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