Science 10: Chemistry Handbook part 1
Name: ___________________________________________
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Block: _______________
Chemistry 10 2021 Information
Here are the major course topics for the assessment:
CHEM A Chemical Reactions
C.1 Explain differences, name, and determine formulae for ionic and covalent compounds.
C.2 Construct diagrammatic models to illustrate the structure of atoms and molecules.
C.3
Classify reactions as being one of the six main types and predict the products of a reaction based on
its type.
CHEM B Acid and Bases
C.4 Name, classify, and determine the formula for various acids and bases.
CHEM C Mass conservation in Chemical Reactions
Balance chemical equations (including compounds as formulae or names) using appropriate state
C.5
symbols and notations.
CHEM D Energy changes in Chemical Reactions
C.6 Justify whether a chemical reaction is endothermic or exothermic by evaluating the evidence.
Topic 2.2 What happens to atoms in a chemical reaction?
 Atoms bond together to form ionic and covalent compounds.
 Bonds are broken, atoms are rearranged, and new bonds are formed.
 Mass cannot be created or destroyed in a chemical reaction.
 A chemical equation represents what happens to the atoms in a reaction.
Topic 2.3 How is energy involved in chemical processes?
 Matter and energy interact in physical and chemical changes.
 Energy is transferred between chemical reactions (the system) and the surroundings.
Topic 2.4 How do atoms rearrange in different types of chemical reactions?
 A compound forms in a synthesis reaction and breaks down in a decomposition reaction.
 In replacement reactions, elements replace other elements.
 Most combustion reactions release heat and light.
 In a neutralization reaction, an acid reacts with a base.
Textbook and Workbook Sections that may be helpful:
1. What happens to atoms in a chemical reaction?
2. How is energy involved in chemical process? Section 2.3
3. How do atoms rearrange in different types of chemical reactions? Section 2.4
4. Workbook pages 73-129
5. Textbook pages
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CHEM A Chemical Reactions
C.1 Explain differences, name, and determine formulae for ionic and covalent compounds.
 Name ionic and covalent compounds
- including elements that are multivalent
- including polyatomic ions
- go from name to formula and vice versa
C.2 Construct diagrammatic models to illustrate the structure of atoms and molecules.
 Draw Bohr diagrams and Lewis structures for atoms and compounds
 Show using Lewis structures how atoms are reacting in a chemical reaction
Classify reactions as being one of the six main types and predict the products of a reaction based on
its type.
Identify if a chemical reaction is synthesis, decomposition, single replacement, double replacement,
combustion, or neutralization reaction
Predict the products of a chemical reaction when given the reactants
C.3


CHEM B Acid and Bases
C.4 Name, classify, and determine the formula for various acids and bases.





Determine the chemical formula for acids and bases
Understand and differentiate between an acid and a base
Determine the name of a an acid or base when provided the chemical formula
pH scale and examples
Formation from metal or non-metal oxides
CHEM C Mass conservation in Chemical Reactions
Balance chemical equations (including compounds as formulae or names) using appropriate state
C.5
symbols and notations.
 Balance chemical questions
 Understand the Conservation of mass
 Denote whether a substance is solid, liquid, aqueous, or gas
CHEM D Energy changes in Chemical Reactions
C.6 Justify whether a chemical reaction is endothermic or exothermic by evaluating the evidence.





Predict whether a reaction is exothermic or endothermic
Understand activation energy
System vs Surroundings
Conservation of Energy
Energy-level graphs
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Vocabulary for Chemistry
activation energy: the minimum amount of energy needed for a reaction to occur (2.3)
endothermic reaction: a chemical reaction in which there is net absorption of energy from the surroundings.
