Unit 6 Notepack:
Chemical Quantities
Chapters 9 &10
NAME_______________________
Period:_______
9.1 Naming Ions
A. ____________________ions: Ions made of single __________________.
B. ___________________ Elements: There is a pattern in predicting how many electrons are lost and
gained for the ___________________ elements, can you guess it?
WRITE ON YOUR PERIODIC TABLES THE CHARGES OF THE REPRESENTATIVE ELEMENTS NOW.
B. _________________ Elements: We cannot determine how many electrons are lost for the
_________________ elements b/c their in their valence electrons can change.
1. How can we tell someone how many electrons lost for the transition metals?
Examples: Name the following transition metal ions: For help, refer to page 144, table 6.3
a. tin (lost 2 electrons)
b. tin (lost 4 electrons)
c. iron (lost 3 electrons)
d. iron (lost 2 electrons)
C. There are 3 exceptions to this rule:
1. Silver is always a _____
2. Zinc and Cadmium are always a ______
*Place their charges in their boxes on your periodic table so you don’t forget this rule.
3. DO NOT USE A ROMAN NUMERAL WHEN NAMING SILVER, ZINC AND CADMIUM IONS.
4. ALWAYS USE A ROMAN NUMERAL WHEN NAMING ANY OTHER TRANSITION METAL ION.
Write the symbol and charge of the
following elements.
a. sulfur
S-2
b. lead (4 electrons lost)
c. strontium
Name the ion
Cation or Anion?
d. argon
e. bromine
f. copper (1 electron lost)
g. selenium
h. silver
i.
cesium
j.
phosphorus
D. ___________________ Ions: Ions made of _________________________________.
1. What endings to polyatomic ions receive when naming them?
2. There are 3 important exceptions, they are:
3. Use chart 6.4 to write the formulas for the following polyatomic ions:
a. ammonium ion
f. permanganate ion
b. sulfate ion
g. hypochlorite ion
c. sulfite ion
h. phosphate ion
d. carbonate ion
i. cyanide ion
e. nitrate ion
j. hydroxide ion
9.2 Ionic Compounds
A. ______________ Ionic Compounds
1. What are Binary Ionic Compounds?
2. What are the “rules” for writing Binary Ionic Compounds?
a. Write the __________ (positive) ion first
b. Write the ______________ (negative) ion last
c. The net charge for the compound must add to _____ (positives + negatives = 0)
d. Use _____________ to indicate how many of each ion you need to “balance” the charge.
Practice:
Write the name &formula for the ionic compound formed b/w magnesium and chlorine.
Write the name &formula for the ionic compound formed b/w sodium and oxygen.
Write the name &formula for the ionic compound formed b/w aluminum and sulfur.
3. Another approach to writing a balanced formula for a compound is to use the crisscross method.
In this method, the numerical charge of each ion is crossed over and used as the subscript for the
other ion. The signs of the numbers are dropped.
Practice
Write the name &formula for the ionic compound formed b/w iron (III) and oxygen.
Write the name &formula for the ionic compound formed b/w calcium and sulfur.
PRACTICE PROBLEMS:
1. Write the formulas for the compounds formed between these pairs of ions. NAME THEM.
a. Ba+2, S-2
b. Li+1, O-2
c. Ca+2, N-3
d. Cu+2, I-1
2. Write formulas for these compounds.
a. sodium iodide
d. rubidium nitride
b. tin (II) chloride (also called stannous chloride)
e. barium fluoride
c. potassium sulfide
f. lithium bromide
3. Write names for these binary ionic compounds
a. ZnS
b. KCl
c. BaO
d. CuBr2
e. CuO (careful!)
f. Ag2S
g. Al2Se3
B. Naming compounds with Polyatomic Ions.
A. Write the symbol for the _______________ followed by the formula of the ___________
_________ _______, then balance charges.
B. Write the formula for lithium nitrate.
C. Sometimes, we need to take more than one polyatomic ion to _____________ the charge to 0. If
this happens, place the polyatomic ion in _____________________ and the subscript
