Intermolecular Forces:

Intermolecular Forces:
Intermolecular Forces: What holds everything together (Chapter 14)

Intramolecular forces (bonds)
Hold atoms together in molecules Have high energy associated with them it’s difficult to break molecules into their individual atoms Different types based upon what is going on with the electrons (electron clouds)

Types of bonds: Ionic Covalent
Ionic attraction between fully charged molecules/ atoms NaCl, made from Na+ and Cl-; or Ca(OH)2, made from Ca2+ and 2OH- Covalent electrons are shared between atoms, water (H2O) and sugar (C6H12O6) Can be polar or nonpolar Based on electronegativity VSEPR geometry (shape)

Intermolecular forces (IMFs)
Hold molecules together Much weaker than intramolecular forces Intramolecular bonds are usually 100x or even 1000x stronger *(kJ are units of energy like Calories; 1Cal= 4.184kJ) 1000cal= 1Cal 1cal =4.184J

Figure 14. 2: Intermolecular forces exist between molecules
Figure 14.2: Intermolecular forces exist between molecules. Bonds exist within molecules.

Why do we care? The strength of the IMFs determine the state of matter
The strength of the IMFs determine the state of matter Solid, liquid, or gas* *Not plasma- intramolecular bonds are broken to get plasmas

level of organization*
Solids, Liquids, and Gases shape volume density energy* motion with Energy* level of organization* strength of IMFs* Gas indefinite variable with P and T variable with volume change high high; molecules freely moving with great distance compared to molecular size between them very low low Liquid constant** moderate high; molecules freely moving past each other but in close proximity to each other Solid definite low; vibration only as molecules are basically fixed in place *all at room temperature, ~25C **small variations occur due to temperature changes, very little variable with pressure changes

Things with strong IMFs tend to be solids at room temperature
Things with strong IMFs tend to be solids at room temperature Things with weak IMFs tend to be gases at room temperature Medium IMFs tend to be in between- liquids, yes, but with varying characteristics Amorphous solids: long transition between solid and liquid states- gets soft, then melts (like wax) Crystalline solids: definite, clear melting point (no soft transition- ie: ice)

Types of IMFs In order of increasing strength:
In order of increasing strength: London dispersion forces Dipole- dipole Hydrogen bonds

London dispersion forces
LDFs occur in all molecules, but are the only forces that are present in nonpolar molecules such as diatomic molecules and atomic substances CO2, N2, He They occur because the electron clouds around molecules are not always evenly distributed. When the electron clouds are unevenly distributed, temporary partial charges result

Figure 14.6: Atoms with spherical electron probability.
Figure 14.6: Atoms with spherical electron probability. 14.6: The atom on the left develops an instantaneous dipole.

LDFs, con’t These temporary partial charges are called induced or temporary dipoles This temporary dipole forming in a nonpolar substance is strong enough to cause a dipole to occur in a neighboring molecule

Figure 14.3: (a) Interaction of two polar molecules. (b) Interaction of many dipoles in a liquid.

LDfs, con’t Basically, everything lines up temporarily, but long enough to keep everything together Common in gases

See LDFs at work here These dipoles fluctuate; they do not last very long, but they do occur frequently enough to have a significant effect overall

Dipole- dipole forces:
Are stronger than LDfs because they occur in polar molecules that already have permanent dipole moments (in other words, partial charges already exist) Are AKA as van der Waals interactions at times, but in actuality both induced dipole attractions and dipole-dipole attractions are van der Waals forces

Examples HCl and other acids* HCN NH3
HCl and other acids* HCN NH3 *except HF, which does something else

What would happen between polar and nonpolar molecules? (Do forces of attraction exist? Do the molecules repel?) Explain!

