ap chem unit 10
TRANSCRIPT
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Liquids and Solids
AP Chem Unit 10
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SectionsIntermolecular Forces
Liquid state
Solid Structures
Metal Structures
Carbon and Silicon Networks
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SectionsMolecular Solids
Ionic Solids
Vapor pressure and State Change
Phase Diagrams
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States of Matter
When considering the three states of matter, properties of gases are strikingly different than solids and liquids. Liquids and solids share many similar characteristics
compressibility
density
intermolecular forces
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States of Matter
H2O(s)H2O(l) ΔH°fus = 6.02 kj/mol
H2O(l)H2O(g) ΔH°vap = 40.7 kj/mol
Water densities:25°C and 1atm .99707g/cm3
25°C and 1065 atm 1.046g/cm3
400°C and 1atm 3.26x10-4 g/cm3
400°C and 242 atm .157g/cm3
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Intermolecular Forces
10.1
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Intermolecular Forces
Electrons shared within the molecule are called intramolecular bonding.
In the condensed states of matter the attraction between molecules are called intermolecular forces.
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Intermolecular Forces
It is important to realize that when a molecule changes state, the molecule stays intact. The changes in state are due to the change in forces surrounding the molecule not from changes within the molecule.
40.7kj needed to vaporize water934kj to break the O-H bond
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Dipole–Dipole Forces
Dipole-dipole forces occur when polar molecule (molecules with dipole moments) electrostatically attract each other by lining up the positive and negative ends of the dipoles.
Dipole-dipole forces are about 1% as strong as a covalent or ionic bond and rapidly become weaker when distances between the dipoles increases. The distances in a gas make these attractions relatively unimportant
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Dipole-Dipole Forces
In a condensed state, molecules line up dipoles to minimize repulsions and maximize attractions.
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Dipole-Dipole Forces
Some dipole-dipole forces are unusually strong. These usually form between H and another very electronegative atom.
These are stronger due to the high polarity of the bond and the closeness of the dipoles between the atoms.These strong attractions have a
strong impact on melting points and boiling points.
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Boiling Points of Covalent Hydrides
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Hydrogen bonds
Hydrogen bonds are the strongest in the smallest and lightest of the covalent molecules. This is primarily due to two factors:
large difference in electronegativities
small size of the atoms allows for close dipole interactions.
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Hydrogen bonds
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Hydrogen Bonds and Organics
Methanol (CH3OH) and ethanol (CH3CH2OH) have much higher boiling points than would be expected from their molar masses because of the O-H bonds that produce hydrogen bonding.
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London Dispersion Forces
Even without dipoles, molecules exert forces on each other.
The forces that exist among noble gas atoms and nonpolar molecules are called London dispersion forces.
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London Dispersion Forces
Usually it is assumed that electron dispersion is uniform throughout the molecule, but this is not always the case.
Since the movements of the electrons around the nucleus are somewhat random, a momentary nonsymmetrical electron distribution can develop that creates a temporary dipolar arrangement of charge.
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London Dispersion Forces
This temporary change in polarity can, in turn, temporarily change the distribution of the neighboring molecule.
This phenomenon leads to an inter-atomic attraction that is relatively weak and short-lived, but can be significant in larger atoms at lower temperatures.larger atoms have more electrons and
increases the probability of a temporary dipole.
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London Dispersion Forces
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London Dispersion Forces
Polarizability is the ease at which an electron cloud can be distorted into a temporary dipole.
large atoms have a larger polarizability than smaller atoms
This also applies to molecules like H2, CH4, CCl4 and CO2; smaller molecules, but nonpolar.
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The Liquid State10.2
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Liquid Characteristics
lack of rigidity
low compressibility
high density
rounded droplets
capillary action
viscosity
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Rounded Droplets
Occur due to the intermolecular forces of the liquid. The liquid molecules are subject to attraction from the side and from below, so liquid tends to form a shape with the minimum surface area – sphere.
The resistance of a liquid to increase surface area is from the energy that it takes to overcome intermolecular forces. This resistance is called surface tension.
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Rounded Droplets
Molecules that are polar and have stronger intermolecular forces have stronger surface tensions.