___________________ of the parentheses.
1. Write the formula for aluminum carbonate:
2. Write the formula for potassium sulfate:
3. Write the formula for ammonium phosphate:
4. Write the formula for calcium nitrate:
D. To name a compound with a polyatomic ion, state the __________ first, then the anion just like
you did in binary ionic compounds. Remember the list of polyatomic ions on the back of your
periodic table. Most end in _____ or _____. Remember the three exceptions
____________________ & ______________________&________________________
Practice Problems
1. Write the name & formulas for ionic compounds formed from these pairs of ions:
a. NH4+1, SO32b. Calcium ion, phosphate ion
c. Al 3+, NO3 -1
d. Potassium ion, chromate ion
2. Write formulas for these compounds
a. lithium hydrogen sulfate
b. chromium (III) nitrite
c. mercury (II) bromide
d. ammonium dichromate
3. Name these compounds:
a. LiCN
b. (NH4)2C2O4
c. Fe(ClO3)3
d. Sr(H2PO4)2
e. CaC2O4
f. KClO
g. KMnO4
h. Li2SO3
9.3 – 9.4 Molecular Compounds and Acids
A. Molecular Compounds
1. What is a binary molecular compound?
2. We use prefixes when naming binary molecular compounds: Refer to Table 6.5 and fill in the
following:
Prefix
Number
Prefix
Number
1
6
2
7
3
8
4
9
5
10
a. How do we name binary molecular compound??
b. Say the name of the first element, say the name of the second element, ending in –IDE, and put the
appropriate prefix in to indicate how many of each element there are in the formula:
**If the prefix for the first element in a binary molecular compound is ________, it may be dropped.
However, it must be said if it is for the second element.
**Don’t reduce the subscripts (like you did for binary ionic compounds)
Name these binary molecular compounds
a. N2O ___________________________
b. SF6__________________________________________
c. OF2_________________________________________
d. Cl2O8______________________________________
Write formulas for the following binary molecular compounds
a. nitrogen trifluoride ______________________
b. dinitrogen tetroxide_______________________
c. octoxygen dichlorice_____________________
d. trinitrogen pentoxide_____________________
NAMING ACIDS
A. Acids always are a positive hydrogen cation with a negative anion attached. Charges are crossed just like
ionic compounds.
B. To name acids, look at the anion name ending. Use the anion to name the acid but follow these rules..
There are only three possible anion endings.
1. If the anion ends in –ide, the acid name is _________- ________________ - ______
HCl = __________________, HCN = ____________________, Hydrobromic acid =_________
2. If the anion ends in –ate, the acid name is ________________-______
HNO3=____________________, H2SO4 = ____________________, arsenic acid =_________
3. If the anion ends in –ite, the acid name is __________________-_______
HNO2=____________________, H2SO3 = ____________________, Phosphorus acid =_________
Chapter 10.1: Mole
What is a mole?
How do we measure quantities of
matter?
What are examples of counting
units?
What can we count in chemistry?
How much is a mole?
What is molar mass?
What are the four representative
particles that we can calculate the
molar masses of? Give examples.
10.1 Sample Calculations: We must show our work using dimensional analysis.
1. How many moles of magnesium is 1.25 x 1023 atoms of magnesium?
2. How many moles is 2.80 x 1024 atoms of silicon?
3. How many molecules is 0.360 mole of water?
4. How many moles are equal to 2.41 x 1024 formula units of sodium chloride (NaCl)?
5. How many atoms are in 2.12 moles of propane (C3H8)?
6. How many atoms are there in 1.14 mole of sulfur trioxide?
7. How many moles are there in 4.65 x 1024 molecules of nitrogen dioxide?
8. How many atoms of Carbon are in 2.0 moles of C12H22O11, sucrose sugar?
9. What is the molar mass of carbon?
10. What is the molar mass of hydrogen?
11. What is the molar mass of sulfur?
12. What is the molar mass of sulfur trioxide?
13. What is the molar mass of hydrogen?
14. What is the molar mass of carbon dioxide?
15. What is the molar mass of sodium chloride?
16. What is the molar mass of ammonium carbonate?
17. What is the molar mass of potassium oxide?
Review Problems:
1. Find the gram formula mass of each compound:
a. lithium sulfide
b. iron (III) chloride
c. calcium hydroxide
2. How many oxygen atoms are in a representative particle of each substance?
a. ammonium nitrate
b. acetylsalicylic acid (C9H8O4), the chemical name of aspirin
c. ozone (O3), a disinfectant and natural molecule found in the atmosphere
3. How many moles in each of the following:
a. 1.50 x 1023 molecules of ammonia, NH3?
b. 1 billion (1 x 109) molecules of oxygen, O2?
c. 4.81 x 1024 atoms of lithium, Li?
Chapter 10.2: Mole mass and Mole Volume Relationships.
How are moles and mass related?
What is STP?
Why is STP important when
quantifying gases?
What is the volume of a Mole of
Gas at STP?