Hydrogen bonding Are stronger than dipole-dipole forces or LDFs
Are stronger than dipole-dipole forces or LDFs Occurs in only the most polar bonds between molecules containing H-F, H-O and H-N bonds only Are the reason that water is so different from any material from similar atoms, like H2S

Figure 14.4: Hydrogen bonding among water molecules. Norton Interactive: IMFs tutorial Select Hydrogen bonding in water from bottom of list

(note: I am not responsible for the music on the above web site) Polarity and hydrogen bond formation Ice at different temperatures

Which is ice? Which is liquid water? Explain.
Which is ice? Which is liquid water? Explain. Ice at different temperatures

Water is special because…
It has a high specific heat, meaning that it takes a lot of energy to raise the temperature of a sample of water by even 1 degree Specific heat of water (c)= 1 cal/ g°C or 4.184J /g°C The solid phase is LESS dense than the liquid phase, so ice floats on water It’s a good solvent for many substances due its polarity H2O is liquid at RT, where H2S is a gas

Figure 14.5: The boiling points of covalent hydrides.

Water is special And water would not be special without hydrogen bonding H bonding plays vital roles in DNA (holding together the chains of DNA) Protein shape (and therefore the protein’s function; think hair!)

H bonding in dna1

H bonding in DNA

Amino Acids- they make proteins

Protein Structures

Protein Structure and H Bonding

For the next slides: Determine polarity of group
Determine polarity of group Determine type of IMFs are possible in group Determine if the group will be highly soluble in water

IMFs in proteins

Sickle Cell Anemia Glu (glutamic acid) replaced by Val (valine)

What would happen if a molecule capable of H-bonding comes into contact with: A nonpolar substance A polar substance that does not H-bond

Strength increases from left to right; when ions are involved, attractive forces are greater than when they are not involved.

Dealing with this pic… Ion- dipole forces Ionic Bonding
Dealing with this pic… Ion- dipole forces Ionic Bonding Basically electrostatic attractive forces between positive and negative charges Strong

IMFs influence… Boiling point/ Melting Point Viscocity Surface Tension
Boiling point/ Melting Point Viscocity Surface Tension Capillary Action Vapor pressure/ rate of evaporation State of Matter (at room temp) Density falls here, but can vary even within state

IMFs and mass The mass of a material makes a difference, so yes, mass (size) matters Larger molecules have stronger forces than similar molecules that are smaller (in terms of mass)

Figure 14.5: The boiling points of covalent hydrides.

Boiling points and masses of noble gases
Helium: -269°C g/mol Neon: °C g/mol Argon: °C g/mol Krypton: -152°C g/mol Xenon: °C g/mol radon -62°C ~222 g/mol Larger atoms have larger e- clouds, which lead to greater polarizability

Saturated Hydrocarbons, or Alkanes
Name Molecular Melting Boiling State at Formula Point 25oC (oC) methane CH4 -183 -164 gas ethane C2H6 -89 propane C3H8 -190 -42 butane C4H10 -138 -0.5 pentane C5H12 -130 36 hexane C6H14 -95 69 heptane C7H16 -91 98 octane C8H18 -57 125 nonane C9H20 -51 151 liquid decane C10H22 -30 174 undecane C11H24 -25 196 dodecane C12H26 -10 216 eicosane C20H42 37 343 triacontane C30H62 66 450 solid Saturated Hydrocarbons, or Alkanes As melting point increases, boiling point increases (saturated hydrocarbons are hydrocarbons with as many Hs as possible)

Shape also matters Butane, bp -0.5 degrees C
Butane, bp degrees C 2-methylpropane degrees C Butane has a higher boiling point because the dispersion forces are greater. The molecules are longer (and so set up bigger temporary dipoles) and can lie closer together than the shorter, fatter 2-methylpropane molecules. Also, the molecules can stack with each other better H

Butane and 2-methylpropane
Butane and 2-methylpropane Compare the properties of these two compounds: n-butane methylpropane relative density (liquid) refractive index (liquid) boiling point (oC) melting point (oC) It is clear that the different carbon skeletons make a difference to the properties, especially the melting and boiling points.

