1. CCl4 MgCl2
Guess at the names for these two compounds
Predict whether each is ionic or molecular compound
Understanding the difference in bonding, explain what the formula means for each. Draw pictures to help.

2. Ch. 7: Chemical Formulas and Compounds
7.1 Chemical Names and Formulas

3. Chemical Formulas molecular compound
number of atoms of each element contained in a single molecule of the compound
NO2, CH2Cl2,

4. Chemical Formulas ionic compound represents one unit
simplest ratio of the compound’s anions and cations
MgO, Mg(OH)2, NH4Br

5. Monatomic Ions ions formed from a single atom
Not all representative elements easily form ions
Some atoms form covalent bonds instead
Others form ions without noble gas configurations
d and p-block metals
many form 2+ or 3+, some +1 or +4

6. Monatomic Ions
type of ion can be determined by looking at the number of valence electrons in the neutral atom
atom is more stable with a full shell
zero valence
eight valence
determine which is easier for the atom to achieve

7. Monatomic Ions cation anion positive if the atom looses electrons B MgRb Al Na
anion
negative
if the atom gains electrons N F S O Br

8. Naming Monatomic Ions cations anions positive ion
written first in formula
by element’s name
add a Roman Numeral if it can form more than 1 type of ion
that is called the STOCK SYSTEM
anions
negative ion
written second in formula
by element’s root name + -ide ending

9. Practice: Monatomic Ions
Identify the ion created by each atom and give the name for it: Ga Se Ca I N Li

10. Binary Ionic Compounds
compounds made of two different ions
Crossing Over
write two ion types with correct charges
make the anion’s charge, the cation’s subscript
make the cation’s charge, the anion’s subscript
simplify the ratio if possible
combine the names of the two ions

11. Practice Naming
Al and O
Mg and Br
Cu and Br

12. Practice Writing Formulas
iron (III) sulfide
cadmium oxide
potassium nitride
tin (IV) sulfide

13. Cmpds with Polyatomic Ions
the cation or anion could be a group of covalently bonded atoms instead of one atom
most polyatomic ions are oxyanions:
polyatomic ions containing oxygen
more than one polyatomic anion can exist for one element
sulfate ion: SO42- sulfite ion: SO32-

14. Practice Naming
KNO3
Cu(OH)2
NH4Br
CaSO4

15. Practice Writing Formulas
sodium permanganate
iron (II) chlorite
ammonium acetate
calcium carbonate

16. Write the formula for
barium nitrate
zinc phosphate

17. Traditional vs. Stock Stock System
using Roman Numerals in names of metal ions with more than one charge
more important and used more often
Traditional
add –ic or –ous ending to metal name
lower charge: -ous, higher charge: -ic
Shown on the chart below Stock System chart

18. Examples stannous chloride SnCl2 chromic sulfate Cr2(SO4)3
plumbic nitrate
Pb(NO3)4
aurous oxide
Au2O
ferric phosphate
FePO4
cuprous chlorate
CuClO3

19. Examples CuCl2 cupric chloride PbO plumbous oxide Pt3(PO4)2
plantinous phosphate
AuBr
aurous bromide
PbO
plumbous oxide
SnO2
stannic oxide
Fe(NO3)2
ferric nitrate

20. Acids Acids will be focused on greatly later on in the year
Can be identified by H being first in the formula
All acids contain hydrogen and an anion
Can be polyatomic anion : HNO3
Can be monatomic anion: HBr
Name is based on the anion root

21. Binary Acids Only H and monatomic anion Naming Example: Hydro- prefix
Root of element name
-ic ending
Example:
HBr, hydrobromic acid
H2S, hydrosulfic acid

22. Ternary Acids Naming Examples: Polyatomic ion name
If it has –ate ending, change to –ic
If it has –ite ending, change to –ous
Examples:
HNO3: from nitrate, so nitric acid
H2SO3: from sulfite, so sulfurous acid

23. Practice HBr HClO3 H2C2O4 H3AsO3 H2Se H2S Binary Ternary ternary

24. Writing Formulas for Acids
Determine if it is binary (has hydro-prefix) or ternary (no hydro- prefix)
Find the anion contained in the acid (binary- element, ternary- polyatomic ion)
add enough H+1 ions to balance the charge of the anion to the front

