Chemistry 30
Review of Chemistry 20
Name: ______________
The Atom
An atom is made up of the following subatomic particles:
Electrons: negatively charged particles that orbit around the nucleus.
Protons: positively charged particles contained in the nucleus.
Neutrons: particles in the nucleus that aren’t charged but affect the mass of an atom.
To Do: Please label the subatomic particles on the diagram of an atom above.
Elements On The Periodic Table
The Periodic Table provides information about each element. Consider the first
box on the upper left side of the Periodic Table. This is the box for the element
Hydrogen and contains information about that element.
To Do: label the atomic number, atomic symbol, the name of the element and the
average atomic mass.
1
1.01
H
hydrogen
Atomic Number: indicates the number of protons in the nucleus. The atomic number defines an
atom. All atoms of an element have the same atomic number. No two elements have the same
atomic number.
Mass Number: is the sum of the protons plus the number of neutrons in the nucleus. Neutrons
add mass to an atom but do not change the atomic number. The mass of a proton is the same as a
neutron.
Chemistry 30
Counting Subatomic Particles from the Periodic Table



The number of protons in an atom will NEVER change. An atom cannot gain or lose
protons!
The number of electrons in an atom CAN change.
o In a neutral atom, the number of electrons must be equal to the number of protons.
o If an atom gains or loses electrons, the number of protons and electrons are not
the same so the atom will no longer be neutral – it will now possess a charge. We
call a charged atom an ion.
 If an atom loses electrons it becomes a positively charged ion. (cation)
 If an atom gains electrons it becomes a negatively charged ion. (anion)
When examining the periodic table you will see the atomic mass is also reported. The
atomic mass is the combined mass of the protons and neutrons in the atom. Therefore,
the number of neutrons in an atom is the mass number – the atomic number. (Note:
when counting protons and neutrons, we always round the mass to a whole number!)
o Isotopes are atoms of the same element having different atomic masses. That is,
one is “heavier” than the other because it has more neutrons. The reason the
periodic table gives the mass as a decimal number is that not all atoms have the
same number of neutrons so we have to take an average atomic mass of all the
different forms of the same kind of atom.
To Do: Complete the following chart on counting subatomic particles of atoms.
Name and
symbol of
element
Atomic
Mass
Atomic
number
14
7
Protons
Electrons
10
Neon
20
10
Ne
11
10
Argon
42
18
18
Ar
17
18
Neutrons
Neutral
atom,
cation or
anion?
Chemistry 30
Bonds



Remember that atoms can be combined to make molecules or compounds. Bonds
connect atoms together.
The valence shell electrons actually determine how many of each element will combine
to make a compound. They represent the number of electrons that will be shared,
transferred or accepted between elements to form a stable compound.
There are 2 different types of bonds:
1. Ionic Bonds: formed when metals transfer their electron(s) to non-metals
2. Covalent Bonds: formed when electrons are shared between two non-metals.
Naming Compounds
Ionic compounds:
 The first element (the metal ion) in the compound does not change its name.
 The second element (the non-metal ion) drops its ending and adds “-ide”
 Metals with more than one charge use a roman numeral in brackets to indicate the charge
 Use a table for the names of polyatomic ions.
o Examples:
LiCl = lithium chloride
FeCl2 = iron (II) chloride
NaOH = sodium hydroxide
Covalent/Molecular compounds:
 The second element still drops its ending and adds “-ide”
 Use prefixes to indicate the number of each element in the compound. Mono is never
used on the first element.
o Examples:
N2O3 = dinitrogen trioxide
SiO2 = silicon dioxide
Naming acids:
Acids are covalent compounds formed when hydrogen combines with non-metal or polyatomic
ions. The hydrogen is written first in the compound and the naming is based on what the anion’s
(negative ion) name ends in: -ide, -ite or -ate.
Naming acids with anions that end in –ide (for the most part, anions from the PTE):
 Use the prefix “hydro” followed by the name of the anion which has dropped
its ending and added “ic” to it.
 The second word is “acid”.
o Example: HF is hydrofluoric acid (from the anion fluoride)
o Example: HCN is hydrocyanic acid (from the anion cyanide)
Naming acids with polyatomic anions that end in –ate or –ite:
 The first word is the polyatomic anion which has dropped its ending and
added either
o –ic if the anion name ends in –ate
o –ous if the anion name ends in -ite
 the second word is “acid”
o Examples: H2SO4 is sulfuric acid (from the anion sulfate)
HClO3 is perchlorous acid (from the anion perchlorite)
Chemistry 30
To Do: Name the following compounds and indicate if each is an ionic compound, an acid or a
covalent compound.
Name
Ionic, Acid or Covalent?
1. KNO3
________________________
________________________
2. CO2
________________________
________________________
3. H2SO4
________________________
________________________
4. Pb(NO3)4
________________________
________________________
5. Ca(OH)2
________________________
________________________
6. FeCl3
________________________
________________________
7. CaCl2
________________________
________________________
8. HCl
________________________
________________________
9. P2O3
________________________
________________________
10. CCl4
________________________
________________________
11. H2CO3
________________________
________________________
12. H2CO2
________________________
________________________
13. (NH4)2S
________________________
________________________
14. MgBr2
________________________
________________________
15. SeF2
________________________
________________________
Chemistry 30
Writing Formulas
1. Ionic Compounds (metals and non-metals):







