Agenda:

1. Agenda: Homework: New Element Project Thermochemistry • Heating & Cooling Curves – a review • Graphing temperature change • Phase Change Diagrams • Review: Energy in Chemical Reactions • Intermolecular Forces • Impact of polarity, shape and size • VSEPR - shape

2. 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.

3. Characteristics of Solids, Liquids & Gases • Sort the terms into 3 columns: S, L, G • (Hint: Look for 3 cards with similar wording and determine which best fits solid, liquid or gas) • Solid Liquid Gas

4. 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.

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

6. 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)

7. States of matter, energy & phase changes Energy level Energy change Phase changes MP/BP Entropy= degree of disorder

8. 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 = -209.9 °C (-345.81998 °F) M.P. of Silver = 961.93 °C (1763.474 °F) M.P. of Carbon = 3500.0 °C (6332.0 °F)

9. Melting Point Particles of a solid vibrate so fast that they break free from their fixed positions. Increasing Thermal Energy Solid Liquid Melting point

10. 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

11. 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.

12. 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 = -195.79 °C (-320.42 °F) B.P. of Silver = 2162 °C (3924 °F) B.P. of Carbon = 4027 °C (7281 °F)

13. States of matter, energy & phase changes Energy level Energy change Phase changes MP/BP Entropy= degree of disorder

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

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

16. 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:

17. 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

18. 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

19. 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

20. 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

21. 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

22. 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?

23. 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?

24. How much energy is needed to heat water from a solid to a vapor? Graph the data – using most of the graph paper Time (when heat energy is added) Resulting temperature See Textbook

25. Review: Heating curve with heat formulas Scroll down http://www.kentchemistry.com/links/Matter/HeatingCurve.htm

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

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

28. Phase diagrams http://www.youtube.com/watch?v=fLOPaJ8lcr8&feature=endscreen&NR=1

30. For Water A = B= C= D=

31. PHET States of Matter http://phet.colorado.edu/en/simulation/states-of-matter

32. 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?

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

34. 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 -1000 C, what phase change(s) would occur if I increased the pressure to 1.00 atm? At what temperature(s) would they occur?

35. Water: Connecting Phase Diagram and Heating Curve

36. Vapor Pressure – Physical Equilibrium

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

38. Resources for S, L, G http://www.kentchemistry.com/links/Matter/HeatingCurve.htm http://www.middleschoolchemistry.com/

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

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

41. 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

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

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

44. Let us review – covalent bondsIntramolecular 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

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

46. Review

47. 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 ___________.

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

49. Non-Polar Covalent Bond Review

50. Polar Covalent Bond Review