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Surface Tension
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Capillary Action
Capillary action is the spontaneous rising of a liquid in a narrow tube. This action is due to two forcescohesive forces- the intermolecular
forces among the molecules.adhesive forces – the attractive
forces between the liquid and the container.
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Adhesive forces
Adhesive forces happen when bonds within the container have polar bonds
For example: glass has O atoms that carry a partial negative charge that attracts the partial positive charge of the hydrogen in water. This balance between the strong cohesive forces and the strong adhesive forces produce a meniscus.
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Adhesive forces
A nonpolar substance, such as mercury, has a convex meniscus because the cohesive forces are stronger than the adhesive forces.
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Meniscus: Water vs. Mercury
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ViscosityViscosity is a fluids resistance to flow.
liquids with strong cohesive forces tend to be highly viscous. Example: glycerol is highly viscous
because of its ability to create hydrogen bonds.
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Viscosity
Molecular complexity also can affect viscosity because they can become entangled in each other.Example: Gasoline has carbon chains
from 3-8C long and is nonviscous. Grease is 20-25C long and is very viscous.
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Introduction to Structures and Types of
Solids10.3
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Types of Solids
Crystalline solids
Amorphous solids
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Crystalline Solids
Crystalline solids have a regular arrangement of components at a microscopic level and produce beautiful, characteristic shapes of crystals:
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Crystalline Solids
The positions of components are usually represented by a lattice.
lattice is a three dimensional system of units repeating in a pattern. The smallest repeating unit of the lattice is called the unit cell.
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Three types of Crystalline Solids
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Amorphous Solids
Amorphous solids have considerable disorder in their structures.
Example: Common glass looks like a solution frozen in place. It has a rigid shape but a great deal of disorder within its structure.
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X-ray Analysis of Solids
The structures of crystalline solids are commonly determine by X-ray diffraction.
This type of diffraction occurs when beams of light are scattered as they go through spaces between substances. Light scatters when the size of the spaces are similar to the wavelength of light.
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X-ray Analysis of Solids
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X-ray Analysis of Solids
A single wavelength is directed at the crystal and a diffraction pattern is obtained. The diffraction pattern is a series of light and dark areas on a photographic plate from constructive and destructive interference from waves of light.
The diffraction pattern can then be used to determine the interatomic spacings.
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X-ray Analysis of Solids
A diffractometer is a computer-controlled instrument used for carrying out the X-ray analysis of crystalsIt rotates the crystal with respect to
the X-ray beam and collects the data produced by the scattering. The techniques have been refined to the point that very complex structures can be determined, such as large biological enzymes.
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X-ray Analysis of Solids
The Bragg equation combines trigonometry and physics to determine the atomic spaces between crystals:
nλ = 2d sin θ
d is the distance between atoms and θ is the angle of incidence and reflection of the light. n is an integer, most commonly 1. (n is usually given)
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X-ray Analysis of Solids
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Example Problem
X-rays of wavelength 1.54 Â were used to analyze an aluminum crystal. A reflection was produced at θ = 19.3°. Assuming n=1, calculate the distance d between the planes of atoms producing this reflection
2.33 Á
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Types of Solids
Ionic solidsionic solids are made of ions
Molecular solidsMolecular solids have small units of
covalently bonded molecules.
Atomic solidsAtomic solids are made of elements such
as carbon (graphite, diamond and the fullerenes), boron, silicon, and all metals.
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Fullerenes
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Types of Solids
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Atomic SolidsAtomic solids are broken down into
subgroups depending on the bond that exists in the solid:
Metallic solidsHas delocalized nondirectional covalent
bonding.
Network solidsatoms bond with strong directional
covalent bonding that lead to giant molecules and networks
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Atomic SolidsGroup 8A solids
noble gases are attracted to each other with London dispersion forces.
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Classification of Solids
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Structure and Bonding in Metals
10.4
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Metal Characteristics
Most of the properties that we see in metals is due to the nondirectional covalent bonding found in metal crystals.
High thermal conductivity
Electrical conductivity
Malleability
Ductility
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Metallic Crystals
Metallic crystals can be pictured as containing spherical atoms packed together that can be bonded to each other equally in all directions.
