What are intermolecular forces? Forces that exist between molecules (intramolecular forces exist within molecules ex. covalent bonds) Intra- = within Inter- = between

4.4 Intermolecular Forces London (dispersion) forces < dipole-dipole forces < hydrogen bonds.

London (dispersion) forces These intermolecular forces are present in all substances. They are the only intermolecular force present in nonpolar substances. At a given moment in time, the electrons in a molecule can be unevenly spread. In this case, a temporary dipole is formed. This temporary dipole can induce a dipole in a nearby particle, resulting in a weak force of attraction between them.

More Electrons = More attractive force

Dipole-dipole attraction These intermolecular forces are present in polar substances that have permanent dipoles. They are stronger than London dispersion forces but are still quite weak. Dipole-dipole forces are due to the attraction of one dipole by its surrounding dipoles.

Van der Waals’ forces Another commonly used term to refer to intermolecular forces Includes London dispersion forces, induced-dipole and dipole-dipole attractions

Hydrogen Bonding These intermolecular forces are present when hydrogen is directly bonded to fluorine, oxygen, or nitrogen. (Hydrogen bonding is FON!) Temporary and permanent dipole attraction will also be present, but hydrogen bonding is much stronger. Hydrogen bonding was discovered because certain hydrogen compounds, such as water, had an unexpectedly high boiling point that temporary and permanent dipole forces could not account for.

Melting and boiling points More/stronger intermolecular forces = higher melting and boiling points What happens when you boil water?

Melting and Boiling Points More (or stronger) intermolecular forces = higher melting and boiling points It takes more energy (as heat) to overcome the attractive forces between molecules

Solubility “like dissolves like” Polar substances are soluble in polar substances Nonpolar substances are soluble in nonpolar substances Larger molecules are not as soluble if the polar bond is only a small part of the overall structure Giant molecular substances are generally insoluble

Electrical Conductivity Covalent compounds = do not conduct electricity (unless they can break into ions!) Ionic compounds = conduct when molten or dissolved in water Giant covalent = graphite and graphene conduct electricity What about metals?

4.5 Metallic Bonding A metallic bond is the electrostatic attraction between a lattice of positive ions and delocalized electrons. The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion. Alloys usually contain more than one metal and have enhanced properties.

Metallic Bonding When solid metals bond together, one or more valence electrons detach from each atom to become delocalized. Metallic bonding consists of the attraction between the metal cations (positively charged ions) and the delocalized electrons. Metallic bonding: the electrostatic attraction between a lattice of positive ions and delocalized electrons

“A lattice of cations surrounded by a sea of electrons”

Alloys Solid solution of metals with enhanced properties Ex. Bronze

Chart pg. 184

14.1 Further aspects of covalent bonding and structure Covalent bonds result from the overlap of atomic orbitals. A sigma bond (σ) is formed by the direct head-on/end-to-end overlap of atomic oribitals, resulting in electron density concentrated between the nuclei of the bonding atoms. A pi bond (π) is formed by the sideways overlap of atomic orbitals, resulting in electron density above and below the plane of the nuclei of the bonding atoms.

https://www.youtube.com/watch?v=lQ3oDKYYL7k

Identifying sigma and pi orbitals

Some molecules contain a central atom with an expanded octet Usually, the central atom forms bonds to have a “full octet” However, if the central atom is from period 3 or below, sometimes there are more than eight electrons around the central atom Possible because the d orbitals are used to bond There are four distinct molecular geometries in this domain: Trigonal bipyramidal Seesaw T-shaped Linear

Formal charge Treats covalent bonds as if they were purely covalent with equal electron distribution Low formal charge means that less charge transfer has taken place, and usually represents the most stable and preferred structure Formal charge = the number of valence electrons in unbonded atom (V) – the number of electrons assigned to atom in Lewis (electron dot) structure Number of electrons assigned = ½ number of electrons in bonded pairs (½B) + number of electrons in lone pairs (L)

Formal charge Overall formula FC = V – (½B + L)

Which of the following Lewis structures for XeO3 is preferred? Formal charge = the number of valence electrons in unbonded atom (V) – [½ number of electrons in bonded pairs (½B) + number of electrons in lone pairs (L)] FC = V – (½B + L)

The structure with the lowest formal charge is preferred.

Ozone Calculate the formal charge of Ozone

Formal charge distribution