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Ch. 12 – States of Matter I. Intermolecular Forces.

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Presentation on theme: "Ch. 12 – States of Matter I. Intermolecular Forces."— Presentation transcript:

1 Ch. 12 – States of Matter I. Intermolecular Forces

2 A. Definition of IMF  Attractive forces between molecules.  Much weaker than chemical bonds within molecules.  a.k.a. van der Waals forces

3 C. Johannesson B. Types of IMF

4  London Dispersion Forces View animation online.animation

5 B. Types of IMF  Dipole-Dipole Forces + + - - View animation online.animation

6 B. Types of IMF  Hydrogen Bonding

7 C. Determining IMF  NCl 3 polar = dispersion, dipole-dipole  CH 4 nonpolar = dispersion  HF H-F bond = dispersion, dipole- dipole, hydrogen bonding

8 II. Physical Properties Ch. 12 - Liquids & Solids

9 A. Liquids vs. Solids LIQUIDS Stronger than in gases Y high N slower than in gases SOLIDS Very strong N high N extremely slow IMF Strength Fluid Density Compressible Diffusion

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

11 B. Liquid Properties  Capillary Action attractive force between the surface of a liquid and the surface of a solid

12 B. Liquid Properties  Viscosity Measure of the resistance of a liquid to flow

13 B. Liquid Properties  Cohesion Force of attraction between identical molecules  Adhesion Force of attraction between different molecules watermercury

14 C. Types of Solids  Crystalline - repeating geometric pattern covalent network metallic ionic covalent molecular  Amorphous - no geometric pattern decreasing m.p.

15 C. Types of Solids Ionic (NaCl) Metallic

16 C. Types of Solids Covalent Molecular (H 2 O) Covalent Network (SiO 2 - quartz) Amorphous (SiO 2 - glass)

17 C. Types of Solids  Allotrope In a covalent network, different forms in the same state Carbon allotropes watermercury

18 C. Types of Solids  Amorphous Solid No geometric pattern watermercury

19 Ch. 12 - Liquids & Solids III. Changes of State

20 A. Phase Changes

21  Evaporation molecules at the surface gain enough energy to overcome IMF  Volatility measure of evaporation rate depends on temp & IMF

22 A. Phase Changes Kinetic Energy # of Particles Boltzmann Distribution tempvolatilityIMFvolatility

23 A. Phase Changes  Equilibrium trapped molecules reach a balance between evaporation & condensation

24 A. Phase Changes  Vapor Pressure pressure of vapor above a liquid at equilibrium IMFv.p.tempv.p. depends on temp & IMF directly related to volatility temp v.p.

25 A. Phase Changes  Boiling Point temp at which v.p. of liquid equals external pressure IMFb.p.P atm b.p. depends on P atm & IMF Normal B.P. - b.p. at 1 atm

26  Which has a higher m.p.? polar or nonpolar? covalent or ionic? A. Phase Changes  Melting Point equal to freezing point polar ionic IMFm.p.

27 A. Phase Changes  Sublimation solid  gas v.p. of solid equals external pressure  EX: dry ice, mothballs, solid air fresheners

28 B. Heating Curves Melting - PE  Solid - KE  Liquid - KE  Boiling - PE  Gas - KE 

29 B. Heating Curves  Temperature Change change in KE (molecular motion) depends on heat capacity  Heat Capacity energy required to raise the temp of 1 gram of a substance by 1°C

30 B. Heating Curves  Phase Change change in PE (molecular arrangement) temp remains constant  Heat of Fusion (  H fus ) energy required to melt 1 gram of a substance at its m.p.

31 B. Heating Curves  Heat of Vaporization (  H vap ) energy required to boil 1 gram of a substance at its b.p.  EX: sweating, steam burns, the drinking bird

32 C. Phase Diagrams  Show the phases of a substance at different temps and pressures.

33  The following slides… same information, different explanation and examples

34 Phase Changes

35 Why do liquids and solids form at all? KMT postulates – A gas is a collection of small particles traveling in straight-line motion and obeying Newton's Laws. – The molecules in a gas occupy no volume. – Collisions between molecules are perfectly elastic no energy is gained or lost during the collision – There are no attractive or repulsive forces between the molecules. – Kinetic energy is proportional to temperature X

36 What is a phase? Region of matter that is: – chemically uniform – physically distinct – mechanically separable. Often synonymous with (same meaning as) “state of matter”

37 Changing phases Distinguish liquid vs. solid?

38 Properties of … Gas phase – Like/unlike soccer players on field Liquid phase – Like/unlike crowd at a rally – Like/unlike gases Solid phase – Like/unlike movie theatre Compare intermolecular interactions – gases vs. liquids vs. solids

39 Phases of matter Add more energy

40 Phase change vocabulary A - melting B - freezing C – boiling/ evap D - condensation E - sublimation F - solidification

41 Phase changes and IMF’s As InterMolecular Forces increase, melting and boiling temperatures _________? (increase or decrease)? Metallic bonding Network covalent bonding Larger sphere, higher melting point

42 Strength of interactions Which simulation has stronger intermolecular interactions? A or B – How do you know? AB Same temperature

43 Ranking of intermolecular interactions Water Wood Iron Air Gold Mercury Carbon dioxide Oxygen Gasoline Lead vs. Why?

44 Why is water special? Periodic trends – Boiling and melting points of hydrides

45 KMT, energy and phase changes

46 Total energy

47 Water phases present?

48 Heating curves Why are b and d flat?

49 Boiling vs. melting Which takes more energy? (same mass)

50 Consider liquid  gas Evaporation removes energy

51 Evaporative Cooling

52 Evaporation and volatility Evaporation – molecules at the surface gain enough energy to overcome attractive intermolecular forces (IMF) Volatility – measure of evaporation rate – depends on temperature and IMF

53 Volatility Kinetic Energy # of Particles Boltzmann Distribution tempvolatilityIMFvolatility

54 Boiling – Fixed temperature

55 Why do liquids boil? Boiling Point – T at which liquid vapor pressure = atmospheric pressure depends on P atm & IMF – Normal boiling point - b.p. at 1 atm IMFb.p.P atm b.p.

56 Temperature/pressure phase diagram for water

57 Transition to Supercritical CO 2

58 Interpret phase diagram for water

59 Draw a phase diagram for NH 3 (used as a refrigerant in RV’s and solar cooling systems)

60 Ammonia phase diagram

61 Phase Changes Need to know… Know why liquids and solids exist at all – Failure of KMT postulate Know phase names and changes – molecular views of phases and changes – distinguish S vs. L on molecular level Explain heating curve origins and parts – Why no change in T during phase change? Explain differences and similarities between evaporation, volatility and boiling Explain evaporative cooling on a molecular level Interpret temperature-pressure phase diagrams

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