This arrangement is called closest packing.
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Closest Packing
The spheres pack in layers. Each sphere is surrounded by six others. These layers do not lie directly over those in the first layer, instead they fill the indentations of the layer below. The third layer is in the same position as the first. This is called aba arrangement.
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Closest Packing
The aba arrangement has the hexagonal unit cell and the resulting structure is called the hexagonal closest packed (hcp) structure.
The abc arrangement has a face-centered cubic unit cell and the resulting structure is called the cubic closest packed (ccp) structure. This has a repeating vertical placement every fourth layer.
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Closest Packing
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Closest Packing: Hexagonal
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Closest Packed: Cubic
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Closest Packing
Knowing the net number of atoms in a particular unit cell is important for many applications involving solids.
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Closest Packing
Example: A face centered cube (unit cell) is defined by the centers of the spheres on the cube’s corners. Therefore 8 cubes share a given corner sphere, so 1/8 of this sphere lies inside the unit cell. (8 corners x 1/8 sphere = 1sphere). The sphere at the center of each face is shared by two cubes. (6 faces x ½ sphere = 3 spheres). The total number of spheres for a face centered cube is 4.
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Closest Packing
Face – Centered Cubic Unit Cell
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Cubic Substances
Metals that form cubic closest packed solids are:
aluminum
iron
copper
cobalt
nickel
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Hexagonal Substances
Metals that form hexagonal closest packed solids are:
magnesium
zinc
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Other Metal Solids
Calcium and certain other metals can crystallize in either cubic or hexagonal solids.
Some metals, including many alkali metals, have structures that are characterized by a body-centered cubic (bcc) unit cell. In this structure, each sphere has 8 neighbors.
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Example Problem
Silver crystallizes in a cubic closest packed structure. The radius of a silver atom is 144pm. Calculate the density of solid silver?
10.6 g/cm3
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Bonding Models for Metals
In order to determine bonding for metals, one must account for the typical properties: durable, high melting point, malleable, ductile, and efficient in uniform conduction of heat and electricity in all directions.
These characteristics indicate that the bonds are strong and nondirectional. In other words, it is not easy to separate metal atoms but easy to move them.
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Electron Sea Model
Metal cations ‘swim’ in a sea of valence electrons that are mobile and shared.
This accounts for conduction and malleability and ductility.
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Electron Sea Model
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Band Model (MO Model)
In this model, the electrons are assumed to travel around the metal crystal in molecular orbitals formed from the valence atomic orbitals of the metal atoms.
When metals atoms interact, the large number of resulting molecular orbitals become more closely spaced and finally form a virtual continuum of levels, called bands.
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Band Model (MO Model)
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Band Model
The electrons in partially filled MO’s are mobile. These conduction electrons are free to travel throughout the metal crystal. The MO occupied by these conducting electrons are called conduction bands.
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Band Model
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Metal Alloys
An alloy is best defined as a substance that contains a mixture of elements and has metallic properties. There are two types of alloys:
Substitutional alloy– some of the host metal atoms are replaced by other metal atoms of similar size.
Interstitial alloy – is formed when some of the interstices (holes) in the closest packed lattice are occupied by smaller atoms.
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Metal Alloys
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Substitutional Alloys
Example: brass: 1/3 of copper metal atoms are replaced by zinc atoms
Sterling silver- 93% silver and 7% copper.
Pewter- 85% tin, 7% copper, 6% bismuth and 2% antimony.
Plumbers solder – 95% tin and 5% antimony
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Interstitial Alloy
Example: Steel contains carbon atoms in the holes of an iron crystal. The presence of the interstitial atoms changes the properties of the host metal. Iron is relatively soft, ductile and malleable, but when carbon (which forms directional bonds), is introduced into the crystal, it makes the iron bonds stronger and less ductile.
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Interstitial Alloy
The amount of carbon directly affects the properties of steel:
Mild steels- contains less than .2% carbon: nails chains and cables.
Medium steels- contain .2-.6% carbon: rails and structural steel
High-carbon steel – .6-1.5% carbon: springs, tools and cutlery.
