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States of Matter

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States of Matter States of Matter Heating Curve Factors Affecting State Vapor Pressure IMF There are 3 types of IMF of different strengths. London Dispersion Forces ... – PowerPoint PPT presentation

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Title: States of Matter


1
States of Matter
  • States of Matter
  • Heating Curve
  • Factors Affecting State
  • Vapor Pressure

2
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3
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4
Defining States
Has Own Shape? Has Own Volume?
Solid
Liquid
Gas
5
Elements and State
6
Changes of State
7
Changes of State
8
Changes of State
  • Endothermic
  • Heat H2O(s) ? H2O(l)
  • Heat H2O(l) ? H2O(g)
  • Heat H2O(s) ? H2O(g)
  • Exothermic
  • H2O(g) ? H2O(l) heat
  • H2O(l) ? H2O(s) heat
  • H2O(g) ? H2O(s) heat

9
States of Matter
  • States of Matter
  • Heating Curve
  • Factors Affecting State
  • Vapor Pressure

10
Heating Curve
11
Heating Curves
  • Added heat will be used in one of two ways
    NEVER BOTH
  • To raise the temperature of a sample graph line
    goes UP!
  • To overcome IMF so that particles separate from
    each other graph line remains FLAT!

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13
Change of State
  • Since the curve plateaus when a change of state
    occurs, the flat line can be used to find the
    temperature at which the change occurs.
  • Melting / Freezing occur at the SAME temp.
    (freezing point, FP)
  • Boiling / Condensation occur at the SAME temp.
    (boiling point, BP)

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15
Temperature Points
  • Boiling Point (BP) is a temperature!
  • - boiling begins if heating
  • - condensation begins if cooling
  • Freezing Point (FP) is a temperature!
  • - freezing begins if cooling
  • - melting begins if heating

16
Phase Equilibrium
  • Each plateau represents a region of equilibrium
  • At that temperature, if no heat is added or
    removed, the rates of.
  • freezing and melting are equal.
  • Vaporization and condensation are equal.

17
Cooling Curve
18
Heat of Fusion
  • Heat of fusion (Hfus) the amount of heat
    required to melt a sample
  • q m Hfus
  • How much heat is required to melt 9.32 g of solid
    water?

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20
Heat of Vaporization
  • Heat of vaporization (Hvap) the amount of heat
    required to vaporize a sample
  • q m Hvap
  • How much heat is required to evaporate 9.32 g of
    water?

21
Heat and Phase Change
  • The heat/gram given off when a sample is frozen
    is the same amount required to melt the same
    sample.

22
Heating Curve
  • There are 5 line segments
  • Solid temp. rising (q mC?T)
  • Solid-Liquid crystal structure breaking apart (q
    mHfus)
  • Liquid temp. rising (q mC?T)
  • Liquid-Gas molecules escaping (q mHvap)
  • Gas temp. rising (q mC?T)

23
Heat Calculations
  • What amount of heat is needed to change 5 grams
    of ice at -20oC to steam at 125oC?
  • (Hint!)

24
States of Matter
  • States of Matter
  • Heating Curve
  • Factors Affecting State
  • Vapor Pressure

25
Temperature
  • As temperature increases, motion of the particles
    increases.
  • As motion increases, the sample will gain enough
    energy to change
  • Solid ? Liquid ? Gas

26
Pressure
  • A gas is very compressible
  • Under enough pressure it can be liquefied.
  • Even more pressure can create a solid

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28
Phase Diagram
  • A phase diagram combines the effects of
    temperature and pressure

As temperature increases, the phases
change Solid ? Liquid ? Gas
As pressure increases, the phases change Gas ?
Liquid ? Solid
29
  • Triple point all 3 phases exist
  • Critical point substance cannot exist as liquid
    at any pressure past this temperature

30
Phase Diagram for Water
  • Unique because solid changes to liquid as
    pressure increases.

31
Solid Water
Water crystals take up more space than the liquid
form
Increased pressure crushes the crystal to liquid
32
Unusual Phase Diagrams
33
Intermolecular Forces
  • Intermolecular forces (IMF) attraction that
    occurs between molecules
  • Intramolecular forces the attraction that occurs
    between atoms IN the molecule

