CHAPTER 7 – CHEMICAL FORMULAS AND CHEMICAL COMPOUNDS
Naming Ions and Binary Ionic Compounds
CATIONS
• Lose electrons, positive charge
• Use element name as cation name
• d-block cations (except Ag+, Zn2+, and Cd2+) and some lower left p block cations need Roman
numeral in name
Examples:
Write the following cations when given the name:
• Sodium _________
• Hydrogen _________
• Vanadium (III) _________
• Tin (II) _________
Give the name of each of the following cations:
• Sr2+ __________________
• Cd2+ __________________
• Cu2+ __________________
• Pb3+ __________________
• Pb2+ __________________
• K+ __________________
ANIONS
• Gain electrons, negative charge
• Drop element name ending and add –ide
Examples:
Write the following anions when given the name:
• Bromide _________
• Oxide _________
• Nitride _________
Give the name of each of the following anions:
• S2- __________________
• Cl- __________________
• P3- __________________
POLYATOMIC IONS
• Molecule with an overall charge
• MEMORIZE names
• Combine as complete ions to form ionic compounds
Examples:
Write the following polyatomic ions when given the name:
• Oxalate _________
• Chlorate _________
• Acetate _________
• Ammonium _________
MS. JENSEN GENERAL CHEMISTRY – NOTES – CHAPTERS 7
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Give the name of each of the following polyatomic ions
• O22- __________________
• CH3COO- __________________
• Cr2O72- __________________
• AsO43- __________________
• MnO4- __________________
NAMING IONIC COMPOUNDS: NAME à FORMULA
1. Write cation first (include charge)
2. Write anion second (include charge)
3. “Cross over” charges to determine number of each atom or polyatomic ion needed
4. Make sure total charge on compound is zero
5. Rewrite formula without charges, without subscript 1
Examples:
• Lithium oxide
•
Barium nitrate
•
Aluminum sulfate
More practice:
Write the formulas for the following ionic compounds:
• Sodium chloride
•
Aluminum oxide
•
Ammonium nitrate
•
Calcium hydroxide
•
Barium chloride
•
Magnesium oxide
•
Lead (II) Nitride
NAMING IONIC COMPOUNDS: FORMULA à NAME
1. (Determine charge on cation)
2. (Determine charge on anion)
3. Write name of cation (Include Roman numeral if necessary: d-block except Ag+, Zn2+, and Cd2+)
4. Write the name of the anion
Examples:
• Sr(ClO2)2
CuCl
•
Pb(NO3)2
•
Pb(NO3)3
KI
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Some Common Monatomic Ions
You do not need to memorize the monotomic ions, except Ag+, Cd2+, and Zn2+. The charge on main
group (s and p block) ions can be determined from the periodic table. These elements consistently form
monatomic ions of only one charge. The d-block elements always form cations (positive charge).
However, most of the d-block elements can form more than one type of monatomic ion (different
charges). The amount of charge on a d-block ion must be determined from the Roman numeral of the
cation name or by looking at the chemical formula of an ionic compound. There are three d-block
elements that consistently form ions of only one charge: Ag+, Cd2+, and Zn2+.
Some Common Polyatomic Ions
You need to memorize these at least one time. You will be tested specifically on polyatomic ions.
Following the final polyatomic ions quiz, you will be expected to be familiar with these ions and use them
throughout the remainder of the school year, but during tests you will be provided with a chart of ions.
Polyatomic Ion Quiz Dates:
Quiz 1 (1+ and 1- ions):
Quiz 2 (2+, 2-, and 3- ions):
Final Quiz (all ions):
Weds Feb 27 (A) and Thurs Feb 28 (B)
Mon March 4 (A) and Tues March 5 (B)
Friday March 8 (A) and Monday March 11 (B)
Retakes: You may retake one of the quizzes one time only. All retakes must be completed by the last day
of the term (Friday, March 22rd).
