Unit 2: Bonding & Lewis Structures

Question 1

What is a Chemical bond ?

Answer 1

the attraction between positive and negative ions, or two nonmetal atoms.

Question 2

Ionic Bonds

Answer 2

are due to the electrostatic attractions between oppositely charged ions.

• solids with mp > 400C;
• soluble in polar solvent;
• conductive,
• high ΔEn

Question 3

Covalent Bonds

Answer 3

molecules are formed by the sharing of electrons between atoms

• covalent – gases, liquids, solids rt, mp usually <300 C;
• many insoluble in polar solvent,
• nonconductive,
• low ΔEn diff

Question 4

Lewis Formula

Answer 4

Chemical symbol of element surrounded by dots representing valence electrons

Lewis symbol: inner electrons/nucleus as symbol; valence as dots

Primary contribution = valence electrons

G.N. Lewis – express valence electrons

Question 5

Octet Rule

Answer 5

Atoms often gain, lose or share electrons to achieve the same number of electrons as the noble gas closest to them in the periodic table.

Noble gases have very stable electron arrangements, because all noble gases except He have eight valence electrons, many atoms undergoing reactions end up with eight valence electrons. This observation led to octet rule.

Octet Rule: Atoms tend to gain, lose or share electrons until they are surrounded by 8 valence electrons

Question 6

Ionic Bonding

Answer 6

Ionic substances generally result from the interaction of metals on the left side of the periodic table with nonmetals on the right side

For Example:

One element oxidized, another reduced – electron transfers

Ex: NaCl

2Na (s) + Cl2 (g) à 2 NaCl (s)

Question 7

Ionic Bonding (continue example)

Answer 7

Question 8

Ionic Bonding

Answer 8

Question 9

Ionic Bonding

Answer 9

Question 10

Electron configuration for Zn2+ ion:

Electron configuration for Co2+ and Co3+ ion:

Question 11

Covalent Bonding

Answer 11

A chemical bond formed by sharing a pair of electrons is a covalent bond.

Represented by a line between atoms in a Lewis structure: X-X

Ionic compound – strong intermolecular, weak intramolecular

Covalent – weak intermolecular, strong intramolecular

Covalent bond formation follows energy curve w/ minimum distance at bond distance

Question 12

Covalent Bonding

Answer 12

Question 13

Covalent Bonds & Lewis Structures

Answer 13

if a polyatomic ion is formed, the charge is included as in ionic compounds

NH4+

Question 14

Writing Lewis Formulas

Answer 14

• Obey octet rule
• Differentiate bonding and nonbonding electrons

Determine bonding electrons:

S = N - A

S = shared

N = needed

A = available

• For F₂ as an example
  
  • A = 2 F = 2x7 available;
  • N = 2 F = 2x 8 needed
  • S = 16 - 14 = 2 shared
  • n2 electrons = 1 bond

Question 15

Guide to Dot Structures

Answer 15

Question 16

What is a Oxidation State (number)

Answer 16

a means devised for keeping track of electrons being gained or lost by a substance.

Track electronic charges

Single atom ions = actual charge

Polyatomics spread over ion; assign # to each atom

General rules but no absolutes

Question 17

Oxidation Number Rules

(1-4)

Answer 17

1. The oxidation number of an element in its free (uncombined) state is zero
2. The oxidation number of a monatomic (one-atom) ion is the same as the charge on the ion.
3. The sum of all oxidation numbers in a neutral compound is zero.
4. The sum of all oxidation numbers in a polyatomic (many-atom) ion is equal to the charge on the ion.

Question 18

Oxidation Number Rules

(5-7)

Answer 18

5. The oxidation number of fluorine is always –1.
6. The oxidation number of oxygen in a compound is usually –2, except in peroxides (- 1), superoxides (- ½), or combined with fluorine (+1 or +2)
7. The oxidation state of hydrogen in a compound is usually +1, except as a binary compound with a metal (-1)

Question 19

Oxidation Number Rule 8

Answer 19

8. Periodic position contributes to oxidation state:
   a. Group IA = 1+
   b. Group IIA = 2+
   c. Group IIIA = 3+ (some exceptions)
   d. Group VA = 3- w/ metal, H, NH4+; apply 3 & 4 to element to the right of group V
   e. Group VIA = 2- w/ metal, H, NH4+; apply 3 & 4 to element to O or halogen in lower n
   f. Group VIIA = 1- in binary compounds; apply 3 & 4 when combined w/ O or halogen in lower n

If En are equal, negative to element closest to F

Question 20

Determine ox state of N in N2O4, NH3, HNO3, NO3-

Question 21

Resonance

Answer 21

Resonance – situations where a single Lewis structure does not present the actual configuration for a molecule, and it is capable of representation by multiple valid Lewis structures.

Resonance structure – different electron arrangement

Resonance hybrid – real molecule

Ex: CO32-

Question 22

Formal Charges

Rules for assigning formal charge for main group elements:

Answer 22

Formal Charge – the charge an atom would have if each bonding electron pair in the molecule were shared equally between its two atoms.

Rules for assigning formal charge for main group elements

1. FC on an atom is = group number – [number bonds + number unshared electrons].
2. If # bonds = group number, F.C. = 0
3. a. in molecule, sum charge = 0
   b. in polyatomic, sum charge = charge on ion

Question 23

Keep in mind for Formal Charges

Answer 23

Keep in mind for formal charges:

a. most likely formula charges minimize charge – closest to 0
b. put negative charges mostly on more electronegative atom
c. If adjacent atoms have same sign (+ or -) usually incorrect

EX: NOCl

Question 24

Octet Exceptions

Answer 24

Some instances octet rule fails: must modify S = N - A formula for these:

A. covalent compounds of Be – use 4 for N

B. group 3A elements – use N = 6

C. Some compounds w/ unpaired electrons NO and NO2

D. Expanded Valence – > 8 electrons; usually S and P

Question 25

Octet Exceptions

Expanded Valence

Answer 25

Question 26

Write Lewis symbols for Na, F, Si, Ar, S

Write Lewis structures for NaF, SiF4, BeCl2, BCl3, PF5,

Write resonance structures for NO3-