34
IMF
  • There are 3 types of IMF of different strengths.
  • London Dispersion Forces (weakest)
  • Dipole-Dipole
  • Hydrogen Bonding (strongest)

35
London Dispersion Forces
  • Uneven distribution (dispersion) of the electrons
    causes an
  • Instantaneous dipole a particle has a positive
    pole and negative pole only for an instant

36
London Dispersion
  • One dipole can induce another dipole for a
    ripple effect.
  • Very weak occurs mostly in gases

37
London Dispersion
  • Gas molecules tend to
  • Have nonpolar covalent bonds
  • Have symmetry
  • There is no existing polarity in the molecule

38
Dipole-Dipole
  • Dipole-dipole the force of attraction that
    occurs between 2 polar molecules
  • In this case the poles are permanent, not
    instantaneous.

39
Dipole-Dipole
  • Substances that experience dipole-dipole
    attraction tend to be gases or liquids.

40
Hydrogen Bonding
  • Hydrogen bonding exists whenever hydrogen is
    bonded toN, O, or F
  • It is a stronger dipole than most polar molecules
  • Therefore the attraction
  • between molecules is
  • stronger

41
Hydrogen Bonding
  • Water is the most common example of hydrogen
    bonding.
  • Attraction occurs between
  • the negative oxygen in
  • one molecule and the
  • positive hydrogen in
  • another.

42
Substances that experience strong
H-bonding tend to be liquids.
43
Special Properties of Water
  • Surface tension occurs due to the strong
    H-bonding between water molecules

44
Special Properties of Water
  • Capillary action due to the attraction of polar
    H2O to polar glass molecules

45
Hydrogen Bonding
  • NH3 and H2O experience hydrogen bonding.

46
  • Which has hydrogen bonding and which has
    dipole-dipole?

Hint Look for H bonded to O!!
47
Recap - IMF
  • There are 3 types of IMF of different strengths.
  • 1) London Dispersion Forces (weakest)
  • substances are gases
  • 2) Dipole-Dipole (moderate)
  • substances are gases/liquids
  • 3) Hydrogen Bonding (strongest)
  • substances are liquids

48
More Forces
  • There are 2 more forces that exist between
    particles.
  • Van der waals forces attraction between positive
    nucleus of one atom and electrons of another
  • (exceedingly weak and fleeting)
  • London Dispersion ? Dipole-Dipole ? H-bonding
  • Ionic bonding attraction between cation and
    anion
  • (exceedingly strong substances tend to be
    solids)

49
Intermolecular Forces
50
Intermolecular Forces
51
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53
States of Matter
  • States of Matter
  • Heating Curve
  • Factors Affecting State
  • Vapor Pressure

54
Vapor Pressure
  • Vapor pressure the pressure of a vapor over its
    liquid

55
Vapor Pressure Equilibrium
  • Rate of vaporization Rate of condensation

56
VP and Temperature
  • Vapor pressure varies exponentially with
    temperature.
  • More heat applied
  • provides the energy
  • needed for molecules
  • to escape their
  • liquid.

57
VP and IMF
  • Molecules that have STRONG intermolecular forces
    cannot easily escape their liquid.
  • ? they have LOW vapor pressure
  1. Which substance has the greater VP?
  2. Which substance has the stronger IMF?

58
Vapor Pressure
  • Vapor pressure can
  • be considered with an
  • open container
  • Then, it can be
  • thought of as the force
  • with which molecules
  • escape their liquid

59
Vapor Pressure and BP
  • The atmosphere exerts a pressure on the surface
    of the earth of about 15 lbs./in2
  • This pressure works to prevent molecules from
    escaping their liquid when in an open container

60
Vapor Pressure and BP
  • When VP exerts as much force as atmospheric
    pressure, a liquid will boil.
  • Boiling point the temperature at which
  • VP atm pressure
  • A liquid is in equilibrium with it gas

61
Vaporization
  • The difference between evaporation and boiling.

62
Boiling Point
  • Atmospheric pressure is lower at higher
    elevations.
  • ?it takes less energy to reach the boiling point
  • The BP is lower
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