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SCH3U - CHEMISTRY 2020-21

SCH3U - Homework Questions 2020-2021

Unit Topic Suggested reading from

textbook

Practice

Review of SNC2D

Grade 9 review including nomenclature and chemical reactions

11-22,74-81, 152-177

Handout - Compounds and Reactions Review Optional: Pg. 4 and 5 questions 1-16 excluding #10b, 12 and 15; pg. 148-149 #1-9

Unit 1 Periodic Trends Chapter 1.7 pg. 41 #1-7,9

Type of Bonding: Ionic 56-60 pg. 58 #1; pg.60 #2-5,8

Lewis Structures 61-69 Pg. 65# 1;pg. 67 #2; pg. 69 # 1-6 Handout – Extra Lewis Dot Diagrams Practice

Types of Bonding: Electronegativity 70-73 Pg. 73 # 1-7

Molecular Compounds and Intermolecular Forces

102-118 Pg. 107 #1; pg.108 #4,5; pg. 115 #1-5 Handout – Chemical Forces and Bonding (follow-up questions on last page)

Intermolecular Forces: Hydrogen Bonding

116-118 Pg. 118 # 1-3 Handout – Chemical Forces and Bonding (follow-up questions on last page)

UNIT 1 REVIEW Self-Quiz pg. 136-137 #1,3-5,7-12,16-19,24,25,27-31,37-40 Review pg. 138-145 #6-12,14-16,18-20,22,25,28,29,30,34,37-42,46-49,54,62-66,71-73,79-85,87,91,93,94,99-105

Unit 2 Synthesis and Decomposition Reactions

156-161 Pg. 161 #1-4,7

Oxides 200 - 204 Pg. 204 #1,3-6

Single Displacement Reactions 164 - 168 pg. 166 #1; pg. 167 #2; pg. 169 #1-4

Double Displacement Reactions 172 - 177 Pg. 177 #4-6

Neutralization Reactions 205 - 206 Pg. 211 #1

Combustion Reactions 192 - 195 Pg. 197 #1-4

UNIT 2 REVIEW Self Quiz pg. 244-245 #1-4,6-10,13-16,20,21,23,25, 26,29,30,35 Review pg. 246 #1,3,4,7-11; pg. 248 #54,56,64-66,71

Unit 3 Measurement and Significant Figures 649-650 Handout – Significant Figures Rules

Average Atomic Mass 25-27 Pg. 29 #1-9

The Mole 266-270 Pg. 270 #1-3,4,6

Molar Mass 271-277 Pg. 277 practice #1,2 & questions #4-8 Handout – Working with the Mole Worksheet

Mass and Number of Entities 278-282 Pg. 280 #1-4; pg. 282 #1-3 Handout – Working with the Mole Worksheet

Percent Composition 284-288 Pg. 288 #1-4 Handout – Percent Composition, Empirical and Molecular Formulas (practice questions on last page)

Empirical Formula 289-293 Pg. 292 #1; pg. 293 #3-6 Handout – Percent Composition, Empirical and Molecular Formulas (practice questions on last page)

Molecular Formula 296-300 Pg.298 1; pg. 300 practice #2 & question #2,6,7 Handout – Percent Composition, Empirical and Molecular Formulas (practice questions on last page)

Stoichiometry part 1 316-320 Pg. 319 #1-3; pg. 320 #3,4

Stoichiometry part 2 321-325 Pg. 323 #1-3 Handout – Stoichiometry Worksheet

Limiting Reactants 326-335 Pg. 330 #2-4; pg. 332 #1-3; pg. 334 #1-3 Handout – Extra Practice Stoichiometry Practice Including Limiting Reactants

Percentage Yield 336-339 Pg. 338 #1,2; pg. 339 #1-13

UNIT 3 REVIEW Self Quiz p. 345 #1-20 Review p. 346 #1,2,4,6,7,9,10,27,30,33,36,39,42,44-48,62,63

Unit 4 Gases: Introduction 516-521, 541-546

Pg. 546 #1-4

Charles’s Law 547-553 Pg. 553 #1-4,10,11

Boyle’s Law 554-556 Pg. 559 #1-2

Gay-Lussac’s Law; Combined Gas Law

557-561 Pg. 559 #3; pg. 560 #1-3; pg. 562 #2-15

Dalton’s Law of Partial Pressures 592-597 Pg. 594 #1-4; pg. 596 #1-3

Avogadro’s Law and Molar Volume 576-581 Pg. 579 #1-3; pg. 580 #1-3

Ideal Gas Law 582-588 Pg. 587 #1-4

Gas Stoichiometry 598-602 Pg. 599 #1-3; pg. 603 #1-7

UNIT 4 REVIEW Self Quiz pg. 618 #1,5-17,24-29,31,32,34,35 Review pg. 620 #8-10,12-14,16,17,23-26,28,30,54-62,65,66,68,69