(2.3)
exothermic reaction: a chemical reaction in which there is net release of energy to the surroundings. (2.3)
law of conservation of energy: the law that states that the total energy of the universe if constant; law stating
that energy cannot be created or destroyed, but is transformed from one form of energy to another or transferred
from one object to another. (2.3, 3.1)
surroundings: everything else in the universe outside of the system. (2.3, 3.1)
system: anything that is under observation; for example, a chemical reaction. (2.3, 3.1)
anion: a negatively charged ions (2.2)
balanced chemical equation: a complete description of a chemical reaction that provides the chemical formulas
for the reactants and products and the coefficients. (2.2)
cation: a positively charged ion. (2.2)
chemical equation: a representation of a chemical reaction using words or chemical formulas. (2.2)
coefficient: a number placed in front of a chemical formula in a balanced chemical equation to show the ratios
of substances in reaction. (2.2)
covalent bond: a strong attraction between atoms that forms when atoms share valence electrons. (2.2)
covalent compound: a compound that results when atoms of two or more elements bond covalently. (2.2)
electrostatic attraction: the attraction between cations and anions that form the ionic bond. (2.2)
formula unit: the chemical formula for an ionic compound that represents smallest repeating part of the crystal
lattice. (2.2)
iconic bond: a strong attraction that forms between oppositely charged ions. (2.2)
iconic compound: a compound made of oppositely charged ions. (2.2)
law of conservation of mass: in a chemical reaction, the total mass of the substance used is equal to the total
mass of the substances produced. (2.2)
molecule: a particle made up of two or more atoms bonded by covalent bonds. (2.2)
product: any new substance that is formed from a chemical reaction. (2.2)
reactant: a substance that undergoes a chemical change. (2.2)
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skeleton equation: a description of a chemical reaction that provides the chemical formulas for the reactants and
products. (2.2)
word equation: words that describe what happens to reactants and products in a chemical reaction. (2.2)
acid: compound that forms H+ ions when dissolved in water (2.4)
acid-base indicator: a chemical that changes colour in response to the concentration of hydrogen ions in a
solution (2.4)
base: a compound that forms OH- ions when dissolved in water. (2.4)
combustion reaction: a chemical reaction in which an element reacts with oxygen to produce an oxide of the
element and heat; also refers to the burning of hydrocarbons to produce carbon dioxide and water. (2.4)
decomposition reaction: a chemical reaction in which a compound is broken down into two or more elements or
simpler compounds. (2.4)
double replacement reaction: a chemical reaction in which solutions of two ionic compounds react to produce
two other ionic compounds. (2.4)
neutralization reaction: a chemical reaction in which an acid reacts with a base to form salt and water. (2.4)
pH scale: a numbered scaled between 0 and 14 that indicates the acidity or basicity of a solution. (2.4)
single replacement reaction: a chemical reaction in which an element and a compound react to produce another
element and another compound. (2.4)
synthesis reaction: a chemical reaction in which two or more reactants combine to produce a single product.
(2.4)
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Class 1: Safety and Intro
Demo 1: Sketch and describe what happened
Demo 2: Sketch and describe what happened
After the day 1 chemistry demo is done, consider what you observed, and record 3 questions that you have
about what you saw.
1.
2.
3.
Then, choose 1 of them and turn it into a question that could be answered with an experiment. Talk in general
terms (a few sentences) about what you might have to do to start answering that question. What would you
physically to do with the equipment?
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Class 2: Review Chemistry 8/9
1. Atoms: A historical development
John Dalton – solid sphere
JJ Thompson – plum pudding
Rutherford – nuclear model
Niels Bohr – planetary model
2. Subatomic particles: In the last two years, we have looked at three subatomic particles and their properties.
Complete this table:
Particle
Location
Mass
Charge
Proton
Neutron
Electron
What units are we using to measure the mass and charge? Why do you think this might be the case?
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3. Different elements are simply different combinations and arrangements of subatomic particles. What
information can we get from each element’s entry on the Periodic Table? Use sodium as an example.
4. Use a periodic table (inside back cover of your book) to complete this chart:
Hints: atoms not ions, round off the mass to the nearest whole number
Name of the
element
Symbol
Magnesium
Atomic
Number
Atomic
Mass
Number of
protons in an
atom (not ion)
Number of
neutrons in an
atom (not ion)
Number of
electrons in an
atom (not ion)
12
Nb
41
Argon
18
U
238
6
11
Pb
207
100
157
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5. Summarize:
In a neutral atom, the number of electrons is always equal to the number of ____________ or the
____________ number.
The atomic mass is equal to the number of ____________ plus the number of ____________.
