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Solids and Liquids

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Title: Solids and Liquids


1
Solids and Liquids
  • Chapters 16 and 17

2
Kinetic Theory of Matter
  • Gas Solid Liquid

3
Kinetic Theory of Gases
  • Particles are independent from one another and
    move in straight lines.
  • Change of direction occurs only when particles
    strike each other or container.
  • Movement is random.

4
Kinetic Theory of Liquids
  • Particles much closer together.
  • Particles do not act independently from one
    another.
  • Have a vibratory type of motion.
  • Movement occurs when particles shift between the
    spaces of particles.

5
Kinetic Theory of Solids
  • Particles close together and do not act
    independently.
  • Particles occupy a relatively fixed position in
    relation to surrounding particles.
  • Vibration around a fixed point.

6
Properties of Solids
  • All true solid substances are composed of
    crystals.
  • Solids have low kinetic energy.
  • They have a definite volume and assume their own
    shape.
  • High densities and Melting points.

7
Crystal Structure
  • Crystals are made from small units called unit
    cells.
  • These units are repeated over and over again as
    the crystal grows.
  • The arrangement of these unit cells is determined
    by the bonds between the particles.
  • Early crystallographers classified crystals on
    the basis of their external shapes into seven
    crystal systems.

8
Crystal Systems
  • Refer to page 398 in text.
  • EX Cubic Hexagonal

9
  • Cubic
  • Hexagonal

10
Allotropes
  • Allotropes are different chemical states of same
    element in same physical state.
  • EX Allotropes of carbon are diamonds and
    graphite.
  • Diamond cubic lattice held together by strong
    covalent bonds.
  • Graphite layers of hexagonal units held
    together by weak forces.

11
  • Diamond Graphite

12
Amorphous Substances
  • Amorphous substances, sometimes referred to as
    amorphous solids, appear to be in the solid state
    but have no crystal structure.
  • Glass is the best known example.
  • Glass does not have a sharply defined melting
    point (like other solids).
  • Viscosity (resistance to flow) is used to
    describe the texture of glass.
  • High viscosity, high resistance to flow (thick).

13
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14
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15
Properties of Liquids
  • Liquids occupy a definite volume but they do not
    have their own shape.
  • Liquid movement is considered fluid.
  • Bodies of nonmoving liquids are always flat on
    top.

16
Surface Tension and Capillary Action
  • A liquid molecule is attracted to other liquid
    molecules. This is called cohesion.
  • Surface tension is a function of cohesion.
  • A liquid molecule can also be attracted to other
    materials. This is called adhesion.
  • Capillary action is a function of adhesion.

17
Surface Tension
  • Surface tension is the force at the surface of a
    liquid due to the cohesive forces of the
    molecules of liquid for each other and adhesive
    forces of the liquid molecules for the walls of
    the container.

18
  • A paper clip will float on water despite having a
    higher density.
  • The cohesive forces of the water molecules keep
    the paper clip from pushing through.

19
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20
Gerris remigis takes a walk
21
Capillary Action
  • Capillary action is related to the adhesive
    properties of liquids.
  • Liquid molecules are attracted to the straw. As
    they climb up the straw, other molecules
    follow.

22
  • Capillary action is limited by gravity and the
    size of the straw.
  • The thinner the straw or tube the higher up
    capillary action will pull the liquid.

23
Intermolecular Forces
  • Intermolecular forces are called Van Der Waals
    forces .
  • These forces are strongest in liquids and solids.
    (Remember gases have little or no attraction for
    other gas molecules)
  • There are three such types of Van Der Waals
    forces

24
  • London Dispersion Forces (LDF) which are forces
    that exist between elements or nonpolar molecules
    as a result of positive nuclei of one molecule
    attracting the electrons of another molecule. All
    molecular substances exhibit London Forces which
    are the weakest of the Van Der Waals forces.
  • EX O2, CH4, Kr

25
  • 2. Dipole-Dipole interactions which are forces
    that exist between polar molecules where the
    positive end of one molecule attracts the
    negative end of another molecule. Only polar
    molecules can exhibit this type of Van Der Waals
    forces which is considered the second strongest.
  • EX HBr, CH3Cl, CO, PCl3, C12H22O11

