chapter 3-chemical bonds
TRANSCRIPT
Chemical Bonds
Chapter 3
Lewis Model of Bonding In 1916, Gilbert N. Lewis pointed out that the
lack of chemical reactivity of the noble gases indicates a high degree of stability of their electron configurations.
He 1s2
Ne
Ar
Kr
Xe
[He]2s2
2p6
[Ne]3s2 3p
6
[Ar]4s2 4p
6
[Kr]5s2 5p
6
Noblegas
Noble gasnotation
2 Chapter 3- Chemical Bonds
The Octet Rule
Octet rule: The tendency of group 1A-7A elements to react in ways that achieve an electron configuration of eight valence electrons.
– An atom that loses one or more electrons becomes a positively charged ion called an cation.
– An atom that gains one or more electrons becomes a negatively charged ion called a anion.
3 Chapter 3- Chemical Bonds
Na Chlorine atom (17 electrons): 1s2 2s2 2p6 3s2 3p5
The Octet Rule- Ion formation
Cl
Sodium atom (11 electrons): 1s2 2s2 2p6 3s1
Nonmetals Gain electrons to become anions
Metals Lose electrons to become cations
4 Chapter 3- Chemical Bonds
Ionic Bond Formation
5 Chapter 3- Chemical Bonds
+ 2+
3+
Charges of Common Ions
-
2-
2+
+
Label your Periodic Table with these ionic charges!
6 Chapter 3- Chemical Bonds
Naming Cations
+ 2+
3+ -
2-
2+
+
7 Chapter 3- Chemical Bonds
• Elements of Groups 1A, 2A, and 3A form only one type of cation.
• The name of the cation is the name of the metal followed by the word “ion”.
• Most other metals require a roman numeral to indicate charge
Chapter 3- Chemical Bonds 8
Do Not Need to Know Common Names
Naming Cations- Examples Ca2+______________ Na+______________
Al3+_________________________ Potassium ions __________
Magnesium ions ________________ Cesium ions _____________
Fe3+ ________________ Co2+_______________
Mn4+ ___________________ Gold (I) ions ______________
9 Chapter 3- Chemical Bonds
Naming Anions- Monoatomic
10 Chapter 3- Chemical Bonds
add “ide” to the stem name.
Anio n
F-
Cl-
Br-
I-
O2-
S2-
Ste m
name
fluor
chlor
bro m
iod
o x
sulf
A nion
name
fluoride
chloride
bro mide
iodide
oxide
sulfide
Naming Polyatomic Ions
NH4+
OH-
NO2-
NO3-
CH3 COO-
CN-
MnO4-
CO32 -
HCO3-
SO32-
HSO3-
SO42-
PO43 -
HPO42 -
H2PO4-
HSO4-
CrO42 -
Ammonium
Hydroxide
Nitrite
Nitrate
Acetate
Cyanide
Permanganate
Carbonate
Hydrogen carbonate (Bicarbonate)
IonName
Sulfite
Hydrogen sulfite (Bisulfite)
Sulfate
Phosphate
Hydrogen phosphate
Dihydrogen phosphate
Name
Hydrogen sulfate (Bisulfate)
Chromate
Ion
These must be memorized!!! 11 Chapter 3- Chemical Bonds
Predicting formulas of Binary Ionic Compounds
• Metal (+)is always first, nonmetal (-) last
• Must have charge balance- Ex. Aluminum chloride
Chapter 3- Chemical Bonds 12
Al3+ 3Cl--
-
-
-
+
+
+ AlCl3
Try: Sodium Oxide Aluminum Sulfide
• Positive named first – May need a roman
numeral (you will have to figure out what the roman numeral is!!)
• Negative named second – Ending of
monoatomic anions is “-ide”
+ 2+
3+
- 2-
2+
+
13 Chapter 3- Chemical Bonds
CaBr2 Ba3N2
ZnS AgCl
Cd3As2 FeI3
FeF2 Au2Se
Naming Binary Ionic Compounds
• Positive named first – May need a roman
numeral (you will have to figure out what the roman numeral is!!)