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Mixed AlloysSome steels contain elements in addition
to iron and carbon. These are called alloy steels and are viewed as being mixed interstitial and substitutional alloys.
Bicycle frames are usually constructed from a wide variety of alloy steels.
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Carbon and Silicon Network Atomic Solids
10.5
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Network Solids
Many atomic solids contain strong directional covalent bonds to form a solid that might be viewed as a “giant molecule.” These materials are typically brittle and do not efficiently conduct heat and electricity. Two examples of these network solids are carbon and silicon.
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CarbonTwo most common forms of carbon are
diamond and graphite. They are typical network solids.
Diamond is the hardest naturally occurring substance.
Graphite is slippery, black and a conductor.
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DiamondEach carbon is surrounded by a tetrahedral
arrangement of other carbon atoms to form a large molecule. Diamond is an insulator not a conductor. Each carbon is sp3 hybridized with localized bonding and therefore does not conduct.
Diamonds are often used for industrial cutting implements.
The application of 150,000 atm at 2800°C can break graphite bonds and rearrangement into a diamond structure.
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Graphite
The structure of graphite is based on layers of carbon atoms arranged in fused 6 C rings. The unhybridized p orbitals allow for delocalized electrons and therefore conductivity.
Graphite is used as a industrial lubricant. Because graphite has strong bonds within the layers and weak bonding between the layer, the layers slide past one another readily.
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Carbon
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Carbon: Graphite layers
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Silicon
Silicon is an important constituent of the compounds that make up the earth’s crust. Silicon is to geology what carbon is to biology and is fundamental to most rocks, sands and soils found in the earth’s crust.
Carbon compounds typically have long strings of C-C bonds
Silicon compounds typically involve chains of Si-O bonds.
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SilicaThe fundamental
silicon-oxygen compound is silica, which has the empirical formula SiO2. The structure that is formed is based on a network of SiO4 tetrahedra with shared oxygen atoms rather than smaller SiO2 units.
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SilicaWhen silica is heated
above its melting point (1600°c) and cooled rapidly, an amorphous solid called glass results. Glass has a lot of disorder as opposed to the crystalline nature of quartz. Glass, also homogeneous, more closely resembles a very viscous solution than it does a crystalline solid.
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GlassThe properties of glass can vary greatly
depending on the additives.
Common glass results when substances like Na2CO3 are added to the silica melt.
B2O3 produce borosilicate glass which does not expand and contract during large temperature changes. (Pyrex)
K2O produces especially hard glass that can be ground into shapes for lenses and contacts.
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Glass
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Silicates
Compounds closely related to silica and found in most rocks, soils and clays are the silicates. Like silica, the silicates are based on interconnected SiO4 tetrahedra, but instead of a O/Si ratio of 2:1, the ratio is typically higher. This higher ratio tends to make silicon-oxygen anions.
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Silicates
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Silicates
Ceramics are typically made from clays (which contain silicates) and hardened by firing at high temperatures. They tend to be strong, brittle and heat and chemical resistant.
Ceramic is heterogeneous and contain two phases: minute crystals of silicates that are suspended in a glassy cement.
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ClaysClay comes from the weathering
of feldspar, an Aluminosilicate (Na2O/K2OAl2O36SiO2). This weathering produces kaolinite, that consists of tiny thin platelets of Al2Si2O5(OH)4. When dry these platelets cling together and lock into place; when wet they can slide over one another. During firing, these platelets bind and form a glass.
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CeramicsCeramics constitute one of the most important
classes of ‘high-tech” materials. Their stability at high temperatures and resistance to corrosion, make them an obvious choice for constructing jet and car engines.
Organoceramics are taking form by the addition of organic polymers to ceramics. This reduces some of the brittle nature of ceramics and allows them to be used for things such as flexible superconducting wire, microelectronic devices, prosthetic devices and artificial bones.
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Semiconductors
Elemental silicon has the same structure as diamond. The structure is different in that the energy gap between filled and empty MO’s is not as large and electrons can delocalize and make silicon a semi-conductor. At higher temperatures, more electrons get excited in the conduction bands and the conductivity of silicon increases.