Positive ion written first, negative ion second.
The sum of the charges will be zero because ions add on until a neutral molecule
is formed
SHORTCUT: You can criss-cross charge values by writing the number behind
and below the opposite element. If only one atom is needed, you do not write the
subscript 1. Charges of the same value just cancel out. (+1/-1, +2/-2 etc).
Roman numerals are used with metals that make more than one charge. The
number indicates the charge NOT how many atoms are in that compound!
Use the tables provided for charges of ions and for polyatomic ions (ions that, as a
group of atoms, hold a charge).
If there are more than one polyatomic ion needed, use brackets.
Reduce your formula to its lowest form.
Examples:
 Sodium chloride  Na1+, Cl1- = NaCl
 Calcium chloride  Ca2+, Cl1- = CaCl2
 Iron (III) oxide  Fe3+, O2- = Fe2O3
 Aluminum sulfate  Al3+, SO42- = Al2(SO4)3
 Tin (IV) sulfate  Sn4+, SO42- = Sn(SO4)2
(charges cancel)
(criss-cross numbers)
(roman numeral indicates 3+)
(polyatomic ion)
(reduce from Sn2(SO4)4)
2. Covalent/Molecular compounds (two non-metals)



Write the compound by ignoring the charges and writing exactly what the name states
NEVER reduce formulas.
Prefixes are:
1- mono
6-hexa
2- di
7-septa
3-tri
8-octa
NOTE: we never use mono for the
4-tetra
9-nona
first element!
5- penta
10-deca
Examples:
 Nitrogen dioxide  NO2
 Carbon monoxide  CO
 Diphosphorous pentaoxide  P2O5
 Disilicon tetroxide  Si2O4 (don’t reduce!)
Chemistry 30
To Do: Write the formula for the following compounds and indicate if it is ionic or
covalent/molecular.
Formula
Ionic or Covalent
1. lithium fluoride
___________________
____________________
2. zinc carbonate
___________________
____________________
3. magnesium chromate
___________________
____________________
4. aluminum nitride
___________________
____________________
5. dinitrogen trioxide
___________________
____________________
6. ammonium sulfate
___________________
____________________
7. carbon tetrachloride
___________________
____________________
8. aluminum nitrate
___________________
____________________
9. carbon monoxide
___________________
____________________
10. aluminum hydroxide
___________________
____________________
11. cobalt (II) bromide
___________________
____________________
12. silicon tetraiodide
___________________
____________________
13. silver nitride
___________________
____________________
14. zinc acetate
___________________
____________________
15. beryllium sulfite
___________________
____________________
Chemistry 30
3. Formulas for Acids
Acids are covalent compounds formed when hydrogen combines with non-metal or
polyatomic ions. Even though they are covalent, the formulas are calculated the same
way as with ionic compounds; just criss-cross the charges. However, there are three rules
to remember:
I.
If the prefix “hydro” is used in the name, it is hydrogen bonding with a nonmetal off the periodic table (or a polyatomic ion that ends in –ide).
 Ex. hydrofluoric acid is HF; hydrogen 1+ and fluoride 1 Ex. Hydrocyanic acid is HCN; hydrogen 1+ and cyanide 1II.
If the name ends in –ic, a polyatomic anion with the ending –ate was used.
Use your sheet!
 Ex. Sulfuric acid is H2SO4; hydrogen 1+ and sulfate 2III.
If the name ends in –ous, a polyatomic anion with the ending –ite was used.
Use your sheet!
 Ex. Perchlorous acid is HClO3; hydrogen1+ and perchlorite1-
Write the formula for the following acids:
1. hydrobromic acid
__________________________________
2. carbonic acid
__________________________________
3. iodic acid
__________________________________
4. hydrochloric acid
__________________________________
5. chlorous acid
__________________________________
6. nitrous acid
__________________________________
7. nitric acid
__________________________________
8. sulfurous acid
__________________________________
9. hydrosulfuric acid
__________________________________
10. chromic acid
__________________________________
11. hydroiodic acid
__________________________________
12. hydrophosphoric acid
