1. Ch. 8 Covalent Bonds

2. The Covalent Bond Covalent Bond:
The bond that results from the sharing of valence electrons
Both elements have a strong hold on the valence electron
2 nonmetals bonded together
Molecule: a neutral group of atoms joined together by covalent bonds

3. Diatomic Molecules
7 elements that naturally form 2-atom molecules due to increased stability
Found as pairs unless they are in another compound
Five gases  H2, N2, O2, F2, Cl2
One Liquid  Br2
One Solid  I2
**Mr. BrINClHOF

4. Formation of a covalent bond
Sharing of electrons to achieve an octet
Example: Chlorine …7 valence electrons
Notice: neither atom gives up an electron
By sharing, each atom achieves an octet
Sharing

5. Molecular Compounds
Molecular Compounds: a compound composed of molecules
Example: Water and Carbon Monoxide
Molecular Formula: chemical formula for a molecular compound
Shows how many atoms of each element a molecule contains

6. Formulas of Some Molecular Compounds

7. Types of Bonds Single Bond: Two atoms share one pair of electrons
Bonding Pair: The shared pair of electrons can be represented by a pair of dots or a line
H:H or H–H
Lone pairs: non-bonding pairs of valence electrons

8. Types of Bonds Double Bond: Two atoms share two pairs of electrons
Example: O::O O=O
Triple Bond: Two atoms share 3 pairs of electrons
Largest bond possible
Example: N:::N N≡N

9. Bond Length & Strength Bond Length Bond Strength
Distance between two bonding nuclei
Single bond > double bond > triple bond
Bond Strength
Dissociation Energy: The energy needed to break a bond
Indicates the strength of bonds
Triple bond > double bond > single bond

10. Drawing Dot Structures of Molecules
Draw the dot structure for each individual atom. Determine the number of valence electrons
Determine the number of electrons each atom needs to share to become stable
Choose your central atom (if more than one)
- Typically the least electronegative element

11. Properties of Covalent Bonds
Lower melting and boiling points (compared to ionic compounds)
Most are gases or liquids at room temperature
Composed of two or more nonmetals
Poor conductor of electricity

12. Nomenclature for Molecules
Naming Binary Molecular Compounds:
- Composed of 2 nonmetals
- Use Prefixes
Number or atoms – prefix
mono-
hexa-
di-
hepta-
tri-
octa-
tetra-
nona-
penta-
deca-

13. Naming Guidelines First Element: Second Element:
Use a prefix to indicate the number of atoms
Name the element
Note: Don’t use mono- for the first element
Second Element:
Name the root of the element name
Add the –ide suffix
Example: CO2 - Carbon Dioxide

14. Naming Practice
P2O5
SO2
Dinitrogen tetrahydride

15. Naming Acids
Acid: a substance whose molecules yield H+ when dissolved in water (H is first in the formula)
Names are based on the suffix of the anion bonded to the hydrogen
Anion Ending
Formula Example
Anion Name
Naming System
Acid
-ide
HCl
chloride
Hydro-anion root - ic
Hydrochloric Acid
-ate
H2SO4
sulfate
Anion root - ic
Sulfuric Acid
- ite
HClO2
chlorite
Anion root - ous
Chlorous Acid

16. Acid Naming Practice
Phosphoric Acid
Hypochlorous Acid
HBr
H2CO3
HNO3

17. Acid Naming Practice Phosphoric Acid – H3PO4 Hypochlorous Acid - HClO
HBr – Hydrobromic Acid
H2CO3 – Carbonic Acid
HNO3 – Nitric Acid

18. Molecular Structures
Structural Formula: Uses symbols and bond lines to show the relative position and interactions of the elements in the formula
Steps:
Calculate the total number of available valence electrons
For an anion, add e- equal to the negative charge
For a cation, subtract e- equal to the charge
Identify the central atom
Usually the least electronegative
Usually the first atom in the formula
Never hydrogen