Unit 5 Water and Solutions 370-397 Pg. 375 #1-3; pg. 381 #1,5,8,9; Tutorial pg.384 #1; pg.389 #2-10, 14; Tutorial pg.394-5 #1-3; pg. 397 #1,4,6,10

Solutions and their reactions: net ion equations

424-428 Pg. 427 #1; pg. 428 #1-6 Handout – BLM 9-5

Qualitative Analysis 437-440 Tutorial pg. 438-9 #1 Handout – Qualitative Analysis (thought lab)

Concentration and Making Dilutions 398-405 Pg. 400#1-4; pg. 405 practice #1-4

Concentration of Consumer Products 406-411 Pg. 411 #2-7

Stoichiometry of Solutions 444-449 Pg. 447 #1,2; pg. 448 #1-3; pg. 449 #1-9

Acids and Bases 464-485 Pg. 469 #2-7; pg. 475 #4-7,9-11; pg. 484 #1-3; pg. 485 #9-13

UNIT 5 REVIEW Self Quiz pg. 500 #1,3-9,11-19,21,23-29,31-36 Review pg. 502 #1,2,8-12,15,16,19,25-31,36,43,44,47-49,60,61,65,71,74-94

Activity Series of Non-metals

F2 > Cl2 > Br2 > I2

Electronegativities

Shapes

Gas Equations

K = 273 + oC Constants

PV = nRT Avogadro’s number = 6.02 x 1023 particles/mol

𝑉1

𝑇1=

𝑉2

𝑇2 Gas constant, R = 8.314

𝑘𝑃𝑎∙𝐿

𝑚𝑜𝑙∙𝐾

𝑃1𝑉1 = 𝑃2𝑉2 =0.0821 𝑎𝑡𝑚∙𝐿

𝑚𝑜𝑙∙𝐾

𝑃1

𝑇1=

𝑃2

𝑇2 1 atm = 760 mmHg = 760 torr = 101.3 kPa

𝑃1𝑉1

𝑇1=

𝑃2𝑉2

𝑇2 STP = 0.000oC and 1.00 atm

PT = PA + PB + PC …

PA = 𝑋𝐴 x total pressure where 𝑋𝐴 = 𝑚𝑜𝑙 𝐴

𝑡𝑜𝑡𝑎𝑙 𝑚𝑜𝑙

SCH3U

Review of SNC2D – Chemical Compounds and Chemical Reactions

Part A - Chemical Compounds

Pure substances can be either elements or compounds:

An element consists of atoms that have identical numbers of protons in their nucleus. They cannot be broken

down by chemical means. They can be found on the periodic table. There are 7 diatomic elements these are:

A compound is any substance composed of identical molecules (or formula units) consisting of atoms of two or

more chemical elements.

Ionic compounds

Ionic compounds are composed of ions held together by ionic bonds which are electrostatic forces of attraction between

oppositely charged ions. The compound is neutral overall but consists of positively charged ions called cations and

negatively charged ions called anions.

Ionic compounds are formed when:

- Atoms of an element lose electrons from their valence shell so that they have the same electron arrangement as the

nearest noble gas. In doing so, they have become chemically stable. In losing negatively charged electrons, they now

possess a positive charge. These are referred to as cations.

- Atoms of an element gain electrons so that they have the same electron arrangement as the nearest noble gas. In

doing so, they have become chemically stable. In gaining negatively charged electrons, they now possess a negative

charge. These are referred to as anions.

- Opposites attract, so the cations are attracted to the anions. This

electrostatic attraction is referred to as the ionic bond. An ionic

compound has been formed.

Naming ionic compounds:

To name an ionic compound: name of cation followed by name of anion eg. sodium chloride

Notes:

1. Cations have no special suffix, however simple anions always end in -ide.

2. Ions made up of more than one atom are referred to as polyatomic ions. The formulas and names for these can

be found on the data sheet provided.

3. Some metals can form more than one kind of ion. These are called multivalent metals and can be identified on

the periodic table. One must include the charge of the ion as a roman numeral immediately following the cation.

eg. iron (III) chloride indicates that the Fe3+ ion is in this compound. If the metal is not multivalent, then no charge

should be indicated.

