1. Chapter 7
Chemical Formulas and Chemical Compounds

2. caffeine
trichickenfootphos
aspirin
Buckminsterfullerene (buckyball)

3. 7-1 Chemical Names and Formulas
For a molecular compound, formula tells number of atoms of each type in molecule
C8H18
N2O5
C12H22O11

4. 7-1 Chemical Names and Formulas
For ionic compounds, the formula represents the simplest ratio of positive ions to negative ions that will make a neutral compound
Al2(SO4)3
CaCl2
Na2SO4

5. 7-1 Monatomic Ions
Remember, main group elements form ions by gaining or losing electrons to get the electron configuration of a noble gas
Ions formed from a single ion are called monatomic ions
What will be the charge of an ion formed by an element in group 1? Group 2? Group 15? Group 16? Group 17?

6. 7-1 Monatomic Ions
Transition metals – most are +2 or +3, some form +1 or +4
Some transition metals can have more than one oxidation state (charge)
Pb (+2 or +4)
Cu (+1 or +2)
Fe (+2 or +3)
Cr (+2 or +3)

7. 7-1 Naming Monatomic Ions
Monatomic cations – have the same name as the element
Na – sodium, Na+ - sodium
Monatomic anions – drop the ending and add –ide
O – oxygen, O2- - oxide
N – nitrogen, N3- - nitride
Cl – chlorine, Cl- - chloride

8. 7-1 Naming Monatomic Ions
Transition metal ions with more that one oxidation state (charge):
Fe2+ - iron (II), Fe3+ - iron (III)
Cu+ - copper (I), Cu2+ - copper (II)
Pb2+ - lead (II), Pb4+ - lead (IV)

9. 7-1 Binary Ionic Compounds
Binary compounds – compounds with only two elements
Total number of positive charges and negative charges must be equal
Ions Combined: Mg2+, Br-, Br-
Formula:
Ions Combined: Fe3+, Cl-, Cl-, Cl-
Ions Combined: Al3+, O2-

10. 7-1 Binary Ionic Compounds
Write symbols for ions side by side, cation first
Cross over charges, using absolute value of ion’s charge as subscript for the other.
Divide subscripts by largest common factor to give simplest ratio, then write formula.

11. 7-1 Naming Binary Ionic Compounds
Write name of cation first (including Roman numeral if applicable).
Write name of anion second.
AgCl
ZnO
CaBr2
SrF2
BaO
CaCl2
CuO
CoF3
SnI4
FeS

12. 7-1 Compounds Containing Polyatomic Ions
Polyatomic ions – a group of atoms, covalently bonded, with a charge
Only common polyatomic cation is ammonium (NH4+)
Most other polyatomic ions are anions
Many are oxyanions (polyatomic ions that contain oxygen)

13. 7-1 Compounds Containing Polyatomic Ions
If formula requires more than one polyatomic ion to balance charges, use parentheses.
magnesium hydroxide
copper (II) nitrate
aluminum sulfate
potassium phosphate

14. 7-1 Naming Binary Molecular Compounds
Binary molecular compound – usually between two nonmetals
Old system uses prefixes
P4O10
SO3
ICl3
PBr5
N2O3
CI4
mono- 1
di- 2
tri- 3
tetra- 4
penta- 5
hexa- 6
hepta- 7
octa- 8
nona- 9
deca- 10

15. 7-1 Acids binary acids – hydrogen and one of halogens
oxyacids – hydrogen, oxygen and a third element (usually a nonmetal)
Most common acids contain hydrogen
Produce solutions with pH<7

16. 7-1 Acids p. 214 know acids on table 7-5 HF hydrofluoric acid
HCl hydrochloric acid
HBr hydrobromic acid
HI hydroiodic acid
H3PO4 phosphoric acid
HNO2 nitrous acid
HNO3 nitric acid
H2SO3 sulfurous acid
H2SO4 sulfuric acid
CH3COOH acetic acid
HClO hypochlorous acid
HClO2 chlorous acid
HClO3 chloric acid
HClO4 perchloric acid
H2CO3 carbonic acid

17. Lab: Forming and Naming Ionic Compounds
Ionic compounds dissolved in water break apart.
The ions are solvated (surrounded by water molecules).
Solvated ions move around in solution and bump into each other.
Many ionic compounds are soluble in water, but a few are not.

18. Lab: Ionic Compounds
If two ionic compounds are mixed, all of the ions can now interact with each other. New combinations of ions are possible.
If one of the new combinations is insoluble, a precipitate will form.
The insoluble compound will come out of solution as a solid.

19. Lab: Ionic Compounds
AgNO3(aq)  Ag+ + NO3-
NaCl(aq)  Na+ + Cl-

20. 7-2 Oxidation Numbers
Oxidation numbers are assigned to keep track of electrons.
Oxidation numbers are assigned to the atoms in a compound or ion in order to indicate the general distribution of electrons among the bonded atoms.

21. 7-2 Rules for Assigning Oxidation Numbers
Shared electrons are assumed to belong to the more electronegative atom in each bond.
The more electronegative atom has a negative oxidation number, the less electronegative atom has a positive oxidation number.
NaCl
CO2
SF6
LiF

22. 7-2 Rules for Assigning Oxidation Numbers
The atoms in a pure element have an oxidation number of zero. (Na, O2, P4, S8)
The more electronegative element in a binary molecular compound is assigned the number equal to the negative charge it would have as an anion.
The less electronegative atom is assigned the number equal to the positive charge it would have as a cation.
Fluorine has an oxidation number of -1 in all of its compounds.
Oxygen has an oxidation number of -2 in almost all of its compounds. (exceptions: peroxides, in compounds with halogens, OF2)

23. 7-2 Rules for Assigning Oxidation Numbers
Hydrogen is +1 when it is with a more electronegative element and -1 when it is with a metal.
The algebraic sum of all oxidation numbers in a neutral compound is 0.
The algebraic sum of oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion.

