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10.1 – 10.210.8 – 10.9
Intermolecular Forces
Go over Tests and Turkey Questions andRead P. 442, 444-456: Monday 12/1
PPT: Tuesday 12/2 – Wednesday 12/3
In Class Discussion:
• Ch. 10 # 3 – 10, 33, 104
• Homework after 10.1-10.2, 10.8-10.9• Ch. 10 #35-43 odd, 87, 91
• Intermolecular Forces (IM Forces)– The forces that hold one molecule next to another.– Between molecules– Related to heat of fusion and vaporization (phase changes)– Dispersion, Dipole-Dipole (H-bonding), ion-dipole, ion-
induced dipole, dipole-induced dipole,• Ion-ion forces could be called inter and/or intra … arguable
• Intramolecular Forces (Chemical Bonds)– Forces that hold the atoms together inside of a molecule.– Inside a molecule– Related to heat of reaction (bonds breaking and forming)– Ionic, Covalent, Metallic
• Intermolecular Forces are generally weaker than Intramolecular Forces.
– The attraction between an H and an O in one water molecule is greater than the attraction between the H in one water molecule and the O in another.
– The energy needed to boil water is less than the energy needed to break apart a water molecule for a chemical reaction.
• Kinetic Molecular Theory of Gases (Ch. 5) says that we can neglect the interactions between molecules.– Gases experience negligible IM forces.– Remember Ideal Gas vs. Real Gas
• Liquids and Solids properties differ from those of gases due to large intermolecular forces.
• Gas:– KE >> IM Forces– Compressible; expands; volume of container– Fluid (flows quickly, shape of container)
• Liquid:– KE similar to IM Forces– Condensed phase (incompressible; retains volume)– Fluid (flows, although slowly; shape of container)
• Solid:– KE << IM Forces– Condensed phase (incompressible; retains volume)– Not a fluid (doesn’t flow; retails shape)
Phase changes occur due to changes in Temperature (KE), or Pressure
• Increasing KE can cause a change from solid to liquid or liquid to gas, based on relationship between KE and IM Forces.– Example: H2O is solid ice at -10°C and liquid at 20°C.
• Increasing Pressure can cause a change from gas to liquid or liquid to solid, acting with IM Forces against particle’s KE– Example: H2O boils at 1000C at sea level atmospheric
pressure. If you decrease the pressure, it will boil at a lower temperature.
– Less pressure pushing molecules together, less KE needed to overcome the IM Forces.
The Temp and/or Pressure at which phase changes occur depend on the particle’s IM Forces
• The phase a substances is at room temperature (a given KE) depends on the substance’s IM Forces– Greatest IM Forces = solid– Least IM Forces = gas
– Example: I2 has stronger IM forces than Br2,
which has stronger IM forces than Cl2.
At room temp, I2 is solid, Br2 is liquid, and Cl2 is gas.
Boiling Points: I2 (365°F), Br2 (138°F), and Cl2 (-29°F).
• All Intermolecular Forces are electrostatic, involving attractions between positive and negative species or areas.
– Remember: • Electrostatic Forces increase with increased charge.• Electrostatic Forces decrease with increased distance
between charges.
Types of Intermolecular Forces
• (1) Dispersion Forces
• (2) Dipole-Dipole Forces– Hydrogen Bonding being the strongest
• (3) Ion – Dipole Forces
• (4) Induced Forces– Ion induced dipole– Dipole induced dipole
Dispersion Forces (P. 447 Brown)
• Electrostatic attraction between all molecules.
• Electrons move randomly.• Instantaneous dipole moments form during
any given second– More e- on one end than the other, causing partial
positive and partial negative ends– See Fig 11.4 in Brown and 10.5 in Zumdall
Dispersion Forces (P. 447 Brown)• Polarizability: the ease with which the charge distribution is
distorted; how easy it is for the particle to develop a temporary dipole.
• Greater Polarizability = stronger dispersion forces = higher boiling point– See Figure 11.5 in Brown
• more electrons (usually determined by Molar Mass) = greater polarizability– Ex: Br2 has a higher boiling point than Cl2
• More surface area of electrons = greater polarizability– Ex: isomers of C5H12 have different strengths of dispersion forces; See
Figure 11.6 in Brown
Remember, ALL particles experience dispersion forces.
• Larger polarizability (larger particle w/greater surface area) = greater dispersion forces
Dipole-Dipole Forces (P. 448-452 Brown)
• Some particles have a permanent dipole.– These particles are called “polar”.
• The attraction between the partial negative end of one polar molecule and the partial positive end of a second polar molecule.
