1. Conservation of Mass and Stoichiometry
Cartoon courtesy of NearingZero.net
Conservation of Mass and Stoichiometry

2. Ions Cation: A positive ion Mg2+, NH4+ Anion: A negative ion
Cl-, SO42-
Ionic Bonding: Force of attraction between oppositely charged ions.

3. Predicting Ionic Charges
Group 1:
Lose 1 electron to form 1+ ions H+
Li+
Na+ K+

4. Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+

5. Predicting Ionic Charges
Loses 3
electrons to form
3+ ions
Group 13:
B3+
Al3+
Ga3+

6. Predicting Ionic Charges
Lose 4
electrons or gain
4 electrons?
Group 14:
Neither! Group 13 elements rarely form ions.

7. Predicting Ionic Charges
Nitride
Gains 3
electrons to form
3- ions
Group 15:
P3-
Phosphide
As3-
Arsenide

8. Predicting Ionic Charges
Oxide
Gains 2
electrons to form
2- ions
Group 16:
S2-
Sulfide
Se2-
Selenide

9. Predicting Ionic Charges
F1-
Fluoride
Br1-
Bromide
Group 17:
Gains 1
electron to form
1- ions
Cl1-
Chloride
I1-
Iodide

10. Predicting Ionic Charges
Stable Noble gases do not form ions!
Group 18:

11. Predicting Ionic Charges
Many transition elements
have more than one possible oxidation state.
Groups :
Iron(II) = Fe2+
Iron(III) = Fe3+

12. Predicting Ionic Charges
Some transition elements
have only one possible oxidation state.
Groups :
Zinc = Zn2+
Silver = Ag+

13. Writing Ionic Compound Formulas
Example: Barium nitrate
1. Write the formulas for the cation and anion, including CHARGES!
( )
2. Check to see if charges are balanced.
Ba2+
NO3- 2
Not balanced!
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

14. Writing Ionic Compound Formulas
Example: Ammonium sulfate
1. Write the formulas for the cation and anion, including CHARGES!
( )
NH4+
SO42-
2. Check to see if charges are balanced. 2
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!

15. Writing Ionic Compound Formulas
Example: Iron(III) chloride
1. Write the formulas for the cation and anion, including CHARGES!
Fe3+
Cl-
2. Check to see if charges are balanced. 3
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!

16. Writing Ionic Compound Formulas
Example: Aluminum sulfide
1. Write the formulas for the cation and anion, including CHARGES!
2. Check to see if charges are balanced.
Al3+
S2- 2 3
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!

17. Writing Ionic Compound Formulas
Example: Magnesium carbonate
1. Write the formulas for the cation and anion, including CHARGES!
Mg2+
CO32-
2. Check to see if charges are balanced.
They are balanced!

18. Writing Ionic Compound Formulas
Example: Zinc hydroxide
1. Write the formulas for the cation and anion, including CHARGES!
( )
2. Check to see if charges are balanced.
Zn2+
OH- 2
3. Balance charges , if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.
Not balanced!

19. Writing Ionic Compound Formulas
Example: Aluminum phosphate
1. Write the formulas for the cation and anion, including CHARGES!
2. Check to see if charges are balanced.
Al3+
PO43-
They ARE balanced!

20. Naming Ionic Compounds
1. Cation first, then anion
2. Monatomic cation = name of the element
Ca2+ = calcium ion
3. Monatomic anion = root + -ide
Cl- = chloride
CaCl2 = calcium chloride

21. Naming Ionic Compounds (continued)
Metals with multiple oxidation states
- some metal forms more than one cation
- use Roman numeral in name
PbCl2
Pb2+ is cation
PbCl2 = lead(II) chloride

22. Naming Binary Compounds
- Compounds between two nonmetals
- First element in the formula is named first.
- Second element is named as if it were an anion.
- Use prefixes
- Only use mono on second element -
P2O5 =
diphosphorus pentoxide
CO2 =
carbon dioxide
CO =
carbon monoxide
N2O =
dinitrogen monoxide

