1. Unit 11: States of Matter
Ch. 13, Sections 2-4

2. Types of Covalent Bonds
Polar Covalent Bond
e- are shared unequally
asymmetrical e- density
results in partial charges (dipole) + -

3. Types of Covalent Bonds
Nonpolar Covalent Bond
e- are shared equally
symmetrical e- density
usually identical atoms

4. Take the absolute value of the difference!
If ΔEN is:
Bond type is:
< 0.4
Nonpolar covalent
0.4 < Δ EN < 1.7
Polar covalent
> 1.7
Ionic
Given: Electronegativities of these elements
H = 2.2 C = 2.55 N = 3.04 O = 3.44 F = 3.98
Na = 0.93 K = 0.82 P = 2.19 S = 2.58 Cl = 3.16
Determine bond type for the following bonds:
H – H _____________________ H – O ________________________
H – C _____________________ Na – Cl ________________________
C – Cl _____________________ F – F ________________________
K – F _____________________ H – N ________________________
Take the absolute value of the difference!

5. Intermolecular Forces (IMF)
Attractive forces between molecules.
Much weaker than
chemical bonds
within molecules.
a.k.a. van der Waals forces

6. Types of IMF London dispersion forces
Exist in all atoms and molecules (temporary dipoles)
Attraction between two instantaneous dipoles
Weakest
Larger the molecule, larger the dispersion force.

7. Types of IMF
London Dispersion Forces
View animation online.

8. Types of IMF Dipole-Dipole forces
Attraction between two permanent dipoles
Polar molecules
Medium strength
Stronger when molecules are closer together

9. Types of IMF
Dipole-Dipole Forces + -
View animation online.

10. Types of IMF Hydrogen bonding Special kind of dipole-dipole
Occurs between molecules that have an H bonded to either O, N, or F.
Extremely polar bonds
Strongest
Not chemical bonding

11. Types of IMF
Hydrogen Bonding

12. Types of IMF
The hydrogen bonds in water explain its relatively high boiling point, considering that it is a small molecule. The H-bonds hold the water molecules together as a liquid, so you have to heat it a lot before it will change to a gas. Compare boiling points of these molecules:
Molecule
IMF (s) present
Molar Mass (g/mol)
Boiling Point (oC)
CH4
16.05
- 164
HCl
36.46
- 85
H2O
18.02
100
London Disp.
London Disp.
Dipole-Dipole
London Disp./Dipole-Dipole/Hydrogen Bonding

13. Properties of Liquids Density: Compressibility:
Higher in liquids than gases
Intermolecular forces hold liquid particles together
Gas particles are far apart and have minimal attraction between particles.
Compressibility:
Only slightly in liquids
Molecules already closely packed
Gases have a lot of empty space between particles and can be compressed.

14. Liquid Properties
Viscosity
Resistance to flow
Viscosity Demo

15. Liquid Properties Surface Tension
attractive force between particles in a liquid that minimizes surface area

16. Action of detergents:
Detergents are used to clean dirt/oil from clothing, dishes, etc.
How does this work?
The detergent interacts with water in such a way that the hydrogen bonding is disrupted, and this breaks the surface tension of the water, so that the dirt can mix with the water. Detergents are called surfactants, because they are surface active agents.
Detergents have a special structure that aids the cleaning process. A micelle is formed when detergents interact with dirt/oil and water.
Water is polar.
Dirt and oils are nonpolar.
Detergents are long molecules that have polar “heads” and nonpolar “tails”.
“like dissolves like”
Because detergents have a dual nature, they can dissolve grease and carry it away in the water at the same time.

17. Cohesion and Adhesion
Cohesion is the force of attraction between identical molecules in a liquid (cohesion is a result of intermolecular forces).
Adhesion is the force of attraction between liquid molecules and a solid that is touching them.

18. Liquid Properties Capillary Action
attractive force between the surface of a liquid and the surface of a solid
Meniscus of water in a glass tube is concave: adhesion > cohesion
Meniscus of Hg in a glass tube is convex: cohesion > adhesion
water
mercury

19. PROPERTIES OF SOLIDS
Density: In general, the particles of a solid are more closely packed than those of a liquid. So most solids are more dense than most liquids, but there are exceptions (wax, cork). When solid and liquid states of a substance coexist, the solid is more dense than the liquid, so the solid will sink in the liquid. Water is an important exception (ice floats in water). This is because the H2O molecules are less closely packed in ice than in liquid water.

20. Types of Solids decreasing m.p.
Crystalline - repeating geometric pattern
covalent network
metallic
ionic
covalent molecular
Amorphous –
no geometric pattern
decreasing
m.p.

21. Crystalline vs. Amorphous:
A crystalline solid has particles which are arranged in an orderly, geometric, 3-D structure.
Examples: sodium chloride, ice, gems and minerals
In an amorphous solid, the particles are not arranged in any particular pattern.
Examples: rubber, plastics.

22. Types of Solids Covalent Molecular Covalent Network Amorphous (H2O)
(SiO2 - quartz)
Amorphous
(SiO2 - glass)

23. Phase Changes E. Q=mCΔT D. Q= mHv Temperature Tb  C. Q=mCΔT Tm 
B. Q =mHf
C. Q=mCΔT
Tm 
A. Q= mCΔT
Thermal Energy (Heat)

24. Thermal Energy (heat) Phase Changes: Temperature Changes
B and D represent phase changes
Occur at constant temperature
Temperature Changes
A,C, and E
Temp is changing
Sloping portion of the graph

25. Thermal Energy (heat ) What is happening at each part of the graph:
A. Substance is a solid. Can heat it up or cool it down along this line.
B. Phase change: solid-liquid. The temperature at B (Tm) is the melting (freezing) point.
C. Substance is a liquid. Can heat it up or cool it down along this line.
D. Phase change: liquid-gas. The temperature at D (Tb) is the boiling (condensation) point.
E. Substance is a gas. Can heat it up or cool it down along this line.

26. Heat Calculations Q = m c ∆ T
Q = heat or thermal energy in Joules (J) or calories (cal)
m= mass in grams (g)
C = specific heat in J/(g oC) or cal/(g oC)
∆T= change in temperature in oC

27. Heat Calculations Q =mHf
Q = heat or thermal energy in Joules (J) or calories (cal)
m= mass in grams (g)
Hf = Heat of fusion (J/g or cal/g); use for liquid-solid phase change

28. Heat Calculations Q =mHv
Q = heat or thermal energy in Joules (J) or calories (cal)
m= mass in grams (g)
Hv = Heat of vaporization (J/g or cal/g); use for gas-liquid phase change

29. Things to remember
You can move from left to right or right to left along the curve.
If you move left to right:
all the processes are endothermic (heat must be supplied).
Your answer for heat calculations will be positive.

30. Things to remember If you move right to left:
all the processes are exothermic (heat is removed/released).
Your answer for heat calculations will be negative. You need to hand-correct answer!