Handouts
Grade 11 University Level
Unit 1
Nomenclature
Handout 1 – Grade 11 Nomenclature
Handout 2 – Recognizing Patterns 1 and 2
Handout 3 – Combining Patterns 1 and 2
Types of Reactions
Handout 4 – Types of Reactions
Handout 5 – Net Ionic Equations
Atomic Structure
Handout 6 – History of the Atom (Powerpoint)
Handout 7 – The Electromagnetic Spectrum
Handout 8 – The Emission Spectrum of Hydrogen
Handout 9 – Theories of the Atom
Handout 10 – Atomic Orbitals
Handout 11 – Aufbau Diagram
Unit 2
Periodicity
Handout 12 – Atomic Radii
Handout 13 – 1st Ionization Energy
Handout 14 – Successive Ionization Energies
Handout 15 – Electron Affinity/Eletronegativty
Bonding
Handout 16 – Covalent Compounds
Handout 17 – Ionic to Covalent to Metallic Bonding
Handout 1 – Grade 11 Nomenclature
Hydrated Salts
Some solids are crystals that regularly associate with water
- SiO2 placed in shoes to absorb water to protect the leather
- when these compounds are associated with H2O we call them hydrated
- when water is removed we call them anhydrous
-
to name hydrate compounds we use a prefix (same Greek prefixes from
molecular
compound naming) followed by hydrate to indicate the number of H2O molecules
associated to each formula unit.
- i.e. MgSO4∙7H2O is the chemical formula for magnesium sulfate heptahydrate
- could also be called hydrated magnesium sulfate but the above name is better because
indicates number of water molecules.
- CuSO4∙5H2O could be copper (II) sulfate pentahydrate or cupric sulfate pentahydrate
Acids
There are two kinds of acids that you will learn to name Binary acids and Ternary Acids
Binary Acids
-
the term binary indicates that a compound has only two types of atoms
all acids must contain H. Therefore, there is only one other atom that can vary
for each binary acid
Binary acids commonly contain elements from the halogen group, but other
examples include sulfur and selenium.
HCl(aq) is a very common binary acid you have used. Remember all acids are
aqueous because they dissolve in water
HCl(aq) is called aqueous hydrogen chloride by the more modern IUPAC
system, or more commonly, hydrochloric acid by the classical system.
Two naming systems for acids
IUPAC (modern) very simple just put aqueous in front of regular chemical
name.
Classical system follows this general formula
hydro ________ic acid
The blank is filled in with the associated anion name after
removing the “ide”
-
an unusual acid is HCN(aq), which is named using binary acid rules
Chemical
Formula
HCl(aq)
HF(aq)
H2S(aq)
H2Se(aq)
HCN(aq)
IUPAC Nomenclature
Classical Nomenclature
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Ternary Acids
-
-
Ternary means 3, therefore this is an acid containing three types of atoms
After the H the rest of the atoms in the ternary acid are polyatomic ions that
contain oxygen, or also called oxyanions. Therefore, ternary acids can also be
called oxyacids.
HNO2(aq) is called - aqueous hydrogen nitrite by the IUPAC system
-nitrous acid by the Classical system
Conversion from IUPAC to Classical for ternary acids
1) Replace the words “aqueous hydrogen” with the word “acid” at the
end
2) Change the ending: - “ate” to “ic” or “ite” to “ous”
IUPAC
aqueous hydrogen hypo______ite
aqueous hydrogen
______ite
aqueous hydrogen
______ate
aqueous hydrogen per ______ate




