liquids and solids gas –low density –high compressibility –completely fills its container...

Liquids and Solids• Gas

– low density – high compressibility – completely fills its container

• Solid– high density– only slightly compressible– rigid– maintains its shape

Liquids and Solids• Liquids

– properties lie between those of solids and gases• H2O(s) --> H2O(l)

kJ/mol

• H2O(l) --> H2O(g) Hovap = 40.7 kJ/mol

– large value of Hvap suggests greater changes in structure in going from a liquid to a gas than from a solid to liquid

– suggests attractive forces between the molecules in a liquid, though not as strong as between the molecules of a solid

Liquids and Solids

• Densities of the three states of water– H2O(g) D = 3.26 x 10-4g/cm3 (400oC)

– H2O(l) D = 0.9971 g/cm3

(25oC)

– H2O(s) D = 0.9168 g/cm3 (OoC)

• Similarities in the densities of the liquid and solid state indicate similarities in the structure of liquids and solids

Intermolecular Forces

• Bonds are formed between atoms to form molecules– intramolecular bonding (within

the molecule)


• The properties of liquids and solids are determined by the forces that hold the components of the liquid or solid together– may be covalent bonds– may be ionic bonds– may weaker intermolecular forces

between molecules


• During a phase change for a substance like water– the components of the liquid or solid

remain intact– the change of state is due to the

changes in the forces between the components

– e.g., H2O(s) --> H2O (l) …the molecules are still unchanged during the phase change

Dipole-Dipole Forces

• Polar molecules– line up in an electric field

• positive end of molecule will line up with the negative pole of the electric field while the negative end of the molecule will line up with the positive pole

– can attract each other• positive end of one molecule will attract

the negative end of another molecule

Dipole-Dipole Forces

• Dipole-dipole forces – about 1% as strong as covalent or

ionic bonds– become weaker with distance– unimportant in the gas phase

Hydrogen Bonding• A particularly strong dipole-dipole

force• When hydrogen is covalently bonded

to a very electronegative atom such as N, O, or F

• Very strong due to– great polarity of the bond between H

and the N, O or F– close approach of the dipoles due to H’s

small size

Hydrogen Bonding

• H-bonding has a very important effect on physical properties– For example, boiling points are

greater when H-bonding is present

London Dispersion Forces• aka Van der Waals forces• Nonpolar molecules must exert

some kind of force or they would never solidify

London Dispersion Forces

• London dispersion forces (LDF)– due to an instantaneous dipole

moment • created when electrons move about the

nucleus• a temporary nonsymmetrical electron

distribution can develop (I.e., all the electrons will shift to one side of the molecule)


• The instantaneous dipole moment can induce an instantaneous dipole moment in a neighboring molecule, which could induce another instantaneous dipole moment in a neighboring molecule, etc. (like a “wave” in the stands of a football game)


• The LDF is very weak and short-lived• To form a solid when only LDF exists

requires very low temperatures– the molecules or atoms must be

moving slowly enough for the LDF to hold the molecules or atoms together in a “solid” unit


• Element Freezing Point (oC) Helium -269.7 Neon -248.6 Argon -189.4

Krypton -157.3 Xenon -111.9


• Notice that as the MM of the noble gas increases, the freezing point increases– This implies that the LDF between the

atoms is stronger as the MM increases• Large atoms with many electrons have an

increased polarizability (the instantaneous dipole would be larger), resulting in a larger London Dispersion Force between the atoms than between smaller atoms

The Liquid State

• Properties of liquids– low compressibility– lack of rigidity– high density (compared to gases)

The Liquid State

• Surface Tension– results in droplets when a liquid

is poured onto a surface– depends on IMF’s

The Liquid State

– Molecules at the surface experience an uneven pull, only from the sides and below. Molecules in the interior are surrounded by IMF’s• Uneven pull results in liquids assuming a

shape with minimum surface area• Surface tension is a liquids resistance to

an increase in surface area.•Liquids with high IMF’s have high

surface tensions

The Liquid State

• Capillary Action– Exhibited by polar molecules– The spontaneous rising of a

liquid in a narrow tube•due to two different forces involving the liquid

The Liquid State

• Cohesive forces - IMF between the liquid molecules

• Adhesive forces - forces between the liquid molecules and the polar (glass) container– adhesive forces tend to increase the surface area– cohesive forces counteract this

• Concave meniscus (water) - indicates adhesive forces of water towards the glass is greater than the cohesive forces between the water molecules.

• Convex meniscus (nonpolar substances such as mercury) shows cohesive forces is greater than adhesive forces.

