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Intermolecular Forces

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Presentation on theme: "Intermolecular Forces"— Presentation transcript:

1 Intermolecular Forces
The state of matter depends upon intermolecular forces. Intermolecular forces are the electrostatic attractions that molecules feel for one another. To understand intermolecular forces,one must first understand intramolecular forces. Intramolecular force is interpreted using the electronegativity of the atoms combining to form compounds and molecules.

2 Electronegativity Electronegativity is the same property as electron
affinity, except it is for atoms in compounds or molecules. Electronegativity is the desire of atoms to draw electrons when they combine with other atoms. The larger the atomic radius, the harder it becomes to attract electrons to an atom. The larger the atomic radius, the lower the electronegativity of an atom.

3 Electronegativity Atomic Radii Decrease Electronegativity Increase Electronegativity Decrease Atomic Radii Increase

4 Intramolecular Forces – A Review
Intramolecular forces are the forces that bind atoms together resulting in molecular bonding. Intramolecular forces, along with molecular geometry will be used to interpret intermolecular forces. Intramolecular forces fall into one of three categories: 1. Ionic bonding 2. Covalent bonding 3. Polar covalent bonding

5 Ionic Bonding Ionic bonds are formed when the combining atoms
have large differences in electronegativity. Ionic bonds are formed when the combining atoms are separated by large distances in the periodic table. Elements from the left lose electrons (becoming cations) and elements from the right gain electrons (becoming anions). The electrostatically opposites attract to form an ionic compound.

6 The Covalent Bond Covalent bonds are formed by combining atoms with
exactly the same electronegativities. Because only elements of the same type have exactly the same electronegativity, covalent bonds are formed by elements combining with themselves. Because each element in the bond has exactly the same electrongegativity, covalent bonds are non-polar. Covalently bound materials exist as individual entities called molecules. They do not repeat in a lattice like ionic compounds.

7 The Polar Covalent Bond
The polar covalent bond is an intermediate form of bonding. Polar covalent bonds are formed between elements with similar electronegativities. Polar covalent bonds are formed between elements that are nearby each other in the periodic table. Ionic bonds are formed between elements with large differences in electronegativity (far apart in the table). Covalent bond are formed between elements with the same electronegativity (combining with themselves). Polar covalent bonds exhibit some degree of electron sharing with some degree of polarity.

8 Classifying the Type of Bonding
N KBr PCl LiNO H SO2 polar covalent covalent ionic ionic covalent polar covalent

9 The Bond Dipole In polar covalent bonds, electrons are not shared
evenly between the two atoms. The more electronegative element in the bond acquires more of the electron density than the less electronegative element. This leads to an imbalance in charge distribution called a bond dipole.

10 F Cl N O Si S C H The Bond Dipole
Bond dipoles in a molecule are indicated by drawing an arrow pointing from the less electronegative element towards the more electronegative element. F Cl N O Si S C H Atoms bonded to themselves are covalent bonds with no polarity and no bond dipole!

11 The Dipole Moment In polyatomic polar covalent molecules, there
are multiple bond dipoles. When these bond dipoles are in opposition, they cancel leading to a net non-polar molecule. When these bond dipoles are not in opposition, or are in opposition but are not of equal strength, it leads to a net polarity called a dipole moment.

12 Non-Polar Molecules 120o 180o 109.5o Trigonal planar molecule
Linear molecule 120o 180o Tetrahedral molecule Octahedral Molecule 90o 109.5o

13 Polar Molecule Examples
- + O is more electronegative than S - + Total dipole moment is directed tail at more positive head at more negative No bond dipoles along S-S bonds + - - +

14 Total Dipole Examples d- d+ non-polar d- d+ d- d+

15 Intermolecular Force in Non-Polar Molecules
Molecules with symmetric charge distribution have very little attraction for one another. This is the reason that most of them are gaseous at room temperature. However, they must possess some mechanism for attraction, otherwise they would never condense. Their mechanism of attraction is called Dispersion Forces.

16 Dispersion Forces Non-polar molecules have symmetric charge
distribution. Nuclei Electronic charge distributed evenly over the nuclear framework. The net result is a non-polar molecule with little mechanism for intermolecular force

17 Dispersion Forces In the influence of an electric field, this symmetric charge distribution can be induced to distort. Temporary dipole moment + - Positive charge (nucleus) orients toward negative pole. Electron density piles up near positive pole.

