1. Warm-up: Atomic size is one of the many trends of the Periodic Table.
Describe one reason atomic size many vary between the elements on the Periodic Table.
Arrange these elements in descending order: Al, Mg, P, Si, Na, S
Explain why you chose this order.

2. Characteristics of Solids, Liquids & Gases
Sort characteristics of solids, liquids and gases into 3 columns (use workbook page 1 if needed)
Solid Liquid Gas

3. Most substances, like water, can exist in all three states.
An iceberg is made of water in solid form.
This glass contains liquid water.
A cloud is made of water vapor, a type of gas.

4. WHAT ARE THE CHANGES OF STATE?
Which are endothermic?
Which are exothermic?
GAS
Deposition
Boiling / Evaporation
Sublimation
Condensation
Freezing
SOLID
LIQUID
Melting

5. Changing States (Phase changes)
Where on the picture would we place:
Melting Point?
Boiling Point?
Condensing Point?
Freezing Point?
Increase Thermal Energy (Heat up)
Solid
Liquid
Gas
Decrease Thermal Energy (Cool off)

6. Melting point Melting - change from solid to liquid
Melting point - SPECIFIC temperature when melting occurs.
Each pure substance has a SPECIFIC melting point.
Examples:
M.P. of Water = 0°C (32°F)
M.P. of Nitrogen = °C ( °F)
M.P. of Silver = °C ( °F)
M.P. of Carbon = °C ( °F)

7. Melting Point Solid How does melting occur?
Particles of a solid vibrate so fast that they break free from their fixed positions.
Increasing Thermal Energy
Solid
Liquid
Melting point

8. Vaporization Vaporization – change from liquid to gas
Vaporization happens when particles in a liquid gain enough energy to form a gas.
Increasing Thermal Energy
Gas
Liquid
Boiling point

9. Two Kinds of Vaporization
Evaporation – vaporization that takes place only on the surface of the liquid
Boiling – when a liquid changes to a gas BELOW its surface as well as above.

10. Boiling Point Boiling Point – temperature at which a liquid boils
Each pure substance has a SPECIFIC boiling point.
Examples:
B.P. of Water = 100°C (212°F)
B.P. of Nitrogen = °C ( °F)
B.P. of Silver = 2162 °C (3924 °F)
B.P. of Carbon = 4027 °C (7281 °F)

11. Heating and Cooling Curves of a Substance Representing MP, BP, CP, FP
Energy (heat) added
Energy (heat) released:

12. Intermolecular Forces
Forces between molecules (compounds) which helps determine whether a substance is a solid or liquid
Gases have little/no intermolecular forces

13. Energy requirements for water Three formulas : specific heat Q = mCp∆T
Energy requirements for water Three formulas : specific heat Q = mCp∆T heat of fusion Q= mHf heat of vaporization Q= mHv
Heating
Cooling
Energy (heat) added
Energy (heat) released:

14. Energy calculations related heating or cooling specific substances
Specific heat (Cp)
Latent heat
Heat of fusion (Hf)
Heat of vaporization (Hv)
Use reference tables – values for each pure substance

15. Heat calculations – 3 formulas
Specific heat = heat required to raise the temperature of 1 gram of substance 1 °C
Formula: Q = mCp∆T
Specific heat
Specific for each pure substance
Use reference tables

16. Heat calculations – 3 formulas
Heat of fusion -
Amount of heat added to melt a substance
Amount of heat released to freeze a substance
Formula Q= mHf
Specific for each pure substance
Use reference tables

17. Heat calculations – 3 formulas
Heat of vaporization-
Amount of heat added to boil a substance
Amount of heat released to condense a substance
Formula Q= mHv
Specific for each pure substance
Use reference tables

18. Heat energy
In a heat calculation problem, if the problem asks about melting/freezing you would multiply the mass times _____________________.
heat of fusion
heat of vaporization
or specific heat
In a heat calculation problem, if the problem asks about vaporizing/condensing of steam, you would multiply the mass times ________.
Heat of fusion
Heat of vaporization
Specific heat
In a heat calculation problem, if the problem asks about a change in temperature, you would multiply the mass times ___________________ times the change in temperature.
Heat of fusion
Heat of vaporization
Specific heat

19. Thermochemistry Problems related to water
How much heat is required to raise the temperature of 789 g of water from 25oC to 70oC?
2. How much heat is released when 432 g of water cools from 71oC to 18oC?
3. How many joules of heat are given off when 5.9 g of steam cools from 175oC to 125oC?

20. 4. How many joules does it take to melt 35 g of ice at 0oC?
5. How much heat is released when 85 g of steam condense to liquid water?
6. How much heat is necessary to raise the temperature of 25 g of water from 10 oC to 60 oC?
7. How much heat is given off when 50 g of water at 0oC freezes?

21. What factors impact change?
Intermolecular forces
Energy
Conditions: T, P, V, amount,

22. Phase Diagrams: What is added to this diagram? Why?

23. Phase diagrams

25. A = B= C= D= For Water T °C 200 °C -2°C 100 °C -2 °C 30°C P - atm
Phase
Liquid
Vapor
A = B= C= D=

26. PHET States of Matter

27. Review: Interpreting Phase Diagrams
Phase Diagrams. Use the phase diagram for water below to answer the following questions.
Review: Interpreting Phase Diagrams
What is the state of water
at 2 atm and 50 C?
What phase change will
occur if the temperature
is lowered from 80C
to -5C at 1 atm?
You have ice at -10C and 1 atm.
What could you do in order
cause the ice to sublime?

