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Title: Chapter 7 — States of Matter and Changes of State


1
Chapter 7 States of Matter and Changes of State

2
What are the States of Matter?
  • Solids
  • Liquids
  • Gases

3
Solids
  • Solids have a definite volume and a definite
    shape
  • Molecules (or atoms or ions) are closely packed
    together
  • INTERMOLECULAR FORCES are strong enough to hold
    molecules (etc.) rigidly in place with respect to
    each other

4
Liquids
  • Liquids have a definite volume but no definite
    shape
  • Liquids have the ability to flow
  • Molecules are very close together, but can flow
    past each other
  • Intermolecular forces are
  • Strong enough to hold molecules in a condensed
    phase
  • Not strong enough to prevent molecules from
    sliding past each other

5
Gases
  • Gases have neither definite shape nor definite
    volume
  • Gases are ideally independent molecules
  • Intermolecular forces are essentially zero
    between gas molecules

6
Kinetic - Molecular Theory of Matter
  • Gases are well-described by gas laws
  • No such laws exist for solids or liquids.
  • Particles that make up solid and liquid samples
    are touching therefore, solids liquids are not
    easily compressible (both are called the
    condensed states of matter).
  • K M Theory of matter attempts to describe all the
    states of matter and the conversion between
    states
  • To understand the states of matter, we need to
    consider the molecules comprising matter and how
    they interact

7
Ionic Bonds (b/w metals and nonmetals)
  • An ionic bond results from Coulombic attraction
    between oppositely charged ions (like charges
    repel each other and opposite charges attract
    each other).
  • Ions are not molecules
  • Ionic compounds are (almost) always solids under
    normal conditions
  • Metals give up 1 or 2 electrons to achieve a
    noble-gas-like electron configuration (cations)
  • Nonmetals acquire 1 or 2 electrons to achieve a
    noble-gas-like electron configuration (anions)

8
Covalent Bonds (b/w nonmetals)
  • Covalent bonds result from sharing one or more
    pairs of electrons
  • The OCTET RULE states atoms want to be a Noble
    gas (filled valence shell with 8 electrons).
  • Or at least have an noble gas electron
    configuration
  • This is achieved be giving each atom access to 8
    electrons
  • Every atom tends to add, remove, or share
    electrons so as to end up with eight valence
    electrons.
  • Valence electrons electrons in the highest
    energy (outermost) shell (valence shell).
  • Species with the same electron configurations are
    termed isoelectric (every atom strives to be
    isoelectric to the nearest noble gas).
  • When nonmetals combine, neither one can force its
    partner to become an ion so each nonmetal atom
    has access to as many electrons as the nearest
    noble gas (sharing electrons)

9
Valence Bond Theory
  • A covalent bond results from overlap of two
    electron clouds after bringing atoms close enough
  • e- cloud b/w 2 nuclei will shield them from
    each other reducing repulsion
  • This allows the electron on atom 1 to spend time
    around the nucleus of atom 2 and vice versa
    (e.g., 2 hydrogen atoms overlapping)
  • Electron on atom 1 can zip over and bask in the
    positive glow of the nucleus on atom 2 and
    viceversa http//www.chemistryland.com/CHM130W
    /11-Bonds/bonds.html

e1
1

e2
10
Bond Length and Bond Energy
(atoms too close to each other, repulsion b/w
positive nuclei, rapid rise of energy, decreased
stability
(atoms are separated so no covalent bond)
(amount of energy needed to brake the bond)
(energy is at a minimum and stability at maximum)
11
Lewis Dot Structures
  • G.N. Lewis developed the theory of covalent
    bonding
  • Structures showing covalent bonds are called
    Lewis structures
  • Each line represents a shared pair of electrons
  • Lone pairs of electrons are shown by a pair of
    dots

12
Drawing Lewis Structures http//www.whfreeman.com/
chemicalprinciples/content/instructor/sampleproble
ms.pdf
  • Decide on atom connectivity
  • Hydrogen is frequently bonded to oxygen
  • Oxygen is rarely the central atom
  • Count the total number of valence electrons
  • An atoms number of valence electrons is equal to
    its group number
  • Connect the atoms with single bonds
  • A single bond is one shared pair of electrons
  • Use lone pairs and/or multiple bonds to give each
    atom an octet of electrons