25. Practice sulfuric acid nitrous acid hydrofluoric acid
hydrosulfuric acid
phosphoric acid
formic acid

26. Write the formula for sodium sulfate and determine the number of atoms in one unit.

27. Ch. 7: Chemical Formulas and Names
Oxidation States
7.2

28. Oxidation States Also called oxidation numbers
For an individual atom of one type
Used to indicate approximate electron distribution in covalent bonding
What contains covalent bonding?
molecular compounds
polyatomic ions
Not “real” charges since electrons are shared

29. Oxidation States
Shared electron are counted as “belonging” to the more electronegative element
Since it has a greater attraction for the electrons
The electrons are actually located closer to that atom
Periodic Trend for electronegativity: increases as you go up and go to the right

30. Assigning Oxidation States
Atoms in a pure element have an oxidation state of zero.
Na, O2, Fe, etc.
The more electronegative element is assigned a negative oxidation state while the less electronegative has a positive state
CF4:
C is less electronegative: +
F is more electronegative: -

31. Assigning Oxidation States
Fluorine always has a -1 oxidation state
Since it is the most electronegative element
Oxygen usually has a -2
Except in peroxides where it has a -1
Hydrogen usually has a +1
But when it is paired with a metal, it has a -1: ex. LiH

32. Assigning Oxidation States
Sum of the oxidation states must equal zero in a neutral compound
N2O5
O: -2 x 5 = -10
N: +5 x 2 = +10
Sum of states in a monatomic/polyatomic ion must equal the charge of the ion
NO3-1
O: -2 x 3 = -6
N: + 5 x 1 = +5

33. Practice
UF6
H2SO4
ClO31-
CO2

34. Determine the Oxidation States
Br2
NH3
CaSO3
HSO3-
B2H6

35. Using Oxidation States in Naming
Can use Stock System (Roman Numerals) to names molecular compounds too
The Roman Numeral in between names is the oxidation state of the first element
Do NOT simplify the ratio for molecular compounds
Only use either prefixes OR Roman Numerals, not a combination

36. Stock System Naming PCl3 PCl5 N2O Cl: -1, P: +3
Phosphorus (III) chloride
PCl5
Cl: -1, P: +5
Phosphorus (V) chloride
N2O
O: -2, N: +1
Nitrogen (I) oxide

37. Stock System Naming CO
BrF5
CO2

38. Writing Formulas Sulfur (IV) fluoride S: +4 F: -1 SF4
Nitrogen (II) oxide
N: 2+
O: -2 NO
Lead (IV) oxide
Pb: +4
Pb2O4
PbO2
Sulfur (IV) fluoride
S: +4
F: -1
SF4
Chlorine (III) fluoride
Cl: +3
ClF3

39. Writing Formulas Hydrogen (I) oxide carbon (IV) sulfide
Phosphorus (V) oxide
Boron (III) hydride
Hydrogen (I) oxide
carbon (IV) sulfide

40. Name the following compounds
Na2CO3
H2SO4
N2O4

41. Ch. 7: Chemical Formulas and Compounds
7.3 Uses of Chemical Formulas

42. Meaning of Formulas the subscripts in a formula give you:
simplest ratio of atoms OR
the number of atoms in a molecule
What type of compound has a formula that is always the simplest ratio?
What type of compound forms a molecule?
ALWAYS provides you with a ratio of moles of atoms
Example: C13H18O2
13 moles of C : 18 moles of H : 2 moles of O

43. Example
If I have 2.00 moles of C13H18O2, how many moles of each atom would I have?

44. Formula Mass mass of a molecule, ion, or formula unit
sum of mass of all atoms in the chemical formula
in amu
Ex: H2O
amu
formula mass = molecular mass for molecular compound

45. Example
Find formula mass of potassium chlorate.
KClO3

46. Molar Masses mass of one mole of pure substance
numerically equal to formula mass
units: g/mol
Find molar mass of barium nitrate.
Ba(NO3)2

47. Molar Mass in Conversions
can be used as a conversion factor
between grams and moles
What is the mass in grams of 2.50 mol of oxygen gas?

48. Example
Ibuprofen, C13H18O2, is the active ingredient in many pain relievers.
Find molar mass:

49. Example
If the tablets in a bottle contain a total of 33 g of ibuprofen, how many moles are in one bottle?
How many molecules of ibuprofen are in the bottle?