The Liquid State

• Viscosity– Measure of a liquid’s resistance to

flow– Depends on strength of IMF’s

between liquid molecules• molecules with large IMF’s are very

viscous• Large molecules that can get tangled up

with each other lead to high viscosity

The Liquid State

• So what does a liquid “look like?”– A liquid contains many regions where

the arrangements of the components are similar to those of a solid

– There is more disorder in a liquid than in a solid

– There is a smaller number of regions in a liquid where there are holes present

Types of Solids

• Ways to classify solids– Crystalline vs. Amorphous Solid

• Crystalline solids– regular arrangement of components– positions of components represented

by a lattice– unit cell - smallest repeating unit of

the lattice

Types of Solids

• three common unit cells exist– simple cubic– body centered cubic– face centered cubic

Types of Solids

• Amorphous Solids– noncrystalline– glass is an example– disorder abounds

Types of Solids

• X-ray diffraction– used to determine the structures of

crystalline solids– diffraction occurs when beams of light

are scattered from a regular array of points

– obtain a diffraction pattern– Bragg equation: n = 2d sin

Types of Solids

• Where n is an integer is the wavelength of the x-rays• d is the distance between the atoms is the angle of incidence and reflection• Use x-ray diffraction to determine bond

lengths, bond angles, determine complex structures, test predictions of molecular geometry

Types of Solids• Example:• x-rays of wavelength 1.54 A were

used to analyze an aluminum crystal. A reflection was produced at = 19.3 degrees. Assuming n = 1, calculate the distance d between the planes of atoms producing the reflection.

• (D = 2.33 A)

Types of Solids

• Types of Crystalline Solids– Ionic Solids (e.g. NaCl)

– Molecular Solids (e.g. C6H12O6)

– Atomic Solids which include:• Metallic Solids• Covalent Network Solids

Types of Solids• Classify solids according to what

type of component is found at the lattice point (of a unit cell)– Atomic Solids have atoms at the

lattice points– Molecular Solids have discrete,

relatively small molecules at the lattice points

– Ionic solids have ions at the lattice points

Types of Solids

• Different bonding present in these solids results in dramatically different properties

• Element (atomic solid) M.P. (oC)Argon -189C(diamond) 3500Cu 1083

Structure and Bonding in Metals• Properties of Metals

– high thermal conductivity– high electrical conductivity– malleability (metals can be pounded

thin)– ductility (metals can be drawn into a

fine wire)– durable– high melting points

Structure and Bonding in Metals

• Properties are due to the nondirectional covalent bonding found in metallic crystals

• Metallic crystal– contains spherical atoms packed

together– atoms are bonded to each other

equally in all directions


• Closest Packing – most efficient arrangement of these

uniform spheres– Two possible closest packing

arrangements•Hexagonal Closest Packed Structure•Cubic Closest Packed Structure


• Hexagonal Closest Packed Structure (hcp)– aba arrangement– First Layer

• each sphere is surrounded by six other spheres


• Second Layer– the spheres do not lie directly over

the spheres in the first layer– the spheres lie in the indentations

formed by three spheres

• Third Layer– the spheres lie directly over the

spheres in the first layer


• Cubic Closest Packed Structure (ccp)– abc arrangement– First and Second Layers are the same as

in hexagonal closest packed structure– Third Layer

• the spheres occupy positions such that none of the spheres in the third layer lie over a sphere in the first layer


• Finding the net number of spheres in a unit cell– important for many applications

involving solids(when I figure it out, I’ll let you know…

or when it shows up on the ACS or AP test…then I’ll figure it out!)

Structure and Bonding in Metals• Examples of metals that are ccp

– aluminum, iron, copper, cobalt, nickel• Examples of metals that are hcp

– zinc, magnesium• Calcium and some other metals can go

either way


• Some metals, like the alkali metals are not closest packed at all– may be found in a body centered

cubic (bcc) unit cell where there are only 8 nearest neighbors instead of the 12 in the closest packed structures

Bonding Models for Metals

• The model must account for the typical physical properties of metals– malleability– ductility– efficient and uniform conduction of

heat and electricity in all directions– durability of metals– high melting points

Bonding Models for Metals• To account for these physical

properties, the bonding in metals must be– strong– nondirectional

• It must be difficult to separate atoms, but easy to move them (as long as the atoms stay in contact with each other


• Electron Sea Model (simplest picture)– Positive Metal ions (Metal cations) are

surrounded by a sea of valence electrons• the valence electrons are mobile and loosely

held• these electrons can conduct heat and

electricity• meanwhile, the metal ions can move around

easily


• Band Model or Molecular Orbital (MO) model– related to the electron sea model– more detailed view of the electron

energies and motions


• MO model– electrons travel around the metal crystal

in molecular orbitals formed from the atomic orbitals of the metal atoms

– In atoms like Li2 or O2, the space between the energies of the molecular orbitals is relatively wide (big energy difference between the orbitals)


• However, when many metal atoms interact, the molecular orbital energy levels are very close together

• Instead of separate, discrete molecular orbitals with different energies, the molecular orbitals are so close together in energies, that they form a continuum of levels, called bands