Examples:

CaBr2 K2SO3 CuO

Writing formulas for ionic compounds:

Key idea: since the compound has no charge overall, the positive charges of the cations must be equal to the

negative charges of the anions

Notes:

1. Polyatomic ions will need brackets around them if more that one is needed to balance the charge.

2. The cross-over rule may be applied, however using reason to determine the formula is preferential.

3. Always simplify formulas to lowest ratio of cations to anions

:

aluminum oxide magnesium nitrate nickel (II) phosphate calcium sulfide

Molecular Compounds

Many compounds are made of non-metals only. Non-metals are close to filling their

valence shell so they do not readily lose electrons. So instead, when a compound is

formed between non-metals only, valence electrons are shared. The shared

electrons form a covalent bond. Non-metals can share their electrons in a variety of

ways, so naming rules are important.

Naming molecular compounds:

• Prefixes are used to indicate the number of each element.

• Examples The second element ends with -ide.

• If there is only one of the first element, then no prefix is used. A prefix is always used for the

second element.

Examples:

CO CO2 N2O5

Writing formulas for molecular compounds:

• The name indicates the number of each element in a molecular compound.

• Do not simplify formulas.

Examples:

trinitrogen hexoxide carbon tetrachloride

Part B – Chemical Reactions

Note: balance all chemical equations to ensure that they obey the Law of Conservation of Mass

Synthesis reactions:

• Two elements or compounds form one product.

• You were asked to predict the product when a metal reacted with a non-metal to form an ionic compound. Be

careful to determine the correct formula for the product.

Examples: 2Li + F2 → 2LiF 4Fe + 3O2 → 2Fe2O3

Note: if the metal is multivalent, use the more common charge

Decomposition reactions:

• A compound breaks apart into two or more elements or compounds.

• You were asked to predict products for the decomposition of:

o a binary compound into its elements eg. 2NaCl → 2Na + Cl2

o a ternary compound into compounds if one product was given eg. CaCO3 → CaO + ________

Single Displacement reactions:

• A reactive element replaces a less reactive element in a compound.

• The Activity Series of Metals or the Activity Series of Non-metals was

used to determine if the reaction would proceed

Examples:

3Mg + 2FeCl3 → 3MgCl2 + 2Fe Pb + NaNO3 → no reaction

2KI + Cl2 → 2KCl + I2 LiF + Br2 → no reaction

Note: if the more reactive metal is multivalent, use the more common charge

in the new compound

Double Displacement reactions:

• Cations and anions switch partners to form new compounds

Example: Li2CO3 + 2AgCl → Ag2CO3 + 2LiCl

Combustion of hydrocarbons:

• A compound comprising of carbon and hydrogen (and sometimes oxygen) reacts with elemental oxygen (O2)

• The products for a complete combustion are always CO2 and H2O

Example: 2C2H6 + 7O2 → 4CO2 + 6H2O

SCH3U

Success Criteria – Unit 1 - Matter, Chemical Trends, and Chemical Bonding

At the end of the unit, I am able to:

• Use appropriate terminology related to chemical trends and chemical bonding including, but not limited to: atomic radius, effective nuclear charge, electronegativity, ionization energy and electron affinity

• Discuss the structure of the atom and isotopes

• Name and give formulas for ionic and molecular compounds

• Describe how shielding works to determine effective nuclear charge

• Explain how effective nuclear charge can affect atomic size, ionization energies, electron affinities, ion sizes and the trends along a period and down a group of the periodic table

• Analyze data related to the properties of elements within a period and down a group (ex. Ionization energy, atomic radius) to identify general trends in the periodic table

• Draw Lewis structures to represent the bonds in ionic and molecular compounds and polyatomic ions

• Write structural formulae for molecular compounds containing single and multiple bonds

• Identify the type of bonding (ionic, covalent, polar covalent metallic ) based on the nature of the elements involved and electronegativity differences.

• Use VSEPR notation to determine the electron geometry and shape of a molecule

• Determine whether a molecule is polar or non-polar based on electronegativity differences and shape of the molecule

• Discuss all types of intermolecular forces: ionic, London dispersion forces (LDF), dipole-dipole forces, hydrogen bonds

• Given an example, identify the type of intermolecular forces in the substance and make inferences about its physical properties

SCH3U

Trends in the Periodic Table

Learning Goals

• Use appropriate terminology related to periodic trends: valence electron, nuclear charge, shielding, effective nuclear charge, ionization energy, electron affinity, atomic radius, ionic radius

• Explain trends of the periodic table and recognize relationships between them (eg. atomic radius and first ionization energy)

• Using provided data, be able to construct a graph, explain any trends and analyse observed relationships

Some key ideas for this lesson:

Like charges _________________ while unlike charges _________________.