19. 3) Connect the terminal atoms to the central atom(s) with single bonds
4) Complete the octet on the terminal atoms
Hydrogen can only have a maximum of 2 electrons
5) Complete the octet on the central atom with remaining electrons
**For emergency:
Convert one or two lone pairs on the terminal atom to multiple bonds with the central atom if no electrons are left

20. Examples:
CH3Br
- Total # of Valence Electrons: C=4, H=1, H=1, H=1, Br=7
= 14 total valence electrons

21. Practice:
CO2
SO42-
H2CO

22. Practice: O C O O C O O C O O O S O O CO2 SO42-
4+6+6 = 16 total valence electrons
SO42-
(charge) = 32 total valence electrons
O C O X
O C O
O C O
20 valence electrons – too many!
Convert lone pairs to form a double bond
16 valence electrons! O 2-
If the compound has a charge, we put brackets around the molecular structure and indicate the charge on the top right.
O S O O

23. Practice: H C O H H C O H H C O H H2CO
= 12 total valence electrons
H C O H
H C O X H
H C O H
14 valence electrons – too many!
Convert lone pairs to form a double bond
12 valence electron – perfect!

24. Sigma & Pi Bonds
Sigma ( ) bond: the first bond made with any other atom
Pi () Bond: any 2nd or 3rd bond made with any other atom
- 4 bonds
2 sigma bonds
2 pi bonds
- 4 lone pairs
O C O

25. Resonance Structure Resonance Structure:
More than one Lewis structure can be drawn
Atoms remain in the same location
Electrons are moved
Example: Nitrate Ion (NO3-)
The structures vary based on the location of the double bond
All resonance structures are considered equal

26. Octet Rule Exceptions In some cases, the octet rule can be broken
If there is an odd number of valence electrons
Central atom will not have an octet
Examples: ClO2 and NO
Central atom with less than 8 electrons
These compounds tend to be very reactive
B, Al, Be as the central atom
Example: BH3
Central atom with more than 8 electrons
Most common exception
“Expanded octet”
Extra electrons fill in the empty d-sublevel
P, S, Cl, Br, I (3rd row or beyond)
Examples: SF6 and XeF4

27. Molecular Orbits
When two atoms combine, the molecular orbital model assumes that their atomic orbitals overlap to produce molecular orbitals, or orbitals that apply to the entire molecule.
Just as an atomic orbital belongs to a particular atom, a molecular orbital belongs to a molecule as a whole.

28. Molecular Shapes Molecular Shapes…VSEPR
Valence Shell Electron Pair Repulsion Model
Method of predicting molecular shapes
Minimizes repulsion of the electron pairs
According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible.
Predicts bond angles – angles formed by two terminal atoms and the central atom
Lone pairs occupy a slightly large space
Classification is based on a dot structure

30. Molecular Shapes Example: O3 Draw the Lewis (dot) structure
Count the number of electron domains around the central atom
Predict the shape and bond angles of the molecule based on the types and quantities of the domains
3 domains (regions where electrons are found)
2 bonding, 1 nonbonding
Geometric Shape: BENT

31. Molecular Shapes
ClF3
NH3

32. Molecular Shapes ClF3 28 Valence Electrons NH3 8 valence electrons
5 domains (regions where electrons are found)
3 bonding, 2 nonbonding
Geometric Shape: T-Shaped
4 domains (regions where electrons are found)
3 bonding, 1 nonbonding
Geometric Shape: Trigonal Pyramidal

33. Hybridization Hybridization
Process of combining atomic orbitals to make equally shaped bonding orbitals
In hybridization, several atomic orbitals mix to form the same total number of equivalent hybrid orbitals.
This enables the central atom to make identical bonds with terminal atoms
Sigma bonds and lone pairs hybridize
Carbon electron configuration: 1s2 2s2 2p2

