Liquids and Solids: AP Chemistry Unit 10

Sections

Intermolecular Forces Liquid state Solid Structures Metal Structures Carbon and Silicon Networks Molecular Solids Ionic Solids Vapor pressure and State Change Phase Diagrams

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

H2O(s) H2O(l) H°Δ fus = 6.02 kj/mol

H2O(l) H2O(g) H°Δ 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

Electrons shared within the molecule are called intramolecular bonding.

In the condensed states of matter the attraction between molecules are called 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 water 934kj to break the O-H bond

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….

In a condensed state, …

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Dipole-Dipole Forces continued…

Some dipole-dipole forces are unusually strong. These usually form between H and another very electronegative atom.

o These strong attractions have a strong impact on melting points and boiling points.

Boiling Points of Covalent Hydrides

Hydrogen bonds

Hydrogen bonds are the strongest in the smallest and lightest of the covalent molecules. This is primarily due to two factors:

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Hydrogen bonds continued….

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.

London Dispersion Forces

Even without dipoles, molecules exert forces on each other.

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.

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.

o

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London Dispersion Forces continued…

Polarizability is the ease at which an electron cloud can be distorted into a temporary dipole.

This also applies to molecules like H2, CH4, CCl4 and CO2; smaller molecules, but nonpolar.

The Liquid State 10.2

Liquid Characteristics

lack of rigidity

low compressibility

high density

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.

Capillary Action

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Capillary action is the spontaneous rising of a liquid in a narrow tube. This action is due to two forces

cohesive forces-

adhesive forces –

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.

A nonpolar

Meniscus: Water vs. Mercury

Viscosity

Viscosity is a fluids resistance to flow.

Example: glycerol is highly viscous because of its ability to create hydrogen bonds.

Viscosity continued…

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Example: Gasoline has carbon chains from 3-8C long and is nonviscous. Grease is 20-25C long and is very viscous.

Introduction to Structures and Types of Solids 10.3

Types of Solids

Crystalline solids Amorphous solids

Crystalline Solids

Crystalline solids have a regular arrangement of components at a microscopic level and produce beautiful, characteristic shapes of crystals:

The positions of components are usually represented by a lattice.

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.

X-ray Analysis of Solids

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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.

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.

A diffractometer is a computer-controlled instrument used for carrying out the X-ray analysis of crystals

o It 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.

The Bragg equation combines trigonometry and physics to determine the atomic spaces between crystals:

n = 2d sin λ θ d is the distance between atoms and is the angle of incidence and reflectionθ

of the light. n is an integer, most commonly 1. (n is usually given)

X-ray Analysis of Solids continued….

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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 = 19.3°. Assuming n=1, θcalculate the distance d between the planes of atoms producing this reflection

Types of Solids

Ionic solidso

Molecular solidso

Atomic solidso

Atomic Solids

Atomic solids are broken down into subgroups depending on the bond that exists in the solid:

Metallic solidso

Network solidso

Atomic Solids continued….

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Group 8A solidso

Classification of Solids

Structure and Bonding in Metals 10.4

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

Metallic Crystals

Metallic crystals can be pictured as containing spherical atoms packed together that can be bonded to each other equally in all directions.

Closest Packing

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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.

o

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.

Closest Packing: Hexagonal

Closest Packing: Cubic

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Knowing the net number of atoms in a particular unit cell is important for many applications involving solids.

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.

Face – Centered Cubic Unit Cell

Cubic Substances

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Metals that form cubic closest packed solids are:

Hexagonal Substances

Metals that form hexagonal closest packed solids are:

Other Metal 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.

Example Problem: Silver crystallizes in a cubic closest packed structure. The radius of a silver atom is 144pm. Calculate the density of solid silver?

Bonding Models for Metals

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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.

Electron Sea Model

Metal cations ‘swim’ in a sea of valence electrons that are mobile and shared.

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.

Band Model continued…

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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–

Interstitial alloy –

Substitutional Alloy

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Example: brass: 1/3 of copper metal atoms are replaced by zinc atoms

Sterling silver-

Pewter-

Plumbers solder –

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.

The amount of carbon directly affects the properties of steel:

Mild steels-

Medium steels-

High-carbon steel –

Mixed Alloys

Some 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.

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.

