Intermolecular Forces and

Intermolecular Forces and
Liquids and Solids Chapter 10

States of Matter The fundamental difference between states of matter is the distance between particles.

States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases. 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

The States of Matter The state a substance is in at a particular temperature and pressure depends on two antagonistic entities: The kinetic energy of the particles The strength of the attractions between the particles

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

Inter- molecular Forces
Have studied INTRAmolecular forces—the forces holding atoms together to form molecules. Now turn to forces between molecules — INTERmolecular forces. Forces between molecules, between ions, or between molecules and ions.

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.

Types of Bonding in Crystalline Solids

IMF Flowchart

Types of Intermolecular Forces
1. Hydrogen Bond (strongest) 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. IT IS NOT A BOND. A H … B or A & B are N, O, or F

Hydrogen Bond 11.2

Why is the hydrogen bond considered a “special” dipole-dipole interaction?
Decreasing molar mass Decreasing boiling point

Hydrogen Bonding in Biology
H-bonding is especially strong in biological systems — such as DNA. DNA — helical chains of phosphate groups and sugar molecules. Chains are helical because of tetrahedral geometry of P, C, and O. Chains bind to one another by specific hydrogen bonding between pairs of Lewis bases. —adenine with thymine —guanine with cytosine

Double helix of DNA Portion of a DNA chain

Base-Pairing through H-Bonds

Hydrogen Bonding in Biology
Hydrogen bonding and base pairing in DNA.

Nucleotide pairs form H-bonds
* There is no strict cutoff for the ability to form H-bonds (S forms a biologically important hydrogen bond in proteins). * Hold DNA strands together in double-helix Nucleotide pairs form H-bonds DNA double helix

2. Ion-Dipole Forces Attractive forces between an ion and a polar molecule. The magnitude of the attraction increases as either the charge of the ion or the magnitude of the dipole increases Ion-Dipole Interaction

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

4. London Dispersion Forces – van der Walls forces/London forces (weakest) Attractive forces that arise as a result of temporary dipoles induced in atoms or molecules ion-induced dipole interaction dipole-induced dipole interaction

Intermolecular Forces
4. 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.

* Geckos! Geckos’ feet make use of London dispersion forces to climb almost anything. A gecko can hang on a glass surface using only one toe. Researchers at Stanford University recently developed a gecko-like robot which uses synthetic setae to climb walls Jesus Lizard

IMF Flowchart If a compound has Hydrogen Bonding it also has Dipole-Dipole & LDF Forces If a compound has Dipole-Dipole it also has LDF Forces If a compound has LDF as strongest IMF, no other forces present

IMF & Solubility Things with similar IMF tend to be miscible or soluble in one another

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.

Ex. Which of the following is most likely to exist as a gas at room temperature and normal pressure?
P4O10 Cl2 AgCl I2

Ex. Which of the following is most likely to exist as a gas at room temperature and normal pressure?
P4O10 Cl2 (smallest attractive forces) AgCl (largest attractive forces, least likely) I2

Ex. Of Bromine, Neon, Hydrochloric Acid, and Nitrogen-which is most likely to have the largest intermolecular dispersion forces? Ne Br2 HCl N2

Ex. Of Bromine, Neon, Hydrochloric Acid, and Nitrogen-which is most likely to have the largest dipole-dipole forces? Ne Br2 HCl N2

Ex. Rank the following in order of increasing boiling point: barium chloride, hydrogen, carbon monoxide, HF, Ne

Ex. Rank the following in order of increasing boiling point: From smallest to largest: 1. Hydrogen (LDF, smallest MW) 2. Ne (LDF, but larger than H2) 3. Carbon Monoxide (dipole-dipole) 4. HF (hydrogen bonding) 5. Barium Chloride (ionic compound)

BP(⁰C) IMF Explanation HF 20 HCl -85 HBr -67 HI -35
Ex: Identify all IMF present in a pure sample of each substance, then explain the boiling points. BP(⁰C) IMF Explanation HF 20 HCl -85 HBr -67 HI -35 London, dipole-dipole, H-bonds London, dipole-dipole Lowest MW/weakest London, but most polar/strongest dipole-dipole and has H-bonds Low MW/weak London, moderate polarity/dipole-dipole and no H-bonds Medium MW/medium London, moderate polarity/dipole-dipole and no H-bonds Highest MW/strongest London, but least polar bond/weakest dipole-dipole and no H-bonds HF 1.9 1.8 20 deg C

Intermolecular Forces Affect Many Physical Properties
The strength of the attractions between the particles can greatly affect the properties of a substance or solution

Viscosity Resistance to flow is called viscosity High Viscosity=SLOWER
LOW Viscosity=FASTER Viscosity Increases with: Stronger IMF Larger molecules Cooler temperatures

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

attracted to each other
Properties of Liquids Cohesion is the intermolecular attraction between like molecules Adhesion is an attraction between unlike molecules Adhesion attracted to glass Cohesion attracted to each other

Meniscus: curved upper surface of a liquid in a container; a relative measure of adhesive and cohesive forces Ex: Hg H2O (cohesion rules) (adhesion rules)

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. It results from the net inward force experienced by the moelecules on the surface of a liquid Strong intermolecular forces High surface tension

Vapor Pressure Vapor Pressure increases with temperature
When the vapor pressure of a liquid = atmospheric pressure, the liquid boils The normal boiling point of a liquid is the temperature at which its vapor pressure is 760 torr

Vapor Pressure Due to both temperature effects and energy transfers from collisions, molecules on the surface of a liquid are able to gain sufficient kinetic energy to escape into the atmosphere

Vapor Pressure At any temperature, some molecules in a liquid have enough energy to escape. As the temperature rises, the fraction of molecules that have enough energy to escape increases.