Classical
hypo_______ous acid
_______ous acid
_______ic acid
per _______ic acid
Chemical
Formula
IUPAC Nomenclature
Classical Nomenclature
HNO2(aq)
HBrO2(aq)
H3PO4(aq)
H2CO3(aq)
H2SO4(aq)
HClO4(aq)
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_________________________________________________
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-
remember only do classical system rules for binary and ternary acids when
you know it is an acid. For example “(aq)” symbols are a good indication the
compound is an acid if it also contains an H at the beginning. There are some
exceptions such as acetic acid CH3COOH, the last H being the acidic proton.
-
example HBrO2(g) would be called hydrogen bromite gas
Handout 2 - Recognizing Patterns #1
Oxyanions (polyatomic ions containing oxygen) have a pattern in their names to indicate
amount of oxygens. Look at the list of oxyanions that have chlorine in them and see if
you notice the pattern.
ClOClO2ClO3ClO4-
hypochlorite
chlorite
chlorate
perchlorate
Ion charge is (-1), but oxygens are
increasing by 1
- the base ion is the one with “ate” and no prefix
- when the suffix “ite” is used, subtract 1 oxygen atom from the base ion
- when the prefix “hypo” and the suffix “ite” is used, subtract 2 oxygen atoms
- when the prefix “per” and the suffix “ate” is used, add 1 oxygen atom to the
base ion
Use your polyatomic ions list to find the formulas for the following “base” ions:
carbonate, nitrate, phosphate, sulfate, iodate, bromate.
Use the 6 base polyatomic ions above and fill in the boxes below with the 3 other
polyatomic ions that can be known from the base ion.
Name
Formula
Name
Formula
Name
Formula
Name
Formula
Name
Formula
Name
Formula
Recognizing Patterns #2
Acid Anions
- an acid anion is created when one or more H+ ions covalently bond with an
oxyanion (i.e. HCO3-, HPO42-)
-
when acid anions combine with cations, acid salts are created (i.e. CaHPO4)
- using the base polyatomic ions from Recognizing Patterns #1 (carbonate,
phosphate, sulfate) and the pattern below you can create the acid anions
Base Ion
Acid Anion
+H+
carbonate
CO32-
-2 + 1 = -1
hydrogen carbonate
HCO3-
+2H+
phosphate
PO43-
-3 + 2 = -1
dihydrogen phosphate
______
______
sulfate
SO42-
________
hydrogen sulfate
______
______
phosphate
PO43-
________
hydrogen phosphate
______
Handout 3 - Combining Patterns 1 and 2
phosphate
PO43(Pattern 1)
+2H+
Phosphite
PO33-
(Pattern 2)
________
dihydrogen phosphite
________
Sulfate
SO42(Pattern 1)
+H+
Sulfite
SO32-
(Pattern 2)
________
___________________
________
Practicse Questions
1. Oxyanion Pattern: Fill in the table below using the patterns for oxyanions
Chemical Formula
Chemical Name
Calcium hypochlorite
Zn(BrO4)2
Barium phosphate
Aurous nitrite
Mg(IO)2
SnSO3
Lithium persulfate
Iron (III) percarbonate
or
2. Oxyanion Pattern: Fill in the table below using the patterns for oxyanions
Chemical Formula
Sr(HCO3)2
Cu(H2PO3)2
Chemical Name
Sodium hydrogen sulfate
or
Aluminum dihydrogen phosphate
Rb2HPO4
Gold (I) hydrogen sulfite
Handout 4 - Types of Reactions
1) Synthesis Reactions (A + B  AB)
I)
Simple Binary Ionic Compounds
i.e. solid aluminum reacts with chlorine gas
Al(s) + Cl2(g)  AlCl3(s)
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II)
Slightly More complicated Synthesis Reactions
Non-metal oxides
(acidic oxides)
i.e. CO2(g) , SO3(g)
Acid
i.e.H2CO3(aq) ,
H2SO4(aq)
Salts containing
oxyanions
i.e. Li2CO3(s)
H2O
Base
i.e.LiOH(aq),
Ca(OH)2(aq)
Metal oxides
(basic oxides)
i.e. Li2O(s) , CaO(s)
- non-metal oxides such as CO2, SO3, N2O5 react with H2O to form acids
i.e. CO2(g) + H2O(l)  H2CO3(aq)
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___________________________________________
- metal oxides such as Li2O, CaO react with H2O to form bases
i.e. CaO(s) + H2O(l)  Ca(OH)2(aq)
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- non-metal oxides and metal oxides can react to form salts containing oxyanions
i.e. CaO(s) + CO2(g)  CaCO3(s)
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2) Decomposition Reactions (AB  A + B)
-reverse of the above reactions
-often heat is needed; this is called “thermal decomposition”
I)
Simple Binary Ionic Compounds
i.e. aluminum chloride is heated
II)
Slightly More complicated Reactions
- acids will decompose into non-metal oxide and water
i.e. H2CO3(aq)  CO2(g) + H2O(l)
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- bases will decompose into metal oxides and water
i.e. Ca(OH)2(aq) 
CaO(s) + H2O(l)
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- salts containing oxyanions decompose into non-metal oxides and metal oxides
i.e. CaCO3(s)  CaO(s) + CO2(g)
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3) Single Displacement Reactions (AX + B  A + BX)
Create your own activity series mini-lab
Hypothesis: Predict the order of reactivity of the 5 metals in the lab from most to
least reactive. (Hint: use their position on the periodic table)
Most _____ _____ _____ _____ _____ Least
When complete show your teacher the order and obtain the materials
for the lab.
Observation Chart:
Metals
Iron
Magnesium
Copper
Zinc
Calcium
Solutions
Iron nitrate
Magnesium nitrate
Copper nitrate
Zinc nitrate
Calcium nitrate
Indicate a reaction with a checkmark and no reaction with an X.
Conclusion: Using your observation chart order the metals from most to least
reactive
Most _____ _____ _____ _____ _____ Least
From your observations write the products of the following reactions that would
react and if there is no reaction indicate no reaction.
Ca(s) + Mg(NO3)2(aq) 
Cu(s) + Zn(NO3)2(aq)

