Lewis Dot Structures Chapter 9 Lewis Dot Structures

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Lewis Dot Structures Chapter 9

Lewis Dot Structures Chapter 9

Lewis Dot Structures of Ions • ***Remember Lewis Dot structures show the # of

Lewis Dot Structures of Ions • ***Remember Lewis Dot structures show the # of VALENCE ELECTRONS. • FOR IONS 1. Write the symbol of the Ion 2. Put correct number of valence electrons around Ion (cations will have 0, anions will have 8) 3. Put brackets around lewis structure 4. Put charge of ion as superscript outside brackets

Rules for writing ionic compound Lewis structures 1. Determine that a metal is present

Rules for writing ionic compound Lewis structures 1. Determine that a metal is present / first element on left of periodic table(Not counting H). 2. Write the Lewis structures for the elements in the ionic compound. (as atoms) 3. Draw an arrow showing the transfer of the electrons from the metal (low electronegativity element) to the non-metal (high electronegativity. ) 4. Draw a chemical equation arrow. 5. On right side of arrow show all of the Ion Lewis Dot Structures

Draw Lewis Dot structure for Na. Cl

Draw Lewis Dot structure for Na. Cl

Examples of Ionic Compound Lewis Dot Diagrams

Examples of Ionic Compound Lewis Dot Diagrams

Ionic Compounds

Ionic Compounds

Rules for Writing Lewis Dot Structures of Covalent Compounds Remember Covalent Compounds: • Have

Rules for Writing Lewis Dot Structures of Covalent Compounds Remember Covalent Compounds: • Have a difference in electronegativity of less than 2. 0 • Usually composed of two non-metals

Covalent Molecule Lewis Dot Structure Rules • Rule 1: Count the total # valence

Covalent Molecule Lewis Dot Structure Rules • Rule 1: Count the total # valence electrons for entire molecule (v. e. inventory) • Rule 2: Place least electronegative element in center of other atoms (H will always be outer atom) • Rule 3: Add single covalent bond between the centre atom & the outer atoms (subtract 2 v. e. from inventory per bond) • Rule 4: Add lone pairs to the outer atoms & centre atom (subtract from v. e. inventory) so that atoms meet octet rule (H meet duet rule) • Rule 5: If all v. e. from intial inventory (rule 1) are displayed as either a bond or lone pair and all atoms are meeting octet rule (H is meeting duet), then you are done. If not, go to rule 6 • Rule 6: Change 1 single bond to double bond, repeat Rules 3 -5

Covalent Bond When nonmetallic elements react with other nonmetallic elements, they share electrons in

Covalent Bond When nonmetallic elements react with other nonmetallic elements, they share electrons in order to obtain eight valence electrons. Each fluorine atom has seven valence electrons. They each require one more electron to satisfy the Octet Rule.

The left fluorine atom now has a total of eight electrons and the right

The left fluorine atom now has a total of eight electrons and the right fluorine atom now has a total of eight electrons around it. When nonmetallic elements react with other nonmetallic elements, they share electrons in order to obtain eight valence electrons.

The two electrons that form the covalent bond are often Represented by a single

The two electrons that form the covalent bond are often Represented by a single line. The F 2 molecule can be represented using a line and dots to show the bonding pair and the six lone pairs, respectively. This is called a Lewis dot structure.

Multiple Covalent Bond Some atoms have to share more than one electron in order

Multiple Covalent Bond Some atoms have to share more than one electron in order to satisfy the Octet Rule.

Each oxygen atom has six valence electrons. They each require two more electrons to

Each oxygen atom has six valence electrons. They each require two more electrons to satisfy the Octet Rule.

 • The left oxygen atom now has a total of eight electrons around

• The left oxygen atom now has a total of eight electrons around it. The right oxygen atom now has a total of eight electrons around it.

The four electrons shared by the oxygen atoms form a double bond. The double

The four electrons shared by the oxygen atoms form a double bond. The double bond is represented by two single lines. Each line in the Lewis dot structure represents two electrons

The element hydrogen is an exception to the Octet Rule. It only needs two

The element hydrogen is an exception to the Octet Rule. It only needs two electrons, rather than eight, to be stable. The hydrogen atom has one valence electron. It requires one more electron to be stable. The fluorine atom has seven valence electrons. It requires one more to satisfy the Octet rule.

The hydrogen atom now has a total of two electrons around it and is

The hydrogen atom now has a total of two electrons around it and is stable. The fluorine atom now has a total of eight electrons around it and is stable.

 • The Lewis dot structure of the HF molecule shows a line and

• The Lewis dot structure of the HF molecule shows a line and 6 dots to represent the bonding pair and the 3 lone pairs of electrons, respectively.

Rules for writing Lewis Dot structures of Covalent Compounds • Rule 1 Add together

Rules for writing Lewis Dot structures of Covalent Compounds • Rule 1 Add together the number of valence electrons for each atom in the molecule. For example, CF 4 Carbon has four valence electrons and each fluorine has seven valence electrons = 4 + 4(7) = 32

Rule 2 Write out the elements of the molecule so that the least electronegative

Rule 2 Write out the elements of the molecule so that the least electronegative elements is in the center surrounded by the other elements. For example, CF 4

Rule 3 Place a covalent bond between the central atom and the outside atoms.

Rule 3 Place a covalent bond between the central atom and the outside atoms. Remember each covalent bond contains two electrons.

The four covalent bonds use eight of the 32 valence electrons in CF 4

The four covalent bonds use eight of the 32 valence electrons in CF 4 • This uses 24 electrons. There Are no electrons left, so this is The Lewis dot structure for CF 4 • Rule 4 There are 24 valence electrons remaining. Add electrons to the outer atoms as lose pairs to satisfy the Octet Rule.

