Intermolecular Forces and

Intermolecular Forces and
Liquids and Solids Chapter 11

States of Matter Comparison of gases, liquids, and solids.
Gases are compressible fluids. Their molecules are widely separated. Liquids are relatively incompressible fluids. Their molecules are more tightly packed. Solids are nearly incompressible and rigid. Their molecules or ions are in close contact and do not move. 2

A phase is a homogeneous part of the system in contact with other parts of the system but separated from them by a well-defined boundary. 2 Phases Solid phase - ice Liquid phase - water

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.

Intermolecular Forces; Explaining Liquid Properties
The term van der Waals forces is a general term including dipole-dipole and London forces. Van der Waals forces are the weak attractive forces in a large number of substances. Hydrogen bonding occurs in substances containing hydrogen atoms bonded to certain very electronegative atoms. 2

Van der Waals Forces 7 Van der Waals Forces are types of intermolecular forces. Consists of: Dipole-dipole interactions 2 Ion-dipole interactions 3 London forces

Dipole-Dipole Forces H Cl
Polar molecules can attract one another through dipole-dipole forces. The dipole-dipole force is an attractive intermolecular force resulting from the tendency of polar molecules to align themselves positive end to negative end. H Cl d- d+ 2

Dipole-Dipole Forces Attractive forces between polar molecules Orientation of Polar Molecules in a Solid

Ion-Dipole Forces Attractive forces between an ion and a polar molecule Ion-Dipole Interaction

Ion-dipole interaction between ions and water molecules
Why strong here? Higher charge, smaller ionic radius octahedral

London Forces or Dispersion Forces
London forces are the weak attractive forces resulting from instantaneous dipoles that occur due to the distortion of the electron cloud surrounding a molecule. London forces increase with molecular weight. The larger a molecule, the more easily it can be distorted to give an instantaneous dipole. All covalent molecules exhibit some London force. 2

exist between all molecules,
London forces/dispersion forces: exist between all molecules, is the only attractive force between nonpolar atoms or molecules, Electrons are in constant motion. Electrons can be, in an instant, arranged in such a way that they have a dipole. (Instantaneous dipole) The temporary dipole interacts with other temporary dipoles to cause attraction.

Dispersion Forces/London Forces Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules ion-induced dipole interaction dipole-induced dipole interaction

Induced Dipoles Interacting With Each Other

Dispersion Forces Continued Polarizability is the ease with which the electron distribution in the atom or molecule can be distorted. Polarizability increases with: greater number of electrons more diffuse electron cloud Dispersion forces usually increase with molar mass.

What type(s) of intermolecular forces exist between each of the following molecules?
HBr HBr is a polar molecule: dipole-dipole forces. There are also dispersion forces between HBr molecules. CH4 CH4 is nonpolar: dispersion forces. S O SO2 SO2 is a polar molecule: dipole-dipole forces. There are also dispersion forces between SO2 molecules.

Hydrogen Bonding H N O F :
Hydrogen bonding is a force that exists between a hydrogen atom covalently bonded to a very electronegative atom, X, and a lone pair of electrons on a very electronegative atom, Y. To exhibit hydrogen bonding, one of the following three structures must be present. H N O F : Only N, O, and F are electronegative enough to leave the hydrogen nucleus exposed. 2

not considered a Van der Waals Force
Hydrogen Bonding Hydrogen bonding: not considered a Van der Waals Force is a special type of dipole-dipole attraction is a very strong intermolecular attraction causing higher than expected b.p. and m.p. Requirement for hydrogen bonding: molecules have hydrogen directly bonded to O, N, or F

Examples of hydrogen bonding:
H2O NH3 HF

Hydrogen Bond The hydrogen bond is a special dipole-dipole interaction between the hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H … B or A & B are N, O, or F

HCOOH molecule hydrogen bond with two water molecules

Hydrogen Bonding Molecules exhibiting hydrogen bonding have abnormally high boiling points compared to molecules with similar van der Waals forces. For example, water has the highest boiling point of the Group VI hydrides. Similar trends are seen in the Group V and VII hydrides. 2

Decreasing boiling point
Hydrogen bonding made the circled molecules to have higher B.P. as an exception to the trend in the groups. Decreasing molar mass Decreasing boiling point

Review: Hydrogen Bonding
A hydrogen atom bonded to an electronegative atom appears to be special. The electrons in the O-H bond are drawn to the O atom, leaving the dense positive charge of the hydrogen nucleus exposed. It’s the strong attraction of this exposed nucleus for the lone pair on an adjacent molecule that accounts for the strong attraction. A similar mechanism explains the attractions in HF and NH3. 2

Review: Hydrogen Bonding
2

Van der Waals Forces and the Properties of Liquids
In summary, intermolecular forces play a large role in many of the physical properties of liquids and gases. These include: vapor pressure (1) boiling point (2) surface tension (3) Viscosity (4) (2),(3), and (4) increases with high inter molecular forces and (1) decreases. 2

Surface tension increases with increasing intermolecular forces in liquids Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. To increase surface area, it is necessary to pull molecules apart against the intermolecular forces of attraction. 2

Strong intermolecular forces
Properties of Liquids Surface tension is the amount of energy required to stretch or increase the surface of a liquid by a unit area. High surface tension Strong intermolecular forces A molecule within a liquid is pulled in all directions, whereas a molecule on the surface is only pulled to the interior. As a result, there is a tendency for the surface area of the liquid to be minimized.

Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules, water-glass Related to Surface tension, capillary action in water and depression of mercury level Mercury – Cohesion is greater Water - Adhesion is greater Adhesion Cohesion

Properties of Liquids; Surface Tension and Viscosity
This explains why falling raindrops are nearly spherical, minimizing surface area. In comparisons of substances, as intermolecular forces between molecules increase, the apparent surface tension also increases. 2

Review: Surface Tension
Surface tension - a measure of the attractive forces exerted among molecules at the surface of a liquid. Surface molecules are surrounded and attracted by fewer liquid molecules than those below. Net attractive forces on surface molecules pull them downward. Results in “beading” Surfactant - substance added which decreases the surface tension example: soap

Viscosity increases with increasing intermolecular forces because increasing these forces increases the resistance to flow. you know that syrup has a greater viscosity than water. In comparisons of substances, as intermolecular forces increase, viscosity usually increases. 2

Strong intermolecular forces
Properties of Liquids Viscosity is a measure of a fluid’s resistance to flow. Strong intermolecular forces High viscosity

The Solid State Particles highly organized and well defined fashion
9 Particles highly organized and well defined fashion Fixed shape and volume Properties of Solids: incompressible m.p. depends on strength of attractive force between particles Crystalline solid - regular repeating structure Amorphous solid - no organized structure.

Unit cells in 3 dimensions
A crystalline solid possesses rigid and long-range order. In a crystalline solid, atoms, molecules or ions occupy specific (predictable) positions. An amorphous solid does not possess a well-defined arrangement and long-range molecular order. A unit cell is the basic repeating structural unit of a crystalline solid. At lattice points: Atoms Molecules Ions lattice point Unit Cell Unit cells in 3 dimensions

Shared by 8 unit cells Shared by 2 unit cells

How is this calculation?
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)

The primitive cubic system (cP) consists of one lattice point on each corner of the cube. Each atom at a lattice point is then shared equally between eight adjacent cubes, and the unit cell therefore contains in total one atom (1⁄8 × 8). The body-centered cubic system (cI) has one lattice point in the center of the unit cell in addition to the eight corner points. It has a net total of 2 lattice points per unit cell (1⁄8 × 8 + 1). The face-centered cubic system (cF) has lattice points on the faces of the cube, that each gives exactly one half contribution, in addition to the corner lattice points, giving a total of 4 lattice points per unit cell (1⁄8 × 8 from the corners plus 1⁄2 × 6 from the faces).

When silver crystallizes, it forms face-centered cubic cells
When silver crystallizes, it forms face-centered cubic cells. The unit cell edge length is 409 pm. Calculate the density of silver. d = m V V = a3 = (409 pm)3 = 6.83 x cm3 4 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 = 7.17 x g d = m V 7.17 x g 6.83 x cm3 = = 10.5 g/cm3

Determines bond lengths and bond angles in molecules in the solid state

to be in the same phase as the upper wave which is nl
Extra distance = Scattering of X-rays. The lower wave travels extra distance, 2dsinq to be in the same phase as the upper wave which is nl BC + CD = 2d sinq = nl (Bragg Equation)

Calculator has to be in the degree mode when taking the “sin value”
X rays of wavelength nm are diffracted from a crystal at an angle of Assuming that n = 1, what is the distance (in pm) between layers in the crystal? nl = 2d sin q n = 1 q = l = nm = 154 pm nl 2sinq = 1 x 154 pm 2 x sin14.17 d = = pm Calculator has to be in the degree mode when taking the “sin value”

Melting Point and Structure
Physical Properties Many physical properties of a solid can be attributed to its structure. Melting Point and Structure Note that for ionic solids, melting points increase with the strength of the ionic bond. Ionic bonds are stronger when: The magnitude of charge is high. 2. The ions are small (higher charge density). 2

Solid State From this point of view, there are four types of solids.
Ionic (Ionic bond) Covalent (Covalent bond) Molecular (Van der Waals forces) Metallic (Metallic bond) 2

Types of Crystals Ionic Crystals
Lattice points occupied by cations and anions Held together by electrostatic attraction Hard, brittle, high melting point, boiling point Poor conductor of heat and electricity If dissolves in water, electrolytes NaCl CsCl ZnS CaF2

