unit 4

Ionic Compounds form when oppositely charged ions attract ­

ionization= where electrons are transferred between atoms ­

oppositely charged ions resulting from this electron transfer are attracted to each other and are held together by electrostatic forces, called ionic bonds ­

ionic compounds are electrically neutral

Ionic Compounds have a Lattice Structure ­

ionic lattice = predictable 3D crystalline structure, fixed arrangement of ions based on a repeating unit ­

coordination number= number of ions that surround a given ion in the lattice (NaCl lattice, the coordination number is 6 because each Na + ion is surrounded by 6 Cl ­ ions and each Cl ­ ion is surrounded by 6 Na + ions) ­

lattice energy = measure of the strength of attraction between the ions within the lattice (greater for ions that are small and highly charged, as they have a larger charge density)

The Physical Properties of Ionic Compounds Reflect Their Lattice Structure ­

  • tend to have high melting and boiling points as the forces of electrostatic attraction between the ions in the lattice are strong and so require lots of energy to break ­ solids at room temperature ­

  • volatility = tendency of a substance to vaporize ­

  • low volatility ­

  • low odour ­

  • easily hydrated (ions are surrounded by water molecules) ­

  • ionic compounds are generally soluble in ionic polar solvents but not soluble in non­polar solvents ­ dissolved in something other than water, they become solvated ­

  • don't conduct electricity in the solid state, but will when molten or in aqueous solution

Covalent bonds form by atoms sharing electrons

­ 2 or more non­metals sharing electrons

­ shared electrons are concentrated in the region between the 2 nuclei and is attracted to them both

­ held together by electrostatic attraction, covalent bond

­ formation of the covalent bond stabilizes the atoms so energy is released as the bond forms

­ forces of attraction between the nuclei and shared electrons are balanced by the forces of repulsion between the 2 nuclei, holdingthe atoms at a fixed distance apart

octet rule: when atoms react, they tend to achieve an outer shell with eight electrons

­molecules with 2 atoms are diatomic, 3 atoms are triatomic, etc.

­lone pair/non­bonding pair = electrons not involved in forming the bond but important in the shape of the molecule

Short Bonds are Strong Bonds

Every covalent bond is characterized by 2 values.

­bond length: a measure of the distance between the 2 bonded nuclei

­ bond strength: usually described in terms of bond enthalpy a measure of the energy required to break the bond

­ atomic radius increases as we go down a group, we would expect these atoms to form molecules with longer bonds that would be weaker

­ multiple bonds have greater number of shared electrons so have a stronger force of electrostatic attraction between the bonded nuclei (strong bond, pulling atoms closer together so they are shorter than single bonds)

Polar bonds result from unequal sharing of electrons

­ electrons are shared unequally

­caused by differences in the electronegativities of the bonded atoms,

more electronegative atom exerts a greater pulling power and gains more possession of the electron pair making a polar bond

­ dipole = form with a polar bond, 2 separate charges

­ level of polarity depends on how big a difference exists in the electronegativity values of the bonded atoms

pure covalent bond: non­polar bonds between diatomic

presence of polar bonds affects its properties

­partial separation of charges introduces some ionic nature into covalent bonds, the more polar the bond, the more similar to an ionic compound

The Octet Rule isn't always followed

­ some molecules are exceptions to the octet rule

­ small atoms (i.e. Be, B) form stable molecules where the central atom has fewer than 8 electrons in its valence shell (incomplete octet)

­ incomplete octets are electron deficient and have a tendency to accept an electron pair from a molecule with a lone pair (i.e. NH3 or H2O) which forms a coordinate compound where the central atom has now gained an octet

*BCl3 is an important catalyst in several synthetic reactions as a results of this tendency to accept electrons

VSEPR Theory

­ we can use Lewis structures to help us determine the 3­D shape,

which plays a role in its reactivity (i.e. enzymes lock and key)

­ VSEPR = Valence Shell Electron Pair Repulsion Theory ­ because

electron pairs in the same valance shell carry the same charge, they

repel each other and so spread themselves as far apart as possible

­ electron pair isn't a great description, should say electron domain,

which includes all electron locations in the valence shell, occupied by

lone pairs or single/double/triple bonded pair

­ total number of electron domains determines the shape!