Rule 1 Nitrogen has 5 valence electrons and each hydrogen has 1 valence electron

Rule 1 Nitrogen has 5 valence electrons and each hydrogen has 1 valence electron The total number of valence electrons = 5 + 3 (1) = 8 Rule 2 Hydrogen is always an outer atom and is never at the centre of a molecule

Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons.

Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

 • Rule 3 Add the bonding electrons. This uses 6 of the 8

• Rule 3 Add the bonding electrons. This uses 6 of the 8 valence electrons. • Rule 4 The 2 remaining valence electrons are not added to the outer atoms, because each H has its maximum of 2 valence electrons.

Rule 5 This is the Lewis structure For NH 3 Place the remaining 2

Rule 5 This is the Lewis structure For NH 3 Place the remaining 2 Valence electrons on the central nitrogen atom Rule 6 Check all atoms in the molecule to ensure that each has 8 electrons(2 for hydrogen). If an atom has fewer than 8 electrons, create double or triple bonds. (Note: Double bonds only exist between C, N, O and S atoms)

Apply rule 6 to the following; CH 4, CF 4, • Hydrogen : 1

Apply rule 6 to the following; CH 4, CF 4, • Hydrogen : 1 bond = 2 electrons (stable) • Carbon : 4 bonds = 8 electrons (stable) • Fluorine : 1 bond + 3 lone pairs = 2 + 3 (2) = 8 electrons (stable) • Carbon : 4 bonds = 8 electrons (stable)

Example; CH 2 O Apply Rules 1 -5 to the molecule Rule 1: Count

Example; CH 2 O Apply Rules 1 -5 to the molecule Rule 1: Count the valency electrons Rule 2: Place the least electronegative element at the centre, except for H, which is always an outer atom Rule 3: Add covalent bonds between the centre and the outer atoms Rule 4: Add lone pairs to the outer atoms Rule 5: Add lone pairs to the centre atom

Rule 1 Carbon has 4 valence electrons, each hydrogen has 1 valence electron, and

Rule 1 Carbon has 4 valence electrons, each hydrogen has 1 valence electron, and oxygen has 6 valence electrons. Total number of valence electrons : 4 + 2(1) + 6 = 12 Rule 2 Carbon is at the centre of the molecule because it is less electronegative than oxygen. Hydrogen is always an outer atom and is never at the centre of the molecule.

 • Rule 3 Add the bonding electrons. This uses 6 of the 12

• Rule 3 Add the bonding electrons. This uses 6 of the 12 valence electrons • Rule 4 Add the remaining 6 lectrons to the outer atom. Hydrogen does not need any more electrons, but Oxygen needs 6 to complete its octet.

Rule 5 There are no valence electrons left to add to the centre •

Rule 5 There are no valence electrons left to add to the centre • Rule 6 Oxygen shares one of its lone pairs with C and O and give the desired 8 electron total If an atom lacks an octet, use electron pairs on an adjacent atom to form a double or triple bond. This is the Lewis dot Structure for CH 2 O

Bond Number • There are 3 kinds of covalent bonds based on number of

Bond Number • There are 3 kinds of covalent bonds based on number of pairs of electrons shared between the two atoms: 1. single covalent bond - 1 shared pair e 2. double covalent bond - 2 shared pairs e 3. triple covalent bond - 3 shared pairs e • Bond energy: single < double < triple Bond length: single < double < triple

Exceptions to the Octet Rule The Octet Rule applies to Groups IVA through VIIA

Exceptions to the Octet Rule The Octet Rule applies to Groups IVA through VIIA in the second row of the Periodic Table, but there are exceptions to the rule among some other elements. The following two cases are an example Example BF 3 Rule 1 Boron has 3 valence electrons and each Fluorine has 7 valence electrons Total number of electrons = 3 + 3 (7) = 24

Rule 2 Boron is at the centre of the molecule because it is less

Rule 2 Boron is at the centre of the molecule because it is less electronegative than fluorine Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons

Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine has the

Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine has the required 8 electrons Rule 5 This uses the remaining electrons leaving none to add to the Boron central atom

Rule 6 Check the number of electrons around each atom. Each Fluorine atom has

Rule 6 Check the number of electrons around each atom. Each Fluorine atom has 8 electrons, but the Boron Atom has only 6. This is an exception to the Octet Rule. A B=F bond is not an option, because double bonds exist only between C, N, O, and S atoms This is the Lewis dot structure BF 3

Example PF 5 Rule 1 Phosphorus has 5 valence electrons and each fluorine has

Example PF 5 Rule 1 Phosphorus has 5 valence electrons and each fluorine has 7 valence electrons Rule 2 Phosporus is at the centre because it is less electronegative than fluorine Total number of electrons • = 5 + 5(7) = 40

Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons.

Rule 3 Add the bonding electrons. This uses 6 of the 24 valence electrons. Rule 4 Add the remaining electrons to the outer atoms. Each Fluorine requires 6 more electrons

Rule 5 This uses the remaining electrons leaving none to the central P atom

Rule 5 This uses the remaining electrons leaving none to the central P atom Rule 6 Check the number of electrons around each atom. Each Fluorine atom has 8 electrons, but the phoshorus atom has 10. This is an exception to the Octet Rule.

Rule 6 Check the number of electrons around each atom. Each fluorine atom has

Rule 6 Check the number of electrons around each atom. Each fluorine atom has 8 electrons, but the phoshorus atom has 10. This is an exception to the Octet Rule.