1. Chapter 4.1 Notes
Ionic compounds consist of ions held together in lattice structures by ionic bonds

2. Important terms for this section
Chemical bond
Valence electrons
Ions
Cations
Anions
Ionization
Ionic bond
Ionic compound
Lattice
Formula unit
Lattice energy
Volatility
Hydrated
Solvated
Binary
Electronegativity values
Bonding continuum
Note: only outer shell (valence) electrons take part in bonding

3. Ionic bonding
Atoms bonded chemically have different properties than the individual atoms
Ionic bonds are chemical bonds between oppositely charged particles
Ions are created by the gain or loss of valence electrons
Cations: metals – loss of electrons so positively charged (groups 1, 2, 13)
Anions: non-metals – gain of electrons so negatively charged (groups 15, 16, 17)
Stable transition metal ions to know:
Pb2+, Sn4+, Sn2+, Ag+, H+, H-, Cu+, Cu2+, Fe2+, Fe3+
Polyatomic ions to know: see page 142

4. Ionic compounds
Ionic bonding: transfer of electrons from one atom to another
Ionic formula writing and naming:
Mg2+ + F- 
NH4+ + PO4 3- 
Al3+ + S2- 
Ionic formula writing:
Sodium sulfate
Magnesium cyanide
Mercury hydrogencarbonate

5. Ionic compound structure
Ionic lattice is formed:
Electrostatic attraction
Coordination number:
Number of ions around
the given ion in lattice
Lattice can grow indef.
Formula unit:
Ratio of ions present
Lattice energy:
Strength of attraction between ions in same lattice
(smaller = ↑ attraction)

6. Physical props. reflect lattice structure
Melting and boiling points are usually very high (strong electrostatic attraction means stronger bonds)
Industry tends to not extract metals from molten ionics due to the very high temps needed
Low tendency to vaporize = low volatility
Therefore, low odor (think of table salt)

7. Physical props. reflect lattice structure
Solubility: how easy will it dissolve into a solvent (become dispersed through a liquid) to form a solution
Like dissolves like!
Salt is ionic, so it will likely dissolve into….?
Polar solvent
Non-polar solvent
Ionic solvent
Ions dissociate in water and are surrounded by water molecules = hydrated
Ions dissociate in other liquid = solvated

8. Physical props. reflect lattice structure
Electrical conductivity:
Ionic compounds that are molten have moveable charges
Ionic compounds dissolved in solution have moveable charges
Brittleness:
Ionic compounds tend to shatter on impact (not malleable or ductile)
This occurs because the ions moving from impact will put like charges next to each other, so they repulse and split

9. Ionic character (how we know it’s ionic)
How do you know if it’s ionic or covalent? Sometimes the metalloids (for example) show some grey areas…
Si can have 4+ or 4- and H can have 1- or 1+
Remember:

10. Ionic character (how we know it’s ionic)
Determining if a compound is ionic or covalent depends on the difference of electronegativity (section 8 of data booklet)

11. Ionic character
The bonding continuum is the range of ionic to covalent
Based on difference of electronegativity values
Which will likely form ionic bonds? Covalent? Polar covalent?
Be and F Si and O N and Cl K and S

12. Covalent compounds form by the sharing of electrons
Chapter 4.2 Notes
Covalent compounds form by the sharing of electrons

13. Important terms in this section
Molecule
Diatomic
Triatomic
Octet rule
Non-bonding pairs
Double bond
Triple bond
Bond length
Bond strength
Bond enthalpy
Polar
Directionality
Dipole
Partial charge
Pure covalent

14. Covalent bonding Atoms in covalent molecules share electrons
Covalent bond is an electrostatic attraction between a pair of e- and the positive nuclei
Atoms with covalent bonds are called molecules (as opposed to the ionic called formula units)
Hₓ •H  H – H
Two hydrogen atoms bond covalently to form a hydrogen molecule (these are diatomic molecules) ex. of triatomic?
Covalent bonding is favorable (stabilizes the atoms) so forming the bonds releases energy

15. Bonding releases energy

16. Covalent bonding
Octet rule: tendency for atoms to form stable arrangements of 8 e- in outer shell (like Ar)
This is not really a “rule” but a common occurrence as this “rule” has exceptions
Non-bonding pairs or lone pairs: pairs of electrons on an atom that do not participate in the covalent bond

17. Multiple bonds & bond length/strength
Double and triple bonds can be formed when there are not enough electrons to give all atoms a desired octet
Double bond: sharing of 4 e-
Triple bond: sharing of 6 e-
Bonds: use sections 10 & 11 in data booklet
Bond length: measure of distance of bonded nuclei
Bond strength: measure of E needed to break the bond (bond enthalpy)

18. Bond length and strength
Increase of number of bonds increase strength and decreases length

19. Polar covalent bonds These bonds do not share electrons evenly
There is DIRECTIONALITY to the bond and/or the molecule
Dipole: having one end being slightly more positive and one end being slightly more negative

20. Polar covalent bonds
Again, the difference in electronegativity will tell you the extent of how polar the bond

21. Pure covalent bond When the electronegativity difference is zero
H – H F – F are examples
C – H is typically referred to as non-polar, but there is a slight polarity as C is more electronegative (0.4)
Also referred to as non-polar covalent

22. The continuum from ionic to pure covalent demonstrates why it is better to use terms such as “predominately” covalent or ionic or “strongly polar” to describe bonds or substances.