28. Interpreting a Phase Diagram of Water at varying pressures
Example: 100 atm

29. 1) What is the normal melting point of this substance
1) What is the normal melting point of this substance? ________ 3) What is the normal boiling point of this substance? ________ 4) What is the normal freezing point of this substance? ________ 5) If I had a quantity of this substance at a pressure of 1.25 atm and a temperature of 00 C and heated it until the temperature was 7500 C, what phase transition(s) would occur? At what pressure(s) would they occur? 6) At what temperature do the gas and liquid phases become indistinguishable from each other? ________ 7) If I had a quantity of this substance at a pressure of 0.25 atm and a temperature of C, what phase change(s) would occur if I increased the pressure to 1.00 atm? At what temperature(s) would they occur?

30. Water: Connecting Phase Diagram and Heating Curve

31. Vapor Pressure – Physical Equilibrium
The vapor pressure is the pressure measured when there is an equilibrium between the gas and liquid phases. The rates of condensation and vaporization are equal.

32. Vapor pressure http://www.chem.purdue.edu/gchelp/liquids/vpress.html
Discovery Ed video

33. Resources for S, L, G

34. How does the chemical composition of a substance impact whether it is a gas, liquid or solid at room temperature?

35. Overview: Factors that Impact State of Matter
Type of compound – Ionic, Covalent, Metallic
Intermolecular Forces, impacted by
Shape
Size
Polarity

36. Intermolecular Forces
Attractive forces between molecules
Not between individual atoms
Much weaker than the bonds within a molecule
Intramolecular bonds form between 2 atoms in a molecule/compound
_________ , _________, ________
Can determine the state of matter by the number and type of these forces
Lots of forces= liquid Lots and lots = solid

37. What causes these intermolecular forces?
Opposites attract:
In chemistry this means:
How do these attraction between molecules form?
Polarity (partial polarity)
Shape
Size

38. Intermolecular Forces
Three Types
Hydrogen
Dipole – dipole
London Dispersion (Van der Waals)
Based on weak attraction between molecules
partial negative – partial positive

39. Let us review – covalent bonds Intramolecular bond
Type of atoms in covalent bond
Electronegativity Difference
Sharing valence electrons to form bonds
Some share equally = non-polar covalent bonds
Some share unequally = polar covalent bonds

40. Electronegativity Differences
Review
Electronegativity Differences = ∆EN
Covalent bonds
Ionic Bonds
∆ O
∆ 3.2
∆ 1.7
Increasing polar (+ side and – side) characteristics

41. Review

42. Electronegativity Difference
Review
The electronegativity difference must be equal to or less than _______.
It is a polar covalent bond if the difference is between __________.
It is a non-polar covalent bond if the difference is between ___________.

43. Non-Polar Covalent Bond ∆EN= 0 – 0.3
Review
Non-Polar Covalent Bond ∆EN= 0 – 0.3
The Electron pair that makes up the bond is shared evenly.

44. Non-Polar Covalent Bond
Review

45. Polar Covalent Bond
Review

46. Review
Polar Covalent Bond

47. Polar Covalent Bond ∆EN = 0.4 – 1.7
Review
The electron pair that makes up the bond is closer to
the element that has the higher electronegativity.

48. Intermolecular Forces
Three Types
Hydrogen
Dipole – dipole
London Dispersion (Van der Waals)
Based on weak attraction between molecules
partial negative – partial positive

49. Types of Intermolecular Forces
Strongest intermolecular force
Hydrogen “bond” (~ 10% of a covalent bond)
Molecule must be polar (+ and – sides)
H in one molecule is attracted to the N,O,F of another molecule
Hydrogen Bond - bad choice of words – an attractive force , not a bond

50. Hydrogen “bonds”: attraction between H with N, O, F
Hydrogen bonds between water molecules.
Hydrogen bonds give unique properties to water.

51. Types of Intermolecular Forces
Other intermolecular forces
Dipole-dipole – all polar molecules (weaker)
London dispersion forces – all molecules (weakest)

52. Dipole-dipole Based on polarity of molecules
Found with polar covalent compounds
Use with elements other than H attracted to N, O or F

53. London Dispersion Weakest Temporary polarity
Based on movement of the electrons around the nucleus
Impacts all molecules –
non-polar and polar

54. Factors that impact the state of matter
Intermolecular forces
How do these attraction between molecules form?
Polarity (partial polarity)
Shape
Size

55. Polarity of Water
Short Tutorial with animations of polarity

56. Polarity http://phet.colorado.edu/en/simulation/molecule-polarity
Use real molecules tab
Molecular dipoles
Electronegativity
Electron density or electrostatic potential
Note: VSEPR – valence shell electron pair repulsion impacts shape

57. Shapes of Molecules
PHET simulation (?)

58. Predicting the Shape: VSEPR
Valence shell electron pair repulsion
VSEPR theory
Non-bonding pairs of electrons (lone pairs) will push away (repel) from each other
Use Lewis structures to model VSEPR

59. VSEPR: Lewis dot structure shows the pulling away of electron pairs
Water
Ammonia

60. Methane: CH₄

61. Ammonia: NH₃

62. Water: H₂0

63. Why is water unique? Why does ice float on water?
Hydrogen bonds -

64. Review: Energy in chemical reactions
Exothermic chemical reaction

65. Review: Energy in chemical reactions
Endothermic chemical reaction