1. Ionic Compounds
Ch.6 & 7

2. Ionic Bonding – strongest of all bonds
definiton – results from the electrical attraction between cations (+) and anions (-)
“ions” means charges in our case
Calculated by the difference between electronegativity values (over 1.7)
Figure 2 -- p.176

3. Ionic Bonding
Positive ions are formed from the metal elements when they lose electrons
Negative ions are formed from the nonmetal elements when they gain electrons
The two, (+) and (-) come together to form a new compound

4. Ionic Compounds
definition – composed of positive and negative ions that are combined so that the numbers of positive and negative charges are equal
In a neutral compound, the net charge will be zero
Balance the charges from cations and anions until the compound is neutral

5. Forming Ionic Compounds
Na+ + Cl-  NaCl
Ca2+ + Cl-  CaCl2
Mg2+ + O2-  MgO
Fe3+ + S2-  ______
beryllium + iodine  ______

6. Naming Ionic Compounds
Name the metal first (complete name)
Name the root of the nonmetal the change the ending to –ide
If there is a polyatomic (more than one atom carrying one charge) use the name as it is
Transition metals have a Roman numeral after it to tell its charge

7. Naming Ionic Compounds
NaCl ____________________
CaCl2 ____________________
MgO ____________________
Fe2S3 ____________________
BeI2 ____________________

8. Ionic Compound Characteristics
A formula unit is the smallest whole number ratio of cation to anion
ex. NaCl
Ionic compounds mostly have a crystalline structure
Lattice energy is the energy released when one mole of an ionic crystalline compound is formed from gaseous ions
Insert the crystalline structure of NaCl

9. Ionic Compound Characteristics
Generally have a high melting point and boiling point
Hard but brittle
Not electrical conductors as solids
Most can dissolve in water
Strongest bond

10. Polyatomic Ions Definition: a charged group of covalently bonded atoms
Lewis structures can be used to show the electron placement and where the charge of the ion comes from (p.194)

11. Polyatomic Ions Names have a system:
ending -ate means the highest number of oxygen bonded for that nonmetal (sulfate, chlorate, iodate, etc)
ending –ite means one less oxygen than -ate (sulfite, chlorite, iodite, etc)
Chlorine is a special case: ClO4-, ClO3-, ClO2-, ClO-, Cl- (p.226)

12. Polyatomic Ions to Study
Phosphate PO4-3
Chlorate ClO3-
Nitrate NO3-
Sulfate SO3-2
Carbonate CO3-2

13. Oxidation Numbers
Definition: the number of electrons that must be added to or removed from an atom in a combined state to convert the atom into the elemental form
Example: H+ + Cl-  HCl
KMnO4 K = ___ Mn = ___ O = ___

14. Oxidation Number Practice
NaH Na = _____ H = _____
KClO2 K = ____ Cl = ____ O = ____
KClO3 K = ____ Cl = ____ O = ____

15. Metallic Bonding
Definition: the chemical bonding that results from the attraction between two metal atoms and the surrounding sea of electrons
Properties include:
High electrical and thermal conductivity
Strong absorber and reflector of light
Conforms to shape easily
Malleability – hammered into sheets
Ductility – drawn into wires

16. Covalent Molecules
Ch. 6 & 7

17. Covalent Bonding
Definition: bond formed by the sharing of electron pairs between two nonmetals
There are two types of covalent bonds:
Nonpolar: electrons are shared equally
Polar-covalent: unequally shared electrons

18. Electronegativities
Every element on the PT has an electronegativity value
Those values are subtracted to see whether is:
Nonpolar (0-0.3)
Polar-covalent ( )
Ionic (over 1.7)

19. Lewis Structures
Definition – formulas in which atomic symbols represent nuclei and inner-shell electrons, dot-pairs in covalent bonds
Sample Problem C - p.185
Practice - p.186

20. Molecular Geometry
VSEPR theory – repulsion between the sets of valence electrons surrounding an atom causes these sets to be oriented as far as possible
Use VSEPR with Lewis structures to come up with shapes
Shapes are:
Linear, trigonal planar, tetrahedral, bent, and trigonal pyramidal, etc. (p.200)
Sample Problem F – p.201

21. Intermolecular Forces
Definition: forces in between molecules
Weaker than bond strength
Types:
Dipole-dipole: strongest in polar covalent
Hydrogen bonding: hydrogen bonded to highly electronegative atom
London dispersion forces: weakest force due to motion of atoms in compound

22. Naming Covalent Molecules
Use prefixes listed on p.228
Prefixes tell how many atoms of each nonmetal make up the molecule
The nonmetal farthest left on PT is written first, then the most electronegative atom (farthest right on the PT)
Mono does not appear on the first atom to notate one, it is understood
Omit vowels on the prefix if there is a vowel on the element
Second atom ends in -ide

23. Covalent Prefixes mono 1 di 2 tri 3 tetra 4 penta 5 hexa 6 hepta 7
octa 8
nona 9
deca 10
Two Rules of Thumb:
If two vowels end up next to each other, the vowel on the prefix will be deleted.
Mono is never used for the first element.

24. Practice Name the following covalent compounds: CO2 _________________
OF3 _________________
SO2 _________________

25. Acids Formed from H+ + a nonmetal or a polyatomic
If H+ is bonded to a nonmetal, use the prefix hydro- + the root of the nonmetal + ending with –ic acid
If H+ is bonded to a polyatomic, do NOT use the prefix hydro! Use the root of the polyatomic and use ending from –ate to –ic or –ite to –ous acid
Section Review p.231 #4 f-h

26. Practice Name the following acids: HCl ________________
HClO4 ________________
H2SO3 ________________
HI ________________