– OR Ammonium (NH4
+ )
• Negative named second – Polyatomic ions
from table 3.4
+ 2+
3+
- 2-
2+
+
14 Chapter 3- Chemical Bonds
Naming Compounds Containing Polyatomic Ions
NaNO3 NaH2PO4 NH4OH Fe2(CO3)3 CuSO4
More Examples
Chapter 3- Chemical Bonds 15
NH4I CaSO4
LiCH3COO Na2SO3
K3PO4 Mg(CN)2
Zn3(PO4)2 (NH4)2SO4
Cu(OH)2 Fe2(CO3)3
Lithium acetate Sodium sulfite
Potassium phosphate Cobalt (II) nitrite
Vanadium (IV) chlorate Aluminum chlorite
Manganese (II) acetate Ammonium perchlorate
Ammonium dichromate Lithium permanganate
Chapter 3- Chemical Bonds 16
Ionic Compounds
• Between ions
• Cation Named (Roman Numeral)
• Anion Named (“-ide” or polyatomic
• Charge Balance
Covalent Compounds
• Electrons are shared between atoms
• Pair of shared electrons is a covalent bond
• Prefix system used
Forming a Covalent Bond
A covalent bond is formed by sharing one or more pairs of electrons.
– The pair of electrons is shared by both atoms and, at the same time, fills the valence shell of each atom.
– Example: in forming H2
+H H
the s ingle line represents a shared pair of electrons
.. H H
17 Chapter 3- Chemical Bonds
Polarity of Covalent Bonds
Although all covalent bonds involve sharing of electron pairs, they differ in the equality of the sharing: – Nonpolar covalent bond: Electrons are shared equally. – Polar covalent bond: Electron sharing is not equal. – The equality of the sharing depends on the relative
electronegativities of the bonded atoms. Table 3.6 Classification of Chemical Bonds
Type of Bond
Less than 0.5
0.5 to 1.9
Greater than 1.9
Nonpolar covalent
Polar covalent
Ionic
Two nonmetals or a
nonmetal and a metalloid
Electronegativity
Difference Between
Bonded Atoms
A metal and a nonmetal
Most Likely to
Form Between
19 Chapter 3- Chemical Bonds
Polarity of Covalent Bonds
Examples:
H-Cl
BondDifference in Electronegativity Type of Bond
3.5 - 2.1 = 1.4
3.0 - 2.1 = 0.9
4.0 - 0.9 = 3.1
2.5 - 1.2 = 1.3
polar covalent
polar covalent
ionic
polar covalent
2.5 - 2.5 = 0.0 nonpolar covalent
3.0 - 2.1 = 0.9 polar covalent
O-H
N-H
Na-F
C-Mg
C-S
20 Chapter 3- Chemical Bonds
Polar Covalent Bonds
Chapter 3- Chemical Bonds 21
Polar Covalent Bonds
Chapter 3- Chemical Bonds 22
Using delta notation, label each atom in the
following polar covalent bonds
C-O
H-F
N-O
Non-polar Covalent Bonds
–Elements with the same electronegativity value share
the electrons in a covalent bond equally. If the electrons
are shared equally, the bond is considered nonpolar
–Would the covalent bond that occurs in the diatomic
molecules be considered polar or nonpolar?
Chapter 3- Chemical Bonds 23
Interpreting Lewis Structures
Table 3.7 Lewis Structures for Several Small Molecules. (The number of valence electrons is given in parentheses after the molecular formula.)
Carbonic acidFormaldehydeAcetyleneEthylene
Hydrogen chlorideMethaneAmmoniaWater
H
H N H C H H ClH
H
C C
H
C C HH
H
C
HH
O
H
H2O (8) NH3 (8) CH4 (8) HCl (8)
C2H4 (12) C2H2 (10) CH2O (12) H2CO3 (24)
H
HH
OH
O OC HH
O
24 Chapter 3- Chemical Bonds
• Determine the total number of valence electrons present in the molecule – Valence electrons for an atom = position with in a row (i.e. Na = 1, C = 4, He = 2)
– If the species is an ion, add electrons if the charge is negative, subtract electrons if the charge is positive
• Write the symbols of the atoms arranged according to what is bonded to what – If only 2 elements are present, the element of which there is only one atom is central (example PF3)
– If more than 2 elements are present, bonding occurs in the same order as the formula is written (example: HCN)
– Hydrogen and fluorine are never central. (these elements can only form one single bond)
– Carbon is often central (it commonly makes 4 bonds)
• Place single bonds (2 electrons) between each bonding pair of atoms
• Complete octets for all atoms EXCEPT the central atom
• Compare number of electrons drawn in structure to total number of valence electrons calculated in step one. – If there are any remaining valence electrons, place them on the central atom
– use pairs when possible-sometimes this will result in MORE than an 8 electrons for the central atom
• If there is LESS than an octet on the central atom, form multiple bonds by shifting nonbonding pairs from outer atoms. – Don’t throw in double bonds prematurely! This is the final step after every other condition has been met!
