Chapter 10 Liquids and Solids

Chapter 10 Liquids and Solids

Topics 10.5 Section is self study Intermolecular forces
Dipole-dipole forces Hydrogen bonding London Forces The liquid state Surface tension Capillary action Viscosity An introduction to structures and types of solids X-ray analysis of solids Types of crystalline solids Structure and bonding in metals Bonding metals for metals Meta alloys Molecular solids Ionic solids Vapor pressure and changes of state Phase diagrams

Intra- vs. Inter-molecular forces
intramolecular forces inside molecules (bonding) hold atoms together into molecule intermolecular forces These are what hold the molecules together in the condensed states. Forces between molecules They get weaker as phase changes from S – L – G When a substance changes state, molecule stays together but intermolecular forces are weakened

Intermolecular Forces
Gases – fill container, random rapid motion, never coming to rest or clustering together Motion is mainly translational Liquids – fixed volume, flow and assume shape of container, only slightly compressible, stronger forces hold molecules together Motion is mainly translational Solids – fixed volume, definite shape, generally less compressible than liquids, forces hold particles in a fixed shape Motion is mainly vibrational

Intermolecular Forces
Intermolecular forces are attractive forces between molecules Intramolecular forces hold atoms together in a molecule. Intermolecular vs Intramolecular 41 kJ to vaporize 1 mole of water (inter) 930 kJ to break all O-H bonds in 1 mole of water (intra) “Measure” of intermolecular force boiling point melting point DHvap DHfus DHsub Generally, intermolecular forces are much weaker than intramolecular forces.

+ + Dipole – Dipole Foces - -
When molecules with dipole moments line up to minimize repulsion and maximize attraction Very weak compared to covalent and ionic bonds + - Attractions + - Repulsion

Dipole – Dipole Foces 1% as strong as covalent bonds
Molecules that line up in the presence of a electric field are dipoles. The opposite ends of the dipole can attract each other so the molecules stay close together. 1% as strong as covalent bonds Small role in gases. Molecules with these forces possess higher melting points and boiling points than nonpolar molecules of comparable molar mass

The strengths of intermolecular forces are generally weaker than either ionic or covalent bonds.
Polar molecules have dipole-dipole attractions for one another. 16 kJ/mol (to separate molecules) + - + - 431 kJ/mol (to break bond)

.. .. : : : .. .. .. : .. : + S O O - - + S O - -
Intermolecular forces between molecules that posses dipole moment Dipole-dipole forces: (polar molecules) .. + S .. : O O : dipole-dipole attraction : .. - - .. S O : .. + : .. - - What effect does this attraction have on the boiling point?

Effect of polarity on boiling points Polar Nonpolar
BP MM Molecule -192 28 CO -196 N2 -88 34 PH3 -112 32 SiH4 -62 78 AsH3 -90 77 GeH4 97 162 ICl 69 160 Br2 Effect of polarity is usually small enough to be obscured by differences in molar mass HCl -85 BP (oC) HBr -60 HI -30 BP increase although polarity decreases

Hydrogen Bonds A hydrogen bond is an intermolecular force in which a
hydrogen atom covalently bonded to a nonmetal atom in one molecule is simultaneously attracted to a nonmetal atom of a neighboring molecule The strongest hydrogen bonds are formed if the nonmetal atoms are small and highly electronegative – e.g., N, O, F very strong type of dipole-dipole attraction because bond is so polar because atoms are so small

Hydrogen bonding: It is very strong dipole-dipole interaction (bonds involving H-F, H-O, and H-N are most important cases). +H-F- --- +H-F- Hydrogen bonding

Cl(HCl) and S(H2S) do not form hydrogen
Hydrogen bond Cl(HCl) and S(H2S) do not form hydrogen bonding although they have electronegativity similar to N, why? They are of bigger size to approach the hydrogen atom Hydrogen bond is 5-10% as strong as the covalent bond

Hydrogen bonding between water molecules

Water d+ d- d+

Hydrogen Bonding Bonding between hydrogen and more electronegative
neighboring atoms such as oxygen and nitrogen Hydrogen bonding between ammonia and water

Hydrogen Bonding Effects
Solid water is less dense than liquid water due to hydrogen bonding Hydrogen bonding is also the reason for the unusually high boiling point of water

Boiling Points for Some Non Polar Molecules
The larger the molecule the larger the Van der Waals attraction due to more electrons in the molecule. The stronger the attraction, the higher the boiling point.