VSEPR Theory Basics

­ repulsion applies to electron domains (bonds or lone pairs)

­ total number of electron domains around the central atom determines the geometrical arrangement of the electron domains

­ shape is determined by the angles between the bonded atoms

­ lone pairs (non­bonding pairs) have a higher concentration of charge than a bonded pair because they aren't shared between 2 atoms so have higher repulsion

repulsion decreases as follows: lone pair- lone pair > lone pair-bonding pair>bonding pair bonding pair

­ molecules with lone pairs on the central atom have distortions in their structure that reduce angle between bonded atoms

Molecules with Polar Bonds aren't always Polar

­polarity of bonds depends on the charge separation between its 2 bonded atoms, based on their electronegativities

­ polarity of molecule depends on: ­ polar bonds it contains, shape of the molecule

­ if bonds are of equal polarity and are symmetrical, charge separations will effectively cancel each other out = non­polar molecule

­ if bonds of the molecule had different polarity, or bonds are asymmetrical, dipoles won't cancel out = polar molecule (aka. dipole moment, turning force in an electric field)

Giant Molecular Crystalline Solids

­ crystal = single molecule with a regular repeating pattern of covalent bonds, no finite size

­ aka. giant molecular or network covalent structure or macromolecular structure

Allotropes of Carbon

­allotropes are different forms of an element in the same physical state (i.e. oxygen and ozone)

­ different bonding within these structures gives rise to distinct forms with different properties

Carbon Allotropes

graphite: each C atom covalently bonded to 3 others, hexagons in parallel layers, bond angle 120, london dispersion forces

diamond- each C atom covalent bonded to 4 others, tetrahedral, 109.5 degree bond angles

fullerene- sphericaly bonded with 60 C atoms, 12 pentagons, 20 hexagons, closed spherical cage with each C atom bonded to 3 others,

graphene- each C atom bonded to 3 others, hexagon, 120degree, 2D

Intermolecular Forces

­ intermolecular forces = forces that exist between molecules

strength of intermolecular forces determines the physical properties of a substance (i.e. volatility, solubility and conductivity)

London (dispersion) Forces

­ instantaneous/temporary dipole can be forced based on electrons moving around the nucleus

­ induced dipole= lasts for a moment as the electron density is constantly changing, but it may influence the electron distribution in a neighbouring atom

­ London (dispersion) forces = weak attraction between the opposite ends of the 2 temporary dipoles

exist with non­polar molecules and noble gases as well as polar molecules

­ strength increases with increasing molecular size (greater number of

electrons within a molecule increases the probability of temporary dipoles)

­ generally have low melting and boiling points

­ responsible for non­polar substances being able to be condensed into liquids and solids at low temperatures

Dipole­dipole Attraction

­ exists between polar molecules

­ permanent dipoles exist between polar molecules due to one side of the molecules being electron deficient and the other being electron rich

­dipole­dipole attraction = attraction between opposing charges on neighbouring molecules

­strength depends on distance and orientation of the dipoles

­generally stronger than London forces

­cause MP/BP to be higher in polar compounds than those in non- polar compounds

­ for 2 substances of similar molecular mass, the more polar substance will have the higher boiling point

Hydrogen Bonding

­ Hydrogen bonding = when a molecule with hydrogen covalent bonded to a very electronegative atom (F,O,N) are attracted to each other by a

strong intermolecular force

­ type of dipole­dipole attraction

­large EN difference between H and F,O,N causes the electron pair to be pulled away from the hydrogen

­ Hydrogen's small size and lack of electrons to shield the nucleus, H exerts a strong attractive force on a lone pair in the EN atom of a neighbouring molecule

­ strongest form of intermolecular attraction high BP, higher than expected from their molar mass

The Physical Properties of Covalent Compounds are Largely a Result of their Intermolecular Forces

­ stronger the intermolecular force the more energy required to overcome the forces and the higher the MP/BP

­ covalent substances generally have lower MP/BP than ionic compounds strength of the intermolecular forces increases with increasing molecular size and with an increase in the extent of polarity within the molecule

Solubility

­ non­polar substances are generally able to dissolve in non­polar solvents by the formation of London (dispersion) forces between solute and solvent

­ polar covalent compounds are generally soluble in water (highly polar solvent), solute and solvent interact through dipole interactions and H bonding

­ aqueous solubility of polar compound is reduced in larger molecules where the polar bond is only a small part of the total structure

non­polar covalent substances don't dissolve well in water

­ polar substances have low solubility in non­polar solvents as they will remain held to each other by their dipole­dipole attractions

­ giant molecular substances are generally insoluble in all solvents, too much energy is required to break the strong covalent bonds in their structure

**non­polar covalent substances dissolve best in non­polar solvents; polar substances dissolve best in polar solvents**

Electrical Conductivity

­ covalent compounds don't have ions, they don't conduct electricity

­ some polar covalent molecules in conditions where they can ionize will conduct electricity