__________________________________
13. dichromic acid
__________________________________
14. iodous acid
__________________________________
15. selenic acid
__________________________________
Chemistry 30
Types of Reactions
Synthesis or Composition are reactions in which two or more substances combine to form a
more complex substance. The general formula is A + B  AB.
Examples
2Na + Cl2  2 NaCl
SO3 + H2O  H2SO4
Li2O + H2O  2 LiOH
CaO + CO2  CaCO3
(2 elements)
(non-metal oxide + water  acid)
(metal oxide + water  base)
(acidic oxide + basic oxide  salt)
Decomposition is where one substance breaks down to form 2 or more simpler substances.
The general formula is AB  A + B.
Examples
2 HgO  2 Hg + O2
Ca(OH)2  CaO + H2O
H2CO3  CO2 + H2O
ZnSO4  ZnO + SO3
KClO3  2KCl + 3O2
(Simple compound breaks down into its two elements)
(Hydroxides – break down into a metal oxide and water)
(Acids – break down into water and a non-metal oxide)
(Salts – break down into a metal oxide and a non-metal oxide)
(Chlorates – break down into a chloride and oxygen)
Combustion is when a compound or element reacts with oxygen. For hydrocarbons (organic
compounds), they ALWAYS form carbon dioxide and water. This is also known as burning!
Examples
C3H8 + 5 O2  3 CO2 + 4 H2O
OR
2 Mg + O2  2 MgO
Single Replacement is when one element replaces another element of the same charge sign
that is in a compound. The general formula is A + BC  AC + B for metal replacement
or D + EF  ED + F for halogen replacement.
Examples
Li + CuCl  LiCl + Cu
OR
2 KI + Br2  2 KBr + I2
NOTE: a single replacement will only occur if the element that is doing the replacing is more
active than the replaced element. We can use the activity series table to see which elements
replace others.
Double Replacement is where two compounds switch partners.
The general formula is AB + CD  AD + CB
Example
2 Al(NO3)3 (aq) + 3 Ca(OH)2 (aq)  2 Al(OH)3 (s) + 3 Ca(NO3)2 (aq)
H2SO4 (aq) + 2 NaOH (aq)  Na 2SO4 (aq) + 2 HOH (l) (note: HOH = H20)
NOTE: a double replacement will only occur if one of the products is insoluble in water (a
precipitate), is water, is an acid or is a base. We can use the Solubility Rules table to predict
which double replacement reactions will or will not occur.
Chemistry 30
To Do: List what type the following chemical reactions are using the five types of reactions
from the previous page.
Type of Chemical Reaction
1. 4 Fe + 3 O2  2 Fe2O3
______________________________
2. ZnCl2  Zn + Cl2
______________________________
3. Cu + AgNO3  Cu(NO3) 2 + Ag
______________________________
4. CaCO3 + HCl  CaCl2 + H2CO3
______________________________
5. CH4 + 2 O2  CO2 + 2 H2O
______________________________
Writing Balanced Equations
a)
b)
c)
d)
Represent the reactants and products by correct formulas
Check for diatomic elements: H2, O2, F2, Br2, I2, N2, Cl2 as well as P4 and S8
Balance the equation using coefficients (a large number in front of the element or
compound….do not change the subscripts!!!). The coefficient applies to the whole
molecule, not just the element in front! (ex. 2 NaCl means 2 Na AND 2 Cl)
Balancing equations allows for the law of conservation of mass. (Think of the arrow
as an equal sign, everything to the left must equal everything to the right!)
To Do: Balance the following chemical reactions and identify the type of reaction it is.
Balance
Type
1. ___ Na + ___ O2  ___ Na2O
____________________________
2. ___ Hg + ___ O2  ___ HgO
____________________________
3. ___ Ag2O  ___ Ag + ___ O2
__________________________________________
4. ___ C4H8 + ___ O2  ___ CO2 + ___ H2O
____________________________
5. ___ Fe + ___ Cl2  ___ FeCl3
__________________________________________
6. ___ Al + ___ HCl  ___ AlCl3 + ___ H2
__________________________________________
7. ___ F2 + ___ Al2O3  ___ AlF3 + ___ O2
__________________________________________
8. ___ Na2S + ___ FeBr3  ___ NaBr + ___ Fe2S3
____________________________
9. ___ Na2O + ___ H2O  ___ NaOH
____________________________
10. ___ Mg(OH)2  ___ MgO + ___ H2O
____________________________
Chemistry 30
Significant Figures