26
  • 3. Hydrogen bonding interactions are forces that
    exist between molecules that have a hydrogen atom
    bonded to a highly electronegative atom such as
    Oxygen, Nitrogen, or Fluorine. This represents a
    strong polarity that will have the Hydrogen end
    (positive) attracting the negative end (O,N, or
    F) of other molecules. This type of Van Der Waals
    force represents the strongest type.
  • Alcohols are considered Hydrogen bonding
    because of the OH group.
  • EX H20, NH3, HF, CH3OH

27
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28
Identify the IMF Force
  • Must do Lewis Structures (Bonding Unit, Steps
    1-9) to determine polarity of molecules
  • Ar
  • CCl4
  • HF
  • NCl3
  • C6H11OH
  • C6H12O6
  • NH3
  • CO2

29
Ranking Strength of IMFs
  • When determining the strength of IMF forces when
    comparing different substances, follow these
    guidelines
  • Identify type of force for each substance.
  • Ionic bonds are not considered to be an IMF but
    you will often see them. Ionic bonds are stronger
    than IMFs.
  • Make subcategories of hydrogen bonded,
    dipole-dipole, and LDF.

30
  • 4. Rank according to that order. If presented
    with more than one substance in a category follow
    these guidelines
  • Hydrogen Bonded the greater the
    electronegativity difference between atoms, the
    stronger the force.
  • Dipole-Dipole Same as hydrogen bonded. Look at
    electroneg difference.
  • LDF Calculate molar mass of each substance.
    The greater the molar mass, the stronger the
    force.

31
  • Rank the following set of substances in order of
    increasing strength of IMF.
  • H2, H2O, CO, CO2, NH3, SO3, KCl, Br2, NCl3

32
IMFs and Boiling and Melting Points (phase
changes)
  • There is a relationship between Intermolecular
    forces and Boiling and Melting points of
    substances (all phase changes EX evaporation).
  • The stronger the IMF, the higher the Boiling and
    Melting points are.
  • It takes more energy (temperature) to overcome
    the force created by the IMF.

33
  • Identify which substance out of each pair has the
    higher boiling point
  • O2 or CO
  • H2O or HBr
  • HF or CH4
  • C6H12O6 or C12H22O11
  • PCl3 or NO
  • HF or NH3

34
Phase Changes
  • Phase changes are due to changes in temperature,
    which affects kinetic energy, or pressure, which
    establishes how close the molecules of a
    substance are.

35
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36
Heating and Cooling Curves
  • Heating Curve

37
  • Cooling Curve

38
Phase Diagrams
  • Phase diagrams show how the states of matter in a
    system are affected by changes in temperature and
    pressure.

39
  • This phase diagram shows the relative
    temperatures and pressures at which each of the
    three states of matter can exist.
  • Note the point labeled triple point.
  • This is the point at which all three states can
    exist in equilibrium.

40
Phase Diagram for CO2
41
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42
Equilibrium
  • Equilibrium is the state in a system in which
    there is no net change in that system.
  • In a state change system, for example, a beaker
    of water that is closed off will have two phases.
  • 1. Liquid
  • 2. Gas
  • The point at which there is no increase in the
    gas particles or no increase in the liquid
    particles is called equilibrium.

43
  • Molecules sublimating Molecules condensing

44
Le Chateliers Principle
  • If stress is applied to a system at equilibrium,
    the system will tend to readjust so that the
    stress is reduced. (for chemical reactions)
  • The stress may be a temperature, pressure,
    concentration of products or reactants, or other
    external factors.
  • Equilibrium shifts occur after a stress has been
    applied.
  • Stress ? Shift ? New equilibrium

45
H2O (l) H2O (s)
  • If we add heat to this system at equilibrium,
    what will happen?
  • Thats right! We call that an equilibrium shift
    to the left because we will form more of the
    substance on the left side of the yield sign.
  • If we place this system in a freezer, what type
    of equilibrium shift will occur?

46
Shifts
  • Shift to the left to make more reactants.
  • Everything on the left increases, everything on
    the right decreases.
  • Shift to the right to make more products.
  • Everything on the right increases and reactants
    decrease.

47
Stresses to Equilibrium
  • Only gases and aqueous substances can induce a
    stress.
  • Addition or elimination of solids or liquids will
    not affect equilibrium. They can however,
    increase or decrease when other stresses occur.