50. Example What is the number of moles of carbon in that bottle?
What is the total mass in grams of carbon in the bottle?

51. I have a sample of 67.9 grams of Na2CO3
Find the number of moles
Find the number of molecules

52. Percentage Composition
percentage by mass of each element in the compound
divide each part of your molar mass into totals for each element
divide these element subtotals by the total molar mass of compound
multiply by 100

53. Example 1
Find the percent composition for calcium nitrite
Ca(NO3)2

54. Example 1 Find the total molar mass
Divide each mass by total molar mass and multiply by 100

55. Example 2
Find percent composition for C13H18O2

56. Hydrates ionic compound with water molecules attached to it
name of ionic compound + prefix for the number of water molecules + hydrate
Ex. CaSO4.2H2O
calcium sulfate dihydrate

57. Practice Naming BaCl2.2H2O LiClO4.3H2O MgCO3.5H2O
barium chloride dihydrate
LiClO4.3H2O
lithium perchlorate trihydrate
MgCO3.5H2O
magnesium carbonate pentahydrate

58. Finding Percent of Water
find percent composition but group all of the atoms in water molecules together
To find the % of water
find total molar mass (including waters)
divide mass of just waters by total
multiply by 100

59. Example 172.172 g/mol Find percent of water in CaSO4.2H2O
total molar mass =
g/mol

60. Example
Divide mass of water by total and multiply by 100

61. Consider N2O4 and NO2 name both of these using prefixes
find the percent composition of N and O for both compounds

62. Ch. 7: Chemical Formulas and Compounds
7.4 Determining Chemical Formulas

63. Empirical Formula
formula containing the simplest whole-number ratio of atoms
not necessarily the CORRECT molecular formula
Why?
Ex. BH3 and B2H6
have same empirical formula
have different molecular formulas

64. Calculating Empirical Formulas from Percent Composition
convert percentage to grams by assuming there is 100 g total of sample
convert grams to moles for each element using molar mass
identify the smallest mole value
divide each mole value by that smallest value

65. Example 1
Quantitative analysis shows that a compound contains 32.38% Na, 22.65% S, and 44.99% O.
Find the empirical formula.
Assuming you have 100 g of sample total
32.38 g Na
22.65 g S
44.99 g O

66. Example 1 Na2SO4 sodium sulfate Convert each of those to moles
Divide by the smallest
/ ≈ 2
/ ≈ 1
/ ≈ 4
Na2SO4
sodium sulfate

67. Calculating Empirical Formula from mass composition
don’t have to assume you have 100 g since you have an actual amount
follow all other steps the same way

68. Example 2
Analysis of a g sample of a compound containing only P and O, is known to contain g of P. Find the empirical formula.
Find the mass of all components
mtotal = mP + mO so mO = mtotal – mP
mass of O : = 5.667g

69. Example 2 P2O5 diphosphorus pentoxide
Convert all mass values to moles using molar mass
Divide by the smallest mole value
To get rid of the decimal, multiply by integer
/ ≈ 1
/ ≈ 2.5
P2O5
diphosphorus pentoxide

70. Molecular Formulas
the molecular formula is the actual formula for a compound
could be same as the empirical formula but doesn’t have to be same
to find molecular formula:
need empirical formula
need actual molar mass (or formula mass)
compare the molar mass of empirical formula to actual molar mass

71. Example 1
In the last example the empirical formula was found to be P2O5. Experimentation shows that the molar mass is actually g/mol. Find the molecular formula.

72. Example 1 Find molar mass of empirical formula
2( ) + 5( ) =
Divide the actual molar mass by the empirical formula’s molar mass
/ ≈ 2
Multiply this number by each subscript in empirical formula
P4O10

73. Example 2 Use the percent composition to write the empirical formula:
41.39 % C, 3.47 % H, and % O
/ = 1
/ = 1
/ = 1

74. Example 2 empirical formula: CHO
If the actual molecular mass is g/mol, what is the molecular formula?
Find the molar mass of CHO
1(12.011) + 1( ) + 1( ) =
Compare it to the molecular mass
/ = 4
Multiply that by subscripts
4(CHO) = C4H4O4

75. Example 3
The empirical formula of a compound is C2H5 and its formulas mass is 87 amu. What is the molecular formula?