (such as NO3- or SO42-), follow the instructions for “Creating Lewis Structures for Large Covalent Molecules” in the above method.
  2. Fill in the electrons. Generally, atoms are neutral and do not carry a positive or negative charge. However, when an atom loses or gains electrons, the balance of positive and negative charge in the atom is altered. Then the atom becomes a charged particle known as an ion. On you Lewis structure, add any extra electrons and remove any electrons that were given up.[2]
    • When drawing the electrons, keep the octet rule in mind.
    • When electrons are lost, a positive ion (known as a cation) is formed. For example, lithium loses its one and only valence electron during ionization. Its Lewis structure would just be ‘Li’ with no dots around it.
    • When electrons are gained, a negative ion (known as an anion) is formed. Chlorine gains one electron during ionization, giving it a full shell of eight electrons. Its Lewis structure would be ‘Cl’ with four pairs of dots around it.
  3. Designate the charge of the ion. Counting dots on every atom would be a tedious way of determining if that atom had a charge. To make the structures easier to read, you need to show that your structure is an ion with some charge. To show this, draw brackets around the atomic (or polyatomic) symbol. Then, write the charge outside the brackets in the upper right corner.[2]
    • For example, the magnesium ion would have have an empty outer shell, and would be notated as [Mg]2+.

Video

Sources and Citations

  1. ↑ 1.0 1.1 1.2 1.3 http://www.chem.ucalgary.ca/courses/351/Carey5th/Ch01/ch1-3depth.html
  2. ↑ 2.0 2.1 2.2 2.3 2.4 http://web.chem.ucla.edu/~harding/lewisdots.html
  3. ↑ http://web.chem.ucla.edu/~harding/lewisdots.html
  4. ↑ 4.0 4.1 4.2 http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/lewis.html

You may like

Retrieved from "https://kipkis.com/index.php?title=Draw_Lewis_Dot_Structures&oldid=338668"