1. States of Matter and Energy
Dr. Walker

2. States of Matter Phase – A part of matter that is uniform
Solid (lowest energy)
Liquid
Gas
Plasma (highest energy)
Found in stars, lightning

3. State Changes
Enthalpy = Energy

4. Intermolecular Forces
Intramolecular Forces – bonding between atoms to achieve octet
Ionic Bonding
Covalent Bonding

5. Intermolecular Forces
Forces between separate molecules
As energy is added (higher temperature), molecules have more motion (higher KE) and overcome intermolecular forces to change state
These forces differ with the type of compound (polar vs. nonpolar)
Determines physical properties (bp/mp, density, etc.)

6. Intermolecular Forces
Dipole-Dipole Forces
Interactions between polar molecules
Results from unequal sharing
About 1% of strength of ionic bonds
“partial charges” – shows unequal
sharing between the two atoms

7. Review… Polar molecules
Molecule with difference in electronegativity between atoms
Not symmetric
Unequal sharing of electrons (shown with d symbol

8. Hydrogen Bonding Special dipole-dipole attraction
Hydrogen covalently bonded to highly electronegative elements (N, O, F) has a higher than normal d+ charge
Bond strength is higher than other dipole-dipole attractions
VERY IMPORTANT in biology
Links strands of DNA
Involved in protein folding (3D shape)

9. T A C G Thymine hydrogen bonds to Adenine
Cytosine hydrogen bonds to Guanine C G

10. Hydrogen Bonding Gives water unique properties
Solid is less dense than liquid (why ice floats)

11. London Dispersion Forces (aka Van der Waals Forces)
Movement of electrons in an atom create temporary dipole
Attraction between nonpolar molecules
Much weaker than dipole interactions

12. London Dispersion Forces

13. Intermolecular Forces vs. Melting/Boiling Points
Ionic Networks – Strongest interaction
Covalent Bonding
Hydrogen Bonding (dipole forces) – Water, Compounds with N, O
London Dispersion Forces – hydrocarbons (Weakest Interaction)

14. Bonding Strength vs. MP
Stronger interaction = higher boiling/melting point
Order of increasing melting points:
London Dispersion (hydrocarbons) < Hydrogen bonding and dipole interactions (covalent compounds with N, O, halogens) < Ionic bonding compounds

15. Comparing Intermolecular Forces vs. Boiling Point
Material
Symbol/Formula
Boiling Point
(oC)
Nonpolar Covalent
Hydrogen H2
-253
Methane
CH4
-164
Polar Covalent
Ammonia
NH3
-33
Water
H2O
100
Ionic
Sodium Chloride
NaCl
1413
Magnesium Oxide
MgO
2826

16. Vapor Pressure
Pressure created by gas molecules above the surface of a liquid
Volatile liquids (alcohol, gas, nail polish remover) have a high vapor pressure
Nonvolatile liquids (water) have a low vapor pressure
Solutions have a lower vapor pressure (remember colligative properites!
More vapor =
Higher vapor pressure!

17. Vapor Pressure Inversely proportional to intermolecular (IM) forces
Weak intermolecular forces allows more gas to form, resulting in higher vapor pressure
Polar compounds – higher IM forces – less gas formation – higher bp
Nonpolar compounds – low IM forces – more gas formation – lower bp

18. Vapor Pressure
When vapor pressure = atmospheric pressure, a substance will boil
Lower vapor pressure – further from 1 atm – more heat needed to increase pressure
Higher vapor pressure – closer to 1 atm – less heat needed to increase pressure

19. Inc. polarity
Inc. IM forces
Dec. VP
Inc. MP
Notice – as the boiling points increase…
1) ..the vapor pressures start off lower
2) …the polarity increases (propane, ethanol, water, ethanoic acid)
3) …the IM forces increase (with polarity)

21. Vocabulary Enthalpy Entropy Total Energy of a system
Disorder of a system

22. Measuring Energy What do we measure energy in?
The Joule is the SI (metric) system unit for measuring heat:
The calorie is the heat required to raise the temperature of 1 gram of water by 1 Celsius degree
1 BTU is the heat required to raise the temperature of 1 pound of water by 1 F

23. Energy and Heating/Cooling
What is the relationship between heat and temperature?
Heat – transferred energy
Temperature – measure of kinetic energy
If you transfer the same amount of heat in different substances, the temperature change will not be the same
The temperature change is governed by…

24. Specific Heat Capacity
Determines temperature changes vs. heat transfer
The energy required to raise 1 g of material by 1 degree celsius
E or E = mCDT
change in
E = Energy m = mass C = specific heat DT = temperature
Temperature scale doesn’t matter, 1 C = 1 K
Different substances have different specific heat capacities
Metals tend to be low (>1 J/g x K)
Water is very high (4.18 J/g x K)
Reason why a blacktop gets hotter much faster than a pool on a sunny day!

25. Example
A 4.0 g sample of glass was heated from 274 K to 314 K and absorbed 32 J of heat. What is the specific heat?

26. Example
A 4.0 g sample of glass was heated from 274 K to 314 K and absorbed 32 J of heat. What is the specific heat?
Specific Heat = 32 J/(4.0 g x [ ])
32 J/(4.0 g x 40 K)
0.20 J/g x K

27. Example
Water has a specific heat of 4.18 J/g C. How much energy is required to raise the temperature of a 1000 g sample by 5 C?