6. Electrons in atoms
According to the model of the atom proposed by Neils Bohr, electrons move around the atom in orbits or
orbitals. When one orbital is full, the electrons start filling the next one.
The first orbital holds _____ electrons, the second holds _____, the third holds _____, and the fourth holds
_____.
Give the total number of electrons in the following atoms (not ions) and the number of electrons in each orbit:
Element
Total # of
electrons
Electrons in
level 1
Electrons in
level 2
Chlorine (Cl)
Carbon (C)
Neon (Ne)
Calcium (Ca)
Sodium (Na)
Aluminum
(Al)
Krypton (Kr)
Check in here before continuing
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Electrons in
level 3
Electrons in
level 4
7. Atoms are most stable when they have full valence (outer) shells of electrons. If the valence shell is close
to being empty, it will give away the electrons to another atom. If the valence shell is close to being full, it will
accept electrons from another atom to complete it.
On your table on the previous page, write gain next to one element likely to accept electrons and lose next to
one element likely to donate them.
8. Use the periodic table to complete this table. The first row is done for you.
Element
Nitrogen
Total # of
electrons
in an
atom
# of
valence
electrons
in an atom
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5
Ion charge # of
from
electrons
periodic
gained or lost
table
3-
Gains 3
Oxygen
Fluorine
Neon
Sodium
Magnesium
Aluminum
Phosphorus
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Total #
of
electrons
in the
ion
# of
valence
electrons
in the ion
10
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Class 3: Periodic trends & Isotopes
1. Here are three isotopes of an element:
12
6C
13
6C
14
6 C
a. The element is: __________________
b. The number 6 refers to the _________________________
c. The numbers 12, 13, and 14 refer to the ________________________
d. How many protons and neutrons are in the first isotope? _________________
e. How many protons and neutrons are in the second isotope? _________________
f. How many protons and neutrons are in the third isotope? _________________
2. Complete the following chart:
Isotope name
atomic #
mass #
# of protons
# of neutrons
# of electrons
Potassium-37
Oxygen-17
uranium-235
uranium-238
boron-10
boron-11
DIRECTIONS: For the following problems, show your work! Be thorough by showing each step of your
calculations.
3. Naturally occurring europium (Eu) consists of two isotopes with masses of 151 and 153. Europium-151 has
an abundance of 48.03% and Europium-153 has an abundance of 51.97%. What is the average atomic mass
of europium?
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4. Lithium-6 is 4% abundant and lithium-7 is 96% abundant. What is the average atomic mass of lithium?
5. Hydrogen is 99% 1H, 0.8% 2H, and 0.2% 3H. Calculate its average atomic mass.
6. What is the atomic mass of hafnium if, out of every 100 atoms, 5 have a mass of 176, 19 have a mass of
177, 27 have a mass of 178, 14 have a mass of 179, and 35 have a mass of 180.0?
7. Boron exists in two isotopes, boron-10 and boron-11. Based on the atomic mass listed on the Periodic
Table, which isotope is likely to be more abundant? Explain your reasoning. (optional math challenge:
calculate the % breakdown of boron).
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Class 4: Ionic & Covalent Compounds
Ionic Compounds:
When ionic compounds form between a metal and a non-metal, the overall charge is neutral. This is a
consequence of conservation of charge. You start with neutral atoms, electrons are transferred from the metal
to the non-metal, but none are gained or lost. The overall compound must still be neutral. This lets us
determine how many of each ion we need.
In naming compounds, name the metal ion first. Then name the non-metal ion by ending the element name
with the suffix –ide.
For example: Calcium + chlorine.
 the cation is calcium, it has a charge of 2+ (convention is to write the number first, then the sign)
 the anion is chlorine, it has a charge of 1 the name is calcium chloride
one calcium
 two chloride ions are needed to balance the charge of
ion so the formula is CaCl2
Practice: Determine the formulas and names of the following:
 Fluorine + Magnesium

Scandium + Oxygen

Sodium + Nitrogen
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Translate between names and formulas:
Magnesium bromide
Na3P
Potassium chloride
BeCl2
Aluminum sulfide
GaN
Calcium oxide
Mg3P2
We can also use the charges to determine the charge of a metal ion found in a compound:
For example: SnCl4
The chloride ion has a charge of 1- and there are four of them, for a total of 4-. Since there is only one tin ion,
its charge must be 4+.