The larger the charge, the __________ it will either attract or repel.

The closer in distance unlike charges are to each other, the ______________ the force of attraction between them while

the farther they are apart, the ______________ the force of attraction.

Also, the closer like charges are the ________________ the force of repulsion and vice versa.

The above ideas collectively are called Coulomb’s Law.

Overall Periodic Trends

As you go across a period:

• The charge on the nucleus gets ____________________.

• The outermost electrons are in the _______________ energy level.

• So, the outermost electrons are pulled more ______________ across a period.

Overall Group Trends

As you go down a group:

• There are ______________ energy levels.

• The outermost electrons are ______ attracted to the nucleus for two reasons:

o Distance from the nucleus

o Shielding

Shielding

“Electrons in filled energy levels between the nucleus and outer shell electrons shield the outer shell electrons somewhat

from the effect of protons in the nucleus.”

In the same energy level, the shielding is the _______________.

The more energy levels between the nucleus and the outer electrons, the __________ shielding.

The outer electrons are _________ attracted to the nucleus.

Effective Nuclear Charge, Zeff

Effective nuclear charge, Zeff, is the net force of attraction between the electrons and the nucleus of the atom. It

considers both the attractiveness of the nucleus, the energy level of the valence electrons and the shielding of valence

electrons by core electrons. We say that an atom such as Cl has a _____________ effective nuclear charge than Si since,

even though the valence electrons are on the _______ energy level and amount of shielding is the ___________, Cl has

_________ protons in the nucleus.

Sizes of Atoms (Radius)

Group trend: as we go down a group, the size of the atoms _________________

• Each atom has another energy level

• The distance between the nucleus and the valence electrons becomes ___________ so the attraction is ______________.

• ____________ shielding also results in ____________ attraction.

• The atoms get _______________

Periodic trend: as we go across a period, the size of the atoms _________________

• Each atom has the same number of energy levels and the _________ shielding.

• ___________ nuclear charge as you go across a period.

• More nuclear charge means ____________ attraction.

• The outermost electrons are pulled ____________ making the atom _______________.

Ionization Energy of Atoms (IE)

IE is:

Overall Trend: IE _______________ across a period and ________ a group

Periodic trend: as we go across a period, the IE _________________

• Valence electrons are in the __________ energy level and have the _________ shielding.

• ___________ nuclear charge as you go across a period.

• More nuclear charge means a ____________ force of attraction to the nucleus

• This makes more _______________ to remove an electron

• Therefore, it takes ______________ energy to remove an electron

Group trend: as we go up a group, the IE _________________

• Each atom has one ________ energy level

• The distance between the nucleus and the valence electrons becomes ___________ so the attraction is ______________.

• ____________ shielding also results in ____________ attraction.

• It is _______________ to remove an electron.

• So it takes ____________ energy to remove an electron.

Electron Affinity (EA)

EA is:

EA is basically a measure of:

Overall Trend: EA _______________ across a period and ________ a group.

Except:

Size of Ions: Cations

Cations are ___________ ions formed when atoms _____________ electrons in order to have filled outer shells.

_______________ form cations.

Cations are ____________________ than the atom they came from.

To explain why, we’ll look at magnesium:

The nucleus has not changed, so the nuclear charge felt by the electrons is the _____________.

The outermost electrons are now in a ________________ energy level. The shielding of the outermost electrons is now

________________________. So, the size __________________ significantly.

Size of Ions: Anions

Anions are ___________ ions formed when atoms _____________ electrons in order to have filled outer shells.

_______________ form anions.

Anions are ____________________ than the atom they came from.

To explain why, we’ll look at chlorine:

The nucleus has not changed, so the nuclear charge felt by the electrons is the _____________.

The outermost electrons are still in the ________________ energy level. However, as the energy level fills up, the

electrons begin to repel each other more and ____________ apart. So, the size ______________________ .

