Lewis Structures: Non-Charged

Sections

Overview
Full Text

Overview

Lewis structures show:

• Atoms, as periodic symbols
• Bonds, as lines
• Lone pairs, as dots

Definitions

• Bonds – shared electrons
• Lone pairs – nonbonding electrons

The octet rule:

• Atoms form bonds that will give them the same number of electrons as a noble gas

General predictions:

• Hydrogen – forms one bond
• Oxygen – forms two bonds
• Nitrogen – forms three bonds
• Carbon – forms four bonds
• Halogens – form one bond

Steps for drawing Lewis Structures:

• Form one bond from each H
• Form bonds to connect the other atoms
• Check to see if each atom has achieved an octet

Full Text

Here, we will learn how to draw Lewis structures, which combine individual atoms into molecules by showing how they share electrons.

Lewis structures: three main constituents

Atoms

• The atoms that we will encounter most often in organic chemistry are hydrogen, oxygen, nitrogen, and carbon.

Bonds

• Atoms each donate one electron to make a bond of two shared electrons.

Lone pairs

• Lone pairs are made of two nonbonding electrons.

Periodic Table

The Periodic Table, as it relates to anatomic structure

Draw an outline of the periodic table...

• The outline should have two columns on the left, an empty space to represent transition metal elements in the center, and 6 columns on the right.
• Label the 8 columns that we have created with the numbers 1 through 8, starting with column 1 on the far left and ending with column 8 on the far right.
• Indicate that the numbers are referred to as group numbers. Atoms within the same group have similar characteristics.
• Write that the number of valence electrons in an atom is the same as its group number in the periodic table. Only valence electrons (rather than core electrons) are represented in Lewis structures.

Lewis dot structure of oxygen

To illustrate how we represent the electrons of an atom, let's draw the Lewis dot structure of oxygen.

• Show that oxygen is in row 2 and group 6 of the periodic table.
  Since its group number is 6, oxygen has 6 valence electrons in one atom.
• Draw an O and place one valence electron by itself on each side of the O.
  Since oxygen has more than 4 electrons, pair each of the two remaining two electrons with an electron that has already been drawn.

Lewis structures of molecules: the octet rule

The Lewis structures of molecules are made by combining Lewis dot structures of individual atoms. When atoms combine to form bonds, they tend to form a number of bonds that will give them the same number of electrons as a noble gas; this is known as the octet rule.

• Label group 8 on the periodic table as "noble gases."
• Show that helium is in row 1 of group 8, and neon is in row 2.
• Now, show that hydrogen is in row 1 of group 1.
  Hydrogen therefore has one valence electron, but will tend to bond so that it has two electrons like helium.

General predictions

Start a table. Across the top, write "general predictions."

• Denote that hydrogen will form one bond.
• Write that carbon will form 4 bonds.
  Carbon is in row 2, group 4 of the periodic table. Carbon has four valence electrons, but wants 8 electrons like neon. Thus, it will tend to form four bonds and access two electrons for each bond. Hence, a total of 8 electrons are available.
• Write that nitrogen will form 3 bonds.
  Next to carbon, show that nitrogen is in group 5 of row 2. Nitrogen has 5 valence electrons and will bond to gain three more for a configuration like neon.
• Write that halogens (abbreviated as X) will form 1 bond.
  Next to oxygen, show that fluorine is in group 7 of row 2. Fluorine, like all the halogens, have seven valence electrons and only needs one more to achieve 8 like neon.

Lewis structure of formaldehyde: formula CH2O.

In the same way that made the Lewis dot structures of oxygen, draw the structures of carbon and hydrogen.

• Draw a carbon and surround it by four dots, one on each side.
  Carbon is in group 4, so it has 4 valence electrons.
• Draw a hydrogen with one dot.
  Hydrogen is in group 1, so it has 1 valence electron.

To create the molecular Lewis structure, combine the individual Lewis dot structures as follows...

• On the right side of the carbon with its four dots, draw the oxygen with its six dots.
• On the left side of the carbon, draw one hydrogen and its dot slightly above the carbon, and one slightly below.
• Now, connect one electron from each of the hydrogens to an electron on the carbon with a line to indicate that they have formed a bond.
  Note that the first step of drawing Lewis structures is to form one bond on hydrogen atoms.
• Connect an electron on the carbon and one on the oxygen with a line.
  Note that the second step of drawing Lewis structures is to connect atoms that form more than one bond.
• Connect the last electron on the carbon with another electron on the oxygen with a second line. ** This represents a double bond.**
  Since carbon has achieved an octet (as a reminder, the octet rule tells us that atoms bind to achieve the configuration of a noble gas), it cannot form any more bonds. Therefore, the remaining two pairs of electrons remain on oxygen as lone pairs.
• Note that the final step of drawing Lewis structures is to continue forming bonds from unpaired electrons to achieve an octet for each atom.

Validation

As a validation of our structure, let's count the electrons used to ensure they match the original number of valence electrons.