Practice: Find the ion charge of each of the metals.
BiBr5
PoO2
PbI2
HgO
PdF4
Os2O3
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This process is necessary for writing the names of ionic compounds involving multivalent metals, or metals
that have more than one possible ion charge. We have to determine which one is involved.
Writing the formula from the name: Cobalt(III) oxide
 Cobalt(III) means we are using cobalt ions with a charge of 3+
 Oxide ions have a charge of 2 To balance charges, we need two Co3+ and three O2 The formula is Co2O3
Writing the name from the formula: Au3N
 The charge on the nitride ion is 3 Gold could have an ion charge of 1+ or 3+
 Since there are three of them, and the total charge must be 3- to balance the negative charges, we must
be using the Au+ ion
 The name is Gold(I) nitride
We use Roman Numerals to indicate the ion charge on multivalent metals. These must be included in the name
when naming the cation. Even if the ion charge is 1+, such as with Copper(I), they still need to be used for
clarity. Aside from the addition of the numerals, the naming rules for ionic compounds don’t change.
 Counting to 10: I, II, III, IV, V, VI, VII, VIII, IX X
Translate between names and formulas:
Copper(II) oxide
Mn2O3
Mercury(I) oxide
VBr4
Gold(III) chloride
Nb2O5
Thallium(III) bromide
TiO2
Bismuth(V) oxide
TiN
Check your answers now. If you are finding these difficult, it would be wise to practice more before we add
new rules to this for grade 10. Extra ionic naming practice is available on Canvas.
Covalent Compounds
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Covalent compounds are ones where electrons are shared. Electrons are not transferred, and ions are not
formed. The products are not as specifically determinable (as in, there aren’t set combinations based on
charges), so the information contained in names and formulas needs to be more specific.
For example, H2O2 (hydrogen peroxide) is very different than HO, and you never reduce to lowest terms like
you would for an ionic compound (eg with ionic - germanium oxide is GeO2, not Ge2O4). Subscripts in ionic
compounds show the smallest whole-number ratio of ions, but with covalent compounds, they show the actual
number of atoms of each element in the molecule.
Covalent compounds are generally more complex chemically, but it is actually simpler to translate between
names and formulas. We use prefixes to indicate exactly how many atoms of each element are present.
 Mono- = 1
 Hepta- = 7
 Di- = 2
 Octa- = 8
 Tri- = 3
 Nona- = 9
 Tetra- = 4
 Deca- = 10
 Penta- = 5
 Hexa- = 6
Name the first element first, then the second element changing its end to –ide, and add prefixes
to the start of each word to indicate how many of that type are present. NOTE: if there is only
one of the first atom, you don’t need to include mono-. If there is only one of the second atom,
you do include it.
Writing the formula from the name: dinitrogen trioxide
 Dinitrogen means two nitrogen atoms
 Trioxide means three oxygen atoms
 The formula is N2O3
Writing the name from the formula: CO
 There is one carbon atom: carbon (not monocarbon since it’s the first one)
 There is one oxygen atom: monoxide (take out the double o)
 The name is carbon monoxide
More examples:
CS2 = carbon disulfide
CCl4 = carbon tetrachloride
P4O10 = tetraphosphorus decaoxide
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Translate between names and formulas:
Oxygen dichloride
SO3
Carbon tetrabromide
OF2
Carbon dioxide
CI4
Nitrogen triiodide
N2O
Chlorine monofluoride
N2O4
Check your answers now. If you are finding these difficult, continue with more practice in the workbook on
page 78.
Summary:
Describe how to distinguish ionic compounds from covalent compounds by looking only at the compound’s
name. Be specific by mentioning the information that is only included in one type of name and not the other.