34. Electronegativity and Bond Polarity
Electronegativity: the ability to attract electrons in a chemical bond
Trend: increases LR, decreases TB
F is the most electronegative, Fr is the least
Bond Character
The bonding electron pairs are usually unequally shared
The bonding pairs of electrons in a covalent bond are pulled (like tug-a-war) between the nuclei of the atoms sharing the electrons.
Bond Polarity
Describes the sharing of electrons

35. Bond Polarity Nonpolar Covalent Bond: Electrons are shared equally
Example: two identical atoms (Cl2)
Polar Covalent Bond: unequal sharing of electrons
Also known as a dipole: one of the atoms exerts a greater attraction for the bonding electrons than the other
The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge.
Example: HCl

36. Bond Character
Bond Character: based on the electronegativity difference

37. H – Cl H – Cl Polar Bonds Polar Bonds can be represented 2 ways:
#1 : Using () delta (meaning there is a partial charge)
Example: HCl
Electronegativity of Cl: 3.0
Electronegativity of H: 2.1
= Polar Covalent Bond
**The more electronegative atom attracts electrons more strongly and gains a slightly negative charge. The less electronegative atom has a slightly positive charge.
#2 : Using an arrow point to the more electronegative atom +
 -
H – Cl
H – Cl

38. Practice
CH4
CO2
H2CO
HCN

39. Practice: Key CO2 C – H 2.5 – 2.1 = 0. 4 Nonpolar O – C 3.5 – 2.5 = 1
 - +
C – H
2.5 – 2.1 = 0. 4
Nonpolar
O – C
3.5 – 2.5 = 1
Polar

40. Practice: Key HCN C – H N – C 2.5 – 2.1 = 0. 4 3.0 – 2.5 = 0.5
H2CO
 - +
C – H
2.5 – 2.1 = 0. 4
Nonpolar
N – C
3.0 – 2.5 = 0.5
Polar
 - +
C – H
2.5 – 2.1 = 0. 4
Nonpolar
O – C
3.5 – 2.5 = 1
Polar

41. Molecular Polarity All Nonpolar Bonds Polar Bonds If bonds are:
Shape Keeps Polarity
Shape Cancels Polarity
Nonpolar Molecule
Nonpolar Molecule
Polar Molecule

42. Practice: Molecular Polarity
H2O
CS2
CO2

43. Practice: Molecular Polarity
H2O
CS2
O – H
3.5 – 2.1 = 1.4
Polar Bond
Shapes Keeps Polarity –
Polar Molecule
C – S
2.5 – 2.5 = 0
Nonpolar Bond
Nonpolar Molecule

44. Practice: Molecular Polarity
O – C
3.5 – 2.5 = 1
Polar Bond
Nonpolar Molecule
CO2

45. Attractions between Molecules
Intermolecular forces
The attractive forces that causes the interactions between two molecules
Weaker than ionic and covalent bonds
Van der Waals Forces
The weakest attractions between molecules
Dipole Interactions
Occurs when polar molecules are attract on another
The electrical attraction involved occurs between the oppositely charged regions of polar molecules
Dispersion Forces
The weakest of all molecular interactions
When the moving electrons happen to momentarily more on the side of a molecule closest to a neighboring molecule – attracts to the neighboring molecule

46. Forces of Attraction
Hydrogen Bonds: The attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom.

47. Characteristic Ionic Covalent
Bond Formation
Transfer of Electrons
Sharing Electron
Types of Elements
Metal & Nonmetals
Two or more nonmetals
Internal Structure
Crystalline Solid
Geometric Shapes
Melting & Boiling Points
High
Low
Physical State
Solids
Liquids & Gases
Solubility in Water
Usually High
High to Low
Conductivity
Naming techniques
Cation (+) – Element name
Anion (-) – Element root plus –ide suffix
*Unless the anion is a polyatomic ion, then it ends in –ate or –ite
Use PREFIXES based on the number of atoms
* If mono is attached to the first element, drop it off.
Add the suffix –ide to the end of the second element