18 Systems with Dispersion Force Interactions
Group VIIIA Elements (Noble Gases) He Ne Ar Kr Xe Rn Homonuclear diatomics O O Other homonuclear diatomics: H2 N2 F2 Cl2 Br2 I2 Non-polar polyatomics

19 Dispersion Force and Size
In symmetric charge distributions, dispersion forces are the only mechanism of attraction. The charge distributed around larger systems feels less electrostatic attraction, and is more polarizable. Dispersion forces increase with the size of system Increasing Strength of Dispersion forces He < Ne < Ar < Kr < Xe < Rn F2 < Cl2 < Br2 < I2

20 The Hydrogen Bond Because of their relatively small size and
extremely high electronegativity, the elements F, O, and N bond to H and form an extremely polar bond. The resulting polarity creates unusually strong intermolecular force between molecules. This intermolecular force is called the Hydrogen Bond.

21 Water is a “really bent”
Hydrogen Bonding in Water Water is a “really bent” molecule Lone pairs + -

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23 Hydrogen Bonding and Intramolecular Force

24 Hydrogen Bonding in the Alpha - Helix
The alpha-helix is a corkscrew strucutre held in place by hydrogen bonding between N-H groups and C=O groups in the next turn of the helix four amino acids away. This creates a “telephone cord” affect.

25 Hydrogen Bonding in the Beta-Pleated Sheet
In a beta-pleated sheet, strands of polypeptide chains are held together side by side by hydrogen bonds between the peptide chains.

26 Hydrogen Bonding in the Double Helix
A double helix is formed by intertwining strands of polypeptide chains that hydrogen bond. DNA is a double helix

27 Hydrogen Bonding in Starch
Hydrogen bonding between glucose monomers makes the conformation of the amylose polymer coil into a helical structure known as starch.

28 Hydrogen Bonding and the Triple Helix
Collagen makes up about 1/3 of all protein in vertibrates Collagen is a triple helix that forms the connective tissue (ligands and tendons, etc) in animals.

29 Hydrogen Bonding and Tertiary Structure
The tertiary structure of a protein involves attractions and repulsions between the side chain groups of the amino acids in the polypeptide chain. This creates twists and bends in the in the protein until it acquires a specific three-dimensional shape that determines its biological function (or structure-activity relationship).

30 Summary of Intermolecular Forces
There are various types of distinct intermolecular force which molecules can employ: 1. Ionic Forces - Forces exerted by cations and anions in ionic compounds 2. Hydrogen bonding – Exerted by polar molecules which contain H bonded to F, O, or N. 3. Dipole forces – Exerted by polar covalent molecules with unsymmetric charge distributions such that they possess a permanent dipole moment. 4. Dispersion forces – Exerted by non-polar molecules when they acquire temporary dipoles.

31 Hierarchy of Intermolecular Forces
Strongest Weakest Ionic Forces Hydrogen Bonding > Dipole Forces > Dispersion Forces >

32 Summary of Intermolecular Forces
There are various types of distinct intermolecular force which molecules can employ: 1. Ionic Forces - Forces exerted by cations and anions in ionic compounds 2. Hydrogen bonding – Exerted by polar molecules which contain H bonded to F, O, or N. 3. Dipole forces – Exerted by polar covalent molecules with unsymmetric charge distributions such that they possess a permanent dipole moment. 4. Dispersion forces – Exerted by non-polar molecules when they acquire temporary dipoles.

33 Hierarchy of Intermolecular Forces
Strongest Weakest Ionic Forces Hydrogen Bonding > Dipole Forces > Dispersion Forces >

34 Classifying Intermolecular Forces
The interaction in the ionic lattice is characterized as ion-ion interaction. Na+ Cl- The interaction in pure water is best characterized as hydrogen bonding interaction. The interaction between water molecules and ions in solution is best characterized as ion - dipole interaction. Na+

35 Classifying Intermolecular Forces
Ammonia - Water Solution Hydrogen Bonding PCl3 - Water Solution Dipole-Dipole Interaction Gas mixture (such as air) - Dispersion Forces

36 Consequences of Polarity
The more polar a molecule the more intermolecular forces exerted among a bulk sample of molecules. Physical properties of bulk sample of these molecules are affected by their degree of attraction to one another. These properties include: Bulk Phase Vapor Pressure Boiling Point Melting Point Solubility

37 Intermolecular Force and Bulk Phase
Because of their highly polar bonds, ionic compounds are solids at ambient temperatures. CaCl2 Na2S MgSO4 LiNO3 K3PO4

38 Intermolecular Force and Bulk Phase
Because of their non-polar charge distributions the covalent diatomics only exert dispersion forces and are mostly gases at ambient temperatures. H2 (g) N2 (g) O2 (g) F2 (g) Cl2 (g) Br2 (l) Larger size, higher dispersion forces. Even larger size, even more dispersion Forces. I2 (s)

39 Intermolecular Force and Bulk Phase
In polar covalent molecules, the geometry and dipole moment play a role in the polarity and phase of matter. Of course, temperature and pressure play significant roles in the bulk phase of matter as well. At atmospheic pressure, water becomes a solid at temperatures of 0 oC and below. At pressures over the atmospheric value, the freezing point of water is lowered below 0 oC. At atmospheic pressure, water exists as a vapor at temperatures of 100 oC and above. At pressures over the atmospheric value, the boiling point of water is raised above 100 oC.

