2.2.8 Intermolecular Forces

London ( dispersion ) forces, dipole-induced forces, dipole-dipole forces, Hydrogen Bonding, first 3 called van Waals forces

Bonding --> refers to attraction that holds atoms ahd ions together

Intermolecular forces --> are forces that hold molecules together

• Molecular forces apply to molecular structure substances

• What forces keep water close together? Whats the fore that is overcome when the water molecules in ice break away and join the surrounding liquid water.

• Despite inter-molecular forces being weaker than chemical bonds --> affect physical properties such as Volatility, Solubility, and boiling point of molecular substances

• When Molecular Substances boil, melt, or sublime --> intermolecular forces are overcome

  • Therefore melting and boiling points are indicators of intermolecular strength

• Note that covalent bonds do not break during phase changes

• Note that the ideal Gas law states assumes that there are no intermolecular forces in gases. This is usually true for low pressure and high temperature gases, ( Ideal gas equation can be used )

• But at lower temperatures it tends to deviate from it, and thus should be adjusted using a real gas model.

  • Take into account of actual volume of gas particles

    • Considered negligible in ideal gas model

• 

London ( dispersion ) Forces ( LDFs )

• result from Instantaneous dipoles

  • Dipole when one side of molecule has partial charge and other the opposite charge

• Involves induced dipoles

  • One molecule causes / induces a temporary dipole moment in another moelcule

    • Due to random movement of electrons around the molecule

Take a simple non-polar example $H_2$. Note that the electron distribution is on average equal. Due to the random movements of electrons, taking a picture of electrons in an atom would likely result in an unequal distribution of the electrons.

• This results in a slight negative charge ($-\delta$) on one side and a slight positive charge ( $+\delta$) in the other

  • Which results in a dipole moment in surrounding molecules

    • $\because$ the region of negative charge will repel other electrons, which creates a region of positive charge in the surrounding molecules, electrostatically attracting each other

      • These surrounding molecules then will have a partial negative charge on the other side and thus attract more surrounding molecules

      • Note this is temporary, and the next moment will have a different pattern of induced dipoles

      • This is the London ( Dispersion ) force

    

    

• Two factors affected LDFs

1. Number of Electrons

2. Molecular Shape

$\because$they affect the polarizability ( how easily the electron distribution is distorted by an electric field / Tendency to generate induced electric dipole moments ) of the molecule

• Greater the polariability--> stronger the LDF

  • $P \propto LDF$
• From the table, we see that when going down the group, more electrons --> Electron clouds become larger and less attracted to necleus --> More chance for electrons to get polarized and thus strength of LDF increases.

• For the Homologous series of Organic Compounds ( 3.2 ), the boiling point increases for successive members

  • Each successive compound has one more $CH_2$molecule than the one before. Hence greater molecular size and larger number of electrons.
• Note that the molecular size can quantified in terms of mass, so its often refer to molecular mass when comparing the LDFs of molecules.

  • Note that electrons have negligible mass so how does this work?

    • Well its not causal ; instead a larger amount of electrons are often accompanied by a greater atomic mass --> each proton an electron to stabilise.

• Compare two isomers ( each of two or more compounds with the same formula but a different arrangement of atoms in the molecule and different properties ): pentate, $CH_3CH_2CH_2CH_2CH_3$ and 2,2-dimethylpropane $(CH_3)_4C$

  • $\because$same chemical formula, and are non-polar, only differing factor is shape
• Pentane is long --> better SA to interact with others--> Higher boiling point --> Liquid at room temperature

• 2,2 dimethylpropane is gas room temperature

Dipole - induced dipole forces

• LDFs are forces of attraction between temporary, or instanteous dipole moments in the molecule

• Dipole-induced dipole forces occur between a Polar molecule and surrounding non-polar molecules

  • A permanent dipole in a molecule makes temporary charges on surrounding non-polar molecules

    • $O_2$, a non-polar, attract to water molecules which are polar

      • This is weak as seen from the low aqueous solubility

Dipole-dipole forces

• where as LDFs involved temporary dipoles, dipole-dipole forces involve permanent ones ( polar molecules )

  • $HCl$, hydrogen Chloride, and Flourine, $F_2$have similar sizes and comparable molecular masses --> Experience similar LDFs force

    • However $\because$$HCl$molecules are polar, it experiences dipole-dipole forces in addition to LDFs, which makes the boiling forces higher than Flou]rine as a result of the stronger intermolecular force

Hydrogen Bonding

• Strong inter-molecular forces with molecules that forms when a molecule containing a strong dipole involving hydrogen

  • When Hydrogen covalently with an atom with high EN ( I.e. Oxygen, Nitrogen, or Flourine), the electron distribution will lean towards the more electronegative atom, leaving a big $\delta+$ partial charge on the hydrogen

    • This hydrogen atom forms a strong attraction ( Hydrogen bonds ), with electrons of another electronegative atom

• Usually found in different molecules but intramolecular hydrogen bonds exist as well

• Hydrogen bonds occur in :

1. water Molecules

2. ammonia ( $NH_3$) molecules
3. hydrogen fluoride ( $HF$) molecules
4. water molecules and dimethy either molecules $(CH_3)_2O$

• Stronger than other types of intermolecular forces, but much weaker than covalent

  • Not actually a chemical bond

$LDFs < DiD < DD < HB$

• A way to remember this is to note the boiling of water

  • At $100\degree$C it is sufficient to remove the hydrogen bonds but not the covalent bonds between oxygen and hydrogen atoms

• Looking at the group 14 Hydrides ( $CH_4, SiH_4, GeH_4, SnH_4$)

  • The molecular mass increases $\therefore$ more electrons and a strong polarization and $\therefore$ a higher boiling point

    • Seen in curve

• Looking at group 15 hydrides ( $NH_3, PH_3, AsH_3, SbH_3$

  • For the most part the latter pattern works except for Ammonia, $NH_3$

    • This is due to hydrogen bonds

  • Similar trends with group 16 and group 17 hydrides with the outliers of high boiling points: $H_2O$and $HF$respectively

• A single water molecules can form up to 4 hydrogen bonds and therefore its unexpectedly higher than the Flourine hydride!

• Hydrogen bonds give water its notable properties

  • Normally the solid form should be more dense than the liquid form as a liquid is more free-moving

• But Ice can float on water

  • Due to the excessive hydrogen bonding the ice contains well-ordered spaced out open-cavity network of molecules

    • This is less dense than the liquid water, as liquid water's free-moving property makes the molecule locations much more random and $\therefore$ more dense

• Note that there are many important applications of this unusual property --> for the formation of life

  • Due to ice being less dense than water, it floats up when water freezes

    • This ice acts like an insulator which keeps the liquid water underneath liquid..

      • allowing aquatic ecosystems to survive

Summary

TOK

• DNA ( deoxyribonucleic acid ) molecules store genetic information

  • A key feature for DNA has to be an accurate genetic copying mechanism to pass on to future generations

• DNA has a double-helix structure that contains two strands of organic molecules held together by hydrogen bonds.

  • When DNA needs copying, the hydrogen bonds are broken or "unzipped" to allow copies to be made

What other information storage system exists in the natural world? In Science? In other areas of knowledge??