1. Chemical Formulas Types of formulas
Ionic bonding-transfer of electrons
Covalent bonding-sharing of electrons

2. Determining the types of bonds
To determine the bond types look at values. (Electronegativity value)
On chart. (Given)
Bond type electronegativity difference
Covalent 0.0 – 0.5
Polar covalent
Ionic bonds

3. Example Determine the type of bond that exist between the following:
1. Na and Cl
2. H and O
3. C and O

4. Ionic bonds Often formed between metals and nonmetals
Covalent bonds are formed between similar elements
Binary compounds are those with only two elements.

5. Writing Ionic chemical formulas
1. Write the symbol and charge of the positive ion.
2. Write the negative ion and charge.
3. DO THEY EQUAL ZERO?
YES –You’re done. Rewrite or erase charges
NO – Criss –cross or find the lowest number both charges go into and multiply charges to get that number. The number of times you multiply becomes the subscript.

6. Polyatomics Polyatomics –elements that as a group have 1 charge
Ex. SO₄-² = sulfate
If you need more than 1 polyatomic ion, you must put it in parentheses so subscript multiplies out.
Ex.
** Roman numerals give the positive charge.

7. Examples

8. Examples

9. Examples calcium bromide Iron (III)sulfide Potassium fluoride
Lithium carbonate

10. Formula Mass
The formula mass of a compound is the total mass of all the atoms.
The mass of one molecule is expressed in amus.
The mass of whole moles worth is expressed in grams.

11. Problems
Determine the formula mass of

12. Percent composition
Determine the percent of each element in the whole formula
Weight of I kind of element
Total weight of the formula X 100 =
Example : Determine the % composition of iron(III)chloride.

13. Practice
Iron(III)chloride

14. Terms/additional notes
Chemical formula is a shorthand way to represent composition of a substance
Ex. C₁₂H₂₂O₁₁
Empirical formula- simplest ratio of elements in a compound
Ex. CH₃ is empirical
C₂H₆ is molecular (actual) formula
Subscript – Small number that sits low, behind a symbol. It tells how many atoms of the element are present.

15. More info
In forming bonds, electrons are transferred or shared. This creates ions with resulting charges (positive or negative). These are also called oxidation numbers.
The following shows the changes as a chemical bond forms.
Na Cl → NaCl

16. Na + Cl →NaCl
The oxidation number of a free (uncombined) element is zero.
If an element becomes more positive (sodium) it is “oxidized”. (Oxidation)
If an element becomes more negative (Chlorine) it is “reduced”. (Reduction)
These are often referred to as redox reactions

17. Info cont. Transfer of electrons always involve energy changes.
1 mole Cl mole e- → 1 mole Cl kcal
Is this exothermic or endothermic?

18. Rules for oxidation numbers
1. The oxidation number of an atom of a free element is zero.
2. The oxidation number of a monatomic ion is its charge.
3. The algebraic sum of the oxidations numbers of all atoms in a formula must equal zero.

19. Rules cont.
4. In compounds, the oxidation number of hydrogen is +1 except in metal hydrides where it is HCl (+1) AgH (-1)
In compounds, the oxidation of oxygen is -2, except in peroxides where it is -1 or when it is with very electronegative elements and it is +2.
In compounds composed of nonmetals, the less electronegative element is positive and the more electronegative element is negative.
The sum of the oxidation numbers of the atoms in the formula of a polyatomic ion is equal to its charge.

20. Writing names of ionic formulas
1. Write the name of the positive ion(metal). If the metal has more than 1 possible charge, you must include the charge as a Roman numeral. Ex. Iron(II)
Write the name of the negative ion (nonmetal). Change the ending to “ide”.
A polyatomic keeps its name.
Ex. CaBr is Calcium bromide
CuF is copper(I)fluoride
Na₂SO₄ is sodium sulfate

21. Molecular compounds
Covalent bonds are formed when elements share electrons. Often there are more compounds formed between the same two elements. (usually two nonmetals).
Covalent = molecular
Ex. CO = carbon monoxide
CO₂ = carbon dioxide
*Use prefixes to name the number of atoms not charge.

22. Molecular compound prefixes
Don’t use mono on first word.
1. Mono Hexa
2. di Hepta
3. tri Octa
4. Tetra Nona
5. penta deca

23. Examples
CF₄
N₂O
Sulfur hexafluoride
Dinitrogen pentaphosphide

24. THE MOLE A mole is a quantity.
It is equal to 6.02 x 10 ^23 number of anything. Usually representing very small things like atoms, molecules electrons, formula units etc.
6.02 x 10^23 is also called Avogadro’s number.
Where does this number come from?

25. Where does the mole number come from
1 amu= 1/12 th mass of carbon -12.

26. More mole stuff
The mass in grams of this number of atoms is called gram atomic mass.
The gram atomic mass of an atom is numerically the same as the atomic mass of the element.
So…
The mass of a He atom = 4.0 amus
The mass of a mole of He atoms = 4.0 grams.

27. Mole stuff cont.
Example: Argon 1 atom argon = 39.9 amu 1 mole argon = 6.02 x 10^23 atoms 1 mole argon = 39.9 grams 1 mole argon gas = 22.4 liters (*gases only at STP)

28. MOLE CONVERSIONS
Practice and chart

29. More conversions
Convert 0.9 moles to atoms