Chapter 7 — States of Matter and Changes of State

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Chapter 7 States of Matter and Changes of State

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What are the States of Matter?

• Solids

• Liquids

• Gases

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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

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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

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Gases

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

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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

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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)

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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)

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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
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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)
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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

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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

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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.

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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.

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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

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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

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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

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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

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Trigonal Planar
Urea (NH2)2CO

• AX3 (no lone pairs of electrons)
• Bond angle 120
• BF3
• CO32

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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

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Tetrahedral

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

CHCL3 (chloroform)
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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

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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!)

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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)
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Molecular Polarity

• Bond dipoles are vectors
• The vectoral sum of the bond dipoles gives the
  molecular dipole

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Example

• Predict the molecular dipole for water

Bond dipole vectors act together to give a net
molecular dipole
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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)
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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

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Dipole Dipole Attraction (aka dipolar
attraction)

• This is the attraction between the opposite
  (partial) charges of polar molecules

(dipolar attractive force)
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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.

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London Forces Between Helium Atoms
For the merest fraction of time, there is a
dipole-dipole attraction between the atoms.
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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)

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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

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A Sodium Ion and a Chloride Ion Hydrated by
Water Molecules (very polar)
Na
Cl
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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

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Explain the trend in boiling points of the
halogens(temperature at which the vapor pressure
is equal to the ambient pressure (1 atm)
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Which of These Will Be Soluble in Water?

• HCl(g) (polar)
• O2(g) (nonpolar)

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Which of These Is a Gas?

• N2O or NaN3 ?
• Cl2 or I2 ?

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Changes of State
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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

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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

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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

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Heating Curve for 1 Mole of Water
Water is boiling Heat of vaporization
Ice is melting Heat of fusion
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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

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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

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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.

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Some Heats of Vaporization Fusion
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Distribution of Energy

• In a sample of material, the kinetic energies of
  the molecules (etc.) follow a Boltzmann
  Distribution

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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

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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

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Question

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

Tlow
Molecules
Kinetic Energy
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Vapor Pressure of Water and Ethanol
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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

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Volatility

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

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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

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Question

• How do BP and VP relate to IM forces?

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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?

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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

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Phase Diagram of Water
liquid
solid
pressure
gas
temperature
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Thank you!