Fats v Oils: Saturated v. Unsaturated
Molecular size, bond order, and bond orientation: How different IMFs result in differences in food molecules

A carbon exists where two lines intersect
A carbon exists where two lines intersect Atoms other than C and H are written in Hs are not usually written out- They fill in to complete octets on other atoms

Random cis trans fats (Omega 3 and Omega 6 fats have the double bonds on the 3 or 6th carbon)

fatty acids and triglycerides
3 Fatty acid chains (above) join with a glycerol molecule (top right) to form a triglyceride (right, saturated)

Triglyceride formation

Triglycerides Oils Fats More unsaturated FAs Liquid at RT
Oils More unsaturated FAs Liquid at RT Fats More saturated FAs Solid at RT

Fatty acid pamlmitic stearic oleic
16C sat MP= 62.9°C C sat MP= 69.6°C C unsat MP= 13°C

Oleic v linoleic acid Melting Points °C: Oleic acid: 13
Melting Points °C: Oleic acid: 13 Linoleic Acid: -5

Why? Why do the melting points differ between
Why do the melting points differ between Palmitic Acid (16 C, sat) Stearic Acid (18 C, sat) Oleic Acid (18C, mono unsat) Linoleic Acid (18C, polyunsat. 2=) (Explain the impact of number of carbons and the number of double bonds)

WHY DOES THIS HAPPEN? Proximity of atoms; regular shape allows the IMFs to hold everything in place (to “stack”) molecules rather than have the irregular shapes slide past each other

Triglycerides In unsaturated triglycerides, the molecules can not stack In the saturated molecules, the fatty acids are tightly packed and stacked

More carbons, higher MP The more double bonds, the lower the MP

Of the following, which would have the highest MP? The lowest?
Of the following, which would have the highest MP? The lowest? Lauric 12C, unsat, MP +44°C Stearic 18C, sat, MP 70°C Arachodonic 20, unsat, MP -50°C Elmhurst

Which fats are saturated? Unsaturated? What type of IMF would predominate? Rank the molecules in order from lowest to highest MP.

Percent Fatty Acids in Percent Fatty acids in selected triglycerides

Cis and trans fats

Cis- v Trans- fats Cis- fats are naturally occurring fats from animal products Trans- fats occur from modifying oils chemically Partially hydrogenating oils Adding H’s causes double bonds to convert to single bonds Unsaturated to saturated conversion Due to steric hindrance, when the H is added, they convert some cis bonds to trans bonds Why do manufacturers make trans fats for use in foods? Trans fats cost less (vegetable sources v. animal sources) Fats in foods are usually more desirable that oils- Less greasy Can control how solid the fats are by controlling the number of double bonds Better/ easier to cook with (especially in baked goods)

Saturated Fats: These are considered to be the “bad” fats. They are called “saturated” because their carbon structures are completely filled (saturated) with hydrogen atoms. Their chemical structure is very linear which allows for a “stacking” effect to occur. This is what promotes the solidifying effect of most saturated fats (butter, lard, most animal fats). This solidification may also occur in the body which partly explains the artery-clogging effects linked to saturated fats. Examples of saturated fats include myristic acid, palmitic acid, stearic acid, arachidic acid, and lignoceric acid. These fats may raise cholesterol levels in the body and should be used in moderation

Why are trans- fats bad? The trans- double bonds
The trans- double bonds Are more reactive in the body Promote free radical formation Leads to destruction of biomolecules Are more likely to clog arteries Due to shape, get caught in body Promote cholesterol levels to increase, since they can be used to make cholesterol in the body We don’t have the enzymes to process the trans- fats (we can process cis- fats)

Good fat/ Bad fat?

