6.6 Lewis Structures for Molecules

6.6 Lewis Structures for Molecules

• Draw the Lewis structures for compounds with more than one bond.

• More complex chemical bonds and how they contribute to the structure of a molecule can now be investigated.

• H2 is the simplest molecule.
• There is no attrac tion between the two H atoms when they are far apart.
• The positive charge of each nucleus attracts the electron of the other atom.
• The H atoms are more stable than two individual H atoms.

• A bond is formed when H atoms share electrons.

• Two fluorine atoms are in Group 7A, each with seven exits as diatomic molecules.

• There are elements that have electrons.
• Each F atom can achieve an octet as Diatomic Molecules by sharing its unpaired electron in the Lewis structure.

• Hydrogen and fluorine are examples of nonmetal elements that are diatomic.
• The number of electrons that a nonmetal atom shares and the number of covalent bonds it forms are usually equal to the number of electrons it needs to achieve a stable electron arrangement.

• We draw the Lewis symbols for carbon and hydrogen first.

• The carbon atom is the central atom and the hydrogen atoms are on each side of the Lewis structure.

• There are single lines between the carbon atom and each of the hydrogen atoms.

• The Lewis structure is used to make PCl3.

• The arrangement of atoms can be determined.
• There is only one P atom in PCl3.

• Determine the number of electrons.
• The number of valence electrons for each of the atoms is determined by using the group number.

• Attach each atom to the central atom with a pair of electrons.

• A bond line can be used to represent each bonding pair.

• The central P atom is bonded to three Cl atoms.
• Twenty electrons are left.

• To complete octets, use the remaining electrons.
• The remaining 20 electrons are placed around the outer Cl atoms and on the P atom, so that all the atoms have octets.

• The Lewis structure can be drawn.

• We have looked at bonding in single bonds in the past.
• atoms share two or three pairs of electrons to complete their octets

• Double and triple bonds form when the number of electrons in the molecule isn't enough to complete the octets.
• One or more lone pairs of electrons from the atoms attached to the central atom are shared with the central atom.
• Multiple bonds are formed by the atoms of carbon, oxygen, nitrogen, and sulfur.

• Double or triple bonds are not formed by the atoms of hydrogen and the halogens.
• Sample Problem 6.13 shows the process of drawing a Lewis structure with multiple bonds.

• The Lewis structure for CO2 has the central atom C.

• The arrangement of atoms can be determined.

• Determine the number of electrons.
• The number of valence electrons for each of the atoms is determined by using the group number.

• Attach each atom to the central atom with a pair of electrons.

• Four electrons are used to attach the central C atom to the two O atoms.

• If needed, use multiple bonds to complete octets.

• Six lone pairs of electrons are placed on the outside O atoms.
• This doesn't complete the octet of the C atom.

• To get an octet, the C atom must share electrons with the O atoms.

• A double bond is a pair of bonds between atoms.

• The Lewis structure has a triple bond.

• There are exceptions to the octet rule.
• There is a hydrogen molecule that only requires two electrons or a single bond.
• The nonmetals form numbers.
• The B atom has three electrons to share.
• Boron compounds have six electrons on the B atoms and form three bonds.
• Although we will usually see compounds of P, S, Cl, and I with octets, they can form a molecule in which they share more of their electrons.
• Their valence electrons can be expanded to 10, 12 or 14 electrons.
• The P atom in PCl3 has an octet, but in PCl5 it has five bonds with 10 electrons.
• In H2S, the S atom has an octet, but in SF6, there are six bonds to sulfur.

• Draw the Lewis structure for each molecule.