1. Chapter 8 notes Covalent bonding

2. Covalent Bonding Why do atoms bond?
Atoms held together by sharing electrons
Molecules – a neutral group of atoms joined together by covalent bonds
Diatomic molecules – a molecule consisting of two atoms
Some elements exist in nature as diatoms
H,N, O, F, Cl, Br, I
Healthy Nerves Originate From Clear Brown Iodine
H O F Br I N Cl
Lucky 7
Molecular compound – compound composed of molecules
Lower melting and boiling points than ionic compounds

3. Covalent Bonding
Molecular formula - the chemical formula of a molecular compound (how many atoms of each element a molecule contains)
Tells you nothing about the structure of the molecule
Structural formula
Space-filling molecular formula
Perspective drawing
Ball-and-stick molecular model

4. Structural Formulas
The shared pair of electrons count as valence electrons for both atoms, so you can see that each atom has a filled outer level.

5. Covalent Bonding
Single covalent bond – 1 pair of electrons are being shared between two atoms or 2 shared electrons
Example: Methane (CH4)
Double covalent bond – 2 pairs of electrons are being shared between two atoms
Example: Ethene (C2H4) or Carbon dioxide
Triple covalent bond – 3 pairs of electrons are being shared between two atoms
Example: N2
dkreutz.basd.k12.wi.us

6. Covalent bonding # of electrons needed = # of electrons shared
Covalent bond – bond in which electrons are shared
Atoms overlap their valence orbitals  therefore leading to the “sharing” of valence electrons between the two atoms.
Key idea: How many electrons are shared between atoms?
# of electrons needed = # of electrons shared

7. Something to think about…
You’ve been told that molecular compounds generally have lower melting and boiling points than ionic compounds. What do you think causes this difference?
We will return to this question for the next few days to answer it completely.

8. Drawing Lewis Dot Structures for Compounds
Example: Water molecule, H2O

9. Model Building Activity
To learn how atoms covalently bond with each other, we will complete a model building lab with model kits.
Different elements are represented by different colored atoms.
Each atom has a different number of “holes” in it. The holes equal the number of bonds that atom makes (there are some exceptions so use the rule we learned about the # of bonds an atom makes.
Use the short springs for the bonds, unless you need to have 2 bonds going between 2 atoms, then use the longer, more flexible springs.
Identify the shape and molecular polarity.
Let’s do HBr together.

10. Molecular Shapes

11. Bond Polarity
Individual bond polarities can result in partial charges on the molecule!
These two dipoles do not cancel each other out, so the molecule has a molecular dipole (or molecular polarity).
The overall molecule is neutral, but each “side” of the molecule has a partial charge.
Big Idea: Molecular polarity is about determining if individual bond polarities cancel out. The geometry of the molecule, and the relative strengths of the bond dipoles, determine if they cancel.

12. Bond Polarity Dipole – a charge imbalance between two bonded atoms;
a molecule that has two poles also called a dipolar molecule
Polar molecule placed between two oppositely charged plates, usually align with the plates (+ to – and – to +)

13. Shapes and polarity Pyramidal molecules – dipoles never cancel
Bent molecules – dipoles never cancel
Dipoles in trigonal planar and tetrahedral molecules, do cancel out.
No net dipole

14. Bond Angles Linear – 180° Trigonal planar - 120° Tetrahedral -109.5°
Pyramidal - 107°
Bent -105°

15. Intermolecular Attraction Forces
Weaker than ionic and covalent bonds
Why is water a liquid at room temperature, but CO2 is a gas?
Why is I2 a solid, but F2 a gas?
The strength of the attractions between the molecules, determines the state!
Inter = between
Van der Waals forces – weakest attractions
Dipole Interactions
Dispersion Forces

16. Van der Waals Forces
Dispersion: caused by random motions of electrons, leading to "induced dipoles."
Occur between nonpolar
substances.
The larger the molecule, the stronger they are!
Weakest of the forces

17. Van der Waals Forces
Dipole interactions: attractions between polar molecules,
due to partial charges.
Very common, since polar molecules are very common.
Similar to ionic bonds but weaker.