23. Calculating Formula Mass
Calculate the formula mass of magnesium carbonate, MgCO3.
24.31 g g + 3(16.00 g) =
84.32 g

24. Calculating Percentage Composition
Calculate the percentage composition of magnesium carbonate, MgCO3.
From previous slide:
24.31 g g + 3(16.00 g) = g
100.00

25. Formulas molecular formula = (empirical formula)n [n = integer]
Empirical formula: the lowest whole number ratio of atoms in a compound.
Molecular formula: the true number of atoms of each element in the formula of a compound.
molecular formula = (empirical formula)n [n = integer]
molecular formula = C6H6 = (CH)6
empirical formula = CH

26. Formulas (continued)
Formulas for ionic compounds are ALWAYS empirical (lowest whole number ratio).
Examples:
NaCl
MgCl2
Al2(SO4)3
K2CO3

27. Formulas (continued)
Formulas for molecular compounds MIGHT be empirical (lowest whole number ratio).
Molecular:
H2O
C6H12O6
C12H22O11
Empirical:
H2O
CH2O
C12H22O11

28. Empirical Formula Determination
Base calculation on 100 grams of compound.
Determine moles of each element in 100 grams of compound.
Divide each value of moles by the smallest of the values.
Multiply each number by an integer to obtain all whole numbers.

29. Empirical Formula Determination
Adipic acid contains 49.32% C, 43.84% O, and 6.85% H by mass. What is the empirical formula of adipic acid?

30. Empirical Formula Determination (part 2)
Divide each value of moles by the smallest of the values.
Carbon:
Hydrogen:
Oxygen:

31. Empirical Formula Determination (part 3)
Multiply each number by an integer to obtain all whole numbers.
Carbon: 1.50
Hydrogen: 2.50
Oxygen: 1.00
x 2
x 2
x 2 3 5 2
Empirical formula:
C3H5O2

32. Finding the Molecular Formula
The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?
1. Find the formula mass of C3H5O2
3(12.01 g) + 5(1.01) + 2(16.00) = g

33. Finding the Molecular Formula
The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?
2. Divide the molecular mass by the mass given by the emipirical formula.
3(12.01 g) + 5(1.01) + 2(16.00) = g

34. Finding the Molecular Formula
The empirical formula for adipic acid is C3H5O2. The molecular mass of adipic acid is 146 g/mol. What is the molecular formula of adipic acid?
3. Multiply the empirical formula by this number to get the molecular formula.
3(12.01 g) + 5(1.01) + 2(16.00) = g
(C3H5O2) x 2 =
C6H10O4

35. Synthesis (Composition) Reactions
Two or more substances combine to form a
new compound.
A + X  AX
Reaction of elements with oxygen and sulfur
Reactions of metals with Halogens
Synthesis Reactions with Oxides
There are others not covered here!

36. Decomposition Reactions
A single compound undergoes a reaction that
produces two or more simpler substances
AX  A + X
Decomposition of:
Binary compounds H2O(l )  2H2(g) + O2(g)
Metal carbonates CaCO3(s)  CaO(s) + CO2(g)
Metal hydroxides Ca(OH)2(s)  CaO(s) + H2O(g)
Metal chlorates 2KClO3(s)  2KCl(s) + 3O2(g)
Oxyacids H2CO3(aq)  CO2(g) + H2O(l )
Peroxide H2O2 (l)  2H2O(l) + O2 (g)

37. Single Replacement Reactions
A + BX  AX + B
BX + Y  BY + X
Replacement of:
Metals by another metal
Hydrogen in water by a metal
Hydrogen in an acid by a metal
Halogens by more active halogens

38. The Activity Series of the Metals
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
Metals can replace other metals
provided that they are above the
metal that they are trying to
replace.
Metals above hydrogen can
replace hydrogen in acids.
Metals from sodium upward can
replace hydrogen in water (cold)
Metals between iron upward to
magnesium can replace hydrogen
in hot water (steam)