24. 7-2 Oxidation Numbers
Assign oxidation numbers to each atom is the following compounds or ions.
HCl
CF4
PCl3
SO2
HNO3 KH
P4O10
HClO3
N2O5
GeCl2
N2O
BF3

25. 7-3 Using Chemical Formulas
Formula Mass – sum of the average atomic masses of the atoms in a formula
Ex. NH4NO3

26. 7-3 Using Chemical Formulas
Percentage composition – percentage by mass of each element in a compound
mass of element in sample of compound x 100 = % element
mass of sample of compound in compound

27. 7-3 Molar Mass
Remember, the molar mass of a compound is numerically equivalent to the formula mass.
Find the sum of the atomic masses of all the atoms in the compound
Use the unit g/mol

28. 7-3 Molar Mass as a Conversion Factor
mass (g)
moles
# of particles

29. 7-3 Sample Problem
What is the mass in grams of 2.5 moles of oxygen gas?

30. 7-3 Sample Problem
Ibuprofen, C13H18O2, is a nonprescription pain reliever. Its molar mass is g/mol.
If a bottle of ibuprofen contains 33g, how many moles of ibuprofen are in the bottle?
How many molecules of ibuprofen are in the bottle?
What is the total mass in grams of carbon in 33g of ibuprofen?

31. 7-3 Hydrates
As some salts crystallize, they bind water molecules in their crystal structure. These are called hydrates.
Water combines with the compound in fixed whole number ratio.

32. 7-3 Hydrates
Sodium carbonate forms a hydrate in which 10 water molecules are present for every formula unit of sodium carbonate.
It is called sodium carbonate decahydrate.
The formula is Na2CO3●10H2O.

33. 7-3 Sample Problem
Find the mass percentage of water in sodium carbonate decahydrate, which has a molar mass of g/mol.

34. 7-4 Determining Chemical Formulas
A substance can be analyzed quantitatively to determine its percentage composition.
From this data, chemists can determine the empirical formula of the substance.
An empirical formula shows the simplest whole number ratio of elements in the compound. H C C H

35. 7-4 Determine the empirical formula…
C6H12O6
NaCl
H2C2O4
H2SO4
C5H10
PCl3
H2O2

36. 7-4 How do chemists calculate empirical formula from percentage composition?
Percentage composition is the percent by MASS of each element in the compound.
The empirical formula is the simplest MOLE ratio of the elements in the formula.
A MASS to MOLES conversion must be involved.

37. 7-4 Solving process 1. Start with the percentage composition.
2. Assume a sample of g.
3. Determine the mass of each element in a g sample of the substance.
4. Determine the moles of each element.
5. Simplify the mole ratio.

38. 7-4 Example
A compound contains 32.38% sodium, 22.65% sulfur and 44.99% oxygen. Determine the empirical formula of the compound.

39. 7-4 Example
Analysis of a g sample of a compound known to contain only phosphorus and oxygen indicates a phosphorus content of g. What is the empirical formula of this compound?

40. 7-4 Example
A compound is known to contain 63.52% iron and 36.48% sulfur. Find its empirical formula.

41. 7-4 Example
Analysis of 20.0 g of a compound containing only calcium and bromine indicates that 4.00 of calcium are present. What is the empirical formula of the compound?

42. 7-4 Determination of Molecular Formula
An empirical formula shows the simplest ratio of elements in a compound.
A molecular formula shows the actual number of atoms of each type in one molecule of a compound.
What additional information is needed to determine the molecular formula if the empirical formula is known?

43. 7-4 Relationship between empirical formula and molecular formula.
The molecular formula is some whole number multiple of the empirical formula.
x(empirical formula) = molecular formula
For example, the molecular formula for benzene (C6H6) is 6 times its empirical formula (CH).

44. Therefore…
x(empirical formula mass) = molecular formula mass
In order to determine the molecular formula of a substance from its empirical formula, you must also know its formula mass.

45. 7-4 Example
A compound has the empirical formula P2O5. Experimentation reveals that its molar mass is 284 g/mol. What is its molecular formula?

46. 13-3 Solution Concentration (p. 412-415)
The concentration of a solution is a measure of the amount of solute in a given amount of solvent or solution.
The terms “concentrated” and “dilute” are relative terms used to describe solutions with a lot of solute or a little solute, respectively.

47. Molarity A way to express the concentration of a solution.
Molarity is the number of moles of solute in one liter of solution.
A “one molar” solution of sodium hydroxide contains one mole of sodium hydroxide in every liter of solution.
The symbol for molarity is M, and such a solution would be written as “1 M NaOH”

48. Molarity Molarity is a ratio. molarity (M) = moles solute (mol)
volume solution (L)
How many moles are there in 2.5 L of a 0.5 M solution?
You have 5 moles of a solute in 2.0 L of solution. What is the molarity of the solution?

49. Molarity: Example Problem
You have 3.50 L of solution that contains 90.0 g of sodium chloride, NaCl. What is the molarity of that solution?

50. Molarity: Example Problem
You have 0.8 L of a 0.5M HCl solution. How many moles of HCl does this solution contain? How many grams of HCl does it contain?

51. Molarity: Example Problem
Describe how you would prepare 2.0 L of 3.2 M KOH.