• From the molecular formula of CH2O, calculate that the total number of valence electrons in formaldehyde is:
  4 + (2 x 1) + 6 = 12 electrons
• Look at the Lewis structure we have drawn to calculate the number of electrons involved:
  (4 bonds x 2 electrons/bond) + 4 unpaired electrons = 12 electrons
• Since the hydrogens in formaldehyde have one bond, the carbon has four bonds, and the oxygen has two bonds, all the atoms follow the trend that we predicted.

Lewis structure with a molecule that is not neutral: methoxide

Let's repeat the steps for drawing a Lewis structure with a molecule that is not neutral.

• When atoms do not bond according to the predicted trend, they display either a positive or negative charge. This formal charge is given by the formula:
  Formal charge = # valence electrons - (# unpaired electrons + # bonds)
• The formula for methoxide, a negatively charged molecule, is CH3O-.
• To draw its Lewis structure, we will start with the individual Lewis dot structures of carbon, hydrogen, and oxygen...*
• First, see that oxygen has six valence electrons by locating it under group 6.
• Draw a dot on each side, then pair the two remaining dots with two dots already drawn.
• Next, determine that carbon has 4 valence electrons by seeing that it is in group 4 of the periodic table.
  Draw one dot on each of its four sides.
• Now, show hydrogen, which is under group 1, with its one valence electron.

Let's draw the Lewis structure for methoxide by combining the structures of the atoms.

• First, show carbon with its 4 valence electrons.
• Next, place an oxygen with its six valence electrons above the carbon.
• Now, put the three hydrogens, each with one valence electron, below and on either side of the carbon.

We still have to account for the negative charge on methoxide.

• The negative charge is represented by a single electron, which will go to the most electronegative element in the molecule. A high electronegativity indicates that an atom will attract electrons.
• In regards to electronegativity: On the periodic table, draw an arrow pointing from the bottom left corner to the upper right corner to indicate the direction of increasing electronegativity. Electronegativity increases from left to right across a row in the periodic table, and it decreases from top to bottom down a column.
• Show the negative charge by adding another electron to the oxygen.
• Combine an electron on each of the hydrogens with a corresponding electron on the carbon with lines to form three C-H bonds.
• Now, combine an electron from carbon and an electron from oxygen to form a C-O bond.
  Since carbon has now satisfied the octet rule, the remaining electrons will remain on oxygen. Oxygen does not have the appropriate number of valence electrons to satisfy the octet rule, and will exhibit a formal charge.

Formula for formal charge

• Use the formula for formal charge to count the number of electrons on each atom in the Lewis structure of methoxide.
• The hydrogens have 1 electron from their shared bond with carbon.
• Since this is the same as their 1 valence electron, they have a formal charge of:
  1 – 1 = 0
• The carbon also has a charge of 0, since it has 4 electrons from its bonds:
  4 – 4 = 0

Calculate that the formal charge of oxygen is:

• 6 valence electrons – (6 unpaired electrons + 1 bond) = -1
  Indicate the negative formal charge on the oxygen in this Lewis structure by drawing a negative sign within a circle.
  The charge of methoxide is the same as the -1 formal charge on the oxygen.
• Write that the charge of a molecule is the same as the sum of all the formal charges of the atoms.

Electronegativity

• Now that we have seen the effect of electronegativity in a molecular structure, we can relate the electronegativity of atoms to the type of bonds that they create. Start a second table and label the top as: Categories of bonds
• Write the three categories of bonds as three separate columns:
  Covalent
  Polar covalent
  Ionic

In the first column of covalent bonds, write that: Electrons are shared equally

Next, put two examples: C-C and C-H

• These are covalent bonds because there is little or no electronegativity difference (<0.5) between the atoms carbon (electronegativity 2.5) and hydrogen (electronegativity 2.1).

Under the second column of polar covalent bonds, write that: Electrons are not shared equally.

• As an example, write: C-O
  There is a large difference in electronegativity (0.5 - 1.7) between carbon and oxygen (electronegativity 3.5), so the atoms are not shared equally. This difference is known as a dipole moment.

Dipole moment

• To represent the dipole moment on our structure of formaldehyde, draw a small arrow pointing from the carbon to the oxygen. This indicates that the direction of the dipole goes from the less electronegative carbon to the more electronegative oxygen.
• Draw another arrow on the structure of methoxide, also pointing from the carbon to the oxygen. Since oxygen has a negative charge, it will draw the electrons away from the carbon.
• Indicate that carbon has the lesser share of electrons by drawing a partially positive charge (+) on the atom.

Dipole moment of formaldehyde

The same concept applies to formaldehyde; even though the oxygen in this case does not have a formal negative charge, it is still attracting more electrons than the carbon.

• Draw a partially positive charge on the carbon in formaldehyde, and a partially negative charge on the oxygen.
• Write in the third column of ionic bonds that: electrons are not shared
  Put as an example: the bond between sodium (Na+) and methoxide (CH3O-)
  Sodium has such a low electronegativity (0.9) that the electrons of the bond are possessed solely by methoxide (in an ionic bond, the electronegativity difference > 1.7).