100 H2O HF 0ºC H2Te H2Se NH3 SbH3 H2S HI AsH3 HCl HBr -100 PH3 SnH4
Boiling Points CH4 SiH4 GeH4 SnH4 -100 Molar mass 200

Hydrogen Bonding in other molecules
Many organic acids can form dimers due to hydrogen bonding Certain organic molecules can also form an intramolecular hydrogen bond

Ethanol shows hydrogen bonding
Do these compounds show hydrogen bonding?

Do these compounds show hydrogen bonding?

Hydrogen bonding and solubility
Some compounds containing O, N & F show high solubilities in certain hydrogen containing solvents. NH3 & CH3OH dissolves in H2O through the formation of H-bonds

London Dispersion Forces
Non - polar molecules also exert forces on each other. Otherwise, no solids or liquids. Electrons are not evenly distributed at every instant in time. Thus, they will have an instantaneous dipole. Atom with instantaneous dipole induces a dipole in the atom next to it. As a result, induced dipole- induced dipole interaction would take place. London forces increase with the size of the molecules.

attraction + - + - “Electrons are shifted to overload one side of an atom or molecule”.

They exist in every molecular compound They are significant only for nonpolar molecules and noble gas atoms They are weak, short-lived Caused by formation of temporary dipole moments The ease with which electron “cloud” of an atom can be distorted is called polarizability.

In general big molecules are more easily polarized
Polarizability: the ease with which an atom or molecule can be distorted to have an instantaneous dipole. “squashiness” In general big molecules are more easily polarized than little ones.

Weak, short lived. Lasts longer at low temperature. Eventually long enough to make liquids. More electrons, more polarizable. Bigger molecules, higher melting and boiling points. Much, much weaker than other forces. Also called Van der Waal’s forces.

Relative Magnitudes of Forces
The types of bonding forces vary in their strength as measured by average bond energy. Strongest Weakest Covalent bonds (400 kcal/mol) Hydrogen bonding (12-16 kcal/mol ) Dipole-dipole interactions (2-0.5 kcal/mol) London forces (less than 1 kcal/mol)

Which one(s) of the above are most polarizable?
Hint: look at the relative sizes.

London Forces in Hydrocarbons

Properties of Liquids 10.2 The Liquid state
Low compressibility Lack of rigidity High density compared to gases Beading (beads up as droplets) Surface tension Capillary action Viscosity Stronger intermolecular forces cause each of these to increase.

Surface tension The resistance to an increase in its surface area
Polar molecules and liquid metals show high surface tension

Surface tension Molecules at the the top are only pulled inside. Molecules in the middle are attracted in all directions. Minimizes surface area.

All liquids have surface tension
water is just higher than most others How can we decrease surface tension? Use a surfactant - surface active agent Also called a wetting agent, like detergent or soap Interferes with hydrogen bonding

Adhesive forces are intermolecular forces between unlike molecules
Cohesive forces are intermolecular forces between like molecules

If a polar substance is placed on a non-polar surface.
Beading If a polar substance is placed on a non-polar surface. There are cohesive, But no adhesive forces. And Visa Versa

Capillary Action Capillary action results from
intermolecular interactions Liquids spontaneously rise in a narrow tube. Glass is polar. It attracts water molecules (adhesive forces)

Glass has polar molecules.
Glass can also hydrogen bond. This attracts the water molecules. Some of the pull is up a cylinder.

Water curves up along the side of glass.
This makes the meniscus, as in a graduated cylinder Plastics are non-wetting; no attraction

Meniscus In Glass In Plastic

Viscosity is a measure of a liquid’s resistance to flow
strong inter molecular forces  highly viscous large, complex molecules  highly viscous Cyclohexane has a lower viscosity than hexane. Because it is a circle- more compact.