Measured quantities are generally reported in such a way that only the last digit is uncertain. All digits
including the uncertain one are called significant figures. The number of significant figures indicates
the exactness of a measurement.
All place-holding zeros are NOT significant!
The following guidelines apply to determining the number of significant figures in a measured
quantity:
a) Do not count LEADING zeros on a decimal number.
Ex) 0.000201  3 sig. figs. 2.01  3 sig. figs.
0.00300  3 sig. figs.
b) Do not count ENDING zeros on a number without a decimal.
Ex) 20010  4 sig. figs.
1000  1 sig. fig.
543  3 sig. figs.
NOTE: To indicate the actual number of sig figs, we can use scientific notation:
Ex: if you had the number 130 g but it should have 3 sig figs, you could change it to 1.30 x 102 g
To Do: Count the number of Significant Figures in the following measurements.
1) 0.35 g
_______
4) 0.00250 mL _______
2) 3500 mg
_______
3) 0.0035 mol
_______
5) 3501 L
_______
6) 0.000009 moles
_______
Calculations Involving Significant Figures
a) In multiplication and division the result must be reported with the same number of
significant figures as the measurement with the fewest significant figures. When your result
contains more than the correct number, it must be rounded off.
Ex) 6.221cm x 5.2cm = 32.3492cm2 - the least significant number is 5.2cm with
2 sig figs. Our answer should therefore be 32cm2
To Do: For each of the following, calculate and state the answer in the proper number of significant
figures.
1) 15.2 × 3.5 =
3) 8500 ÷ 9 =
2) 0.0025 x 0.03 =
4) 15.3528976 ÷ 2.5 =
b) In addition and subtraction, the result cannot have more digits to the right of the decimal
point than any of the original numbers.
Example: 20. 4
1. 322
83
104. 722
 rounds off to 105 because of 83 having no decimal places.
To Do: For each of the following, calculate and state the answer in the proper number of
significant figures.
1) 129.9 + 2.345 + 1.2 =
3) 67.009 – 45.98 =
2) 23.9 + 12 + 45.500 =
4) 10.00 – 9.9867 =
Chemistry 30
Mole Concept (1 mol = 6.02 × 1023 particles)  units: particles/mol
The mole is the standard unit for measuring the amount of a substance. It is a name for a very
big number: 602,000,000,000,000,000,000,000 (or 6.02 × 1023). It is equal to 6.02 × 1023 atoms
of an element or molecules of a compound (Avogadro’s number). Just like we like to count eggs
by the dozen (12), we like to count particles by the mole (6.02 × 1023).
Molar Mass (1 mol = ? grams)  units: g/mol
Molar mass is the mass of one mole of an element or compound. In chemistry, the unit is g/mol.
The atomic mass listed for the elements on the periodic table can be interpreted in two ways: first
as the mass of a single average atom of the element in atomic mass units (amu), or secondly as
the mass of one mole of the element in grams. For example, the mass of one atom of carbon is
6amu and the mass of 1 mole of carbon is 6 g.
The molar mass of a compound is the sum of the molar masses of the elements that make up that
compound. For example, 1 mole of carbon dioxide, CO2, has a molar mass of:
1 mole C + 2 mole O = 12.01 g/mol + 2(16.00 g/mol) = 44.01 g/mol
Molar Volume (1 mol = 22.4L)  units: L/mol
At standard temperature and pressure, the volume of 1 mole of ANY substance in the gaseous
state will always be 22.4 L.
Mole Conversion Guidelines
Often in chemistry, we need to change moles into other units, or other units into moles to help
facilitate chemistry calculations. Using the mole triangle/bow-tie can help you do these
calculations.
Mass (g)
Moles
# of particles
or molecules
Volume of
a gas (L)
Concentration/
Molarity (mol/L)
Chemistry 30
When doing mole conversions, include your units throughout the whole calculation to ensure
that your final answer has the correct units!!! (this is called dimensional analysis). You should
also round your answers to the correct significant figures during all chemistry calculations.
Sample Problems Involving the Mole
a) Converting from Moles to Mass (g) – What is the mass of 2.00 moles of oxygen?
 Moles × Molar Mass = Mass