13
How many valence electrons are in the following?
  • N
  • Nitrogen is in group 5A. It has five valence
    electrons.
  • H2S
  • Hydrogen has one valence electron, and sulfur has
    six. The total for the molecule is 2(1)  6  8.
  • CO32
  • Carbon has four valence electrons oxygen has
    six then two for the charge. 4  3(6)  2  24.
  • NH4
  • Nitrogen has five valence electrons hydrogen has
    one, minus one for the charge. 5  4(1)  1  8.

14
Example
  • H2O
  • Atom connectivity H-O-H
  • Total Valence electrons?
  • 8 (6 electrons from oxygen, group 6A 1
    electron per hydrogen, group 1A)
  • 6 2 8 (divide by 2 4 electron
    pairs)
  • -use 2 of the pairs to give single bonds
    b/w the oxygen and each hydrogen.
  • -place two remaining electron pairs on oxygen
    as lone pairs of electrons.
  • CO2
  • Atom connectivity O-C-O
  • Total Valence electrons?
  • 16 (6 electrons per oxygen, group 6A 12
    4 electrons from carbon, group 4A)
  • 12 4 16 (divide by 2 8 electron
    pairs)
  • -use 4 of the pairs to give double bonds
    b/w the carbon and each oxygen.
  • -place 2 of the remaining electron pairs on one
    oxygen as lone pairs of electrons and the other
    2 pairs on the other oxygen.

15
Example (Isoelectric Species Triple Bonds)
  • CO
  • CN
  • Cyanide is CN-, which means that there are 4
    shell electrons for carbon, 5 for nitrogen and
    one more to make the anion. This makes a total of
    10 electrons. Nitrogen naturally wants to make
    three bonds, carbon wants four, but two atoms
    can't have a quadruple bond between them. So they
    have three covalent bonds, which uses up 6
    electrons. This leaves four electrons, which
    divides up nicely into two lone pairs (one for
    each atom). So the lewis dot structure is

16
VSEPR Theory
  • Valence Shell Electron Pair Repulsion Theory
    (VSEPR)
  • Atoms are bound into molecules by shared pairs of
    electrons
  • Electrons repel each other (like charges repel
    each other)
  • Therefore, the groups bonded to a central atom
    try to get as far apart from each other as
    possible
  • The goal is to minimize electron pair repulsions
    around a given central atom

17
Molecular Geometry describes spatial
arrangement of the atoms in a molecule
(intermolecular forces are caused by electrons
are arranged in a molecule)
Linear Geometry
  • AX2 (A is central Atom Xs are stick-on
    groups bonded to central Atom) e.g., CO2
  • Bond angle b/w C O 180
  • HCN
  • XeF2

18
Bent
  • AX2E2 (A is central Atom Xs are stick-on
    groups Es are lone pairs of electrons on the
    central Atom
  • Bond angle 105
  • H2O
  • SO2

19
Trigonal Planar
Urea (NH2)2CO
  • AX3 (no lone pairs of electrons)
  • Bond angle 120
  • BF3
  • CO32

20
Pyramidal
  • AX3E (three bonding pairs and one lone pair all
    point to the vertices of a tetrahedron but we
    only consider the bonding pairs in molecular
    geometry)
  • Bond angle 108
  • NH3

21
Tetrahedral
  • AX4 (central Atom with 4 stick-on groups)
  • Bond angle 109.5
  • CH4 (methane)

CHCL3 (chloroform)
22
Electronegativity
  • Linus Pauling was the first to develop an
    electronegativity scale
  • Electronegativity is the tendency of an atom in a
    molecule to attract shared electrons to itself.1
    (Atoms pulling of electrons to themselves)
  • Fluorine is the most electronegative element (EN
    4.0)
  • The closer an atom is to fluorine, the more
    electronegative it is (O is more electronegative
    than N Cl is more electronegative than Br)
  • 1Chemistry, 7th Edition, Zumdahl Zumdahl