Spider silk monomer (amino acid) Amino acid R groups Kevlar monomer

Silk Silk and proteins

Viscosity Viscosity is the resistance to flow
Viscosity is the resistance to flow The greater the viscosity, the greater the resistance to flow Determined: How quickly a fluid flows through a tube under gravitational force (slower= more viscous) Or by Determining rate at which steel sphere fall through the liquid (more viscous= more slowly) Changes as temperature changes

What is surface tension?
Resistance of a liquid to an increase in it’s surface area (Zumdahl) Free energy per unit surface area (Tinoco, Sauer, Wang and Puglisi) Force per unit length (mNm-1, or dyne/cm) Layman’s terms: How much something spreads out on a surface Beading up= high surface tension Spreading out= low surface tension

Surface Tension (High surface tension) (Low surface tension) Cohesion: Molecules sticking (due to IMFs) to the same molecule in a pure compound Adhesion: Molecules sticking (due to IMFs) to other molecules adjacent to the pure compound (not a mixture- at a surface interface) The molecules of water have more adhesion to the (polar) glass than to each other (cohesion); The Hg has more cohesive forces than attraction to the glass

Wetting and Dewetting http://www.mpikg-golm.mpg.de/gf/1

Surface tension of water is 73 dyne/ cm;
Wetting is how water (in this case) adheres to a surface; when the surface tension is lowered, the material becomes wetter. Surface tension of water is 73 dyne/ cm; Water droplet on lotus leaf, with adhering particles

What causes surface tension?
Surface tension is a result of the imbalance of forces at the surface (or interface)

Surface tension of… mNm-1 Temperature (˚C) Platinum 1819 200 Mercury
mNm-1 Temperature (˚C) Platinum 1819 200 Mercury 487 15 Water 71.97 25 58.85 100 (liquid) Benzene 28.9 20 Acetone 23.7 n- Hexane 18.4 Molten Iron 17 1600 Silicon Oil 16.9 Neon 5.2 -247 (Tinoco, Sauer, Wang, and Puglisi);

Does temperature matter?
temperature ( C) surface tension (mNm-1) -8 77 -5 76.4 75.6 5 74.9 10 74.22 15 73.49 18 73.05 20 72.75 25 71.97 30 71.18 40 69.56 50 67.91 60 66.18 70 64.4 80 62.6 100 58.9 Yes- part of the reason that we wash in warm water (at times), not cold – the fabric gets “wetter” As temperature increases, surface tension decreases (Surface tension given for water against air)

Temperature and IMFs IMFs in a substance change in strength in a substance as temperature changes This influences certain properties of the substances Surface tension Viscosity Capillary action Vapor pressure (but not BP, MP)

Capillary Action Capillary action: a phenomenon associated with surface tension and resulting in the elevation or depression of liquids in capillaries (from The molecules of water have more adhesion to the (polar) glass than to each other (cohesion); The Hg has more cohesive forces than attraction to the glass (glass is polar)

Vaporization and Vapor Pressure
The molecules in a sample of a liquid move at various speeds (average speed is the temperature; some have more energy, some have less, but the overall KE is temperature) Sometimes molecules at the surface have sufficient speed to overcome the attractive forces and leave the liquid surface (thus vaporizing)

Figure 14.9: Microscopic view of a liquid near its surface.

Dynamic equilibrium Dynamic equilibrium is the state where there is simultaneous and equal vaporization and condensation of the substance In a closed container, at some pressure, the amount that vaporizes will equal the amount condensing on the surface of the liquid This is the equilibrium vapor pressure

VP and IMFs Stronger IMFs equal LOWER vapor pressures
Weaker IMFs equal high vapor pressures Substance with very low IMFs and therefore high vapor pressures evaporate very quickly and easily Called volatile substance Mass and shape important, just like with boiling point Heavier = lower VP ex: oil Lighter= higher VP ex: alcohol More volatile Think propane (C3H8) v. gasoline (C8H18)

VP and Boiling This vaporization occurs at any temperature, but occurs more rapidly as temperature increases Molecules at the surface would have to have more speed to overcome the IMFs Boiling is the point at which the vapor pressure equals the external pressure on the surface of the liquid

Boiling and VP, con’t Liquids have some air dissolved in them in tiny invisible bubbles As water vaporizes in the liquid, it is added to the bubbles Also, the gas bubbles are expanding because they are being heated; this causes an increase in volume, but not mass At this point, 2 things are going on: This decreases density, causing the bubbles to float to the surface Also, as gas expands, the pressure increases When the pressure of the bubble increases to greater than the vapor pressure at the surface, the liquid is boiling All molecules must be vaporized before a further increase in temperature can occur

Boiling Point and Elevation
As elevation on the Earth’s surface increases, the atmospheric pressure decreases (smaller column of air pushing down on the area; therefore less pressure) Boiling point changes as the atmospheric pressure changes If you could decrease the pressure without changing temperature, the substance would boil at a lower temperature A decrease in pressure results in a decrease in BP

Figure 14.14: The formation of the bubble is opposed by atmospheric pressure.