Overall:

Success Criteria

I can:

• Use appropriate terminology related to chemical trends and chemical bonding including, but not limited to: atomic radius, effective nuclear charge, electronegativity, ionization energy and electron affinity

• Describe how shielding works to determine effective nuclear charge

• Explain how effective nuclear charge can affect atomic size, ionization energies, electron affinities, ion sizes, isoelectronic series, successive ionization energies and the trends along a period and down a group of the periodic table

• Analyze data related to the properties of elements within a period and down a group (ex. Ionization energy, atomic radius) to identify general trends in the periodic table

SCH 3U

Types of Bonding and Compounds

Learning Goals

• Know the different types of chemical bonds, understand why they are formed and why there are differences

• Know by looking at the elements involved what type of chemical bond is formed

• Be able to write names for compounds

• Be able to write chemical formulae

Activity:

You have been given pictures of samples of different ionic and molecular compounds as well as a substance held together

by metallic bonding. You also have been given a list of possible names and formulas. Match the names and formulas

with the compounds. Record your work in the chart below and provide a physical description of each substance. Identify

the compound as ionic, molecular or metallic.

Substance # 1 2 3 4 5 6

name

formula

physical characteristics

type of bonding

Substances:

# 1 #2 #3 #4 #5 #6

Names: copper (II) sulfate iron sucrose (sugar)

sodium chloride (table salt) oxygen water

Formulas: O2 C12H22O11 Fe CuSO4 H2O NaCl

Definitions:

Ionic compound: Ionic bonding occurs when two or more atoms _________________ electrons, forming

___________ and anions. The oppositely charged ions are then held together by forces of

_________________. The electrons are transferred so that the atoms can gain _____________.

Covalent compound: Covalent bonding occurs when two or more atoms ____________________ electrons. The

atoms share electrons in order to gain ____________________. Compounds held together by

covalent bonds are called __________________________.

note: elements that exist as diatomic molecules are: ___________________________________

Metallic bonding: When metal atoms ____________electrons. See description on next page for full details.

Properties of Ionic and Covalent Compounds:

Property Ionic compound Covalent compound

State at room temperature

Melting point

Electrical conductivity as a liquid

Solubility in water

Electrical conductivity when dissolved in water

Explaining the high conductivity of ionic compounds:

Ionic compounds do not conduct electricity in their solid state. They are very good conductors in the liquid state

or when dissolved in water. Why?

An electric current can flow only if _______________ particles

are available to move and carry the current. In solid state, ionic

compounds are arranged in a rigid lattice so they can’t move

much. In molten form, the rigid lattice is broken, allowing the

ions that make up the compound to ________. When the

compound is in the water, the ________ are also free to move,

thus allowing for the conduction of electricity.

Explaining the low conductivity of covalent compounds:

Covalently bonded compounds whether in solid, liquid or gaseous state, never break their ___________. They do

not break up into _________ when melted, boiled or dissolved but rather stay together as _______________.

Since there are no charged particles formed, no electrical current can be carried.

Metallic Bonding:

How do metal atoms bond to each other?

Metals do not form ionic bonds with other metals since neither wants to gain electrons to form anions. Evidence

proves this… Na can be cut with a knife; pure gold and copper can be bent and hammered to a sheet, whereas

ionic compounds are hard and brittle.

Metals do not form covalent bonds with

each other either since they do not have

enough valence electrons to share to make

a full octet.

They DO, however, share electrons, but in a

different manner. In metallic bonding,

atoms release their valence electrons to a

shared pool of electrons. Because of this

the electrons are said to be

_________________. The force that holds

metal atoms together is called a

________________ bond. Because the

electrons are free to move, the ions are not

in a rigid lattice. Thus, metals can change shape when bent, etc.

Activity: p.85 Investigation 2.4.1 #1*b-d

Do not copy the full table. Just fill construct a table with the headings:

Compound Ionic or covalent Possible identity*

A

* choose from: wax, CuSO4, sucrose, hydrogen, table salt, calcium carbonate, water, gasoline, ammonia

Success Criteria

I can:

• Explain, compare and contrast ionic, covalent and metallic bonding

• Be able to name and give formulas for ionic and molecular compounds

SCH 3U

Lewis Structures

Learning Goals

• Use appropriate terminology related to bonding including valence electrons, chemical stability, noble gas electron configuration, bond, ionic, covalent, molecule, formula unit, metal, non-metal, Lewis dot diagram, structural formula, electronegativity, polarity, polar bond, non-polar bond, dipole

• Be able to represent bonding and compounds with diagrams

In the past we have used Lewis structures to represent the bonding in molecules. It is possible to construct these

diagrams through trial and error as we did in the past. Here we will also look at a systematic approach for a wider range

of molecules and polyatomic ions. It is up to you which method you use.