28. Example
Water has a specific heat of 4.18 J/g C. How much energy is required to raise the temperature of a 1000 g sample by 5 C?
E = mCDT
E = (1000 g)(4.18 J/g C)(5 C) = J or 20.9 kJ

29. Energy of Phase Changes
When compounds change phase, the enthalpy (energy) and entropy (disorder) changes
Plasma, Gas have highest entropy
Solid has lowest entropy
Disorder – think of freedom of molecular movement
Gases have highest entropy
Solids have lowest entropy

30. What About During A Phase Change?
When a substance is at the melting or boiling point, the temperature does not change until the phase change is complete
For example, water boils at 100 oC. Temperature can’t rise to 101 oC until the entire sample in steam!
The energy added is breaking intermolecular forces, not to moving the molecules.

31. Energy of Phase Changes
Notice that during a phase change, energy is added, but the temperature does not change

33. Molar Heat of Fusion
The energy required to melt one mole of a solid at its melting point
Molar Heat of Fusion for ice = 6.00 kJ/mole
Molar Heat of Fusion for methane (CH4) = 1592 J/mole
These are per mole, not per gram. You could convert from grams if necessary…

34. Molar Heat of Fusion
How much energy is required to change 6 g of ice to water at 0 C (melting point)?

35. Molar Heat of Fusion
How much energy is required to change 6 g of ice to water at 0 C (melting point)?
You have to find the moles of water first by dividing by molar mass
6 g/(18 g/mole) = 0.33 moles
0.33 moles x 6.00 kJ/mole = 1.98 kJ of energy

36. Molar Heat of Vaporization
Energy required to convert 1 mole of liquid to 1 mole of gas at its boiling point
Water = kJ/mole
Methane (CH4) = 8.19 kJ/mole
Ethanol (C2H6O) = 38.6 kJ/mole
More polar compounds have a higher heat of vaporization (more energy required to break intermolecular forces)

37. Molar Heat of Vaporization
How much energy is required to change 90 g of ethanol (C2H6O) from liquid to gas at its boiling point?

38. Molar Heat of Vaporization
How much energy is required to change 90 g of ethanol (C2H6O) from liquid to gas at its boiling point?
90 g/(46 g/mole) = 1.96 moles
1.96 moles x 38.6 kJ/mole = 75.7 kJ

39. Energy of a Heating Curve

40. Relationships Between Physical Properties

41. Phase Diagrams Definition
Graph of pressure vs. temperature
Can tell the phase of a substance at a given point
Every compound has a different phase diagram

42. Phase Diagram of Water Point A – Triple Point
All three phases at equilibrium at this temperature and pressure

43. Phase Diagram of Water Point C – Critical Point
Above this temperature and pressure, the substance CANNOT exist as a liquid

44. Phase Diagram of Water Atmospheric Pressure = 1.0 atm
The temperatures on the phase lines are the melting point and boiling point

45. Interpret A Phase Diagram

47. Energy and Reactions
All Chemical Reactions either release or absorb energy – it just isn’t always noticeable
Explosions – chemical reaction with a large amount of energy released

48. Energy and Reactions
Exothermic – Reaction process that gives off energy
Reactants  Products + Energy
Combustion reactions – give off a great deal of energy in the form of fire
Neutralization reactions – acid/base reactions typically get fairly hot

49. Energy and Reactions
Endothermic – Reaction process that absorbs energy
Reactants + Energy  Products

50. Reaction Coordinates Exothermic Reaction
Reaction Progress starts with reactants
Ends with products
Energy (H) at beginning higher than Energy at end
Negative Enthalpy
(-DH)
Add acid + base as a demo, feel the beaker

51. Reaction Coordinates Endothermic Reaction
Energy (H) at end higher than Energy at beginning
Positive Enthalpy (+ DH)

52. Quick Example What kind of reaction is this?
Is the DH positive or negative?

53. Activation Energy
Activation Energy – The minimum amount of energy required to start a reaction

54. Catalysts
A catalyst lowers the activation energy necessary to start a chemical reaction
Changes the rate of a chemical reaction without being consumed
Point 2 demonstrates a catalyst on a reaction coordinate

55. Quick Example Is this reaction endothermic or exothermic?
What is the DH of this reaction?

56. Common Catalysts Catalytic Converter Enzymes
Made of precious metals (Pt, Rh, Pd)
2 CO + O CO2
Enzymes
Proteins which speed up metabolic processes
Cells could not function without enzymes
Enzyme malfunctions result in disease
Lactase (lactose intolerance)

57. What Can Affect Reaction Rates?
Concentration – More molecular collisions, faster rate
Temperature – Higher temperature, faster reactions
Molecules move faster, more molecular collisions

58. What Can Affect Reaction Rates?
Surface Area – More exposed area (think sugar cube vs. granulated sugar), faster reactions
Pressure (gas) – More pressure = more molecular collisions, faster reactions