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Class 5: Yet More Practice
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Class 6: Reactivity of Metals Lab
Aim:
The purpose of this lab is to determine the relative reactivity of metals using single replacement reactions.
Background:
Single replacement reactions happen between a single element (a metal, in this case) and an ionic compound.
The single metal element might swap places with the metal in the compound. The general form of the reaction
is A + BX  B + AX where A and B are metals, and X is a non-metal.
Whether or not this happens depends on the chemical reactivity of the metals. Differences in chemical
reactivity are caused by how easily the metals give up electrons. You can measure the relative reactivity of two
metals by placing a pure sample of one metal (A in the general form above) in a solution containing ions of the
other metal (B). If metal A is more reactive than metal B, electrons will move from the solid metal sample into
the solution. For example, a piece of iron placed in a solution containing copper(II) ions will react, producing
copper particles on the surface of the iron.
Fe + CuCl2  Cu + FeCl2 (iron is more reactive than copper)
However, no reaction occurs when copper is put into a solution containing iron(II) ions
Cu + FeCl2  nothing happens (copper is less reactive than iron)
The purpose of this experiment is to rank a selection of metals from most reactive to least reactive.
Pre-lab Questions:
1.
a) You put a piece of lead into a copper(II) chloride solution. There is a colour change, meaning a
reaction happened. Write the equation for this reaction.
b) Based on these results, which metal (Pb or Cu) is more reactive?
2.
a) Suppose we have metals H, M, J, and W. You add them to chloride solutions of the same metals.
Using the following results, determine the relative reactivity of the metals.
Metal
H
H
H
J
J
M
Solution
JCl
MCl
WCl
MCl
WCl
WCl
Observations
Colour change
Bubbles
Bubbles
Colour change
No reaction
No reaction
Relative Reactivity
H>J
b) Based on the relative reactivities, rank the 4 metals from most reactive to least reactive.
_____ > _____ > _____ > _____
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Procedure:
You will need to conduct a series of tests to determine the activity series (ranking) of Cu, Mg, Zn, Pb, Ag, and
Na. You will have the following solid metals: Cu, Mg, and Zn. Make sure the copper is cleaned with steel
wool. You will have the following aqueous solutions: Pb(NO3)2, AgNO3, CuSO4, MgCl2, ZnSO4, and NaCl.
You will need to use a large well plate to set up and observe the possible reactions. About 5-10 drops of
solution work well in the plates. Be very specific with your observations – look for gas production (aka
bubbles), colour changes in the solution or on the solid metal, temperature changes, and/or or solid precipitates
forming. When you are done, dry off and return unreacted solid metals. Rinse the contents of your well plate in
the sink with LOTS of water.
Data Collection
1. Record your observations of each combination of metal + solution. If there is no reaction, write NR.
2. Determine the relative reactivity of the two metals like you did in the pre-lab questions.
3. Write the chemical equation for each reaction that occurred (not the NR ones!).
Metal
Solutio
n
Cu
Pb(NO3)
Cu
AgNO3
Cu
NaCl
Cu
Mg(NO3)2
Cu
ZnSO4
Mg
Pb(NO3)2
Mg
AgNO3
Mg
CuSO4
Mg
NaCl
Mg
ZnSO4
Zn
Pb(NO3)
Zn
AgNO3
Zn
CuSO4
Zn
Mg(NO3)2
Zn
NaCl
Observations
Relative
Reactivit
y
2
2
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Chemical Equation
Analysis + Conclusion:
1. Using your relative reactivities from above, rank the six metals (Cu, Mg, Zn, Pb, Ag, and Na) from most
reactive to least reactive. If you could not distinguish between the activity of two metals, put them together,
such as A > B > (C = D) > E.
2. Compare your results to theory – search online for the reactivity series of metals. Do your results match? If
no, what was different, and suggest a reason why.
3. Explain how the reactivity series helps us predict the products of chemical reactions.
Application:
1. Based on metal reactivity, why are coins made out of copper and nickel instead of iron?
2. Many plumbing pipes are made out of a plastic called PVC. Before the invention of PVC, most plumbing
pipes were made of copper instead of the less expensive option like iron. Explain why.