Note: polyatomic ions are ions made up of covalently bound atoms

Steps for Drawing Lewis Structures

1. Position the least electronegative atom (other than H) in the centre of the molecule or polyatomic ion. Write the other atoms around this central atom, with each atom bonded to the central atom by a single bond. Always place a hydrogen atoms or a fluorine atom at the end position in the structure.

2. A) Determine the total number of valence electrons in the molecule or ion. For polyatomic ions, pay close attention to the charge. For example, if you are drawing a polyatomic anion such as CO3

2-, add two electrons to the total number of valence electrons calculated fort the structure CO3. For a polyatomic ion such as NH4

+, subtract one electron from the total number of valence electrons calculated for the structure NH4. B) Once you have the total number of valence electrons, determine the total number electrons needed for each

atom to achieve a noble gas configuration.

C) Subtract the first total from the second total to get the number of shared electrons. Then divide this number

by 2 to give the number of bonds. Double or triple bonds may be needed to account for this number of bonds.

Double bonds count as two bonds and triple bonds count as three bonds.

3. Subtract the number of shared electrons from the number of valence electrons to get the number of non-bonding electrons. Add these electrons as lone pairs to the atoms surrounding the central atom so that you achieve a noble gas configuration for each atom.

4. Structural formulas are the Lewis structures without the electron dots.

Examples:

compound Lewis structure structural formula

NH3

SO2

CH2O

CO32-

Success Criteria

I can:

▪ Draw Lewis structures to represent the bonds in ionic and molecular compounds ▪ Write structural formulae for molecular compounds containing single and multiple bonds

SCH 3U

Electronegativity Worksheet

Learning goals:

• Use appropriate terminology related to bonding including bond, ionic, covalent, molecule, metal, non-metal, electronegativity, polarity, polar bond, non-polar bond, dipole

• Know what electronegativity is and be able to describe its trend on the Periodic Table

• Know the different types of chemical bonds and be able to identify them by calculating electronegativity differences • Know by looking at the elements involved what type of chemical bond is formed

Electronegativity

Read p70-74 of your textbook to answer the following questions:

1. Fill in the blanks:

The electronegativity is the ability of an _______________________ atom, when ________________ to attract

______________________________ to itself.

2. How is electronegativity similar to periodic trends such as atom size, ionization energy and electron affinity? How is it different?

3. Circle the correct word and strike out the incorrect word in the following sentences:

a. Electronegativity increases / decreases from left to right on the periodic table. b. Electronegativity increases / decreases down a group.

4. Draw 2 arrows on this periodic table to show the trends for increasing electronegativity.

Using Electronegativity Values to Determine Bond Types

• One can use electronegativity to predict the type of bonds that will form between two atoms.

• EN stands for the difference between two electronegativity values. When calculating, always substract the smaller value from the larger one (no negatives)

Examples: Using the electronegativity values from the periodic table at the back of your textbook (refer to

legend)

KF EN = 4.0 - 0.8 = 3.2

O2 EN = 3.4 - 3.4 = 0

HCl EN = 3.2 - 2.2 = 1.0

The values for EN can help you determine the type of bond. There is an established range for each type of bond:

To determine if a bond is ionic:

➢ check that there are metals and non-metals present

➢ usually but not always, EN is between 1.7 - 3.3 for ionic

To determine if a bond is covalent and non-polar:

➢ check that only non-metals are present

➢ EN is between 0 - 0.5 for non-polar molecules

To determine if a bond is polar covalent:

➢ check that only non-metals are present

➢ EN is between 0.5 - 1.8 for polar molecules

➢ A bond that is polar covalent is said to have a dipole.

Fill in the following chart using the EN values from page 71. The first one is done for you.

Bond EN ionic covalent polar covalent

O --- H 3.4 – 2.2 = 1.2 x

Al --- F

C --- H

Mg --- Cl

Ni --- O

Na --- Cl

S --- O

Ni – Ag (be careful!)

Success criteria

I can:

• Explain what electronegativity is, and describe its trend on the periodic table

• Identify the type of bonding (ionic, covalent, polar covalent, metallic) based on the nature of the elements involved and electronegativity differences

• Use appropriate terminology in my explanations including bond, ionic, covalent, molecule, metal, non-metal, electronegativity, polarity, polar bond, non-polar bond, dipole

SCH 3U

Polar Molecules

Polar Covalent Bonds

When two bonding atoms have an ________________________________ difference that is greater than _____ but less

than ______, they are considered to have a __________________ _________________ bond. The difference between

the electronegativities is not great enough for a transfer of _________________ electrons, but it is great enough for the

bonding pair to spend _________ time near the _____________ electronegative atom. In other words, the atoms share

the electrons but not ___________________. This results in an __________ _________________ of charge occurring

across polar covalent bonds. The bond will have a __________ ____________ side, indicated by a _____ sign and a

__________ ____________ indicated by a ____ sign.

For example: HCl

Molecular Shapes

Electron pairs in the valence shell, both bonding pairs (BP) and non-bonding pairs (lone pairs, LP), ___________ each

other and will spread apart as far as possible to ____________________ repulsions, ie. the angle between two adjacent

electron clouds will be maximized.