b) Converting from Mass (g) to Moles – How many moles are in 25.0 g of carbon?
 Grams ÷ Molar Mass = Moles

c) Converting from Moles to Volume (L) – What is the volume of 2.00 moles of lithium?
 Moles × Molar Volume = Volume (L)

d) Converting from Volume (L) to Moles – How many moles are in 44.8 L of CO2?
 Volume ÷ Molar Volume = Moles

e) Converting from Moles to # of Particles – How many particles are in 2.00 moles of
NaCl?
 Moles × Avagadro’s Number = # of Particles

f) Converting from # of particles to Moles – How many moles are in 1.81 × 1024
molecules of NaCl?
 # of Particles ÷ Avagadro’s Number = Moles

Chemistry 30
Ham Sandwich Stoichiometry
Stoichiometry: the study of relationships between the amounts of reactants used and products
formed by a chemical reaction; is based on the Law of Conservation of Mass.
To make the perfect ham sandwich, we would need:
2 slices bread + 1 slice ham + 3 pieces lettuce  1 ham sandwich
Mol  Mol
How do you find the moles of one substance given the moles of another?
Moles known × Moles unknown = Calculated moles unknown
Moles known
Given
ratio from coefficients in the balanced equation
Ex) If you had 6.00 moles of lettuce, how many moles of ham could you use?
Mol  Mol  Grams
What happens if you know that 1 ham sandwich has a mass of 15.0 g/mol.
Moles known × Moles unknown × ? grams unknown = Calculated mass unknown
Moles known
1 mol
Given
ratio from balanced equation
Molar mass
Ex) If you had 4.00 moles of bread, determine the mass of ham sandwiches you could make.
Mass  Mol  Mol  Mass
What happens if you only know the masses of the reactants and products?
bread = 5.0 g/mol ham = 2.0 g/mol lettuce = 1.0 g/mol ham sandwich = 15.0 g/mol
Mass known × 1 mol known × Moles unknown × ? grams unknown = Calculated mass unknown
? grams
Moles known
1 mol
Given
Molar mass ratio from balanced equation
Molar mass
Ex) If you had 10.0 g of bread, determine the mass of ham sandwiches you could make.
Chemistry 30
Stoichiometry Worksheet
Mol known  Mole unknown
1. Hydrogen and nitrogen react to produce ammonia and energy. Balance.
___H2 (g) + ____N2 (g) ____ NH3 (g) + energy
a) How many moles of NH3 (g) are formed when 3.5 mol of H2 (g) are reacted?
b) How many moles of NH3 (g) are formed when 19.6 mol of N2 (g) are
reacted?
c) How many moles of H2 (g) will completely react with 1.2 mol of N2 (g)?
Mol known  Mol unknown  Mass unknown
2. A solution of potassium chromate reacts with a solution of lead (II) nitrate to
produce a yellow precipitate of lead (II) chromate and a solution of
potassium nitrate.
K2CrO4 (aq) + Pb(NO3)2 (aq)  PbCrO4 (s) + 2 KNO3 (aq)
a) Determine the mass of PbCrO4 that can be obtained if you start with
0.250 mol of K2CrO4.
b) Determine the mass of KNO3 that can be obtained if you start with 0.500
mol of K2CrO4.
Mass known  Mol known  Mol unknown  Mass unknown
3. Suppose iron (II) sulfide is treated with enough hydrochloric acid to
complete the following reaction.
FeS (s) + 2HCl (aq)  H2S (g) + FeCl2 (aq)
a)
b)
If 10.0 g of iron (II) sulfide reacts, how many grams of H2S could be
produced?
If 15.0 g of FeCl2 was formed, what mass of hydrochloric acid was
used?