18. Hydrogen Bonds
Hydrogen bond - a special type of Dipole interaction. Very strong.
Occurs between H bonded to a very Χ atom on one molecule (N, O, F), and another very Χ atom on another molecule.

19. Hydrogen Bonding Has about 5% the strength of an average covalent bond
Strongest of the intermolecular forces

20. Check for Understanding
Place the intermolecular attractions in order of increasing strength.
Dispersion forces < dipole interactions< hydrogen bonding

21. Covalent Network Solids
Covalent Network Solids - a covalently-bonded crystal. Very strong! (all atoms are covalently bonded to each other.
Carbon (diamond allotrope) SiO2 (quartz)
Diamond vaporizes at 3500 degrees C (does not melt)
Quartz melts at 2700 degrees C

22. Ionic Forces in a Crystal
Attractions between formula units in an ionic crystal are very strong! This explains why ionic compounds are solids at room temperature.
NaCl – table salt

23. Bond Dissociation Energy
The energy required to break the bond between two covalently bonded atoms
Large BDE = a strong covalent bond.
Expressed as the energy needed to break 1 mole of bonds (6.02 x 1023 bonds)
Units - kJ/mol

24. Characteristics of Ionic and Covalent Compounds
Ionic Compound
Covalent Compound
Representative Unit
Formula unit
Molecule
Bond formation
Transfer of one or more electrons between atoms
Sharing of electron pairs between atoms
Types of elements
Metallic and nonmetallic
Nonmetallic
Physical state
Solid
Solid, liquid, or gas
Melting point
High (usually above 300˚C)
Low (usually below 300˚C)
Solubility in water
Usually high
High to low
Electrical conductivity of aqueous solution
Good conductor
Poor to nonconducting

25. Coordinate Covalent Bonds
A bond in which one atom contributes both bonding electrons.  
The atom needs a lone pair of electrons to do this.
Examples
CO (triple bond)
NH4 N, H, H, H, and H+

26. Polyatomic Ions
A tightly bound group of atoms that has a positive or negative charge and behaves as one unit
Most polyatomic have both covalent and coordinate covalent bonds

27. Resonance
Resonance structure – a structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion
Use double headed arrows to connect resonance structures
Generally you find that the bond strength is an average.
There is no back and forth between the resonance structures.

28. Comparing Bonds
As the number of bonds increases, bond length decreases
As the number of bonds increases, bond strength increases
Comparing bde in increasing order
Single bonds < double bonds < triple bonds
As bde increases, chemical reactivity decreases

29. Exceptions to the Octet Rule
The octet rule will not be satisfied in molecules whose total number of valence electrons is an odd number.
Some molecules will have fewer or more than a complete octet
Example NO2 – two possible structures
Remember the number of bonds formed will equal the number of electrons needed
There is an unpaired electron on nitrogen no matter how the structure is drawn
Some with even numbers will also be exceptions to the octet rule.
Example BF3
Expanded octets
PCl5 and SF6

30. Bonding Theories VSEPR Valence-Shell Electron-Pair Repulsion Theory
The repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible.
Bond angles
Tetrahedral angle– 109.5º
Bent - 105º
Linear - 180º

31. Molecular Orbitals Quantum mechanic model
Produced when two atoms combine their atomic orbitals overlap to produce molecular orbitals
Bonding orbital
A molecular orbital that can be occupied by two electrons of a covalent bond
Sigma bonds
Two atomic orbitals combine – the resulting molecular structure is symmetrical around the bond axis

32. Sigma Bonds

33. Pi Bonds Side by side overlapping
Weaker than sigma bonds because there is less overlapping

34. Hybridization
Occurs when several atomic orbitals mix to form the same number of hybrid orbitals
Gives information about molecular bonding and molecular shape