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N-type Semiconductor
When small fraction of silicon atoms are replaced by arsenic atoms (one more valence electron), extra electrons become available for conduction and produce an n-type semi-conductor. These can conduct an electric current.
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P-type Semiconductor
When small fraction of silicon atoms are replaced by boron atoms (one less valence electron), an electron ‘vacancy’ is made. As electrons move, the fill the ‘hole’ and make a new one. This movement of electrons can therefore carry a current. This type of conductor (less electrons) is called a p-type semiconductor.
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Energy Level Diagrams for N-type and P-type Semiconductors.
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P-N Junction
Most important applications of semiconductors involve connection of a p-type and an n-type to form a p-n junction.
The red dots represent excess electrons in the n-type semiconductor and the white circles represent holes (electron vacancies.
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P-N Junction
At the junction a small number of electrons migrate from the n-type region into the p-type region. The effect of these migrations is to place a negative charge on the p-type region and a positive charge on the n-type region.
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P-N Junction
This charge buildup, called the contact potential or junction potential, prevents further migration of electrons. This transfer of electrons is therefore a ‘one-way’ transfer and under an external battery source will allow flow of electrons from the n to the p type regions.
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P-N Junction
When current is opposed it is said to be under reverse bias. When current flows easily, the junction is said to be under forward bias.
A p-n junction is a good rectifier, a device that produces a pulsating direct current from an alternating current.
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P-N Junction
When placed in a circuit where the current is constantly reversing, a p-n junction only transmits current under forward bias. Radios, computers and other electronic devices all use this rectifiers. This p-n junction revolutionized electronics.
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Molecular Solids10.6
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Molecular Solids
Sometimes Molecular solids can have large discrete molecular units in a lattice-type position. These molecules have strong bonds within the molecules but relatively weak between the molecules.
Common examples: Ice, dry ice (solid carbon dioxide), Sulfur (S8), Phosphorus (P4)
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Molecular
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Molecular Solids
When molecules do have dipole moments, their intermolecular forces are significantly greater, especially when hydrogen bonding is possible.
Water not only has polar bonds, a dipole moment, has hydrogen bonds, but it also can have a total of four hydrogens associated with every oxygen atom.
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Ionic Solids10.7
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Ionic Solids
Ionic solids are stable, high melting substances held together by the strong electrostatic forces that exist between oppositely charged ions.
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Ionic SolidsMost binary ionic solids can be explained
by the closest packing of spheres. Typically the larger ions, usually anions, are packed in one of the closest packed arrangements (hcp and ccp).
The smaller cations fit into the holes among the closest packed anions. This packing maximizes the electrostatic attractions among oppositely charged ions and minimizes the repulsion of like charges.
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Ionic SolidsThere are three
types of holes in closest packed structures:
1.Trigonal holes are formed by three spheres in the same layer
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Ionic SolidsThere are three
types of holes in closest packed structures:
2.Tetrahedral holes are formed when a sphere sits in the dimple of three spheres in an adjacent layer.
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Ionic SolidsThere are three
types of holes in closest packed structures:
3.Octahedral holes are formed between two sets of three spheres in adjoining layers of the closest packed structures.
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Ionic Solids
The holes increase in size in the order:
trigonal < tetrahedral < octahedral
The trigonal holes are so small that they are never occupied in binary ionic compounds. Tetrahedral and octahedral holes are occupied if the relative sizes of the ions allow.
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Ionic SolidsExample: Zinc Sulfide (ZnS) creates a ccp
structure. The Zn2+ has a radius of 70pm and the S2- ion has an ionic radius of 180pm. There are 4 spheres (atoms/anions) in a face-centered cubic unit cell and 8 tetrahedral holes. So only half of the holes in the ccp unit are filled with cations.
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Ionic SolidsExample: Sodium chloride can be described in
terms of a ccp structure. Na+ resides in octahedral holes. The locations of the octahedral holes in the face-centered cubic unit is marked by X. The number of spheres (anions) in the structure is the same number of octahedral holes. Since NaCl is a 1:1 binary compound. All octahedral holes are used.
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Example Problem
Determine the net number of Na+ and Cl- ions in the sodium chloride unit cell.