Chapter 3- Chemical Bonds 25
How To Draw An Electron Dot Picture (Lewis Structure)
Simple Lewis structure examples
CCl4
H2O
HCN
NO3-
Chapter 3- Chemical Bonds 26
Lewis Structures Practice problems:
– Draw a Lewis structure for hydrogen peroxide, H2O2.
– Draw a Lewis structure for methanol, CH3OH.
– Draw a Lewis structure for acetic acid, CH3COOH.
27 Chapter 3- Chemical Bonds
Exceptions to the Octet Rule
PCl5 BH3
XeF4 I3
-
Look at OWL for excellent guided learning. Lesson 3.7 (a-k)
28 Chapter 3- Chemical Bonds
Expanded Valence of Sulfur and Phosphorus
Chapter 3- Chemical Bonds 29
Recognize as
Expanded valence
Resonance
Linus Pauling - 1930s – Many molecules and ions are best described by
writing two or more Lewis structures. These molecules or ions are said to exhibit resonance.
– Individual Lewis structures are called contributing structures.
– Double-headed (resonance) arrows are placed between individual contributing structures.
– The molecule or ion is a hybrid of the various contributing structures.
30 Chapter 3- Chemical Bonds
Resonance Examples
NO2-
CH3 COO-
Chapter 3- Chemical Bonds 31
VSEPR Model
Valence-Shell Electron-Pair Repulsion (VSEPR) – Valence electrons of an atom may be involved in
forming single, double, or triple bonds or they may be unshared.
– Each arrangement of electrons creates a negatively charged region of electron density around a nucleus.
– Because like charges repel each other, the various regions of electron density around an atom spread so that each is as far away as possible from the others.
32 Chapter 3- Chemical Bonds
Chapter 3- Chemical Bonds 33
Illustrations of VSEPR
Four regions of electron density
109.5°
Tetrahedral
Tetrahedral
Chapter 3- Chemical Bonds 34
Illustrations of VSEPR
Four regions of electron density
109.5°
Tetrahedral
Trigonal Pyramidal
Chapter 3- Chemical Bonds 35
Illustrations of VSEPR
Four regions of electron density
109.5°
Tetrahedral
BENT
Chapter 3- Chemical Bonds 36
Illustrations of VSEPR
(Trigonal)
Planar 120°
Illustrations of VSEPR
Chapter 3- Chemical Bonds 37
Regions of electrons density surrounding central atom
Polarity
Chapter 3- Chemical Bonds 38
Remember Bond Polarity
POLAR Molecules
POLAR: Molecule has opposite poles (One + and one -)
LOOK FOR
• Polar Bonds
• Unshared e- pairs on central atom
• Unsymmetrical molecules
• Dipole Moments ADD
Chapter 3- Chemical Bonds 39
Polar Molecules
p. 96 40 Chapter 3- Chemical Bonds
Polar Molecules
Chapter 3- Chemical Bonds 41
NON-POLAR Molecules
Chapter 3- Chemical Bonds 42
Non-polar: (No opposite poles.) •Same on all corners •Di-poles cancel out
Polar or Non-Polar?
• H2S
• HCN
• C2H6
Chapter 3- Chemical Bonds 43
• VSEPR: An acronym that stands for V______ S______ E_______ P____ R_______ • Electron pair geometry indicates the arrangement of the electron pairs around the central atom • Molecular geometry (or molecular shape) indicates the arrangement of atoms around the central atom • The bond angle is the angle formed by any two atoms bonded to the central atom
Chapter 3- Chemical Bonds 44
Formula Lewis Structure Electron Pairs e-pair geometry Bond angle Molecular geometry
Polarity
Total Bonding pairs Non-bonding pairs
CO2
BCl3
O3
Chapter 3- Chemical Bonds 45
Formula Lewis Structure Electron Pairs e-pair geometry Bond angle Molecular geometry
Polarity
Total Bonding pairs Non-bonding
pairs
CH4
NH3
H2O