1. Intermolecular Forces © 2009, Prentice-Hall, Inc. A Molecular Comparison of Liquids and Solids The fundamental difference between states of matter is the distance between particles.

2. Intermolecular Forces © 2009, Prentice-Hall, Inc. States of Matter Because in the solid and liquid states particles are closer together, we refer to them as condensed phases.

3. Intermolecular Forces © 2009, Prentice-Hall, Inc. 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.

4. Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces The attractions between molecules are not nearly as strong as the intramolecular attractions that hold compounds together.

5. Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces They are, however, strong enough to control physical properties such as boiling and melting points, vapor pressures, and viscosities.

6. Intermolecular Forces © 2009, Prentice-Hall, Inc. Intermolecular Forces These intermolecular forces as a group are referred to as van der Waals forces. Dipole-dipole interactions Hydrogen bonding London dispersion forces

7. Intermolecular Forces © 2009, Prentice-Hall, Inc. Ion-Dipole Interactions Ion-dipole interactions (a fourth type of force), are important in solutions of ions. The strength of these forces are what make it possible for ionic substances to dissolve in polar solvents.

8. Intermolecular Forces © 2009, Prentice-Hall, Inc. Dipole-Dipole Interactions Molecules that have permanent dipoles are attracted to each other. –The positive end of one is attracted to the negative end of the other and vice- versa. –These forces are only important when the molecules are close to each other.

9. Intermolecular Forces © 2009, Prentice-Hall, Inc. Dipole-Dipole Interactions The more polar the molecule, the higher is its boiling point.

10. Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces While the electrons in the 1s orbital of helium would repel each other (and, therefore, tend to stay far away from each other), it does happen that they occasionally wind up on the same side of the atom.

11. Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces At that instant, then, the helium atom is polar, with an excess of electrons on the left side and a shortage on the right side.

12. Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces Another helium nearby, then, would have a dipole induced in it, as the electrons on the left side of helium atom 2 repel the electrons in the cloud on helium atom 1.

13. Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces London dispersion forces, or dispersion forces, are attractions between an instantaneous dipole and an induced dipole.

14. Intermolecular Forces © 2009, Prentice-Hall, Inc. London Dispersion Forces These forces are present in all molecules, whether they are polar or nonpolar. The tendency of an electron cloud to distort in this way is called polarizability.

15. Intermolecular Forces © 2009, Prentice-Hall, Inc. Factors Affecting London Forces The shape of the molecule affects the strength of dispersion forces: long, skinny molecules (like n-pentane tend to have stronger dispersion forces than short, fat ones (like neopentane). This is due to the increased surface area in n-pentane.

16. Intermolecular Forces © 2009, Prentice-Hall, Inc. Factors Affecting London Forces The strength of dispersion forces tends to increase with increased molecular weight. Larger atoms have larger electron clouds which are easier to polarize.

17. Intermolecular Forces © 2009, Prentice-Hall, Inc. Which Have a Greater Effect? Dipole-Dipole Interactions or Dispersion Forces If two molecules are of comparable size and shape, dipole-dipole interactions will likely the dominating force. If one molecule is much larger than another, dispersion forces will likely determine its physical properties.

18. Intermolecular Forces Practice 1. The dipole moments of acetonitrile, CH 3 CN, and methyl iodide, CH 3 I, are 3.9 D and 1.62 D, respectively. (a)Which of these substances has greater dipole–dipole attractions among its molecules? (b) Which of these substances has greater London dispersion attractions? (c) The boiling points of CH 3 CN and CH 3 I are 354.8 K and 315.6 K, respectively. Which substance has the greater overall attractive forces? 2. Of Br 2, Ne, HCl, HBr, and N 2, which is likely to have (a) the largest intermolecular dispersion forces, (b) the largest dipole–dipole attractive forces? © 2009, Prentice-Hall, Inc.

19. Intermolecular Forces © 2009, Prentice-Hall, Inc. How Do We Explain This? The nonpolar series (SnH 4 to CH 4 ) follow the expected trend. The polar series follows the trend from H 2 Te through H 2 S, but water is quite an anomaly.

20. Intermolecular Forces © 2009, Prentice-Hall, Inc. Hydrogen Bonding The dipole-dipole interactions experienced when H is bonded to N, O, or F are unusually strong. We call these interactions hydrogen bonds.