10.3 An introduction to structures and types of solids
Crystalline Solids: highly regular three dimensional arrangement of their components [table salt (NaCl)] Amorphous solids: considerable disorder in their structures (glass: components are frozen in place before solidifying and achieving an ordered arrangement) The positions of components in a crystalline solid are usually represented by a lattice

Crystalline solids Amorphous solids

Representation of Components in a Crystalline Solid
Lattice: A 3-dimensional system that describes the locations of components (atoms, ions, or molecules) that make up the unit cells of a substance. Unit Cell: The smallest repeating unit in the lattice. There are Three common types of unit cells: simple cubic body-centered cubic face-centered cubic

Crystal Structures - Cubic
Simple Face-Centered Body-Centered

Unit Cells The simple cubic cell is the simplest unit cell and has structural particles centered only at its corners The body-centered cubic (bcc) structure has an additional structural particle at the center of the cube The face-centered cubic (fcc) structure has an additional structural particle at the center of each face

Cubic lattice points At lattice points: Atoms Molecules
Ions Unit cells in 3 dimensions The simple cubic cell is the simplest unit cell and has structural particles centered only at its corners

Body-Centered Cubic Unit cells in 3 dimensions The body-centered cubic
(bcc) structure has an additional structural particle at the center of the cube

Face-Centered Cubic The face-centered cubic (fcc) structure has an additional structural particle at the center of each face

X-Ray analysis of solids
Sample is powdered X-rays of single wavelength is used Distance between planes of atoms in the crystal are calculated from the angles at which the rays are diffracted using Bragg equation

X-Ray analysis of solids
X-Ray diffraction Spots from diffracted X-rays Spot from incident beam

Reflection of X-rays from two layers of atoms
1st layer of atoms 2nd layer of atoms Extra distance traveled by lower ray = BC + CD = n = 2d sin

nl = 2d sinq

Bragg Equation nl 2 = d sin q d = distance between atoms
n = an integer l = wavelength of the x-rays

When silver crystallizes, it forms face-centered cubic cells
When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is pm. Calculate the density of silver. d = m V V = a3 = (408.7 pm)3 = 6.78X107 pm3 ___ atoms/unit cell in a face-centered cubic cell 107.9 g mole Ag x 1 mole Ag 6.022 x 1023 atoms x m = 4 Ag atoms = 1.79X10-22 g d = m V 7.17 x g 6.83 x cm3 = = 1.79X10-22 g / 6.78X107 pm3 = 2.6X10-30 g/pm3

Types of crystalline Solids
Ionic solids (ionic compounds) Ions (held by electrostatic attraction) at point in lattice They conduct electric current when dissolved in water Molecular solids (molecular compounds) Molecules (held by: dispersion and/or dipole-dipole forces) at each point in lattice. Ice is a molecular solid H2O atomic solids (metals, nonmetals, noble gases) Elements (C, B, Si) that are composed of atoms at lattice points. Three types: Metallic– metallic bond Network – strong covalent bonding Group 8A –London Dispersion Forces

10.4 Structure and Bonding of Metals
Physical properties of metals Ionization energy E is small (outer electrons move relatively free); this results in High electrical conductivity High thermal conductivity They are Ductile: can be drawn oust into wires Malleable: can be hammered into thin sheets Electrons act like a glue holding atomic nuclei Crystals of nonmetals break into small pieces if it is hammered (brittle) They have luster (reflect light) They form alloys

Metallic Crystals 6 on the same plane 3 in the plane above
Can be viewed as containing atoms (spheres) packed together in the closest arrangement possible The spheres are packed in layers Closest packing- when each sphere has 12 neighbors 6 on the same plane 3 in the plane above 3 in the plane below

Packing in Crystals “Open” packing has larger voids in between particles compared to close-packed crystals