Energy Changes Accompanying Changes of State
Each change of state is accompanied by a change in the energy of the system Whenever the change involves the disruption of intermolecular forces, energy must be supplied The disruption of intermolecular forces accompanies the state going towards a less ordered state As the strengths of the intermolecular forces increase, greater amounts of energy are required to overcome them during a change in state Takes more energy to go from a liquid to a gas than from a solid to a liquid Removing energy allows the molecules to “self- organize”, and results in an more ordered state

Heat of Fusion The melting process for a solid is also referred to as fusion The enthalpy change associated with melting a solid is often called the heat of fusion (Δ Hfus) Ice ΔHfus = 6.01 kJ/mol Δ H is a change (Δ) in enthalpy (H), a measure of energy that is much like heat, but takes into account a few other factors

Heat of Vaporization The heat needed for the vaporization of a liquid is called the heat of vaporization (Δ Hvap) Water Δ Hvap = kJ/mol Vaporization requires the input of heat energy

Less energy is needed to allow molecules to move past each other than to separate them totally, so ΔHfus < Δ Hvap

The heating/cooling curve for water heated or cooled at a constant rate.

Energy/ disorder diagram

Think of IMFs like magnets: stronger magnets hold things more firmly together The more firm the connections, the less molecular motion can occur with the same amount of Energy added Adding (or removing) energy from the system can overcome (or increase) the IMFs, and cause a change in state Add Energy, move from S -> L -> G Remove Energy, move from G -> L -> S

Air conditioners take advantage of Energy changes to remove heat energy from a warm indoor environment by vaporizing condensed gas On the outdoor portion of the AC unit, the gas is condensed to a liquid, sending the heat energy to the environment

Vacuum chamber demo

Phase Diagram Due to changes in pressure and temperature, a substance can exist in all three states under specific conditions The Triple Point Think foggy icy days

Energy and IMFs Remember
Kinetic energy is the Energy associated with moving particles Heat is the RANDOM KE of an object (as opposed to directional motion) Temperature is the measure of the AVERAGE KE in a substance When IMFs are disturbed due to E changes, the properties of the substance change, even to the point of changing state

Explain how the lava lamp works (not “you plug it in”
Explain how the lava lamp works (not “you plug it in”! On a molecular level, explain what is happening to the materials and their IMFs)

Calculations (I know you are excited about this…)

Heat of fusion

calorimetry

VP and IMFs Stronger IMFs equal higher vapor pressures
Weaker IMFs equal high vapor pressures Substance with very low IMFs and therefore high vapor pressures evaporate very quickly and easily Called volatile substance Mass and shape important, just like with boiling point Heavier = lower VP ex: oil Lighter= higher VP ex: alcohol More volatile Think propane (C3H8) v. gasoline (C8H18)

VP and IMFs Stronger IMFs equal lower vapor pressures
Stronger IMFs equal lower vapor pressures Weaker IMFs equal high vapor pressures Substance with very low IMFs and therefore high vapor pressures evaporate very quickly and easily Called volatile substance Mass and shape important, just like with boiling point Heavier = lower VP ex: oil Lighter= higher VP ex: alcohol More volatile Think propane (C3H8) v. gasoline (C8H18)

IMFs in proteins

Sickle Cell Anemia Glu (glutamic acid) replaced by Val (valine)

Intermolecular Forces:

Similar presentations

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

Similar presentations

About project

Feedback

Log in

Auth with social network:

Intermolecular Forces:

Similar presentations

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

Similar presentations

About project

Feedback