21. Intermolecular Forces © 2009, Prentice-Hall, Inc. Hydrogen Bonding Hydrogen bonding arises in part from the high electronegativity of nitrogen, oxygen, and fluorine. Also, when hydrogen is bonded to one of those very electronegative elements, the hydrogen nucleus is exposed.

22. Intermolecular Forces Practice 1. In which of the following substances is hydrogen bonding likely to play an important role in determining physical properties: methane (CH 4 ), hydrazine (H 2 NNH 2 ), methyl fluoride (CH 3 F), or hydrogen sulfide (H 2 S) ? 2. In which of the following substances is significant hydrogen bonding possible: methylene chloride (CH 2 Cl 2 ), phosphine (PH 3 ), hydrogen peroxide (HOOH), or acetone (CH 3 COCH 3 ) ? © 2009, Prentice-Hall, Inc.

23. Intermolecular Forces © 2009, Prentice-Hall, Inc. Summarizing Intermolecular Forces

24. Intermolecular Forces More practice 1.List the substances BaCl 2, H 2, CO, HF, and Ne in order of increasing boiling points. 2.Given the following: CH 3 CH 3, CH 3 OH, and CH 3 CH 2 OH. (a) Identify the intermolecular attractions present in the following substances (b) select the substance with the highest boiling point: © 2009, Prentice-Hall, Inc.

25. Intermolecular Forces © 2009, Prentice-Hall, Inc. Some Properties of Liquids The strength of the attractions between particles can greatly affect the properties of a substance or solution. With increasing IMF…. Higher Viscosity Higher Surface Tension Higher Boiling and Melting Point Lower Equilibrium Vapor Pressure

26. Intermolecular Forces © 2009, Prentice-Hall, Inc. Viscosity Resistance of a liquid to flow is called viscosity. It is related to the ease with which molecules can move past each other. Viscosity increases with stronger intermolecular forces and decreases with higher temperature.

27. Intermolecular Forces Properties of Liquids Cohesion is the intermolecular attraction between like molecules 11.3 Adhesion is an attraction between unlike molecules Adhesion Cohesion

28. Intermolecular Forces Concave Meniscus of Water Chemistry; The Science in Context; by Thomas R Gilbert, Rein V. Kirss, and Geoffrey Davies, Norton Publisher, 2004, p 458

29. Intermolecular Forces © 2009, Prentice-Hall, Inc. Surface Tension Surface tension results from the net inward force experienced by the molecules on the surface of a liquid.

30. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Changes

31. Intermolecular Forces © 2009, Prentice-Hall, Inc. Energy Changes Associated with Changes of State The heat of fusion is the energy required to change a solid at its melting point to a liquid.

32. Intermolecular Forces © 2009, Prentice-Hall, Inc. Energy Changes Associated with Changes of State The heat of vaporization is defined as the energy required to change a liquid at its boiling point to a gas.

33. Intermolecular Forces © 2009, Prentice-Hall, Inc. Energy Changes Associated with Changes of State 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 a phase change.

34. Intermolecular Forces Practice 1. Calculate the enthalpy change upon converting 1.00 mol of ice at –25 °C to water vapor (steam) at 125 °C under a constant pressure of 1 atm. The specific heats of ice, water, and steam are 2.03 J/g-K, 4.18 J/g-K, and 1.84 J/g-K, respectively. For H 2 O, ΔH fus = 6.01 kJ/mol and ΔH vap = 40.67 kJ/mol. 2. What is the enthalpy change during the process in which 100.0 g of water at 50.0 °C is cooled to ice at –30.0 °C ? © 2009, Prentice-Hall, Inc.

35. Intermolecular Forces © 2009, Prentice-Hall, Inc. 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.

36. Intermolecular Forces © 2009, Prentice-Hall, Inc. Vapor Pressure As more molecules escape the liquid, the pressure they exert increases.

37. Intermolecular Forces © 2009, Prentice-Hall, Inc. Vapor Pressure The liquid and vapor reach a state of dynamic equilibrium: liquid molecules evaporate and vapor molecules condense at the same rate.

38. Intermolecular Forces © 2009, Prentice-Hall, Inc. Vapor Pressure The boiling point of a liquid is the temperature at which it’s vapor pressure equals atmospheric pressure. The normal boiling point is the temperature at which its vapor pressure is 760 torr. * Estimate the boiling point of diethyl ether under an external pressure of 0.80 atm. * At what external pressure will ethanol have a boiling point of 60 °C?

39. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams (prior knowledge) Phase diagrams display the state of a substance at various pressures and temperatures and the places where equilibria exist between phases.

40. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams The circled line is the liquid-vapor interface. It starts at the triple point (T), the point at which all three states are in equilibrium.

41. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams It ends at the critical point (C); above this critical temperature and critical pressure the liquid and vapor are indistinguishable from each other.

42. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams Each point along this line is the boiling point of the substance at that pressure.

43. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams The circled line in the diagram below is the interface between liquid and solid. The melting point at each pressure can be found along this line.

44. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagrams Below the triple point the substance cannot exist in the liquid state. Along the circled line the solid and gas phases are in equilibrium; the sublimation point at each pressure is along this line.

45. Intermolecular Forces © 2009, Prentice-Hall, Inc. 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.

46. Intermolecular Forces © 2009, Prentice-Hall, Inc. 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.

47. Intermolecular Forces © 2009, Prentice-Hall, Inc. Phase Diagram of Carbon Dioxide Carbon dioxide cannot exist in the liquid state at pressures below 5.11 atm; CO 2 sublimes at normal pressures.

48. Intermolecular Forces © 2009, Prentice-Hall, Inc. Structures of Solids We can think of solids as falling into two groups: –crystalline, in which particles are in highly ordered arrangement.

49. Intermolecular Forces © 2009, Prentice-Hall, Inc. Solids We can think of solids as falling into two groups: –amorphous, in which there is no particular order in the arrangement of particles.

50. Intermolecular Forces © 2009, Prentice-Hall, Inc. Attractions in Ionic Crystals In ionic crystals, ions pack themselves so as to maximize the attractions and minimize repulsions between the ions.

51. Intermolecular Forces © 2009, Prentice-Hall, Inc. Bonding in Solids

52. Intermolecular Forces © 2009, Prentice-Hall, Inc. Covalent-Network and Molecular Solids Diamonds are an example of a covalent- network solid, in which atoms are covalently bonded to each other. –They tend to be hard and have high melting points.

53. Intermolecular Forces © 2009, Prentice-Hall, Inc. Covalent-Network and Molecular Solids Graphite is an example of a molecular solid, in which atoms are held together with van der Waals forces. –They tend to be softer and have lower melting points.

54. Intermolecular Forces Metallic Bonding, Alloys & Semiconductors Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

55. Intermolecular Forces © 2009, Prentice-Hall, Inc. Metallic Solids Metals are not covalently bonded, but the attractions between atoms are too strong to be van der Waals forces. In metals valence electrons are delocalized throughout the solid.

56. Intermolecular Forces Alloys Alloys contain more than one element and have the characteristic properties of metals. Solid Solution alloys are homogeneous mixtures. Heterogeneous alloys: The components are not dispersed uniformly (e.g., pearlite steel has two phases: almost pure Fe and cementite, Fe 3 C). Pure metals and alloys have different physical properties. An alloy of gold and copper is used in jewelry (the alloy is harder than the relatively soft pure 24 karat gold). 14 karat gold is an alloy containing 58% gold. Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

57. Intermolecular Forces Metal Alloys-Solid Solutions Substance has mixture of element and metallic properties. 1.Substitutional Alloy: some metal atoms replaced by others of similar size. Electronegativities usually are similar. The atoms must have similar atomic radii. The elements must have similar bonding characteristics. brass = Cu/Zn Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

58. Intermolecular Forces Metal Alloys (continued) 2.Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. Solute atoms occupy interstices “small holes” between solvent atoms. One element (usually a nonmetal) must have a significantly smaller radius than the other (in order to fit into the interstitial site). steel = iron + carbon 3.Both types: Alloy steels contain a mix of substitutional (Cr, Mo) and interstitial (Carbon) alloys. Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

59. Intermolecular Forces Substitutional Alloy Interstitial Alloy Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

60. Intermolecular Forces Alloys vs. Pure Metal The alloy is much harder, stronger, and less ductile than the pure metal (increased bonding between nonmetal and metal). An example is steel (contains up to 3% carbon). mild steels (<0.2% carbon) - useful for chains, nails, etc. medium steels (0.2-0.6% carbon) - useful for girders, rails, etc. high-carbon steels (0.6-1.5% carbon) - used in cutlery, tools, springs. Other elements may also be added to make alloy steels. Addition of V and Cr increases the strength of the steel and improves its resistance to stress and corrosion. The most important iron alloy is stainless steel. It contains C, Cr (from ferrochrome, FeCr 2 ), and Ni. Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

61. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

62. Intermolecular Forces Which two substances are most likely to form an interstitial alloy? Nickel and titanium Silver and tin Tin and lead Copper and zinc Tungsten and carbon Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

63. Intermolecular Forces Which two substances are most likely to form an interstitial alloy? Nickel and titanium Silver and tin Tin and lead Copper and zinc Tungsten and carbon Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

64. Intermolecular Forces Bonding Models for Metals APSI 2014 PWISTA.com

65. Intermolecular Forces Bonding Models for Metals Electron Sea Model: A regular array of metals in a “sea” of electrons. The electron-sea model is a qualitative interpretation of band theory (molecular-orbital model for metals). 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. APSI 2014 PWISTA.com

66. Intermolecular Forces APSI 2014 PWISTA.com Molecular Orbital Theory Recall that atomic orbitals mix to give rise to molecular orbitals.

67. Intermolecular Forces Molecular-Orbital Model for Metals Delocalized bonding requires the atomic orbitals on one atom to interact with atomic orbitals on neighboring atoms. Example: Graphite electrons are delocalized over a whole plane, while benzene molecules have electrons delocalized over a ring. Recall that the number of molecular orbitals is equal to the number of atomic orbitals. Each orbital can hold two electrons. In metals there are a very large number of orbitals. As the number of orbitals increases, their energy spacing decreases and they band together. The available electrons do not completely fill the band of orbitals. APSI 2014 PWISTA.com

68. Intermolecular Forces APSI 2014 PWISTA.com

69. Intermolecular Forces Molecular-Orbital Model for Metals Therefore, electrons can be promoted to unoccupied energy bands. Because the energy differences between orbitals are small the promotion of electrons requires little energy. As we move across the transition metal series, the antibonding band starts becoming filled. Therefore, the first half of the transition metal series has only bonding-bonding interactions and the second half has bonding–antibonding interactions. We expect the metals in the middle of the transition metal series (group 6B) to have the highest melting points. The energy gap between bands is called the band gap. APSI 2014 PWISTA.com

70. Intermolecular Forces APSI 2014 PWISTA.com Molecular Orbital Theory In such elements, the energy gap between molecular orbitals essentially disappears, and continuous bands of energy states result.

71. Intermolecular Forces Formation of Bands APSI 2014 PWISTA.com When atoms come together to form a compound, their atom orbital energies mix to form molecular orbital energies. As more atoms begin to mix and more molecular orbitals are formed, it is expected that many of these energy levels will start to be very close to, or even completely degenerate, in energy. These energy levels are then said to form bands of energy remember each orbital only holds two electrons.

72. Intermolecular Forces The electronic band structure of nickel. The left side of the figure shows the electron configuration of a single Ni atom, while the right-hand side of the figure shows how these orbital energy levels broaden into energy bands in bulk nickel. The horizontal dashed gray line denotes the position of the Fermi Level, which separates the occupied molecular orbitals (shaded in blue) from the unoccupied molecular orbitals. APSI 2014 PWISTA.com

73. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com Types of Materials Rather than having molecular orbitals separated by an energy gap, these substances have energy bands. The gap between bands determines whether a substance is a metal, a semiconductor, or an insulator.

74. Intermolecular Forces Energy bands in metals, semiconductors, and insulators. Metals are characterized by the highest-energy electrons occupying a partially filled band. Semiconductors and insulators have an energy gap that separates the completely filled band (shaded in blue) and the empty band (unshaded), known as the band gap and represented by the symbol E g. The filled band is called the valence band (VB), and the empty band is called the conduction band (CB). Semiconductors have a smaller band gap than insulators. Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

75. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com Metals Valence electrons are in a partially-filled band. There is virtually no energy needed for an electron to go from the lower, occupied part of the band to the higher, unoccupied part. This is how a metal conducts electricity.

76. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com Insulators The energy band gap in insulating materials is generally greater than ~350 kJ/mol. They are not conductive.

77. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com Semiconductors Semiconductors have a gap between the valence band and conduction band of ~50-300 kJ/mol.

