Chapter 6- Intermolecular forces

explain the electron-pair repulsion theory?

• electron pairs surrounding a central atom determine the shape of the molecule or ion

• the electron pairs repel one another so that they are arranged as far apart as possible

• the arrangement of electron pairs minimises repulsion and thus holds the bonded atoms in a definite shape

• different numbers of electron pairs result in different shapes

represent the molecule methane in 3 dimensions

• four bonded pairs of electrons around the central carbon atom

• the 4 electron pairs repel each other as far apart as possible in three-dimensional space

• this results in a tetrahedral shape with 4 equal H-C-H bond angles of 109.5

what are intermolecular forces?

weak interactions between dipoles of different molecules

what are the 3 categories of intermolecular forces?

• induced dipole-dipole interactions (london forces)

• permanent dipole-dipole interactions

• hydrogen bonding

induced dipole-dipole?

weakest type of intermolecular interactions

permanent dipole-dipole interactions?

• stronger than induced dipoles-dipole but weaker than hydrogen bonds

hydrogen bonds?

strongest form of intermolecular forces

• they are still much weaker than INTRAmolecular forces (e.g- covalent bonds)

difference between intermolecular forces and covalent bonds?

intermolecular forces: largely responsible for physical properties such as melting and boiling points

covalent bonds: determine the identity and chemical reactions of molecules

what are the london dispersion forces (induced dipole-dipole interactions)?

• weak intermolecular forces that exist between all molecular substances and noble gases

• they do not occur in ionic substances

• they act between induced dipoles in different molecules

*describe how van der waals forces arise!?

• uneven distribution of electrons/ movement of electrons

• causes an instantaneous dipole/temporary dipole in a molecule

• this then causes induced dipoles in neighbouring molecules

temporary dipoles?

• the electron clouds around molecules are constantly in motion

• one moment, the electron density can be on one side of the molecule

• the next, it can be somewhere totally different

• this is called a temporary dipole, there are partial charges, but they change very rapidly

image of london forces?

induced dipoles?

• if one molecule has a temporary dipole, its partial charges will exert a force on nearby molecules

• the partial charge of one molecule can push away the electrons in another, or attract them towards it

• this means that temporary dipoles will induce dipoles in nearby molecules

what happens once a dipole has been induced?

it will be attracted to the initial dipole

• this is called an induced dipole interaction

what does the strength of london dispersion forces depend on?

• the number of electrons in a molecule

  • this is why london dispersion forces are not all the same strength

question: why does the boiling points of the halogens down group 7 increase?

• number of electrons increases in the bigger molecules

• this causes an increase in the size of induced dipole-dipole interactions between molecules

• this is why iodine is a solid, whereas chlorine is a gas

question: why does the boiling points of the alkane homologous series increase?

• the number of electrons increases in the bigger molecules

• this causes an increase in the size of the induced dipole-dipole interactions between molecules

what else has an effect on the size of the induced dipole-dipole interactions?

the shape of the molecule

example of shape of molecule having an effect on the dipole interactions?

• long chain alkanes have a larger surface area of contact between molecules for induced dipole-dipole interactions to form

• compared to spherical branched alkanes

• therefore, they have induced dipole-dipole interactions

why will molecules with more electrons have stronger london dispersion forces?

• because they have lots of electrons

• this is because they will have larger fluctuations in electron density

• this then leads to larger temporary dipoles and stronger dipole-dipole interactions

where do permanent dipole-dipole interactions exist between?

• they exist between 2 permanently polar molecules

permanent dipole image?

what symbols do permanent dipole-dipole interactions have?

• they have a delta- on one side and a delta + on the other side

• the delta + of one molecule will attract the delta - of another molecule

where are hydrogen bonds found?

• when you have a hydrogen atom bonded to either oxygen, nitrogen or fluorine

• that hydrogen atom will form a strong permanent dipole-dipole interaction with another oxygen, nitrogen or fluorine atom

examples of liquids with hydrogen bonds?

• water

• hydrogen fluoride

• ammonia

why do hydrogen bonds form?

• when hydrogen is bonded to an extremely electronegative element, it develops a strong delta + charge

• hydrogen is a very small atom, so it has a high charge density in this situation

  • this allows it to form a strong bond with any highly delta - charged atom

why do H2O, NH3 and HF have anomalously high boiling points?

because of the hydrogen bonding between the molecules in addition to their london forces

• the additional forces require more energy to break and therefore, they have higher boiling points

why is there a general increase between the molecules?

because there are more electrons in the bigger molecules and therefore, they have stronger london forces

properties of compounds with hydrogen bonding?

• have higher boiling points compared to compounds with other types of intermolecular forces

• tend to be soluble in other compounds with hydrogen bonds

• they have higher viscosity

  • the stronger the hydrogen bonding, the more viscous the liquid

• they have higher surface tension

what acronym is used to remember how to draw hydrogen bonding?

• P- partial charges

• H- hydrogen bond

• I- in line

• L- lone pair

draw the hydrogen bonding between methanol molecules

why is ice less dense than water?

• hydrogen bonds hold water molecules apart in an open lattice structure

• the water molecules in ice are further apart than in water

• solid ice is less dense than liquid water, and floats