23
Polar Covalent Bonds
  • If two atoms of identical electronegativity are
    bonded together, the bond is non-polar (no
    electron hot spots)
  • If two atoms of different electronegativity are
    bonded together, the bond is polar, and the
    electrons spend more time around the more
    electronegative atom
  • This creates partial charges (electron hot
    spots)
  • The greater the difference in EN between two
    atoms, the more polar the bond (the bonding
    electrons spend more time around he more
    electronegative atom since they are pulled to
    such atom therefore the more time the shared
    electron pair spends on that atom)
  • The limiting example of this is the ionic bond
    (dont confuse ionic bonds with polar covalent
    bonds!)

24
Example
  • The bond in hydrogen is non-polar
  • The bond in hydrogen chloride is polar
  • Chlorine is more electronegative than hydrogen so
    the bonding pair of electrons spends more time
    around Cl therefore, the Cl end of the molecule
    has a partial negative charge (delta) and the
    hydrogen end of the molecule has a partial
    positive charge (delta-)

(direction in which the bond is polarized)
25
Molecular Polarity
  • Bond dipoles are vectors
  • The vectoral sum of the bond dipoles gives the
    molecular dipole

26
Example
  • Predict the molecular dipole for water

Bond dipole vectors act together to give a net
molecular dipole
27
Example
  • Predict the molecular dipole for carbon dioxide

Bond dipole vectors cancel each other out to
resulting in a nonpolar molecule (even though the
molecule has polar bonds)
28
Intermolecular Forces
  • These are attractive forces between molecules or
    atoms or ions
  • Immensely important
  • These forces hold DNA molecules in a helix and
    are the mechanism for DNA transcription
  • There are three main varieties of IM forces
  • Dipolar, hydrogen bonding, and London forces

29
Dipole Dipole Attraction (aka dipolar
attraction)
  • This is the attraction between the opposite
    (partial) charges of polar molecules

(dipolar attractive force)
30
London Forces
  • Also called Van der Waals forces, these are
    created by instantaneous dipoles
  • London forces are much weaker than either
    dipole-dipole or H-bonding (ubiquitous)
  • London forces get stronger with larger
    atoms/molecules because larger molecules have
    more electrons.

31
London Forces Between Helium Atoms
For the merest fraction of time, there is a
dipole-dipole attraction between the atoms.
32
Hydrogen Bonding (special type of dipolar
interaction)
  • This is generally stronger than dipolar
    attractions (strongest intermolecular forces)
  • Hydrogen bonding is the attraction between a
    hydrogen bonding directly to an F, O, N (all of
    these 3 are highly electronegative). Hydrogen
    bonding is very important intermolecular force
    (page 166)

33
Ion Dipole Attraction
  • This is the attraction between an ionic charge
    and a polar molecule
  • This attraction allows ionic solids to dissolve
    in water
  • The strength of this force varies widely and
    depends on the magnitude of the dipole moment of
    the polar species and the size of the ion

34
A Sodium Ion and a Chloride Ion Hydrated by
Water Molecules (very polar)
Na
Cl
35
Effects of Intermolecular Forces
  • More intermolecular forces mean
  • Higher boiling and melting points
  • Higher heats of fusion and vaporization
  • Lower vapor pressure
  • More viscous liquids
  • IM Forces also affect solubility
  • like dissolves like

36
Explain the trend in boiling points of the
halogens(temperature at which the vapor pressure
is equal to the ambient pressure (1 atm)
37
Which of These Will Be Soluble in Water?
  • HCl(g) (polar)
  • O2(g) (nonpolar)

38
Which of These Is a Gas?
  • N2O or NaN3 ?
  • Cl2 or I2 ?

39
Changes of State
40
Vapor Pressures
  • The most energetic molecules in a liquid have
    sufficient kinetic energy to escape into the gas
    phase
  • Once the molecules are free as gases, they exert
    a pressure
  • This is called the vapor pressure
  • How does vapor pressure depend on temperature?
  • Vapor pressure INCREASES with INCREASING
    temperature