4 Na+ and 4 Cl-
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Example Problem
Classify each of the following substances according to the type of solid it forms:gold
carbon dioxide
lithium fluoride
krypton
metallic
molecular
ionic
8a
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Types and Properties of Solids
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Vapor Pressure and Changes of
State10.8
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Vaporization
Vaporization, or evaporation, is the process of liquid molecules escaping the liquid’s surface and forming a gas.
Vaporization is endothermic because energy is required to overcome the relatively strong intermolecular forces in the liquid. Water has strong intermolecular forces and
this increases the energy required to vaporize. Also making it a great coolant.
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Vaporization
The energy required to vaporize 1 mole of liquid at 1 atm is called the heat of vaporization or the enthalpy of vaporization.The symbol for this is ΔHvap.
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Vapor Pressure
Condensation is the process by which vapor molecules re-form a liquid.
The evaporation process occurs at a constant rate at a given temperature, and once an equilibrium has been reached, the rate of condensation will equal the rate of evaporation.
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Vapor PressureMolecules in a
given system are constantly escaping from and entering the liquid at high rate. However, there is not net change because the two opposite processes just balance each other.
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Rates of Condensation and Evaporation.
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Vapor Pressure
The pressure of the vapor present at equilibrium is called the equilibrium vapor pressure, or more commonly, the vapor pressure of the liquid.
A simple barometer can measure the vapor pressure of a liquid.
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Vapor PressureLiquid is injected at the bottom of the tube of
mercury and floats to the surface. A portion of the liquid evaporates at the top of the column, producing a vapor whose pressure pushes some mercury out of the tube.
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Vapor PressureWhen the system reaches equilibrium,
the vapor pressure can be determined from the change in the height of the mercury column
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Vapor PressurePatmosphere = Pvapor + PHg column
Pvapor = Patmosphere - Phg column
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Vapor Pressure
The vapor pressure of liquids vary widely. Liquids with high vapor pressures are said to be volatile. They evaporate rapidly in an open dish.
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Vapor Pressure
The vapor pressure of a liquid is principally determined by the size of the intermolecular forces in the liquid.Liquids with strong molecular forces
have relatively low vapor pressures because it takes so much energy for the molecules to escape.
In general, substances with large molar masses have relatively low vapor pressures due of large dispersion forces.
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Vapor Pressure
Vapor pressure increases significantly with temperature. In order to break intermolecular forces,
a sufficient amount of kinetic energy is needed.
As temperature of the liquid increases, so does kinetic energy of the liquid.
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Vapor Pressure of Water
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Vapor PressureThe nature of the temperature, vapor
pressure relationship is quadratic. Pvap vs. 1/T (Kelvin) gives a direct relationship.
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Vapor Pressure
Vapor pressure equation:
R is the universal gas constant (8.3145 J/K). In means natural logarithm.
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Example Problem
Using the graph, determine whether water or diethyl ether has the larger enthalpy of vaporization.
Ether has the smaller slope therefore smaller ΔHvap
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Example Problem
The vapor pressure of water at 25°C is 23.8 torr, and the heat of vaporization of water at 25°C is 43.9 kj/mol. Calculate the vapor pressure of water at 50°C.
93.7 torr
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Sublimation
Like liquids, solids have vapor pressures. When a solid sublimes, it goes directly from the solid to the gaseous state without passing through the liquid state.
Example: Dry Ice.
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Changes of State
Typically when a solid is heated, it will form a liquid and then boil to form a vapor. This process can be represented by a heating curve.
Temperature vs. time when energy is added a constant rate.
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Heating Curve
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Heating Curve
The plateaus in the heating curve represent the positions of phase change.
At the melting point, the temperature remains constant until the solid has completely changed to liquid.
At the boiling point the temperature remains constant as the added energy is used to vaporize the liquid
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Heating CurveThe energy change that occurs at the
melting point when a solid melts is called heat of fusion or enthalpy of fusion.Note that changes of state are physical
changes, although intermolecular forces have been broken, no chemical bonds have been broken.
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Enthalpy of Fusions and Melting Points
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Melting and Boiling
The melting and boiling points for a substance are determined by the vapor pressure of the solid and liquid states.
