Chapter 4 4 Intermolecular forces The physical properties

Chapter 4. 4: Intermolecular forces The physical properties of molecular substances result from different types of forces between their molecules.

Important terms in this section �Intermolecular forces �Temporary/instantaneous dipole �Induced dipole �London (dispersion) forces �Permanent dipole �Dipole-dipole interaction �Van der Waal’s forces �Hydrogen bonding

Intermolecular vs Intramolecular forces �INTERmolecular forces: forces between molecules � Forces where whole molecules interact with each other �INTRAmolecular forces: forces within molecules � Forces that hold a single molecule together �So what are examples of intramolecular forces? �Examples of intermolecular forces?

Inter- and Intra-molecular forces �Just as the strength of the intramolecular forces can determine properties such as melting and boiling points, so can intermolecular forces �The increase in intermolecular forces will � Increase m. p. and b. p. (less volatile) � Can determine the Solubility (like dissolves like) � Conductivity (moving charges = conductivity)

London (dispersion) forces �Non-polar molecules can have: � Weak dipole = temporary/instantaneous dipole � This can influence a neighbor molecule = induced dipole

London (dispersion) forces �LDF strength increases as molecular size increases � More e- in atom = more chances for temporary dipoles occurring �LDF are ONLY forces between non-polar molecules �But exist in all molecules!!! �Generally have low m. p. and b. p. = Easy to break weak LDFs

Dipole-dipole attraction �Polar covalent molecules have a permanent dipole �These permanent charges orient other molecules the same way = dipole-dipole attraction �Strength of the attraction depends on distance & orientation of dipoles similar molecular mass

Van der Waal forces �Note: A polar molecule can induce a dipole in a non- polar molecule = dipole-induced dipole �Van der Waal forces: a term that includes both London dispersion forces and the dipole interactions (both dipole-dipole and dipole-induced dipole) �Note: this term does not include H-bonding!

Hydrogen bonding (H-bonding) �Is a special case of dipole-dipole attraction � Hydrogen bonded to: �Nitrogen �Fluorine �Oxygen �These are the strongest form of intermolecular attraction (so b. p. and m. p. are much higher than predicted)

H-bonding �Can you explain why water is less dense in solid than in liquid form? �H-bonds are the strongest intermolecular force, but is still weaker than covalent and ionic bonds

Intermolecular forces

Physical properties – M. P. & B. P. �What happens when CH 3 CH 2 OH is boiled? � H 2 O and CO 2 are released as gases � CH 3 CH 2 OH is released as gas � CH 3 and O 2 are released as gases � H 2 and CO 2 are released as gases �So are intermolecular forces broken or intramolecular forces broken when something is boiled? �Covalent cmpds have lower values than ionic � Weak intermolecular forces vs electrostatic attraction �Higher intermolecular bond strength, polar strength, and increase in molecular size increases m. p. and b. p.

Physical Properties – Solubility �“Like dissolves like” �Non-polar dissolve in non-polar (such as oil) � Due to the formation of LDF �Polar covalent molecules are usually soluble in water and other polar solvents � Dipole interactions and H-bonding �The larger the molecule with non-polar areas, the less polar the overall molecule �Giant covalent are insoluble in nearly all solvents � Too much E needed to break the bonds for interaction

Physical Properties – Electrical conductivity �Covalent compounds do not conduct as there are no charges particles to carry a charge � However, some polar covalent will conduct when in aqueous solution �Giant covalent molecules � Diamond = no conduction � Fullerene & silicon = semiconductors � Graphite & graphene = conductors

Chapter 4. 5: Metallic Bonding Metallic bonds involve a lattice of cations with delocalized electrons.

Delocalized electrons in metals �Metals have: � Few valence electrons (typically) � Low ionization energy � Form positive ions by losing electrons when reacted �When in elemental state, who can accept the e- s? �These electrons tend to “wander off” or be delocalized � These e- are no longer associated with 1 atom!

Metallic bond strength �The strength is determined by: � Number of delocalized e- (more e-, stronger bond) � Charge on cation (higher charge, stronger bond) � Radius of the cation (smaller radius, stronger bond) �Write the full e- configuration for Na and Mg: �Compare the above points and determine which will have stronger bonding. Why?

Metallic bonding strengths �Melting point of Na = 98 °C �Melting point of Mg = 650 °C �Which atom will have a higher melting point: � Na, K or Rb? �How did you determine this? �Melting point of Na = 98 °C �Melting point of K = 63 °C �Melting point of Rb = 39 °C �Transition metals: very strong metallic bonds. Why? � Many delocalized e- from 3 d and 4 s sub-shells

Alloys �Alloys: � Solutions of metals � Enhanced properties �Typically solid but made when molten �Metallic bonds can accommodate other cations of different sizes into lattice �Bound by the delocalized e- in metal (thus metallically bonded) �Have different properties than original components � More chemically stable � Often stronger � Often more resistant to corrosion

Examples of Alloys