Vapor Pressure If the container is open to the atmosphere, the molecules simply escape. This process is called evaporation As molecules escape from the surface, they take energy with them resulting in a cooling effect on the liquid

Vapor Pressure If the container is closed to the atmosphere, as more molecules escape the liquid, the pressure they exert increases.

Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

The four compounds are all polar liquids containing an –OH group
The four compounds are all polar liquids containing an –OH group. The amount of hydrogen bonding differs as the molecular weight of the compound increases. Evaporation for the alcohols will occur according to molecular weight, with the lowest molecular weight compound evaporating first. So methanol will be first, followed by ethanol, and 2-propanol. Water will usually evaporate last. Why? The strong hydrogen bonding between water molecules affects evaporation (and boiling point) more than molecular weight.

Ice is less dense than water
Water is a Unique Substance Maximum Density 40C Density of Water Ice is less dense than water

Phase Changes

Energy Changes Associated with Changes of State
Heat of Fusion: Energy required to change a solid at its melting point to a liquid.

Heat of Vaporization: Energy required to change a liquid at its boiling point to a gas.

The heat added to the system at the melting and boiling points goes into pulling the molecules farther apart from each other. The temperature of the substance does not rise during the phase change.

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

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

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.

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.

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

Molar heat of fusion (DHfus) is the energy required to melt 1 mole of a solid substance.

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)

A phase diagram summarizes the conditions at which a substance exists as a solid, liquid, or gas.
The triple point is where all 3 phases meet. Phase Diagram of Water

Where vaporization occurs?
Can you find… The Triple Point? Critical pressure? Critical temperature? Where fusion occurs? Where vaporization occurs? Melting point (at 1 atm)? Boiling point (at 6 atm)? Carbon Dioxide

Phase Diagrams Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

Phase Diagrams The AB line is the liquid-vapor interface.
It starts at the triple point (A), the point at which all three states are in equilibrium.

Phase Diagrams It ends at the critical point (B); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

Phase Diagrams The AD line is the interface between liquid and solid.
The melting point at each pressure can be found along this line.

Phase Diagrams Below A the substance cannot exist in the liquid state.
Along the AC line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

Phase Diagram of Water Note the high critical temperature and critical pressure: These are due to the strong van der Waals forces between water molecules.

Phase Diagram of Water The slope of the solid–liquid line is negative.
This means that as the pressure is increased at a temperature just below the melting point, water goes from a solid to a liquid.

Phase Diagram of Carbon Dioxide
Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO2 sublimes at normal pressures.

Phase Diagram of Carbon Dioxide
The low critical temperature and critical pressure for CO2 make supercritical CO2 a good solvent for extracting nonpolar substances (such as caffeine).

Solids We can think of solids as falling into two groups:
Crystalline—particles are in highly ordered arrangement.

Solids Amorphous—no particular order in the arrangement of particles.

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

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

Types of Crystalline Solids
Ionic Solids – ions at the points of the lattice that describes the structure of the solid. Ion-ion is strong than all IM forces Atomic Solids – atoms at the lattice points that describe the structure of the solid. Stronger than IM forces but generally weaker than ion-ion Molecular Solids – discrete covalently bonded molecules at each of its lattice points. Held together with only IM forces

Examples of Three Types of Crystalline Solids

Three crystalline solids -- a) atomic solid, b) ionic
solid, and c) molecular solid.

Types of Crystals Ionic Crystals – Ion-Ion interactions are the strongest (including the “intermolecular forces” (H bonding, etc.) Lattice points occupied by cations and anions Held together by electrostatic attraction Hard, brittle, high melting point Poor conductor of heat and electricity CsCl ZnS CaF2

Electrostatic attractions
Ionic Anions and cations Electrostatic attractions Hard & brittle High MP Poor conductors Some solubility in H2O NaCl, Ca(NO3)2

Ionic Solids Generally have low vapor pressure due to the strong Coulombic interactions of positive and negative ions arranged in a 3D array Tend to be brittle due to the repulsion of like charges caused when one layer slides across another Do not conduct electricity. However, when melted they do conduct electricity because the ions are free to move When dissolved in water, ions are free to move so will conduct electricity Tend NOT to dissolve in nonpolar solvents because the attractions among ions are much stronger than the attractions among the separated ions and nonpolar solvent molecules

Types of crystalline solids (Table 11.6)
Particles Forces Notable properties Examples Atomic Atoms London dispersion Poor conductors Very low MP Ar (s), Kr (s)