Activity: • Draw Lewis diagrams for each of the following. Put your drawings in Table 2 on page3 of this booklet:

H2O CO2 SO2 NH3 CH4 BF3

• Think of each one as a molecule with a central atom (A) with different numbers of other atoms (X)

bonded to it and different numbers of electron pairs around the central atom (E).

For example: H2O

The central atom is O, it has 2 H’s attached to it and 2 electron pairs unbounded on O. It would be AX2E2

Now do this notation for all of the molecules and include it in Table 2 under the column VSEPR notation.

• Look at the list of possible factors listed below that may affect the shape of a molecule.

Learning Goals:

• Use appropriate terminology related to molecular shape and polarity including Lewis dot diagram, valence shell electron pair repulsion theory (VSEPR), bonding pairs, lone pairs, electronegativity, polarity, polar bond, non-polar bond, dipole, polar molecule, non-polar molecule

• Use the localized electron bonding model to describe and predict molecular geometry using Lewis dot diagrams and the VSEPR model

• Understand and be able to explain why molecular shapes can be predicted using the VSEPR model • Identify polar molecules and assign partial charges to the poles of the molecule and be able to explain their reasoning

• Fill in the Predict column of Table 1 by thinking whether this will be a factor in determining the shape of

the molecule; state your reasons in the column to the right.

• DO NOT FILL IN THE FACT COLUMN AT THIS TIME!

Table 1

factor

Predict: is it important in

determining the shape of the molecule?

Reason

Fact: is it important in

determining the shape of the molecule?

Reason

Number of atoms (X) bonded to A

Number of non-bonded electron pairs on A (E)

Size of X

Size of A

Number of non-bonded electrons on X

Presence of multiple bonds

• Go to https://phet.colorado.edu/en/simulation/molecule-shapes and download program.

• Choose Real Molecules

• Use the modelling program to complete the following chart.

• Using the results from the chart, review your predictions from Table 1. Fill in the Fact column and state

your reasons in the column to the right.

Table 2

Compound name and formula

Lewis Structure # single bonds

# double

or triple bonds

# lone pairs

around central atom

Sketch of shape Including angles between bonds

VSEPR Electron geometry

Molecular geometry (shape)

H2O water

CO2 carbon dioxide

SO2 sulfur dioxide

NH3 ammonia

CH4

methane

BF3 boron trifluoride

VSEPR The most common arrangement of sets of electron clouds (of a molecule) around a central atom may be predicted using Valence Shell Electron Pair Repulsion Theory (VSEPR). See the chart below.

The electron geometry of a molecule indicates the arrangement of the electron pairs around the central atom.

The molecular geometry (or shape) of the molecule describes the overall shape made by the atoms in the molecule.

To determine the molecular geometry and shape of a molecule:

1. Draw a Lewis diagram.

2. Determine the VSEPR notation.

3. Look up the electron geometry and shape using the chart.

Examples:

a) PBr3 b) H2CO

electron geometry shape

linear

trigonal planar

tetrahedral

trigonal bipyramidal

octahedral

Polar Molecules

To determine if a molecule is polar or not you have to consider both:

1. the

AND

2. the

For example: H2O and CO2

More examples: SH2 and SF2

Polar molecules are also said to have a ______________________ because they have a negative and a positive pole. A molecule with polar bonds but without a dipole is a ____________________________ molecule.

Example: Why is CF4 non-polar while CH2F2 is polar?