Mg(s) + Cu(NO3)2(aq) 
Activity Series: is an arrangement of metals in order of their relative reactivities.
Knowing the order allows you to predict if a single displacement reaction will
take place or not. Any metal higher on the list can displace any metal lower on
the list.
Metal
Displaces hydrogen Displaces hydrogen
from acids
from cold water
Lithium
Potassium
Barium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Cobalt
Nickel
Tin
Lead
Hydrogen
Copper
Mercury
Silver
Platinum
Gold
Most Reactive
Least Reactive
Using the activity series which of the following reactions will occur?
a) Au(s) + CuSO4(aq) 
Co(s) + HgClO2(aq) 
Na(s) + Sn(IO3)2(aq) 
b) How could the metal activity series be used to predict reactions with acids?
Give 2 examples of reactions between acids and metals that would occur.
What gas is produced?
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c) How is the metal activity series used to predict reactions with water? Give 2
examples of reactions with metal and water. What gas is produced?
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d) There is also a Halogen Activity Series. Give an example of 2 reactions that
could be predicted by the Halogen Activity Series below.
Halogen Series: Most _____
_____
_____
_____ Least
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4) Double Displacement (AX + BY  AY + BX)
- always occur between two soluble ionic compounds
- there are three possible outcomes a) precipitate forms
b) gas is produced
c) water is produced
a) Precipitate Forms
- know how to use the solubility chart below to identify if a solid is produced
Rule
Nitrates (NO3-) are soluble
Halides (Cl-, Br-, I-) are soluble
Sulfates (SO42-) are soluble
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2-
Sulfides (S ) are insoluble
Carbonates (CO32-) are insoluble
Phosphates (PO43-) are insoluble
Hydroxides (OH-) are insoluble
Exception
None
Ag+, Hg22+, Pb2+
Ca2+, Ba2+, Pb2+, Hg22+, Ag+
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NH4+ and ions of groups 1 and 2 elements
NH4+ and ions of group 1 elements
NH4+ and ions of group 1 elements
2+
2+
2+
Ba , Sr , Ca and ions of group 1 elements
Use the solubility chart above to identify if the following reactions will produce a
precipitate or not. If they do write the products and indicate the precipitate by
using a subscript (s).
K2CO3(aq) + CaCl2(aq) 
Pb(NO3)2(aq) + KI(aq) 
b) Gas is Produced
- there are four cases in which a gas is formed. The first 3 cases are a double
displacement reaction followed by a decomposition.
i)
ii)
iii)
acids + carbonates 
acids + sulfites 
bases + ammonium 
- the double displacement reactions produce products such as carbonic acid,
sulfurous acid and ammonium hydroxide, which then decompose into gas
and water. Try completing the following reactions.