Closest packed structures
Hexagonal closest-packed (hcp) structure It has (aba) arrangements that occur when the spheres of the third layer occupy positions so that each sphere in the third layer lies directly over a sphere in the first layer

Closest packed structures
Cubic closest-packed (ccp) structure It has (abc) arrangement that occurs when the spheres of the third layer occupy positions that NO sphere lies over one in the first layer An atom in every fourth layer lies over an atom in the First layer

Face centered cubic cell
Net number of spheres (atoms) in a unit cell, length of the edge of the cell, density of the closest packed solid Face centered cubic cell Atoms occupy corners and centers of the faces Atoms at the corners do not touch each other Atoms contact is made at the face diagonal 74% of the space is occupied Ca, Sr, transition metals An atom at the center of the face of cube is shared by another cube that touches that face. Only atom is assigned to a given cell An atom at the center of the cube is a part of 8-different cubes touching that point. Only corner atom belongs to the cell

Density of closest packed solid
# of spheres (atoms) per unit cell = Density of closes packed solid

Body centered cubic cell
Atoms contact along a body diagonal Atoms occupy the corners and one at the center In a unit cell, 8 atoms occupy the corners plus one in the center 68% of the space is occupied Available in Group I elements + Ba

Number of atoms assigned to each type of cell
Simple cube Body centered cube Face centered cube

1 atom/unit cell 2 atoms/unit cell 4 atoms/unit cell (8 x 1/8 = 1) (8 x 1/8 + 1 = 2) (8 x 1/8 + 6 x 1/2 = 4)

Electron Sea Model The highest energy level for most metal atoms
does not contain many electrons These vacant overlapping orbitals allow outer electrons to move freely around the entire metal Metallic crystal is an array of positive ions (cations) in a sea of roaming valence electrons

Metal Alloys An alloy is a mixture of elements and has metallic properties substitutional alloy host metal atoms are replaced by other metal atoms host and other metal atoms are similar in sizes interstitial alloy metal atoms occupy spaces created between host metal atoms metal atoms have large difference in size

1/3 of Cu atoms replaced by Zn Steel interstitial
Examples Brass substitutional 1/3 of Cu atoms replaced by Zn Steel interstitial Fe with C atoms in between makes harder and less malleable

10.6 Molecular solids Molecules occupy the corners of the lattices.
Common examples: ice, dry CO2, S8, P4, I2 Different molecules have different forces between them (H-bonds, or dipole-dipole or London forces, or a combination of all these forces) These forces depend on the size of the molecule. They also depend on the strength and nature of dipole moments.

Molecular solids with nonpolar molecules (without dipoles): H2, CCl4
Most are gases at 25ºC. The only forces are London Dispersion Forces. These depend on size of atom. Large molecules (such as I2 ) can be solids even without dipoles.

Molecular solids with polar molecules (with dipoles): HCl, NH3
Dipole-dipole forces are generally stronger than L.D.F. Hydrogen bonding is stronger than Dipole-dipole forces. No matter how strong the intermolecular force, it is always much, much weaker than the forces in bonds. Stronger forces lead to higher melting and freezing points.

Water –Molecular solid
Each molecule has two polar O-H bonds. Each molecule has two lone pair on its oxygen. Each oxygen can interact with 4 hydrogen atoms. H O d+

10.7 Ionic Solids In most of binary ionic compounds, larger ions are arranged in closest packing arrangement, hexagonal (hcp) or cubic (ccp) closest packing smaller ions fit in the holes created by the larger ions Atoms are actually held together by electrostatic attractions of opposite charges. They possess huge melting and boiling points. Atoms are locked in lattice so they are hard and brittle. Every electron is accounted for so they are poor conductors-good insulators.

Closest Packing Holes The hole is formed by 3 spheres in the
Same layer The hole formed when a sphere occupies a dimple formed by three spheres in an adjacent layer The holes formed between two sets of three spheres in adjoining layers of closest packed structure

Closest packing holes Tetrahedral holes: are located above a sphere in the bottom layer Octahedral holes: are located above a void in the bottom layer

Examples Trigonal holes are so small that they are never occupied in binary ionic compounds The type of the hole whether tetrahedral or octahedral depends mainly on Relative sizes of cations and anions In ZnS, S2-, ions are arranged in ccp with the smaller Zn2+ ions in the tetrahedral holes In NaCl, ions are arranged in ccp with Na+ ions in the octahedral holes.