78. Intermolecular Forces An intrinsic semiconductor is a semiconductor in its pure state. For every electron that jumps into the conduction band, the missing electron will generate a hole that can move freely in the valence band. The number of holes will equal the number of electrons that have jumped. The higher the temp more electrons into conduction band. Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

79. Intermolecular Forces The following pictures show the electron populations of the bands of MO energy levels for four different materials: Classify each material as an insulator, a semiconductor, or a metal. Arrange the four materials in order of increasing electrical conductivity. Explain. Tell whether the conductivity of each material increases or decreases when the temperature increases. (a) (b) (c) Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

80. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com Semiconductors Among elements, only silicon, germanium and graphite (carbon), all of which have 4 valence electrons, are semiconductors. Inorganic semiconductors (like GaAs) tend to have an average of 4 valence electrons (3 for Ga, 5 for As).

81. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com Doping By introducing very small amounts of impurities that have more valence electrons (n-Type) or fewer (p- Type) valence electrons, one can increase or decrease the conductivity of a semiconductor.

82. Intermolecular Forces The addition of controlled small amounts of impurities (doping) to a semiconductor changes the electronic properties of the material. Left: A pure, intrinsic semiconductor has a filled valence band and an empty conduction band (ideally). Middle: The addition of a dopant atom that has more valence electrons than the host atom adds electrons to the conduction band (i.e., phosphorus doped into silicon). The resulting material is an n-type semiconductor. Right: The addition of a dopant atom that has fewer valence electrons than the host atom leads to fewer electrons in the valence band or more holes in the valence band (i.e., aluminum doped into silicon). The resulting material is a p- type semiconductor. Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

83. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

84. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

85. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

86. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

87. Intermolecular Forces Which of the following is a p-type semiconductor? Sulfur-doped carbon Boron-doped germanium Phosphorus-doped silicon Ultra-pure silicon Carbon-doped copper Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

88. Intermolecular Forces Which of the following is a p-type semiconductor? Sulfur-doped carbon Boron-doped germanium Phosphorus-doped silicon Ultra-pure silicon Carbon-doped copper Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

89. Intermolecular Forces Which of the following elements, if doped into silicon, would yield an n-type semiconductor? Ga; As; C. Suggest an element that could be used to dope silicon to yield a p-type material. Practice Exercises Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

90. Intermolecular Forces Diode- Used to switch and convert between electromagnetic radiation and electric current Semiconductor created that has p-type on one half and n-type on the other half Known as “p-n junction” Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

91. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

92. Intermolecular Forces Light emitting diodes. The heart of a light emitting diode is a p-n junction where an applied voltage drives electrons and holes to meet. Bottom: The color of light emitted depends upon the band gap of the semiconductor used to form the p-n junction. For display technology red, green, and blue are the most important colors because all other colors can be made by mixing these colors. Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com

93. Intermolecular Forces Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com Color λ Voltage DropComposition Red610 < λ < 7601.63 < ΔV < 2.03 Aluminium gallium arsenideAluminium gallium arsenide (AlGaAs) Gallium arsenide phosphide (GaAsP) Aluminium gallium indium phosphide (AlGaInP) Gallium(III) phosphide (GaP) Gallium arsenide phosphide Aluminium gallium indium phosphide Gallium(III) phosphide Orange590 < λ < 6102.03 < ΔV < 2.10 Gallium arsenide phosphideGallium arsenide phosphide (GaAsP) Aluminium gallium indium phosphide (AlGaInP) Gallium(III) phosphide (GaP) Aluminium gallium indium phosphide Gallium(III) phosphide Yellow570 < λ < 5902.10 < ΔV < 2.18 Gallium arsenide phosphideGallium arsenide phosphide (GaAsP) Aluminium gallium indium phosphide (AlGaInP) Gallium(III) phosphide (GaP) Aluminium gallium indium phosphide Gallium(III) phosphide Green500 < λ < 5701.9 [63] < ΔV < 4.0 [63] Traditional green: Gallium(III) phosphide (GaP) Aluminium gallium indium phosphide (AlGaInP) Aluminium gallium phosphide (AlGaP) Pure green: Indium gallium nitride (InGaN) / Gallium(III) nitride (GaN) Gallium(III) phosphide Aluminium gallium indium phosphide Aluminium gallium phosphide Indium gallium nitrideGallium(III) nitride Blue450 < λ < 5002.48 < ΔV < 3.7 Zinc selenideZinc selenide (ZnSe) Indium gallium nitride (InGaN) Silicon carbide (SiC) as substrate Silicon (Si) as substrate—under development Indium gallium nitride Silicon carbide Silicon Violet400 < λ < 4502.76 < ΔV < 4.0Indium gallium nitrideIndium gallium nitride (InGaN)

94. Intermolecular Forces Solar Cells Presented By, Mark Langella, APSI Chemistry 2014, PWISTA.com