• Greater polarity = stronger dipole-dipole forces = higher boiling point
See Figure 11.8 in Brown
• Lets draw the Lewis Structure of each of these four molecules to analyze their polarity and evaluate what makes one polar than another.
• Remember: more polar = stronger dipole-dipole forces = higher boiling point – Assuming all have similar dispersion forces.
– “For molecules of approximately equal mass and size, the strength of intermolecular attractions increases with increasing polarity … boiling point increases as the dipole moment increases.” Brown P. 449
Greater differences in Electronegativities = greater polarity = stronger dipole-dipole forces
Dipole-dipole Force strength: BrF > ClF > F2
See Figure 11.9 Brown or 10.4 Zumdall
• Each color is its own group on the periodic table.
See Figure 11.9 Brown or 10.4 Zumdall
• Each color is its own group on the periodic table.
• Why do molecules with elements in group 6A have greater boiling points than molecules with elements of similar molar mass in Group 5A?
• Why do all these molecules (containing elements in group 6A and group 5A bonded to H) have higher boiling points than molecules of H bonded to Group 4A elements?
See Figure 11.9 Brown or 10.4 Zumdall
• Each color is its own group on the periodic table.
• Why is the boiling point of H2Te greater than that of H2Se greater, and H2Se’s boiling point greater than that of H2S?
• Why is the boiling point of HI greater than that of HBr, and HBr’s boiling point greater than that of HCl?
See Figure 11.9 Brown or 10.4 Zumdall
• Each color is its own group on the periodic table.
• Why do H2O, NH3, and HF all have higher boiling points than any of the other molecules on this graph?– NOTICE THE DIFFERENCE IN PATTERN WITH THESE 3
Hydrogen Bonding (type of dipole-dipole)
• Attraction between a hydrogen atom that is bonded to a highly electronegative atom (F, O, or N) of one molecule and the lone pair(s) of a highly electronegative atom (F, O, or N) in another molecule.
• See Figure 11.10 of Brown or Figure 10.3 of Zumdall
Hydrogen Bonds
• Causes 3 small molecules (H2O, NH3, and HF) to be liquid when other molecules with similar molar masses are gas at room temperature.
• Gives water its extra high specific heat (changes temperature slower than other molecules)
Hydrogen Bonding plays a major role in biochemistry.
• Stablizing structures of protiens.
• Causes DNA to be double helixed and fold over itself.
Hydrogen Bonding is what causes solid H2O to be less dense than liquid H2O.
• See Figure 11.11 in Brown or Figure 10.12 in Zumdall
• Ice floats on top of liquid water.
• Most solids sink in their own liquid.
Ion-Dipole Forces
• Typically involved in aqueous (water) solutions of ionic solutes.
Ion-induced dipole and
dipole-induced dipole
• A permanent dipole (polarity) can be induced in an otherwise non-polar molecule if it is placed next to a polar molecule or an ion.
• Stronger than dispersion forces of similar sized molecules.
• Weaker than actual dipole-dipole or ion-dipole forces of similar sized molecules.
Properties of Liquids Due to Intermolecular Forces.Greater IM Forces = increase in each
• Surface Tension– A liquid’s resistance to increasing its surface tension.– A liquid’s desire to keep its surface area to a minimum.
• Capillary Action– Spontaneous rising of a liquid in a narrow tube– Causes a concave mensicus (See Figure 10.7)
• Viscosity– A liquid’s resistance to flow– Syrup has greater viscosity than water, causing it to pour
slower than water.
Stronger IM Forces also
• Increases boiling point (bp)• Decreases vapor pressure (opposite of bp)• Increases melting point (mp)• Increases specific heat (c)
Compare IM Forces
• Large molecules have greater dispersion forces than small molecules.
• Polar molecules have dipole-dipole forces when nonpolar molecules do not.
Compare IM Forces
• A nonpolar molecule can have stronger IM forces than a polar molecule if the nonpolar molecule is much larger.
– Ex: C16H34 has stronger dispersion forces than H2O has both dispersion and dipole-dipole forces.
Compare IM Forces
• A smaller molecule can have stronger IM forces than a larger molecule if the smaller molecule is much more polar.
– Ex: 1-propanol (CH3CH2CH2OH) has a boiling point of 97°C when water (H2O) has a boiling point of 100°C. They are both polar, but the relative polarities are very different.
Figure 10.40
Heat Curve
Phase Diagram
• If solid is more dense than liquid, solid-liquid slope is positive.
• If solid is less dense than liquid (water), solid-liquid slope is negative.
In Class Discussion:
• Ch. 10 # 3 – 10, 33, 104
• Homework after 10.1-10.2, 10.8-10.9• Ch. 10 #35-43 odd, 87, 91