Introductory Chemistry 6 – Intermolecular Forces, Bonding, & Properties Callum MacKellar – [email protected] – Office: Ha215a Office hours: 15.00-16.00 on Tuesdays and 10.00-11.00 on Fridays

Intermolecular Forces Intermolecular forces are the attractive forces that exist between molecules. While they are generally much weaker than ionic, covalent and metallic bonds, the interactions between molecules are critical in determining the properties and behaviours of compounds. Molecules can interact through several different types of intermolecular forces, depending on their structures. By understanding the relative strengths of the forces, the properties of a substance can be predicted based on a molecule’s structure. Dipole-dipole forces, dipole-induced dipole forces and dispersion forces are known as Van der Waals forces. These make up the majority of electrostatic attraction forces that molecules experience, and they are critical to understanding the properties of condensed matter. Dipole-Dipole Forces In molecules that have atoms with different electronegativites, bonds are polarised due to the uneven distribution of electrons. In highly symmetrical molecules, these bond polarities cancel out and the molecule is apolar. However, in cases where the polarised bonds do not cancel out, they can give rise to a permanent dipole moment, or an overall polarisation.

CH4 has no dipole moment, it is apolar. NH3 does have a dipole moment, it is polar.

Much like opposite poles of a magnet attracting, the opposite electrostatic charges in molecules also attract one another. Dipole-dipole forces are stronger for molecules with larger dipole moments. They are also stronger in crystalline solids than in liquids due to the alignment of dipoles.

Liquid less organised/aligned and slightly less dense.

Na is less charge dense, so weaker. Ion-dipole is not VdW. Same as larger ΔΧ leads to stronger. Dipole-induced dipole forces come about when a polar molecule and an apolar molecules are adjacent. The dipole moment of the polar molecule can distort the electron cloud of an apolar molecule, creating a temporary dipole that attracts the polar molecule. This interaction is weaker than two permanent dipoles attracting one another, but can still be very important.

Larger, more diffuse, more electrons = more polarisable = stronger interactions.

Weaker permanent dipole = Weaker temporary dipole.

Ions also induce dipoles, not VdW, can be stronger than polar molecules

Dispersion Forces (London Forces) Dispersion forces, also known as London forces, are the attractive forces between all molecules due to the random motion of electrons. At any specific time the electrons in a molecule/atom may not be evenly distributed leading to an instantaneous or temporary dipole moment, which only lasts for a fraction of a second. A temporary dipole moment can induce a dipole on a neighbouring molecule, which leads to an attractive force. As temporary dipoles are constantly rearranging and reorienting, they cancel out to give no net dipole moment. All molecules and atoms have dispersion forces between them.

Dipole changes as time passes

Larger and more diffuse atoms/molecules lead to stronger dispersion forces.

TAR is just CnH2n+2 Molecular/Atomic weight OR electron count gives good sense. Hydrogen bonding Hydrogen bonding (H-bonding) is a specific kind of dipole-dipole interaction that is particularly strong, and as such it gives some compounds and materials anomalous properties. It occurs when a hydrogen atom is bonded to an electronegative atom (N, O, F), and the partial positive charge on the hydrogen atom is attracted to a lone pair on a neighbouring electronegative atom (N, O, F).

Highlight donors (A-H) and Acceptors (B), lone pairs on acceptors, dipoles

Linear alignment is necessary

Water can form 4 H-bonds per molecule, so ice forms an extended 3D network (HF + NH3 form chains or rings).

Structure holds molecules set distance apart. In liquid, less structured, so molecules closer together. Thus water is less dense when it in its solid phase vs liquid.

Intermolecular Force Comparison Force Species Involved Typical Strength Covalent Bond - 200-500 kJ mol-1 Dipole-Dipole Both with permanent dipole 2 kJ mol-1 Ion-Dipole One ion, one permanent dipole 15 kJ mol-1 Dipole-Induced Dipole One with permanent dipole <1 kJ mol-1 Ion-Induced Dipole One ion 10 kJ mol-1 Dispersion All Species dependant Hydrogen Bonding H-bond donor and acceptor 20 kJ mol-1 Bonding Structures With an understanding of chemical bonding and intermolecular forces, the physical properties of materials can be predicted and understood by examining their structure. Simple Molecular Substances that exist as single atoms and small molecules containing covalent bonds only interact through intermolecular forces. Intermolecular forces are relatively weak, and this is reflected in their properties. Mechanical properties: Soft, easily deformed.