10. 7 Vapor Pressure and changes of state
vapor- gas phase above a substance that exists as solid or liquid at 25°C and 1 atm. Vaporization or Evaporation - change from liquid to gas at or below the boiling point . (Endothermic process) Condensation is the change of a gas to a liquid (Exothermic process)

Energy required to vaporize 1 mole of a liquid at 1 atm
Heat or enthalpy of vaporization, ∆Hvap : Energy required to vaporize 1 mole of a liquid at 1 atm water has a large ∆Hvap (40.7 kJ/mol), (because of hydrogen bonding)

Molar Heats of Vaporization for Selected Liquids

Vapor pressure Initially, liquid in a closed container decreases as molecules enter gaseous phase When equilibrium is reached, no more net change occurs Rate of condensation and rate of vaporization become equal Molecules still are changing phase but no net change (Dynamic equilibrium) Gas liquid Vapor pressure is independent of volume of container as long as some liquid is present (liquid-vapor equilibrium)

Evaporation and condensation
H2O (l) H2O (g) Rate of condensation evaporation = Dynamic Equilibrium

Equilibrium vapor pressure or vapor pressure
It is the pressure exerted by the vapor when it is in dynamic equilibrium with a liquid at a constant temperature.

Measurement of Vapor Pressure
The pressure of the vapor phase at equilibrium: Pvap can be measured when using a simple barometer Liquid can be injected under inverted tube part of the liquid evaporates to the top of tube Pvap can be determined by height of Hg Patm = Pvap + PHg

Vapor pressure and nature of liquids
Vapor pressure depends upon the nature of the liquid Liquids with high vapor P (volatile liquids) evaporate quickly weak intermolecular forces Liquids with low vapor P Strong London dispersion forces (large molar masses) or dipole-dipole forces

Vapor pressure and temperature
Vapor pressure increases with T More molecules have enough KE to overcome intermolecular forces

VAPOR PRESSURE CURVES A liquid boils when its vapor pressure = external pressure.

Temperature Effect Energy needed to overcome intermolecular forces in iquid T1 # of molecules Kinetic energy

At higher temperature more molecules have enough energy - higher vapor pressure.
Energy needed to overcome intermolecular forces in liquid T1 T1 # of molecules T2 Kinetic energy

Vapor pressure and Temperature
Molar heat of vaporization (DHvap) is the energy required to vaporize 1 mole of a liquid. P = (equilibrium) vapor pressure Clausius-Clapeyron Equation T = temperature (K) ln P = - DHvap RT + C R = gas constant (8.314 J/K•mol)

Mathematical relationship
DHvap is the heat of vaporization in J/mol R = J/K mol.

Vapor Pressure for solids
Solids also have vapor pressure Sublimination solid gas directly Example: dry ice: CO2 heat of fusion (∆Hfus) enthalpy of fusion enthalpy change at melting point

Heating Curve plot of T vs. time where energy is added at constant rate as energy is added, it is used to increase the T when it reaches melting point, the energy added is used to change molecules from (s) to (l)

Changes of state What happens when a solid is heated?
The graph of temperature versus heat applied is called a heating curve. The temperature a solid turns to a liquid is the melting point. The energy required to accomplish this change is called the Heat (or Enthalpy) of Fusion Hfus

bp mp Heating Curve for Water Steam Time (Heat added) Water and Steam
Temp Water mp Water and Ice Ice Time (Heat added)

Heating Curve for Water
Slope is Heat Capacity Hvap=2260 J/g Heat of vaporization Temp Hfus=334 J/g Heat of fusion Time (Heat added)

Calculate the enthalpy change upon converting 1 mole of water from ice at -12oC to steam at 115oC.
liquid 100oC gas 100oC gas 115oC solid -12oC solid 0oC liquid 0oC H H H H H5 = Htotal Sp. Ht. + Hfusion + Sp. Ht. + HVaporization + Sp. Ht. = Htotal Specific Heat of ice = 2.09 J/g•K Hfus=334 J/g Hvap=2260 J/g Specific Heat of water = J/g•K Specific Ht. Steam = 1.84 J/g•K

Normal melting Point Melting point is determined by the vapor pressure of the solid and the liquid. Melting point is the temp at which the vapor pressure of the solid = vapor pressure of the liquid where the total pressure is 1 atm.