1+
ammonium
2+
NH4
+
dimercury
1acetate
bromate
chlorate
chlorite
cyanide
dihydrogen phosphate
hydrogen carbonate
(bicarbonate)
hydrogen sulfate
hydroxide
hypochlorite
nitrate
nitrite
perchlorate
permanganate
Hg2
2-
CH3COO
BrO3ClO3ClO2CNH2PO4HCO3HSO4OHClONO3NO2ClO4MnO4-
-
32-
carbonate
chromate
dichromate
hydrogen phosphate
oxalate
peroxide
sulfate
CO3
arsenate
2CrO4
phosphate
Cr2O72HPO42C2O42O22SO42-
sulfite
SO32-
MS. JENSEN GENERAL CHEMISTRY – NOTES – CHAPTERS 7
3+
2+
AsO43PO43-
3
Chemical Formulas
C8H18
Al2(SO4)3
Formula Mass
• Mass of __________________, _____________________________, or _____________.
• Calculate by_____________________________________________________________.
• Units:
Practice: Calculate the formula mass of the following compounds:
HCl
Na2O
Molar Mass
•
•
•
Mass of _____________________ (6.022 × 1023) of any molecule, formula unit, or ion.
Calculate by adding ___________________________________ (from the periodic table).
Units:
Practice: Calculate the molar mass of the following compounds:
H2SO4
Ca(NO3)2
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Molar Conversions
(Review from Chapter 3 – see your notes p. 7-8 and HW 3D)
Mass in
grams
Amount in
moles
Number of
particles
Practice:
1. How many moles are in 2.0 g CaCl2?
2. How many grams are in 1.5 mol NH3?
3. How many molecules are in 6.0 mol H2O?
4. How many moles are in 5.0 × 1022 molecules H2?
5. How many formula units are in 3.0 g NaCl?
6. How many grams are in 3.5 × 1024 molecules NH3?
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Percentage Composition
Percentage composition:
The mass percentage of an element in a compound is the same regardless of the sample size.
This is because the ratio of one element to another in a given compound is the same regardless of
sample size.
To calculate percentage composition:
Examples: Find the percentage composition of the following compounds:
Carbon dioxide
Sodium nitrate
PRACTICE :
1. Find the percentage composition of the following compounds
Lead (II) chloride, PbCl2
Barium chlorate, Ba(ClO3)2
2. What is the percentage by mass of hydrogen in methane, CH4?
3. Find the mass percentage of water in ZnSO4⋅7H2O
MS. JENSEN GENERAL CHEMISTRY – NOTES – CHAPTERS 7
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Determining Chemical Formulas
Molecular formula: includes symbols of elements and number of atoms in a molecule
Empirical formula: includes symbols of elements in compounds with subscripts that show
smallest whole-number ratio of atoms
•
•
Ionic compounds – formula unit is already smallest whole-number ratio
Molecular compounds (covalent bonds) – molecule is not always smallest whole-number ratio
Examples:
Molecular Formula
Empirical Formula
Methane
Ethane
Water
Benzene
Calculation of Empirical Formulas
1. Use percentage composition to convert to a mass composition (assume 100 g sample)
2. Convert mass to moles for each element (use molar mass for each element)
3. Find smallest whole-number ratio by dividing each number of moles by the smallest
number of existing ratio. (Look at answer as a simple fraction if necessary.)
Examples:
1. A compound contains 32.38% sodium, 22.65% sulfur, and 44.99% oxygen. Find the
empirical formula.
2. A compound contains 4.43 g phosphorus and 5.72 g oxygen. Find the empirical formula.
MS. JENSEN GENERAL CHEMISTRY – NOTES – CHAPTERS 7
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Calculation of Molecular Formulas
• You MUST know formula/molar mass to calculate molecular formula
x(empirical formula) = molecular formula
x(empirical mass) = molecular mass
𝑥 = molecular mass
formula mass
EXAMPLES:
1. Determine the molecular formula of the compound with an empirical formula of CH and
a formula mass of 78.110 amu.
2. A sample of a compound with a molar mass of 34.00 g/mol consists of 0.44 g H and 6.93
g O. Find its molecular formula.
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