ii) Na2SO3(aq) + HCl(aq) 

iii) NH4Cl(aq) + NaOH(aq) 

i) CaCO3(s) + HCl(aq)
Come up with your own examples of the three kinds of reactions above
i)_____________________________________________________________
ii)_____________________________________________________________
iii)____________________________________________________________
- the 4th case of a reaction that produces a gas involves acids and sulfides.
This case only requires the double displacement, the gas is produced
immediately, H2S(g).
iv) Na2S(aq) + HCl(aq)

iv)____________________________________________________________
c) Water is Produced
- these double displacement reactions are more specifically named
neutralization reactions. They occur when acids are combined with bases
and the products are water and salt. Most times the salt can be labeled
aqueous, but make sure by checking the solubility table. Look at the
following reactions and write the products and their states.
H3PO4(aq) + NaOH(aq) 
H2CO3(aq) + CaOH(aq) 
Handout 5 – Net Ionic Equations
Writing net ionic equations
1. Write balanced equation first
2. To do your total ionic equation
a) Break up all ionic compounds EXCEPT precipitates
b) DO NOT break up molecular compounds EXCEPT acids
3. To get net ionic equation cross out all spectator ions. Because they do not really
take part in the reaction.
Example 1
Write the balanced net ionic equation for the reaction of aqueous sodium carbonate
with aqueous calcium nitrate.
Balanced Chemical Equation
Na2CO3(aq)
+
Ca(NO3)2(aq)

2NaNO3(aq)
+
CaCO3(s)
Total Ionic Equation (break into ions)
2Na+(aq) + CO32-(aq) +Ca2+(aq) + 2NO3-(aq)  2Na+(aq) +2NO3-(aq) + CaCO3(s)
- total charges should be the same on both sides of equation
Cross out spectator ions
2Na+(aq) + CO32-(aq) +Ca2+(aq) + 2NO3-(aq)  2Na+(aq) + 2NO3-(aq) + CaCO3(s)
Net Ionic Equations
CO32-(aq) +
Ca2+(aq)

CaCO3(s)
Balanced Chemical Equation
2NaI(aq)
+

Br2(aq)
2NaBr(aq)
+
I2(g)
Total Ionic Equation (break into ions)
2Na+(aq) + 2I-(aq) + Br2(aq) 
2Na+(aq)
+ 2Br-(aq) + I2(g)
- total charges should be the same on both sides of equation
Cross out spectator ions
2Na+(aq) + 2I-(aq) + Br2(aq) 
2Na+(aq)
Net Ionic Equations
2I-(aq) + Br2(aq) 
2Br-(aq) + I2(g)
+ 2Br-(aq) + I2(g)
Handout 6 – History of the Atom
Handout 7 – The Electromagnetic Spectrum
Handout 8 – The Emission Spectrum of Hydrogen
Handout 9 – Theories of the Atom
Handout 10 – Atomic Orbitals
The first 3 S - orbitals
Shapes of the s, p, and d orbitals
The s and p orbitals around a single atom
Handout 11 – Aufbau Diagram
Handout 12 – Atomic Radii
Handout 13 – 1st Ionization Energy
Handout 14 – Successive Ionization Energies
Handout 15 – Electron Affinity
Handout 16 - Covalent Compounds
Remember:
- an ionic compound has a metal + non-metal
- a molecular compound only has non-metals
Hydrogen - neither a metal nor a non-metal
- when we draw the Lewis/Electron dot diagrams for
compounds containing H, we consider the bond covalent.
- when we name the compound containing H, we use ionic
nomenclature rules
- a compound only needs one ionic bond to be classified as ionic (even if
there are many covalent bonds)
Calculating Polarity of Covalent Bonds
To this point we have described bonds as either ionic or covalent. Due to
electronegativity we can now classify the covalent bonds as non-polar, slightly polar,
polar and very polar.
Step #1 – We need to consider the dipoles of each bond in the molecule by examining the
differences in electronegativities.
Electronegativity
Difference
0
0.5 or less
0.6 – 1
1.1 – 1.6
Greater then 1.7
Polarity
Non-polar, ie/ diatomic molecules: H2, O2, F2
Slightly polar
Polar
Very polar
Usually indicates an ionic bond
Step #2 – Dipoles act as forces that pull electrons toward the more electronegative atom
- Dipoles that are equal in magnitude but opposite in direction will cancel out
Step #3 –Then add labels to your electron dot/structural diagrams to show that the atom
with a greater electronegativity pulls on the electrons better and thus has a partial
negative charge, δ-, the atom that is least electronegativity has a partial positive charge,
δ+.
NOTE: The general rule for writing a chemical formula is to write the most
electronegative atom last. Therefore water’s chemical formula is H2O and not OH2.
Handout 17 – Ionic to Covalent to Metallic Bonding