Unit 2: Topic 5 - Lewis Diagrams Flashcards
Lewis diagrams are used to visualize covalent and ionic bonds.
In a Lewis diagram, the chemical symbols represent the elements of the atoms, and the dots represent their valence electrons. The dots are placed in four groups of one/two electrons each, with eight electrons representing a closed shell/noble gas configuration.
Chemicals bonds between nonmetals are typically covalent bonds, or shared pairs of electrons (AKA a bonding pair). Two atoms each share a single electron with each other. The electron pairs that aren’t part of the covalent bond are known as “lone pairs.”
You can use two dots or a single line to represent a bonding pair. A single bond (one bonding pair) can be represented with one line or two dots. A double bond (two bonding pairs) can be represented with two lines or four dots. A triple bond (three bonding pairs) can be represented with three lines or six dots.
Examples of Lewis diagrams.
Here are examples of Lewis diagrams of covalent bonds.
Chemical bonds between a metal and a nonmetal are typically ionic bonds, electrostatic attractions that form between two oppostively charged ions (atoms that have an electric charge). The metal atom loses electrons to become a postively-charged cation, and the nonmetal atom gains those electrons to become a negatively-charged anion. Each atom prefers to have a full octet, or 8 valence electrons, like noble gases (see octet rule). By convention, you should put brackets around the anion when drawing the Lewis diagram of an ionic compound. Also, make sure to include the charges of the ions in the top right corner of the element symbols and place the coefficients before each.
Octet Rule
Atoms tend to prefer to have eight valence electrons. In nonmetals, these electrons can be shared pairs (covalent bonds) or unshared pairs (lone pairs). Atoms with fewer than eight valence electrons tend to undergo chemical reactions to become more stable; atoms are generally stable when they have eight valence electrons or a “full octet.” This is why noble gases rarely form compounds. Since they already have a stable configuration (full octet), there is no need for them to gain, lose, or share electrons to become more stable. There are some expectations to the octet rule.
What are Odd-electron molecules?
Odd-Electron Molecules:
Some molecules have an odd number of electrons. When this happens, all atoms in the molecule cannot satisfy the octet rule. An example of an odd-electron molecule is nitrogen dioxide or NO₂. NO₂ has a total of 17 valence electrons, as each oxygen atom (there are two of them) contributes 6 electrons, while the nitrogen atom contributes 5 electrons.
What are expanded octets?
Expanded Octets:
Period 2 elements cannot have more than eight valence electrons around the central atom. However, period 3 elements and beyond are able to have more than eight valence electrons around the central atom, as the d sublevel becomes available starting with the third period; atoms can use these orbitals in bonding, leading to an expanded octet. Examples of expanded octets are phosphorus pentachloride and sulfur hexafluoride.
What are incomplete octets?
- Incomplete Octets
Some elements can be stable with less than eight valence electrons. Such elements include hydrogen (stable with only two valence electrons), beryllium (stable with only four valence electrons), and boron and alumnium (stable with only six valence electrons).
Pooled Electron Method
When “simple inspection” (determining the structure of a molecule by comparing the number of bonds each atom/element can form) doesn’t work, we can use the “pooled electron” method to draw a Lewis diagram.
1) Add together the total number of valence electrons of all the atoms in the molecule/compound.
2) Distribute the electrons so that each atom has 8 electrons (either as shared or unshared pairs).
Example: Carbon monoxide (CO)
Carbon monoxide has a total of 10 electrons, as carbon has 4 valence electrons and oxygen has 6 valence electrons. These 10 electrons have to be distributed such that both the carbon and oxygen atoms have 8 electrons (can be shared or unshared). As a result, your diagram will consist of a triple bond (six shared electrons) between the carbon and oxygen atoms, and a lone pair (two unshared electons) for both the carbon and oxygen atoms.
Counting the Number of Electrons for Polyatomic Ions
If a polyatomic ion has a negative charge, the number of electrons it has is equal to the total number of valence electrons of all the atoms have plus the charge. This is because electrons have negative charges. More electrons = more negative.
Example: sulfate (SO₄⁻²)
Both sulfate and oxygen have 6 valence electrons. This means you do 1(6) + 4(6) to get 30 total valence electrons, as there is 1 sulfate atom and 4 oxygen atoms. Then, you add 2 (due to -2 charge) to 30 to get a total of 32 electrons.
If a polyatomic ion has a positive charge, the number of electrons it has is equal to the total number of valence electrons of all atoms minus the charge. As mentioned before, electrons have a negative charge. Therefore, fewer electrons = less negative/more positive.
Example: ammonium (NH₄⁺)
Nitrogen has 5 valence electrons, while hydrogen has 1 valence electron. This means you do 5 + 4(1) to get 9 electrons, as there is 1 nitrogen atom and 4 hydrogen atoms. Then, you subtract 1 (due to +1 charge) from 9 to get a total of 8 electrons.
Central Atom of Lewis Diagram
If a molecule has three or more atoms, the central atom is typically the atom that has the lowest subscript in the molecular formula and the atom that can form the most bonds. If all the atoms form the same number of bonds, then the least electronegative atom typically occupies the central position (you can use periodic trends to determine this).
In CERTAIN cases, the central atom can have more than 8 electrons. That is okay, as long at the total does not exceed 12 (this relates to expanded octets).