Carbon

Two most common forms of carbon are diamond and graphite. They are typical network solids.

Diamond

Graphite

Diamond

Each 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.

The application of 150,000 atm at 2800°C can break graphite bonds and rearrangement into a diamond structure.

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.

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

Silica

The 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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Silica continued…

When 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.

Glass

The properties of glass can vary greatly depending on the additives.

Common glass results when…

B2O3 produces borosilicate glass which…

K2O produces ….

Silicates

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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.

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.

Clays

Clay comes from the weathering of feldspar, an Aluminosilicate (Na2O/K2OŸAl2O3Ÿ6SiO2). 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.

Ceramics

Ceramics 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.

Ceramics continued…

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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.

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.

N-type Semiconductor

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.

Energy Level Diagrams for N-type and P-type Semiconductors.

P-N Junction

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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.

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.

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.

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 Solids 10.6

Network Solids

Sometimes network solids can be considered to be one giant molecule or have large discrete molecular units in a lattice-type position. These molecules have strong bonds within the molecules but relatively weak between the molecules.

Network Solids continued…

When molecules do have dipole moments, their intermolecular forces are significantly greater, especially when hydrogen bonding is possible.

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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.

Ionic Solids 10.7

Ionic Solids

Ionic solids are stable, high melting substances held together by the strong electrostatic forces that exist between oppositely charged ions.

Most 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).

There are three types of holes in closest packed structures:

1. Trigonal

2. Tetrahedral

3. Octahedral

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Ionic Solids continued...

The holes increase in size in the order:

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.

Example: 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.

Example: 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.

Example: Determine the net number of Na+ and Cl- ions in the sodium chloride unit cell.

Example: Classify each of the following substances according to the type of solid it forms:

gold carbon dioxide lithium fluoride krypton

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Types and Properties of Solids

Vapor Pressure and Changes of State 10.8

Vaporization

Vaporization, or evaporation, is the process of liquid molecules escaping the liquid’s surface and forming a gas.

o Water has strong intermolecular forces and this increases the energy required to vaporize. Also making it a great coolant.

o The symbol for this is HΔ vap.

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 Pressure continued…

Molecules 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.

Rates of Condensation and Evaporation

Vapor Pressure

The pressure of the vapor present at equilibrium is called…

A simple barometer can measure the vapor pressure of a liquid.

Liquid 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 Pressure contd

When the system reaches equilibrium, the vapor pressure can be determined from the change in the height of the mercury column

The vapor pressure of liquids vary widely. Liquids with high vapor pressures are said to be volatile. They evaporate rapidly in an open dish.

o Liquids with strong molecular forces have relatively low vapor pressures because it takes so much energy for the molecules to escape.

o In general, substances with large molar masses have relatively low vapor pressures due of large dispersion forces.

o In order to break intermolecular forces, a sufficient amount of kinetic energy is needed.

o As temperature of the liquid increases, so does kinetic energy of the liquid.

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Vapor Pressure of Water

The nature of the temperature, vapor pressure relationship is quadratic. Pvap vs. 1/T (Kelvin) gives a direct relationship.

Vapor pressure equation:

ln (PvapT 1PvapT 2 )=DH vap

R ( 1T2−1T 1 )

R is the universal gas constant (8.3145 J/K), and C is a constant characteristic of a given liquid (y=intercept). In means natural logarithm.

Example: Using the graph, determine whether water or diethyl ether has the larger enthalpy of vaporization.

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Example: 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.

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.

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.

Heating Curve

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Heating Curve

The plateaus in the heating curve represent the positions of phase change.

At the melting point,

At the boiling point,

o Note that changes of state are physical changes, although intermolecular forces have been broken, no chemical bonds have been broken.

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.

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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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.

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.

Melting and Boiling…

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.

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 …

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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.

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.

o 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.

Phase Diagrams

Phase Diagrams

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.

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The solid/liquid boundary has a negative slope. o

o 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.

Example 1: Pressure is 1 atm. Water moves through the changes of state according to the vapor pressure at the corresponding temperatures.

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 3: Pressure 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.

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.

Critical Point

Together, the critical temperature and critical pressure define at the critical point.

Critical Point continued…

The critical point for water is 374°C and 218 atm. Anything beyond this point, involves the intermediate “fluid” region.

Phase Diagrams: CO2

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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