Melting/Boiling point: Generally low as IM forces are easily overcome.

Boiling points of Group 14-17 Hydrides 120 NH3, H2O and HF are higher Group 14 Group 15 than others due to H- bonding. H2O is highest 60 Group 16 Group 17 because of 2 donor 2

acceptor. C

° 0

-60 Increase down periods due to increasing dispersion

Boiling Point / Point Boiling forces some increase in D-D -120 interactions.

-180 2 3 4 5 Period Solubility: Like dissolves like as strong interactions encourage mixing

Polar Solvent Apolar Solute → Mixture D-D + Disp Disp D-ID + Disp

Polar Solvent Polar Solute → Mixture D-D + Disp D-D + Disp D-D + Disp

Electrical Conductivity: Insulator. Conduction requires mobile charge carriers (ions or electrons). In molecules, no ions, and electrons are localised in bonds/lone pairs.

Thermal Conductivity: Crystals and gases are poorly conductive. Liquids can conduct. Giant Covalent Giant covalent structures are held together with covalent bonds as molecules are, but are many orders of magnitude larger. They are extended 3-dimensional networks sometimes known as Covalent Crystals, or Giant Atomic Lattices. As all of their bonding is covalent, the properties of giant covalent structures vary significantly from those of simple molecules. Diamond is a molecular crystal of carbon, and exhibits a lot of the properties that are common in covalent crystals.

Each atom has 4 strong bonds to neighbouring carbon atoms.

Mechanical Properties: Hard, as other covalent crystals tend to be.

Melting/Boiling Point: Very high as covalent bonds are hard to break (sometimes no melting)

Solubility: Insoluble as VdW/H-bonds are too weak to break covalent bonds and separate atoms.

Electrical Conductivity: Insulator. Electrons are localised in bonds.

Thermal Conductivity: Conductor. Bonds transfer vibrations through the material readily. This contrasts with other covalent crystals.

Graphite is a layered from of carbon. It and diamond

are allotropes.

Each atom has 3 bonds to neighbouring carbon atoms in the layer forming edge-sharing hexagons.

Bonds between layers are week.

Mechanical Properties: Soft, as layers can more past each other. Used in pencil lead as layers deposit on page, and as lubricant as layers move past each other.

Electrical Conductivity: Conductor. 4th valence e- on every atom is delocalised within the layer and can move.

Thermal Conductivity: Conductor.

Ionic In ionic materials, electrons have been fully transferred from atoms of one element to another. The electrostatic attraction between the ions holds them together in extended lattices, with each ion surrounded by ions of the opposite charge. Mechanical Properties: Hard and brittle, as disrupting the structure moves ions out of alignment.

Melting/Boiling Point: High as electrostatic attraction between ions is hard to overcome.

Electrical Conductivity: Non-conductive in crystals as ions are immobilised. Can conduct in solution or in melt

Solubility: Soluble in polar solvents (e.g H2O) which can form strong ion-dipole interactions.

Spheres of hydration form around ions.

More charge dense ions attract more water Larger spheres Bigger radius of hydration (Stokes Radius)

Thermal Conductivity: Poor.

Metallic As metals make up the majority of the periodic table, and given the variety in ions that can form and be surrounded by a sea of electrons, metallic materials show a wide variation in properties, with plenty of exceptions. Mechanical Properties: Various, but can be very hard due to strong electrostatic attractions. Ductile (can be drawn in wires) and malleable (deformable) as ions can rearrange within electrons without breaking.

Melting/Boiling Point: Various, from high due to strong attraction to low (gallium melts 30 °C) Dimeric, pseudo molecular structure

Electrical Conductivity: Conductive, as sea of delocalised electrons can carry charge.

Thermal Conductivity: Good conductor