Intermolecular Forces

Section 3.4 Pg. 105-117

1) Explain intermolecular forces, London (dispersion) forces, dipole-dipole attractions and hydrogen bonding

2) Relate properties of substances to the predicted intermolecular bonding in the substance. BACKGROUND • All chemical changes (reactions) are accompanied by energy changes ▫ Energy is mostly heat, light, or electrical energy

▫ Energy can be released slowly (battery) or quickly (fireworks) ▫ Two types of energy changes are possible:

 EXOTHERMIC – energy is released into the surroundings - the product’s bonds have less energy than the reactant’s bonds

 ENDOTHERMIC – energy is absorbed from the surroundings - the product’s bonds have more energy than the reactant’s bonds

▫ Bond Energy – the energy required to break a or the energy released when a bond is formed BACKGROUND • There are three types of forces in matter:

1) Intranuclear force (bond) – bonds within the nucleus between protons and neutrons (very strong)

2) (bond) – bonds between within the or between ions within the crystal lattice (quite strong)

3) (bond) – bonds between (quite weak); are electrostatic (involve positive and negative charges)

There are 3 types of intermolecular bonds:

Weakest a) Dipole-Dipole Forces (a.k.a. Polar Forces)

Medium b) London Force (a.k.a. , Dispersion Force)

Strongest c) Hydrogen Bonding

Note: “Van der Walls force” – includes London and dipole-dipole forces 1) Dipole-Dipole Force • The simultaneous attraction between oppositely charged ends of polar molecules. ▫ Simply put, the attraction between diploes

Dipole: a partial separation of positive and negative charges within a molecule, due to differences

▫ Dipole-dipole forces are among the weakest intermolecular forces, but still control important properties (i.e. Solubility because water is polar)) Note: If a molecule is polar it will be soluble 1) Dipole-Dipole Force in water? Why?

In a liquid, polar molecules can move and rotate to maximize attractions and minimize repulsions. The net effect is greater overall attraction.

The strength of the dipole-dipole force is dependent on the overall polarity of the molecule 1) Dipole-Dipole Forces

In a liquid: In a solid: 2) London Force

• Simultaneous attraction between a momentary dipole in a molecule and the momentary dipoles in surrounding molecules

momentary dipole: an uneven distribution of electrons around a molecule, resulting in a temporary charge difference between its ends

They last for just the instant that the electrons are not distributed perfectly even. 2) London Force

• also showed that momentary dipoles occurring in adjacent molecules would result in an overall attraction

• The strength of the London force is directly related to the number of electrons in the molecule, and inversely related to the distance between the molecules. ▫ Increase electrons = Increase force (directly related).. ▫ Increase distance = Decrease force (inversely related) 2) London Force

• The key point is that: ▫ the more electrons a molecule has, the more easily momentary dipoles will form, and the greater the effect of the London force will be.

• London forces are present between all molecules, whether any other type of attraction is present. Why do we care about intermolecular forces?

• We can use Dipole-Dipole and London Forces to predict Boiling Points

Compound (at SATP) Electrons Boiling Point (°C)

CH4(g) 10 -164

SiH4(g) 18 -112

GeH4(g) 36 -89

SnH4(g) 54 -52

A higher boiling point temperature means more energy has to be added, thus we assume the intermolecular forces are stronger. (see Learning Tip pg. 109) Remember (if all other factors are equal): 1) The more polar the molecule = The stronger the dipole-dipole force 2) Increase the number of electrons = Increase the strength of London Force Example #1 • Use Intermolecular force theory to predict which of the following has the highest boiling point:

▫ methane (CH4), ethane (C2H6), propane (C3H8), butane (C4H10)

1) Are the molecules polar or non-polar? non-polar (no dipole-dipole force)

3) Which has more electrons? butane: greatest # of e-’s = greatest London force

Alkane Boiling Point (°C) Check: methane -162 ethane -89 propane -42 butane -0.5 Example #2 • Use Intermolecular force theory to predict which of the following has the highest boiling point:

▫ (Br2 ) or monochloride (ICl)

1) Which has more electrons? They are isoelectronic: have the same number of electrons (70 e-’s) -Therefore the London force is the same (or nearly the same)

1) Are the molecules polar or non-polar? -Bromine is non-polar (has no dipole-dipole force; only London forces) - Iodine monochloride is polar (has dipole-dipole forces and London forces) - This extra attraction between ICl molecules produces a higher boiling point

Check: Substance Electrons Boiling Point (°C) bromine 70 59 iodine monochloride 70 97 You cannot predict boiling points when:

• One molecule has a stronger dipole-dipole force and the other has a stronger London force

• The two molecules differ significantly in shape

• The central of either molecule has an incomplete octet Practice

• Pg. 109 #1-4

• 4) a) Boron is stable with 6 valence e-

b) chloromethane (CH3Cl) 3) Hydrogen Bonding

• Occurs when a hydrogen atom bonded to a strongly electronegative atom, (N, O and F) is attracted to a lone pair of electrons in an adjacent molecule. ▫ Hydrogen nucleus (proton) is simultaneously attracted to two pairs of electrons; one closer (in the same molecule) and one further away (on the next molecule)

Why do you need a strongly electronegative atom?

It pulls the hydrogen’s electron away making it “unshielded”, so the lone pair on the other side can come much closer 3) Hydrogen Bonding • Hydrogen bonds are momentary attractive forces between passing mobile molecules but are the strongest of the intermolecular forces.

• Hydrogen bonds only act as continuous bonds between molecules in solids, where the molecules are moving slowly enough to be locked into position.

• Hydrogen force would have been a better name. 3) Hydrogen Bonding

• In ice, hydrogen bonds between the molecules result in a regular hexagonal crystal structure.

• The ···H– represents a hydrogen nucleus (proton) being shared unequally between two pairs of electrons 3) Hydrogen Bonding

• Do lakes freeze from the bottom-up or the top-down?

• Top–down, because water is unique in that its solid form (ice) is less dense than its liquid form. Why??

• The hydrogen bonds hold water molecules in a hexagonal lattice with open space in the center, which explains the low density (mass/volume) of ice. Hydrogen Bonding in DNA

• FYI: The double helix of the DNA molecule owes its unique structure largely to hydrogen bonding.

• The red bonds are hydrogen bonds.

• If the helix were held together by covalent bonds, the DNA molecule would not be able to unravel and replicate and life could not continue!! Why do we care about intermolecular forces?

• Explains surface tension, shape of a This shows water adhering to the faucet meniscus, volatility and capillary action gaining mass until it is stretched to a point 1) Surface Tension where the surface tension can no longer ▫ Molecules within a liquid are attracted bind it to the faucet. by other molecules in all directions It then separates and equally, but right at the surface, surface tension forms molecules are only attracted the drop into a sphere. downwards and sideways. This means the net pull is downward so the surface tends to stay intact

▫ The stronger the intermolecular force the stronger the surface tension. Why do we care about intermolecular forces?

2) Capillary Action – due to (attraction between unlike molecules) and cohesion (attraction of like molecules)

▫ The adhesion between water and glass is greater than the cohesion between water molecules.

▫ The cohesion between mercury molecules is greater than the adhesion between mercury and glass

Hg clip In a sense, water is pulled up the tube by the intermolecular forces between water and glass Meniscus Practice

• Pg. 117 # 1, 4, 5 ▫ #1 – use pg. 99 table to determine polarity

▫ #1 – look for NH2, NH, OH2, OH, to determine if hydrogen bonding is possible

• Ex. CH3CHOHCH3 will it have hydrogen bonding?