41
Changes of State
  • Solid to liquid melting
  • Liquid to solid freezing
  • Liquid to gas vaporization (evaporation)
  • Gas to Liquid condensation
  • Solid to gas sublimation
  • Gas to solid deposition

42
Energy Changes and Changes of State
  • Imagine recording the temperature of an 18
    gram(i.e.,1.0 mole) sample of ice at -40C as
    heat is added

43
Heating Curve for 1 Mole of Water
Water is boiling Heat of vaporization
Ice is melting Heat of fusion
44
Molar Enthalpy of Vaporization (?Hvap )
  • ?Hvap is the heat required to convert one mole
    of liquid to a gas at its normal boiling point
  • ?Hvap is an inherently endothermic process
    (amount of energy that most be added to the
    sample for the phase transition to occur)
  • ?Hvap has units of energy/quantity, e.g.,
    kJ/mole
  • ?Hvap represents the energy needed to break
    intermolecular forces and allow molecules to
    escape into the gas phase

45
Molar Enthalpy of Fusion
  • ?Hfus is the heat required to convert one mole
    of solid to a liquid at its normal melting point
  • ?Hfus is an inherently endothermic process
  • ?Hfus can have units of kJ/mole
  • ?Hfus represents the energy needed to break down
    intermolecular forces and allow molecules to
    slide around the liquid phase

46
Question
  • Why does steam at 100C cause more severe burns
    than water at the same temperature?
  • a) When water at 100oC touches your skin, it
    begins to drop its temperature immediately as the
    water cools to your skin temperature.
  • b) When steam at 100oC touches your skin, it
    remains at that temperature while releasing the
    heat of vaporization onto your skin as the gas
    converts to a liquid. Then you have water at
    100oC as a above.

47
Some Heats of Vaporization Fusion
48
Distribution of Energy
  • In a sample of material, the kinetic energies of
    the molecules (etc.) follow a Boltzmann
    Distribution

49
Dynamic Equilibrium
  • At the surface of a liquid, the most energetic
    molecules can escape from the IM forces into the
    gas phase
  • Gas molecules near the surface of a liquid can be
    captured by IM forces into the liquid state
  • When there is a balance between vaporization and
    condensation, a state of dynamic equilibrium
    exists

50
Vapor Pressure
  • Molecules can escape from the surface of a liquid
    into the gas phase
  • The gaseous molecules exert a pressure, call the
    vapor pressure
  • The vapor pressure of a liquid increases with
    increasing temperature
  • This is quantitatively described by the
    Claussius-Claperyon equation

51
Question
  • Use the Boltzmann distribution to explain why
    vapor pressure increases with increasing
    temperature

Tlow
Molecules
Kinetic Energy
52
Vapor Pressure of Water and Ethanol
53
Example
  • Many of the terms in the CC equation are
    constants, so we can recast the equation into a
    simpler form
  • Log P A B/T
  • For enflurane, A 7.967 and B 21678.4. What is
    the vapor pressure of enflurane at 25C and 35C

54
Volatility
  • Volatility is the tendency of a liquid to
    evaporate
  • Does a more volatile liquid have a higher or
    lower vapor pressure?

55
Boiling Point
  • The boiling point is the temperature at which the
    vapor pressure of a liquid equals the ambient
    pressure
  • The normal boiling point is the temperature at
    which the vapor pressure of a liquid equals
    exactly 760 torr

56
Question
  • How do BP and VP relate to IM forces?

57
Evaporation Has a Cooling Effect
  • How does temperature relate to kinetic energy?
  • Which molecules are most likely to escape IM
    forces to the gas phase?
  • What happens to average kinetic energy?

58
Phase Diagrams
  • A phase diagram shows combined effects of
    temperature and pressure on the state of matter
  • Key points
  • Equilibrium lines between states
  • Triple point
  • Critical point

59
Phase Diagram of Water
liquid
solid
pressure
gas
temperature
60
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