Transcript Physiological Chemistry - University of Massachusetts Lowell

Physiological Chemistry
Chapter 5
States of Matter: Liquids and
Solids
Changes in State
• Changes in state are considered to be physical
changes
• During a change of physical state many other
physical properties may also change
• This chapter focuses on the important
differences in physical properties among
– gases
– liquids
– solids
Comparison of Physical
Properties of Gases, Liquids, and
Solids
Changes of State
• A change of state is the process whereby a substance
changes from one physical state to another (e.g., melting,
boiling)
– these changes can be effected by changing temperature
and/or pressure
• endothermic: heat is absorbed, which increases the
kinetic energy of the particles
– melting, evaporation and boiling, sublimation
• exothermic: heat is released, which reduces the
kinetic energy of the particles
– freezing, condensation, deposition
Changes of State
The Liquid State
• Compressibility: liquids are practically
incompressible
• Viscosity: a measure of a liquid’s resistance to
flow
– a function of both attractive forces between
molecules and molecular geometry
– complex or polar molecules tend to have
higher viscosity than simpler or nonpolar
molecules
– viscosity decreases as temperature increases
Surface Tension
• Surface tension: is a measure of the
attractive forces exerted among molecules
at the surface of a liquid
– Surface molecules are surrounded and
attracted by fewer liquid molecules than those
below
– Net attractive forces on surface molecules pull
them downward
• Surfactant: is a substance added which
decreases the surface tension (soap)
Vapor Pressure of a Liquid
• Place water in a sealed container
– Both liquid water and water vapor will exist in the
container
• How does this happen below the boiling point?
– the gas particles that escape the liquid phase are
called the vapor of the liquid
– they would not normally exist in the gas phase at the
temperature/pressure at which evaporation is
occurring
• Kinetic molecular theory: liquid molecules are in
continuous motion, with their average kinetic
energy directly proportional to the Kelvin
temperature
Temperature Dependence of Vapor
Pressure
• As temperature increases, the
fraction of molecules having
the activation energy
necessary for evaporation
increases (shaded area) due
to their higher kinetic energy
• Even though at cold
temperatures some
molecules can be converted
to vapor, molecules with
higher kinetic energy have
more propensity to escape
from the liquid phase
energy + H2O(l)  H2O(g)
Evaporation
• Evaporation is the process by which particles on the
surface of a liquid acquire enough energy to overcome the
attractive forces within the liquid, escaping the liquid phase
and entering the gas phase
– As liquid particles evaporate:
• the volume of the liquid decreases – if the liquid is in
an open container it eventually disappears!
• the temperature of the liquid decreases (as long as
the container does not exchange heat with the
environment) – evaporation has a COOLING effect!
– as particles evaporate they absorb heat from the
liquid
Boiling
• When the vapor pressure is high enough that it equals the
atmospheric pressure, bubbles begin to form throughout the
volume of the liquid and it begins to boil
– boiling is the process of evaporation in which the liquid is
transformed into a gas within the body of the liquid through
bubble formation
– the boiling point of the liquid is the temperature at which its
vapor pressure equals the surrounding atmospheric pressure
• at high altitudes, where the atmospheric pressure is lower
than sea level, water boils below 100°C
• inside a pressure cooker, where the pressure is much
higher than atmospheric, water boils above 100°C
Boiling Point
• Boiling point, and melting point, is dependent on
intermolecular forces
– the stronger the intermolecular (attractive) forces that
exist within the liquid, the more difficult for a liquid to
evaporate, thus it has a LOWER vapor pressure, and
HIGHER boiling point
• liquids that evaporate easily (due to weak
intermolecular forces) have HIGHER vapor
pressures (and LOWER boiling points) and are
said to be volatile
– polar molecules have higher b.p. than nonpolar
molecules of similar molar mass
Intermolecular Forces
• Physical properties of matter are explained in terms of their
intermolecular forces
• Different substances melt or boil at different temperatures
because the strength of the intermolecular forces that hold
particles together within matter varies among different
substances
• There are three major types of intermolecular forces that
affect and determine the behavior of matter
• Understanding the nature of these forces is of fundamental
importance in understanding the physical and chemical
properties of matter, including the multitude of complex
biological molecules that are responsible for life (e.g.,
proteins, carbohydrates, and RNA/DNA)
London Dispersion Forces
• These are the weakest of the
intermolecular forces
• They are not fixed but rather develop
momentarily and intermittently
between molecules as they approach
each other and their electron clouds
become briefly distorted and
instantaneously polarized
• The strength of London forces
depends on the ease of distortion of
the electron cloud, which increases
with size and molar mass
Dipole-Dipole Interactions
• These forces develop
between polar molecules
when the negative end of
one molecule is attracted to
the positive end of another
• The greater the polarity of
the molecules, the stronger
the attraction between them,
and the higher the melting
and boiling point of the
substance
Hydrogen Bonding
• Hydrogen bonding:
– is a special type of dipole-dipole attraction
– is a very strong intermolecular attraction causing
higher than expected b.p. and m.p.
• Requirement for hydrogen bonding:
– a hydrogen atom directly bonded to O, N, or F atom
qualifies a molecule to both donate and accept H in
H-bonding interactions
Hydrogen Bonds
• Based on its molar mass, the calculated boiling point of water
should be around −80°C, making life on Earth impossible!
The strength and number of hydrogen bonds in water is what
makes ice float on water, another reason why life on Earth
flourished and was able to survive the ice ages.
Examples of Hydrogen Bonding
• Hydrogen bonding has an
extremely important influence
on the behavior of many
biological systems
• Water forms FOUR hydrogen
bonds in the solid state, but on
average forms less than four in
the liquid state
• Water molecules in the solid
state are perfectly arranged in
a tetrahedral fashion with
respect to one another, making
ice less dense than liquid
water – thus ice floats on
water!
Hydrogen Bonds
Hydrogen Bonds
Hydrogen Bonds
The Solid State
• Particles are closely packed due to attractive forces
strong enough to resist motion
• Properties of solids:
– fixed shape and volume
– incompressible
– m.p. depends on strength of attractive force between
particles
– a solid may be crystalline (ordered array of particles)
or amorphous (disordered arrangement of particles)
Types of Crystalline Solids
1.
Ionic solids
– made up of positive and negative ions
– high m.p. and b.p.
– hard and brittle
– a common example is NaCl
2.
Covalent solids
– held together entirely by covalent bonds
– high m.p. and b.p.
– extremely hard
– an example is diamond
Types of Crystalline Solids
3. Molecular Solids
– made up of molecules held together by intermolecular
attractive forces
– usually soft with low m.p.
– volatile and poor electrical conductors
– a common example is ice
Types of Crystalline Solids
4.
Metallic Solids
– made up of metal atoms held together by “metallic bonds”
– these bonds are formed by the overlap of metal atomic
orbitals
– there are regions of high electron density, which are very
mobile and move freely from atom to atom
– this results in high conductivity
– examples include Ag and Cu
Sublimation of Solids
• Sublimation: process of conversion of molecules in the
solid state directly to the gaseous state
• An example is dry ice (solid carbon dioxide), which
converts directly to a gas at atmospheric pressure
• Solid water (i.e., ice) slowly sublimes, which is why snow
flurries don’t last for long even on a dry cold day
Sublimation and deposition of iodine (I2)