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Melting and Boiling
At 0°C the vapor pressure of ice is less than that of liquid. Vapor pressure of ice increases more rapidly than water.
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Melting and Boiling
A point is reached when the liquid and solids have identical vapor pressures. This is the melting point.
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Freezing Point
At a temperature at which the vapor pressure of the solid is greater than that of the liquid, the solid would sublime and the vapor would be added to the water.
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Freezing Point
At a temperature at which the vapor pressure of the solid is less than that of the liquid, the liquid would evaporate and the vapor would be added to the ice.
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Freezing Point
At a temperature at which the vapor pressures of the solid and liquid are identical, the vapor is in equilibrium. This is the freezing point of water.
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Melting Point and Boiling Point
The normal melting point of a liquid is the temperature at which the solid and liquid states have the same vapor pressure under conditions where the total pressure is 1 atmosphere.
The normal boiling point of a liquid is the temperature at which the vapor pressure of the liquid is exactly 1 atmosphere.
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Supercooled and Superheated
Changes of state do not always occur exactly at the boiling point or melting point.
Water can be supercooled below 0°C at 1 atm and remain in the liquid state. At some point the correct ordering of molecules occurs and ice forms, releasing energy in the exothermic process and bringing the temperature back up to the melting point.
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Supercooled and Superheated
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Supercooled and Superheated
Changes of state do not always occur exactly at the boiling point or melting point.
A liquid can also be superheated, or raised to temperatures above its boiling point, especially if it is heated rapidly. Boiling requires high-energy molecules to gather in the same vicinity for bubble formation. This may not happen at the boiling point.
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Supercooled and Superheated
Once a bubble does form, when a liquid is superheated, its internal pressure is greater than the atmospheric pressure. This bubble can burst before rising to the surface, blowing the surrounding liquid out of the container. This is called bumping and is a common experimental problem.
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Supercooled and Superheated
Boiling chips are often added to prevent bumping. These are bits of porous ceramic material containing trapped air that escapes on heating, forming tiny bubbles that act as ‘starters’ for the vapor bubble formation. This allows for smooth onset of boiling.
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Phase Diagrams10.9
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Phase Diagram
A phase diagram is a convenient way of representing the phases of a substance as a function of temperature and pressure. It shows which state exists at a given temperature and pressure.
Conditions for these phase diagrams are assumed to be a closed system and is plotted as temperature vs pressure.
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Phase Diagram
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Phase Diagram
The solid/liquid boundary has a negative slope. Melting point of ice decreases as
external pressure increases. This is different for most substances
other than water because the density of ice is less than that of liquid water at the melting point.
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Boiling Points of Water vs. Pressure
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Example 1
Pressure is 1 atm. Water moves through the changes of state according to the vapor pressure at the corresponding temperatures.
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Example 2
Pressure is 2 torr. Water will sublime at -10°C. This is when the vapor pressure of the ice is equal to the external pressure of 2 torr. Vapor pressure of liquid water is always greater than 2 torr and therefore will not form.
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Example 3Pressure is 4.58 torr.
When temperature reaches .01°C (273.16K), water reaches the triple point. Solid and liquid water have identical vapor pressures and all three states of water exist. This is the only condition in a closed system that allows this.
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Example 4
Pressure is 225 atm. Liquid water can be present at this temperature because of the high external pressure. As temperature increases, liquid gradually turns to vapor, but goes through a ‘fluid’ region.
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Example 4
The fluid region is neither true liquid or vapor. This unusual behavior occurs because the conditions are beyond the critical point for water.
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Critical PointThe critical temperature can be defined
as the temperature above which the vapor cannot be liquefied no matter what pressure is applied.
The critical pressure is the pressure required to produce liquefaction at the critical temperature
Together, the critical temperature and critical pressure define at the critical point.
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Critical PointThe critical point for water is 374°C
and 218 atm. Anything beyond this point, involves the intermediate “fluid” region.
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Phase Diagrams
The phase diagram for CO2 shows the liquid state does not exist at a pressure of 1 atm. The solid/liquid line has a positive slope, since the density of solid CO2 is greater than that of liquid CO2
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THE END!!