• Core electrons of metals are localized– the core electrons “belong” to a

particular metal ion

• The valence electrons of metals are delocalized– the valence electrons occupy partially

filled, closely spaced molecular orbitals


• Thermal and Electrical conductivity– metals conduct heat and electricity

because of highly mobile electrons– electrons in filled molecular orbitals

get excited (from added heat or electricity)• these electrons move into higher

energy, empty molecular orbitals


• Conduction electrons– the electrons that can be excited to

empty MO’s

• Conduction bands– the empty MO’s that can accept the

conducting electrons

Metal Alloys

• Alloy– a substance that contains a mixture

of elements and has metallic properties

• Metals can form alloys due to the nature of their structure and bonding

Metal Alloys

• Two types of alloys– Substitutional alloy

• host metal atoms are replaced by other metal atoms of similar size

• ex: brass is an alloy of zinc and copper sterling silver - silver and copper pewter - tin and copper

solder - lead and tin

Metal Alloys

• Interstitial Alloys– formed when some of the holes in the

closest packed structure are filled with smaller atoms

– ex: steel is an alloy with carbon filling the interstices of an iron crystal

Metal Alloys

• Presence of interstitial atoms changes the properties of the host metal

• Iron - soft, ductile, malleable• Steel - harder, stronger, less ductile

than pure iron – due to directional bonds between

carbon and iron – More carbon, harder steel

Covalent Network Solids

• Covalent Network Solids– Macromolecule– A giant molecule containing numerous

covalent bonds holding atoms together– Properties

• brittle• do not conduct heat or electricity• very high melting points


• Typical Covalent Network Solids– Diamond (Cdia) and Graphite (Cgraphite)

– Diamond• each C atom is covalently bonded to four

other C atoms in a tetrahedral arrangement• sp3 hybridization of the C atoms• Using MO model, diamond is a

nonconductor due to the large space between the empty MO’s.

– Electrons cannot be transferred easily to empty MO’s


• Graphite– slippery, black, and a conductor– different bonding than diamond– there are layers of sp2 hybridized C

atoms in fused six member rings• the layers are held loosely with weak

LDF’s• graphite is slippery due to these weak

LDF’s between layers


• Graphite– since the C atoms are sp2 hybridized,

there is one 2p orbital left– the 2p orbitals form molecular

orbitals above the plane of the rings– the electrons are delocalized in these

molecular orbitals• these delocalized electrons allow for

electrical conductivity


• Convert graphite to diamonds– apply pressure…150,000 atm at

2800oC– requires such high pressure and

temperature to completely break the bonds in graphite and rearrange them to yield diamond


• Silicon– makes up many compounds found in

the earth’s crust– silicon:geology as carbon:biology– Even though silicon and carbon are in

the same family, the structures of silicon and carbon compounds are very different


• Carbon compounds usually contain long chains with C-C bonds

• Silicon compounds usually contain chains with Si-O bonds


• Silica– Empirical formula - SiO2

• sand, quartz are composed of SiO2

• Si is the center of a tetrahedron, forming single bonds with four oxygen atoms, which are shared by other Si atoms

• A covalent network solid like diamond


• Silicates– related to silica– found in most rocks, soils, and clays

– based on interconnected SiO4 tetradera

– unlike silica, silicates contain silicon-oxygen anions• silicates need positive metal cations to

balance the negative charge


• Glass– an amorphous solid– formed when silica is heated and

cooled rapidly– more closely resembles a viscous

solution than a crystalline solid– adding different substances to the

melted silica results in different properties for the glass


• Add B2O3 to produce glass for labware (pyrex)– very little expansion or contraction

with large temperature changes

• Add K2O to produce a very hard glass that can be ground for eyeglasses or contacts

Semiconductors• Silicon is a semiconductor

– gap between filled and empty MO’s is smaller than the gap found in diamond (a nonconductor)

– a few electrons can get excited and cross the gap in silicon

– at higher temperatures, more electrons can get across, so conductivity increases at higher temperatures

Semiconductors

• Enhance conductivity of semiconductors by doping the crystal with other atoms

Semiconductors• N - type semiconductor - dope Si

with atoms with more valence e-’s (e.g. with As)– the extra electrons from As can

conduct an electric current

Semiconductors

• analogy: Given a row in a movie theater filled with people. Each person has a bag of popcorn. One person has two bags of popcorn. Passing one bag of popcorn (the extra electron) down the row is like electricity being conducted in an n-type semiconductor

Semiconductors• p-type semiconductor - dope Si

with atoms with less valence e-’s (e.g. with B)– B’s three valence e- leave a hole

in an MO. – Another e- could move into the

hole, but it would leave another hole for another electron to fill

Semiconductors

• Analogy: In a movie theater, a row of seats is filled, except for one seat. One person could get up out of his seat and move into the empty seat. The next person could then move into the newly emptied seat, and so on…

• the p in p-type refers to the positive hole formed with a missing valence electron

Types of Solids

• Ionic Solids– between positive and negative ions– held by ionic bonds

•electrostatic forces between oppositely charged ions

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