Liquids and Solids Gas Solid low density high compressibility

Name: Liquids and Solids Gas Solid low density high compressibility
Uploaded: 2017-10-12T15:44:19+00:00
Duration: PTM23S7
Description: Liquids and Solids Gas Solid low density high compressibility

Liquids and Solids Gas Solid 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) DHofus = 6.02 kJ/mol H2O(l) --> H2O(g) DHovap = 40.7 kJ/mol large value of DHvap 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 = g/cm (25oC) H2O(s) D = g/cm (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 (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 Neon Argon Krypton Xenon

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: nl = 2d sinq

Types of Solids Where n is an integer
l is the wavelength of the x-rays d is the distance between the atoms q 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 q = 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 C(diamond) Cu

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

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!)

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

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
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 p molecular orbitals above the plane of the rings the electrons are delocalized in these p 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

Liquids and Solids Gas Solid low density high compressibility

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Liquids and Solids Gas Solid low density high compressibility

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