Types of Crystals Network Atomic Covalent Crystals – Stronger than IM forces but generally weaker than ion-ion Lattice points occupied by atoms Held together by covalent bonds Hard, high melting point Poor conductor of heat and electricity carbon atoms diamond graphite

Covalent (a.k.a. covalent network)
Atoms bonded in a covalent network Covalent bonds Very hard Very high MP Generally insoluble Variable conductivity C (diamond & graphite) SiO2 (quartz) Ge, Si, SiC, BN Graphite Diamond SiO2

Covalent Network Solids
Only formed from nonmetals: elemental (graphite & diamond) or two nonmetals (silicon dioxide & silicon carbide) High melting points because all the bonds are covalently bonded Tend to be rigid and hard because the covalent bond angles are fixed Generally form in the carbon group because of their ability to form four covalent bonds

Covalent Network Solids-Graphite & Silicon
Has a high melting point because the covalent bonds between the carbon atoms making up each layer are relatively strong Soft because adjacent layers can slide past each other relatively easily; the major forces of attraction between the layers are LDF Silicon is a covalent network solid & a semiconductor Forms a 3D network similar to diamond Conductivity increases as temperature increases See later slide for doping to change from n-type semiconductor to p-type semiconductor

The Structures of Diamond and 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

London dispersion, dipole-dipole, H-bonds Poor conductors
Molecular Molecules (polar or non-polar) London dispersion, dipole-dipole, H-bonds Poor conductors Low to moderate MP CO2 (s), C12H22O11, H2O (s) Carbon dioxide (dry ice) Sucrose Ice

Molecular Solids Consist of nonmetals, diatomic elements, or compounds formed from two or more nonmetals Composed of distinct, individual units of covalently molecules attracted to each other through relatively weak IMFs Not expected to conduct electricity because their electrons are tightly held within the covalent bonds of each constituent molecule Generally have a low melting point because of the relatively weak IMFs present between the molecules Sometimes composed of very large molecules, or polymers, with important commercial & biological applications

Types of Crystals Metallic Crystals – Typically weaker than covalent, but can be in the low end of covalent 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 Cross Section of a Metallic Crystal nucleus & inner shell e- mobile “sea” of e-

Metal cations in a diffuse, delocalized e- cloud
Metallic Metal cations in a diffuse, delocalized e- cloud Metallic bonds Excellent conductors Malleable Ductile High but wide range of MP Cu, Al, Fe

Metallic Solids Can be represented as positive cores consisting of the nucleus and inner electrons surrounded by a sea of mobile valence electrons Good conductors because the electrons are delocalized and relatively free to move Malleable and ductile because deforming the solid does not change the environment immediately surrounding each metal core Often pure substances, but may also be mixtures called alloys Interstitial alloys form between atoms of different radius, where the smaller atoms fill the interstitial spaces between the larger atoms, like steel-carbon occupies space in iron. Makes the lattice more rigid, decreasing malleability and ductility Substitutional alloys form between atoms of similar radius, where one substitutes for the other in the lattice like brass-copper substituted with zinc. Remains malleable and ductile Alloys typically retain a sea of mobile electrons and will continue to conduct electricity In some cases, alloys change the chemistry of the surface. An example is the formation of chemically inert oxide layer in stainless steel

Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called closest packing. Each atom has 12 nearest neighbors. hexagonal closest packed (“aba”) cubic closest packed (“abc”)

Closest packing arrangement of uniform spheres --
aba. This forms hexagonal closest packed -- hcp.

Atoms arranged in aba pattern forming hexagonal
10_218 (a) (b) (a)) Top view Atom in third layer lies over atom in first layer. Atoms arranged in aba pattern forming hexagonal closest packed (hcp) structure -- 2 atoms/cell.

Hexagonal closest packed structure -- central
atom has 12 nearest neighbors.

Face-centered cubic is cubic closest packed
(ccp). The spheres are packed in an abc arrangement.

Bonding Models for Metals
Electron Sea Model: A regular array of metals in a “sea” of electrons. Band (Molecular Orbital) Model: Electrons assumed to travel around metal crystal in MOs formed from valence atomic orbitals of metal atoms. Conduction Bands: closely spaced empty molecular orbitals allow conductivity of heat and electricity.

Representation of the energy levels (bands) in a
magnesium crystal. 1s, 2s, & 2p orbitals are localized, but 3s & 3p orbitals are delocalized to make molecular orbitals.

Substances that have a mixture of elements and metallic properties.
Metal Alloys Substances that have a mixture of elements and metallic properties. 1. Substitutional Alloy: some metal atoms replaced by others of similar size. brass = Cu/Zn

Metal Alloys (continued)
2. Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. steel = iron + carbon 3. Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (Carbon) alloys.

Substitutional Alloy Interstitial Alloy

A substance in which some electrons can cross the band gap.
Semiconductors A substance in which some electrons can cross the band gap. Conductivity is enhanced by doping with group 3a or group 5a elements. n-type semiconductor -- doped with atoms having more valence electrons -- Phosphorus. p-type semiconductor -- doped with atoms having fewer valence electrons -- Boron. See Figure on page 476 in Zumdahl.

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