40 Bulk Phase of Polar Covalent Molecules
Carbon Dioxide is a gas at room temperature d- Water is a liquid at room temperature d+ Sucrose is a solid at room temperature

41 Vapor Pressure Molecules that are loosely held in the liquid or solid
phase can break free to the gaseous phase and exert a vapor pressure. The lower the intermolecular force between molecules in a sample, the more escapable its molecules become. The lower the intermolecular force between molecules in a sample, the higher its vapor pressure.

42 Vapor Pressure H2O(l) H2O(g) H2O(g)  H2O(l) (Evaporation)
Without a lid, water vapor can escape A lid prohibits water vapor from escaping H2O(l) H2O(g) (Evaporation) H2O(g)  H2O(l) (Condensation) When the rate of evaporation matches the rate of condensation we reach equilibrium: H2O(l) H2O(g)

43 Vapor Pressure and Intermolecular Force
Because of the high polarity of their bonds, ionic compounds have high intermolecular force and therefore exhibit no vapor pressure. Nonvolatile substance- A material with no vapor pressure.

44 Vapor Pressure and Intermolecular Force
Because of the non-polar nature of their bonds, covalent molecules have low intermolecular force and therefore exhibit high vapor pressure. Although Br2 is a liquid and I2 a solid at room temperature, they still exhibit very high vapor pressures. Volatility - measures the vapor pressure exerted by a substance.

45 Vapor Pressure and Intermolecular Force
Because of the slight polarity of their bonds, polar covalent molecules may or may not exhibit strong intermolecular forces, depending on their geometry.

46 Vapor Pressure of Polar Covalents
Isomers - molecules with the same chemical formula, but different structures. Polar d- d+ Non-polar Lower Vapor Pressure Higher Vapor Pressure

47 Temperature and Vapor Pressure
As temperature increases, molecules move more rapidly, so that more can escape the liquid phase. As temperature increases, vapor pressure increases. Vapor Pressure of H2O at Various Temperatures 0 o C mm Hg 60 o C mm Hg 20 o C mm Hg 80 o C mm Hg 40 o C mm Hg 100 o C mm Hg

48 The Boiling Point The boiling point of a substance is the
temperature where is vapor pressure matches external pressure. The normal boiling point of a substance is when its vapor pressure matches sea level atmospheric pressure (1 atm or 760 mm Hg). Substances at high altitude boil at lower temperatures. Substances under high pressures boil at higher temperatures.

49 Boiling Point and Intermolecular Force
The higher the intermolecular force between molecules in a bulk sample, the more “stickability” of their molecules The higher the intermolecular force between molecules in a bulk sample, the higher its boiling point.

50 Boiling Points of the Noble Gas Elements
Lowest Highest He < Ne < Ar < Kr < Xe < Rn Source of intermolecular forces – Dispersion forces. Relative size of Group VIIIA elements: Smallest Largest He < Ne < Ar < Kr < Xe < Rn

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52 Boiling Points of Hydrides
Lowest Highest Group IVA: CH4 < SiH4 < GeH4 < SnH4 Group VA: PH3 < AsH3 < SbH3 < NH3 Hydrogen bonding – higher intermolecular force Group VIA: H2S < H2Se < H2Te < H2O Group VIIA: HCl < HBr < HI < HF

53 Boiling Points of Polar Covalent Isomers
d- d+ Polar Non-Polar Higher boiling point Lower boiling point

54 Melting Points X(s)  X(l)
Melting point is the change of phase from solid to liquid – Scientists call this fusion. Melting point trends follow the same pattern as boiling point trends. Materials with high intermolecular force will exert strong attractions to one another, and therefore melt at higher temperatures.

55 Solubility of Polar Covalent Isomers
Polar – Higher solubility in water d- d+ Non-polar – Lower solubility in water

56 Solubility Examples Ammonia (NH3) d- d+ Soluble in water,
ethanol, acetic acid Soluble in toluene, benzene, CCl4 Insoluble in toluene, benzene, CCl4 Insoluble in water, ethanol, acetic acid

57 Soaps and Detergents Polar region
Fats and oils are generally non-polar materials that will not dissolve in water. Soaps and Detergents are designed to make these substances water soluble. Sodium Lauryl Sulfate Non-polar region dissolve in grease Polar region dissolves in water

58 Detergent Action Polar end of detergent points to polar water
molecules Non-Polar end of detergent points toward non-polar grease Polar Water


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