Water Vapor Vapor Solid Water Liquid Water
Apparatus that allows solid and liquid water to interact only through the vapor state Solid Water Liquid Water Water Vapor Vapor

A temp at which the vapor pressure of the solid is higher than that of the liquid the solid will release molecules to achieve equilibrium. Solid Water Liquid Water Water Vapor Vapor

While the molecules of water condense to a liquid to achieve equilibrium. Solid Water Liquid Water Water Vapor Vapor

This can only happen if the temperature is above the melting point since solid is turning to liquid. Solid Water Liquid Water Water Vapor Vapor

A temperature at which the vapor pressure of the solid is less than that of the liquid, the liquid will release molecules to achieve equilibrium. Solid Water Liquid Water Water Vapor Vapor

While the molecules of water condense to a solid. Solid Water Liquid Water Water Vapor Vapor

The temperature must be below the melting point since the liquid is turning to a solid. Solid Water Liquid Water Water Vapor Vapor

Temperature at which the vapor pressure of the solid and liquid are equal, the solid and liquid are vaporizing and condensing at the same rate. This is the Melting (freezing) point (Temp at which solid and vapor can coexist) Solid Water Liquid Water Water Vapor Vapor

Normal Boiling Point Temp when vapor pressure inside the bubbles equals 1 atm If Pvap < 1 atm, no bubbles can form, there is too much pressure on surface

Normal Boiling Point Boiling occurs when the vapor pressure of liquid becomes equal to the external pressure. Normal boiling point is the temperature at which the vapor pressure of a liquid is exactly 1 atm pressure. Super heating - Heating above the boiling point. Supercooling - Cooling below the freezing point.

Exceptions Changes of state do not always occur exactly at bp or bp
supercooling Material can stay liquid below melting point because doesn’t achieve level of organization needed to make solid superheating when heated too quickly, liquid can be raised above boiling point causes “bumping”

10.9 Phase Diagrams A plot Representing phases of a substance in a closed system (no material escapes into the surroundings and no air is present) as a function of temperature and pressure.

Phase changes Gas and liquid are indistinguishable.
Critical temperature and critical pressure (all 3 phases exists here)

Phase Diagrams critical point vapor pressure curve fusion curve
sublimation curve triple point

Critical Point where if the T is increased, vapor can’t be liquefied no matter what P is applied at the end of liquid/gas line after this point, only one fluid phase exists that is neither gas nor liquid called supercritical fluid

Critical temperature Temperature above which the vapor can not be liquefied. Critical pressure pressure required to liquefy gas AT the critical temperature. Critical point critical temperature and pressure (for water, Tc = 374°C and 218 atm).

Phase diagram for Water
Expands upon freezing -ve slope of S/L boundary line means that mp of ice decreases as the external P increases Normal bp Normal mp Critical temp

Phase Diagram for H2O

Most substances have a positive slope of solid/liquid line because solid is usually more dense than liquid water has a negative slope

Pressure 1 Atm Temperature Liquid Solid Gas
This is the phase diagram for CO2 The solid is more dense than the liquid The solid sublimes at 1 atm. Pressure Liquid Solid 1 Atm Gas Temperature

Chapter 10 Liquids and Solids

Similar presentations

Presentation on theme: "Chapter 10 Liquids and Solids"— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

Chapter 10 Liquids and Solids

Similar presentations

Presentation on theme: "Chapter 10 Liquids and Solids"— Presentation transcript:

Similar presentations

About project

Feedback