Mole Map
One step examples:
1. Find the mass of 2.7 moles of
C6H12O6.
2. How many moles are in 6.72g
of Silver Nitrate?
3. Determine the volume, in liters,
of 0.60 moles of sulfur dioxide
gas at STP.
4. Determine the number of
moles of oxygen gas in 11.5 L at
STP.
Multi-step examples:
5. What is the mass of 27 liters of
nitrogen dioxide at STP?
6. If you have 9.64 x 1024 F.U.’s of
sodium oxide, what would be
the volume of that gas at STP?
7. If you weight 300 grams of
Aluminum Oxide, how many
formula units will you have?
Sample Calculations:
1. Find the mass, in grams of each.
a. 3.32 mole of potassium atoms, K.
b. 4.52 x 1021 molecules of C6H12O6
c. 0.0112 liters of carbon dioxide
2. Calculate the number of moles in 75 grams of each substance.
a. dinitrogen trioxide
b. sodium oxide
3. How many grams are in 9.45 liters of dinitrogen trioxide?
4. Find the number of moles in 92.2 grams of iron (III) oxide
5. What is the volume at STP of these gases?
a. 3.20 x 10-3 mol carbon dioxide
b. 0.960 grams of methane, CH4
c. 3.70 x 10 24 molecules nitrogen gas
6. Assuming STP, how many moles are in these volumes?
a. 67.2 liters of sulfur dioxide gas
b. 0.880 liters of helium gas
c. 1,000 liters of neon gas
Chapter 10.3: Percent Composition and chemical formulas
What is percent composition of a
compound?
How do we find the percent of
anything?
What is percent composition used
for and how do we show our work?
How do you calculate the percent
composition of a compound?
Example: Calculate the percent composition of potassium dichromate.
Sample Calculations:
1. 9.30 grams of magnesium combine completely with 3.48 grams of nitrogen gas to form a compound.
What is the percent composition of this compound?
2. 29.0 grams of silver combine completely with 4.30 grams of sulfur to form a compound. What is the
percent composition of this compound?
3. Calculate the percent composition of these compounds:
a. sodium bicarbonate
b. ammonium chloride
c. sulfur trioxide
4. Calculate the percent nitrogen in these common fertilizers:
a. CO(NH2)2
b. NH3
c. NH4NO3
Example: Calculate the mass of carbon in 82 grams of propane (C3H8) by multiplying the 82 grams by the
percent of carbon in the compound. Show your work by using dimensional analysis.
Sample Calculations:
5. Calculate the grams of nitrogen in 125 grams of each fertilizer.
a. CO(NH2)2
b. NH3
c. NH4NO3
A. Using Percent as a Conversion Factor
1. You can use percent composition to calculate the number of __________ of an element contained in a
________________ amount of a compound. To do this, you multiply the mass of the compound by a
conversion factor that is based on the percent composition.
Example: Calculate the mass of carbon in 82 grams of propane (C3H8).
Sample Calculations:
Calculate the grams of nitrogen in 125 grams of each fertilizer.
a. CO(NH2)2
b. NH3
c. NH4NO3
B. Calculating Empirical Formulas
1. Determining the percent composition of a compound has an important application – calculating the
empirical formula:
2. Define empirical formula:
3. Define molecular formula:
4. The empirical formula may or may not be the same as the molecular formula
5. Practice: These are all molecular formulas, in other words, true formulas. Write their empirical
formulas:
H20
H2O2
CO2
N2O4
6. We can use the percent composition of a compound to determine its empirical formula.
Example: What is the empirical formula of a compound that is 25.9% nitrogen and 74.1% oxygen?
Sample Calculations:
1. Calculate the empirical formula of each compound.
a. 94.1% O, 5.9% H
b. 79.8% C, 20.2% H
c. 67.6% Hg, 10.8% S, 21.6% O
d. 27.59% C, 1.15% H, 16.09% N, 55.17% O
C. Molecular Formula Calculations
1. The molecular formula is either the __________ as the empirical formula or it is a simple
_____________ number multiple of an empirical formula. For example, the empirical formula of
glucose is _________________. The molecular formula is six times larger equaling
______________________. Notice that the molar mass of the molecule is __________ times larger
than the empirical formula
2. Calculate the molecular formula of a compound whose molar mass is 60.0 g/mol and the empirical
formula is CH4N.
3. Find the molecular formula of ethylene glycol, which is used as antifreeze. The molar mass is 62 g/mol
and the empirical formula is CH3O