Types of Crystals Covalent Crystals Lattice points occupied by atoms
Held together by covalent bonds Hard, high melting point, boiling point Poor conductor of heat and electricity Extremely Hard Diamond carbon atoms diamond graphite

Types of Crystals Molecular Crystals
Lattice points occupied by molecules Held together by intermolecular forces Soft, low melting point Poor conductor of heat and electricity Often volatile Ice, dry ice

Cross Section of a Metallic Crystal
Types of Crystals Metallic Crystals Lattice points occupied by metal atoms Held together by metallic bonds Soft to hard, low to high melting point Good conductors of heat and electricity Iron, copper, and silver Cross Section of a Metallic Crystal nucleus & inner shell e- mobile “sea” of e-

Crystal Structures of Metals

Melting Point and Structure
Physical Properties Many physical properties of a solid can be attributed to its structure. Melting Point and Structure Metals often have high melting points, but there is considerable variability. Melting points are low for Groups IA and IIA but increase as you move into the transition metals. The elements in the middle of the transition metals have the highest melting points. 2

Hardness and Structure
Physical Properties Many physical properties of a solid can be attributed to its structure. Hardness and Structure Hardness depends on how easily structural units can be moved relative to one another. Molecular solids with weak intermolecular attractions are rather soft compared with ionic compounds, where forces are much stronger. 2

Physical Properties Molecular solids - soft
Covalent network solids are quite hard because of the rigidity of the covalent network structure. Diamond and silicon carbide (SiC), three-dimensional covalent network solids, are among the hardest substances known. 2

Hardness and Structure
Physical Properties Brittle and malleable Hardness and Structure Molecular and ionic crystals are generally brittle because they fracture easily along crystal planes. Metallic solids, by contrast, are malleable. 2

Electrical Conductivity and Structure
Physical Properties Many physical properties of a solid can be attributed to its structure. Electrical Conductivity and Structure Molecular and ionic solids are generally considered nonconductors. Ionic compounds conduct in their molten state, as ions are then free to move. Metals are all considered conductors. 2

Electrical Conductivity and Structure
Physical Properties Many physical properties of a solid can be attributed to its structure. Electrical Conductivity and Structure Of the covalent network solids, only graphite conducts electricity. 2

Types of Crystals

An amorphous solid does not possess a well-defined arrangement and long-range molecular order.
A glass is an optically transparent fusion product of inorganic materials that has cooled to a rigid state without crystallizing Crystalline quartz (SiO2) Non-crystalline quartz glass

Changes of State A change of state or phase transition is a change of a substance from one state to another. gas condensation or deposition condensation boiling sublimation liquid melting freezing solid 2

Freezing Point The temperature at which a pure liquid changes to a crystalline solid, or freezes, is called the freezing point. The melting point is identical to the freezing point and is defined as the temperature at which a solid becomes a liquid. Unlike boiling points, melting points are affected significantly by only large pressure changes. 2

Boiling Point The temperature at which the vapor pressure of a liquid equals the pressure exerted on the liquid is called the boiling point. As the temperature of a liquid increases, the vapor pressure increases until it reaches atmospheric pressure. At this point, stable bubbles of vapor form within the liquid. This is called boiling. The normal boiling point is the boiling point at 1 atm. 2

H2O (s) H2O (l) The melting point of a solid or the freezing point of a liquid is the temperature at which the solid and liquid phases coexist in equilibrium Melting Freezing

(The boiling point lowers.)
What happens when you go to a mountain where the atmospheric pressure is lower than 1 atm? (The boiling point lowers.)

Greatest Order Least T2 > T1 Evaporation Condensation

H2O (s) H2O (g) Molar heat of sublimation (DHsub) is the energy required to sublime 1 mole of a solid. Sublimation Deposition DHsub = DHfus + DHvap ( Hess’s Law)

Vapor Pressure Liquids are continuously vaporizing.
If a liquid is in a closed vessel with space above it, a partial pressure of the vapor state builds up in this space. The vapor pressure of a liquid is the partial pressure of the vapor over the liquid, measured at equilibrium at a given temperature. 2

Vapor Pressure The vapor pressure of a liquid depends on its temperature. As the temperature increases, the kinetic energy of the molecular motion becomes greater, and vapor pressure increases. Liquids and solids with relatively high vapor pressures at normal temperatures are said to be volatile. 2

Variation of vapor pressure with temperature.

The equilibrium vapor pressure is the vapor pressure measured when a dynamic equilibrium exists between condensation and evaporation H2O (l) H2O (g) Rate of condensation evaporation = Dynamic Equilibrium

Before Evaporation At Equilibrium

Once there are molecules in the vapor phase, they can be converted back to the liquid phase
H2O(g)  H2O(l) + energy evaporation - the process of conversion of liquid to gas, at a temperature too low to boil condensation - conversion of the gas to the liquid state.