39. The Activity Series of the Halogens
Fluorine
Chlorine
Bromine
Iodine
Halogens can replace other
halogens in compounds, provided
that they are above the halogen
that they are trying to replace.
2NaCl(s) + F2(g) 
2NaF(s) + Cl2(g)
MgCl2(s) + Br2(g) 
???
No Reaction

40. Double Replacement Reactions
The ions of two compounds exchange places in an
aqueous solution to form two new compounds.
AX + BY  AY + BX
A+ + X- + B+ + Y-  A+ + Y- + B+ + X-
One of the compounds formed is usually a
precipitate, an insoluble gas that bubbles out of
solution, or a molecular compound, usually water.

41. Combustion Reactions
A substance combines with oxygen, releasing a large
amount of energy in the form of light and heat.
Reactive elements combine with oxygen
P4(s) + 5O2(g)  P4O10(s)
(This is also a synthesis reaction)
The burning of natural gas, wood, gasoline
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)

42. Balancing Chemical reactions “Satisfies the Law of conservation of Matter”
Formulas must be written correctly
Balance Polyatomic ions first
Balance elements appearing only once on each side
Balance Hydrogen and oxygen last

43. Stoichiometry
“In solving a problem of this sort, the grand thing is to be able to reason backward. This is a very useful accomplishment, and a very easy one, but people do not practice it much.”
Sherlock Holmes, in Sir Arthur Conan Doyle’s A Study in Scarlet
Stoichiometry - The study of quantities of materials consumed and produced in chemical reactions.

44. Review: Atomic Masses Carbon = 98.89% 12C 1.11% 13C <0.01% 14C
Elements occur in nature as mixtures of isotopes
Carbon = 98.89% 12C
1.11% 13C
<0.01% 14C
Carbon’s atomic mass = amu

45. Review: The Mole
The number equal to the number of carbon atoms in exactly 12 grams of pure 12C.
1 mole of anything = ´ 1023 units of that thing

46. The Mole

47. Using Compound Masses

48. Review: Molar Mass
A substance’s molar mass (molecular weight) is the mass in grams of one mole of the compound.
CO2 = grams per mole
H2O = grams per mole
Ca(OH)2 = grams per mole

49. Review: Chemical Equations
Chemical change involves a reorganization of
the atoms in one or more substances.
C2H5OH + 3O2 ® 2CO H2O
reactants
products
When the equation is balanced it has quantitative significance:
1 mole of ethanol reacts with 3 moles of oxygen
to produce 2 moles of carbon dioxide and 3 moles of water

50. Mole Relations

51. Calculating Masses of Reactants and Products
Balance the equation.
Convert mass to moles.
Set up mole ratios.
Use mole ratios to calculate moles of desired substituent.
Convert moles to grams, if necessary.

52. Working a Stoichiometry Problem
6.50 grams of aluminum reacts with an excess of oxygen. How many grams of aluminum oxide are formed.
1. Identify reactants and products and write the balanced equation. 4 Al + 3 O2 2
Al2O3
a. Every reaction needs a yield sign!
b. What are the reactants?
c. What are the products?
d. What are the balanced coefficients?

53. Working a Stoichiometry Problem
6.50 grams of aluminum reacts with an excess of oxygen. How many grams of aluminum oxide are formed?
4 Al O2  2Al2O3
6.50 g Al
1 mol Al
2 mol Al2O3
g Al2O3 =
? g Al2O3
26.98 g Al
4 mol Al
1 mol Al2O3
6.50 x 2 x ÷ ÷ 4 =
12.3 g Al2O3

54. Limiting Reactant The limiting reactant is the reactant
that is consumed first, limiting the amounts of products formed.

55. Limiting Reagents - Combustion

56. Solving a Stoichiometry Problem
Balance the equation.
Convert masses to moles.
Determine which reactant is limiting.
Use moles of limiting reactant and mole ratios to find moles of desired product.
Convert from moles to grams.