48
Changing Concentrations and the Effect on
Equilibrium
  • N2(g) H2(g) ?? NH3(g)
  • What shift will occur when the NH3 is
    increased?
  • What happens to N2, and H2?
  • What shift will occur when some H2 is taken out
    of the system?
  • What happens to N2, and NH3?
  • What shift will occur when some NH3 is taken out
    of the system?
  • What happens to N2, H2, and NH3?

49
Temperature
  • If temp or energy (J, kJ, Cal, kcal) is a
    reactant then the reaction is endothermic
    (requires heat and has a positive value)
  • Treat temp as any reactant.
  • If temp or energy is a product then the reaction
    is exothermic (releases heat and has a negative
    value)
  • Treat as any product.

50
  • Heats of reactions
  • Seen as ?H where the ? refers to change
  • A positive () ?H is endothermic and energy
    should be written as a reactant.
  • A negative (-) ?H is exothermic and energy should
    be written as a product.
  • EX ?H -53 kJ or ?H 53 kJ
  • ( or let you know which side to place the
    energy on)

51
  • N2(g) H2(g) ?? NH3(g) heat
  • What shift will occur when heat is added to the
    system?
  • What happens to N2, H2, and NH3?
  • What shift will occur when some NH3 is taken out
    of the system?
  • What happens to temperature?
  • What shift will occur when heat is taken away
    from the system?
  • What happens to N2, H2, and NH3?

52
Pressure
  • Applies to gases only.
  • If increasing pressure, shift will occur in the
    direction of the least moles.
  • EX 2 mol ? 1 mol right shift
  • If decreasing pressure, shift will occur in the
    direction of the most moles.
  • EX 2 mol ? 1 mol left shift
  • If moles are equal then there is no shift and
    everything remains the same.

53
  • N2(g) H2(g) ?? NH3(g) heat
  • What shift will occur when pressure is increased?
  • What happens to N2, H2, NH3, and
    temperature?
  • What shift will occur when pressure is reduced?
  • What happens to N2, H2, NH3, and
    temperature?

54
K value
  • The equilibrium constant (K) is a unitless
    quantity characterizing a chemical equilibrium in
    a chemical reaction which is a useful tool to
    determine the concentration of various reactants
    or products in a system where chemical
    equilibrium occurs.
  • It is only affected by temp changes.
  • Use the general equation K products
  • reactants

55
  • Note that during a right shift the concentration
    of products increases while the concentration of
    reactants decrease. Mathematically, this
    increases the K value.
  • During a left shift the concentration of the
    reactants increase while the concentration of the
    products decrease, decreasing the K value.

56
  • Note that during a right shift the concentration
    of products increases while the concentration of
    reactants decrease. Mathematically, this
    increases the K value.
  • K products ? numerator reactants
    gets larger
    making K larger

57
  • During a left shift the concentration of the
    reactants increase while the concentration of the
    products decrease, decreasing the K value.
  • K products
  • reactants ? denominator gets larger
    making K smaller.

58
  • A large K value, over 1, indicates that the K
    value favors products.
  • A small K, less than 1, indicates that K value
    favors reactants.

59
If K gt 1, the reaction is product-favored. There
are more products than reactants at equilibrium
60
If K lt 1, the reaction is reactant-favored. There
are more reactants than products at equilibrium.
61
  • N2(g) H2(g) ?? NH3(g) heat
  • What shift will occur when heat is taken out of
    the system?
  • What happens to the K value?
  • What shift will occur when some NH3 is taken out
    of the system?
  • What happens to the K value?
  • What shift will occur when heat is added to the
    system?
  • What happens to the K value?

62
Reaction Rate (Keq)
  • To express the reaction rate (Keq) for the
    equilibrium reactions, we use the general
    equation
  • mA nB ? sP rQ
  • Keq PsQr
  • AmBn

63
  • Only the phases of matter (g) and (aq) are placed
    into a K expression.
  • Disregard (s) and (l) substances.

64
  • H2 (g) I2 (g) 2HI (g)
  • The K for this reaction is
  • Keq HI2
  • H2 I2
  • If we had values for these molecules, we would
    place the values in this equation.

65
H2 (g) I2 (g) 2HI (g)
  • EX Calculate Keq if
  • H2 0.0050 moles
  • I2 0.0035 moles
  • HI 0.0050 moles
  • Keq 0.00502
    0.0050 0.0035

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