When the rate of evaporation equals the rate of condensation, the system is at equilibrium.
Vapor pressure of a liquid - the pressure exerted by the vapor at equilibrium H2O(g) H2O(l)

Heat of Phase Transition
Molar heat of vaporization-measure of the strength of intermolecular forces in a liquid. The energy needed to vaporize 1 mol of a liquid is called the molar heat of vaporization and denoted DHvap. molar heat of fusion, DHfus energy needed to melt 1 mol of a solid. For ice, the heat of vaporization is kJ/mol. 2

Heating Curve

A Problem to Consider The heat of vaporization of ammonia is 23.4 kJ/mol. How much heat is required to vaporize 1.00 kg of ammonia? First, we must determine the number of moles of ammonia in 1.00 kg (1000 g). 2

A Problem to Consider The heat of vaporization of ammonia is 23.4 kJ/mol. How much heat is required to vaporize 1.00 kg of ammonia? Then we can determine the heat required for vaporization. 2

Clausius-Clapeyron Equation
We noted earlier that vapor pressure was a function of temperature. It has been demonstrated that the logarithm of the vapor pressure of a liquid varies linearly with absolute temperature. Consequently, the vapor pressure of a liquid at two different temperatures is described by: 2

A Problem to Consider Carbon disulfide, CS2, has a normal boiling point of 46oC (vapor pressure = 760 mmHg) and a heat of vaporization of 26.8 kJ/mol. What is the vapor pressure of carbon disulfide at 35oC? Substituting into the Clausius-Clapeyron equation, we obtain: 2

A Problem to Consider Carbon disulfide, CS2, has a normal boiling point of 46oC (vapor pressure = 760 mmHg) and a heat of vaporization of 26.8 kJ/mol. What is the vapor pressure of carbon disulfide at 35oC? Taking the antiln we obtain: 2

The boiling point is the temperature at which the (equilibrium) vapor pressure of a liquid is equal to the external pressure. The normal boiling point is the temperature at which a liquid boils when the external pressure is 1 atm.

Phase Diagrams Below is a typical phase diagram. It consists of three curves that divide the diagram into regions labeled “solid, liquid, and gas”. . B C solid liquid pressure . gas D A temperature 2

Phase Diagrams A phase diagram is a graphical way to summarize the conditions under which the different states of a substance are stable. The diagram is divided into three areas representing each state of the substance. The curves separating each area represent the boundaries of phase changes. 2

Phase Diagrams Curve AB, dividing the solid region from the liquid region, represents the conditions under which the solid and liquid are in equilibrium. . B C solid liquid pressure . gas A D temperature 2

Phase Diagrams Usually, the melting point is only slightly affected by pressure. For this reason, the melting point curve, AB, is nearly vertical. . B C solid liquid pressure . gas A D temperature 2

Phase Diagrams If a liquid is more dense than its solid, the curve leans slightly to the left, causing the melting point to decrease with pressure. . B C solid liquid pressure . gas A D temperature 2

Phase Diagrams If a liquid is less dense than its solid, the curve leans slightly to the right, causing the melting point to increase with pressure. . B C solid liquid pressure . gas A D temperature 2

Phase Diagrams Curve AC, which divides the liquid region from the gaseous region, represents the boiling points of the liquid for various pressures. . B C solid liquid pressure . gas A D temperature 2

Phase Diagrams Curve AD, which divides the solid region from the gaseous region, represents the vapor pressures of the solid at various temperatures. . B C solid liquid pressure . gas A D temperature 2

Phase Diagrams The temperature above which the liquid state of a substance no longer exists regardless of pressure is called the critical temperature. . B C solid liquid pressure . gas A D Tcrit temperature 2

Phase Diagrams The vapor pressure at the critical temperature is called the critical pressure. Note that curve AC ends at the critical point, C. . B Pcrit C solid liquid pressure . gas A D Tcrit temperature 2

The critical temperature (Tc) is the temperature above which the gas cannot be made to liquefy, no matter how great the applied pressure. The critical pressure (Pc) is the minimum pressure that must be applied to bring about liquefaction at the critical temperature.

Phase Diagrams The curves intersect at A, the triple point, which is the temperature and pressure where three phases of a substance exist in equilibrium. . B C solid liquid pressure . gas A D temperature 2

A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas.
Phase Diagram of Water Triple point for water,0.01⁰C and atm. Increasing the pressure, lowers the m.p and increases the b.p

Phase Diagram of Carbon Dioxide
Liquid is not stable below 5.2 atm. Only solid and vapor at normal pressure Dry ice! At 1 atm CO2 (s) CO2 (g)

Effect of Increase in Pressure on the Melting Point
of Ice and the Boiling Point of Water

Homework Ch-11 10,12,39,47,48,41,42,84,76

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