Outcome being graded: Processing and Analyzing Data
Evidence of Exceeding
Expectations
Meeting Expectations
(86%)
Student has compiled data
and used it to anticipate
possible patterns, diligently
search for trends, and
identify outliers.
Analyze patterns, trends,
PA2 and connections in data,
identifying inconsistencies
Final Grade:
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Evidence of not yet
meeting expectations
Class 7: Polyatomic Ions
Polyatomic Ions
◦
◦
◦
Covalently bonded molecules (non-metals, shared electrons) that happen to have a non-zero
charge (are ions)
In bonds, they behave just like simple, single-element ions.
Most are anions
◦ Ammonium is the main exception
Naming – indicates how many oxygens are involved
S2SO32SO42-
Sulfide
Sulfite
Sulfate
Chloride
Hypochlorite
Chlorite
Chlorate
Perchlorate
ClClOClO2ClO3ClO4-
Important: Doesn’t indicate an absolute number, but more/fewer oxygen atoms bonded
Charges on polyatomic ions
Find the charge of Cr in CrO42O is 2-, 4 of them is 8-. Overall ion charge 2-, so Cr must be 6+
Find the charge of P in PO43O is 2-, 4 of them is 8-. Overall ion charge 3-, so P must be 5+
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Ionic Bonds
Polyatomic ions behave just like monoatomic ions when forming bonds:
◦ Charges must balance – neutral compound formed
◦ Subscripts to indicate how many of each ion
The only difference in naming is that you don’t change the polyatomic ion’s name ending (no –ide)
What are the names of:
NaClO4
MgSO4
Na3PO4
Sc(OH)3
Explain the difference:
Ca(OH)2
vs.
CaOH2
Magnesium hydroxide
Sodium hypochlorite
Iron(III) Sulfate
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Class 8: Simplified Bohr Diagrams
Bohr Diagrams – Review and Simplifying
In an earlier class, we reviewed electron orbitals. Go back and look at pages 3+4 of the Grade 9 Review
booklet.
Today we are going to review Bohr diagrams and learn some simplifications of them.
Key points:
 Each shell can only contain a certain number of electrons: 2, 8, 8, 18, 18, 32
 Electrons fill in starting from the lowest shell that has space, and
start the next one once the previous one is full
 Placement of electrons within a shell – the numbers indicate the
which they are added
 (the model is mostly wrong – but we use it since it helps
understand the basics)
1. Draw Bohr diagrams for atoms of the following elements.
Lithium
Sodium
Magnesium
Calcium
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only
order in
Fluorine
Chlorine
Helium
Neon
Describe a pattern that you see. Hint: their placement on the periodic table might help answer this.
2. When atoms gain or lose electrons to become ions, they are looking to have complete valence (outer)
shells, whether that means losing some to empty the shell or gaining some to fill it up. They will always do
whatever is easiest. Which of the elements above will gain electrons and which will lose? How do you know?
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3. When we draw Bohr diagrams of ions, we add square brackets with the ion charge outside like an exponent.
Draw Bohr diagrams of ions of the same elements from #1. Sodium is done as an example.
Lithium
Sodium
Magnesium
Calcium
Fluorine
Chlorine
Helium
Neon
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4. You are probably tired of drawing circles. This year, we are going to simplify these to something called a
Lewis structure. A Lewis diagram omits the inner (complete) shells and only draws the valence electrons.
Examples:
The inner shells are still present, but showing only the valence electrons allows us to focus on how the elements
will behave in bonding. Pairs of electrons like oxygen and fluorine have above are called lone pairs. The solo
electrons that all four examples have are called bonding electrons, since they are the ones that dictate those
behaviours.
Label the lone pairs and bonding electrons on the four examples shown above.
5. When we draw Lewis diagrams of ions, we still add the square brackets and ion charges. The only change
from Bohr diagrams is that since we don’t draw the inner shells, elements that have lost electrons look like they
don’t have any electrons. But they do! Their inner shells are still present.
Examples:
On the next page, draw Lewis diagrams of atoms of the first 20 elements (4 are done for you in the examples!).
On the final page, draw Lewis structures of ions of the first 20 elements.
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1