Success Criteria

I can:

• Explain what a polar bond is

• Use VSEPR notation to determine the shape of a molecule

• Can explain why the VSEPR model can be used to predict molecular shape

• Determine whether a molecule is polar or non-polar based on electronegativity differences and shape of the

molecule

• Can identify the poles (+ and -) of a polar molecule and explain their reasoning

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Chemical Bonding and Forces (Inter and Intramolecular Forces)

• For every type of bonding, there is a that "creates" the bond

• There are two types of forces: ____ and _

• Intramolecular forces are usually called ______________ and occur ___________ a molecule and are

responsible for the _ properties of a substance

• Intermolecular forces (IMF) occur ______ two different molecules and are responsible for the

properties of a substance (ie. boiling point, melting point, solubility, etc)

• The stronger the force, the the bond or IMF, and the the distance between

particles

Forces and Bonds

• There are types of intramolecular forces; these are typically the types of forces as

they need to keep the atoms of a molecule or ionic compound together → chemical bonds

• There are types of intermolecular forces, these help to keep the molecules or formula units

together in a sample of the substance → IMF

Learning Goals:

• Use appropriate terminology related to intermolecular forces and physical properties including intramolecular forces, covalent, polar covalent, ionic, metallic, intermolecular forces (IMF), London dispersion forces (LDF), dipole-dipole forces, hydrogen bonding, boiling point, melting point, solid, liquid, gas

• Know the different types of IMF and be able to describe what they are and why they are formed

• Based on the type of the compound (ionic, metallic, non-polar molecule, polar molecule) determine the types of IMF present

• Based on the IMF predict, compare and contrast physical properties of substances

• Based on boiling points, be able to justify the presence of specific IMRF in a sample of a substance

Substance

Intramolecular Forces

Covalent

Ionic Metallic

Non-polar Molecule Polar Molecule

AND Ionic Metallic

London Dispersion Forces OR H has FON

Dipole-Dipole Forces Hydrogen Bonding

Intermolecular Forces

1. Ionic Substances

• Recall: ionic bonds form between two opposite (cations and anions)

• Ionic compounds are formed between atoms with large differences in their _

very often compounds that include and have ionic bonds

• These ions are strongly attracted together by forces; therefore, the ions are held

together by bonds

• The the charge difference between the ions, the more they are held

together; this is reflected in their physical properties such as melting point and solubility.

Example: Ca and Cl

Example: Ca and O

• Ionic compounds do not exist as discrete ________________________but rather as a lattice of ions; as such

ionic compounds are referred to with respect to __________________________. For example, NaCI is a

formula unit within a lattice of sodium chloride, or Ca(NO3)2 is a formula unit within a lattice of calcium nitrate.

• forces between the ions of different keep ionic compounds

together in a sample of the substance.

Example: NaCI Iattice

3. Molecular Substances

• Molecular substances are discrete molecules with atoms held together by bonds

• Covalent bonding exists when atoms electrons

•

Example: CO2

atoms are involved in covalent bonds

• There are two types of covalent bonding: covalent bonding and covalent

bonding; as such, the type of intramolecular force in a molecule is either or

covalent bonding (note: non-polar covalent bonding is often referred to as just covalent bonding)

• Molecules can be polar or non-polar. To determine if a molecule is polar or non-polar one must consider:

o the of the bonds AND

o the of the molecule

• The type of intermolecular force within a sample of molecules depends on whether it is a or

molecule

Polar and Non-polar Bonds

• Recall: for polar bonds between non-metals, the electronegativity difference is _

for non-polar bonds between non-metals, the electronegativity difference is __________________

Shapes of Molecules

• Recall: steps to determine shape of molecule:

1. draw the _

2. determine the VSEPR notation

3. use the _ to determine the shape

Polar and Non-polar Molecules

• Recall: for a molecule to be polar is must have:

1. At least one _bond

2. A shape that gives the molecule at least distinct "sides"

Example: NH3 vs. CO2

Intermolecular Forces for Molecular Substances

• The type of intermolecular force within a sample of molecules depends on whether it is a ______________ or

_______ molecule; so the first step in determining the type of intermolecular force is to determine

if the molecule is polar or non-polar.

i) London Dispersion Forces (LDF)

• molecules (covalent compounds) are

held together in whole or in part by LDF

• forces that occur between two different atoms

of two different molecules with the same or similar

dipole. They are created by momentary dipole.

• The overall force of attraction between molecules held together by LDF depends on the number of

electrons. involved and the size of the atoms, therefore ______________ atoms and

molecules have _______________overall forces of attraction between them.

• This can be seen in differences in the ________________properties of the substances.

For example:

• natural gas is a gas at room temperature; it consists of non-polar CH, molecules held together by LDF; CH 4has a melting point of _____________________

• octane is a liquid at room temperature; it consists of non-polar C8H18 molecules held together by LDF;

C8H8has a melting point of _____________________

• paraffin wax is a solid at room temperature; it consists of non-polar C20 H42 molecules held together by

LDF; wax has a melting point o f ___________________

I. Normal condition – in a non-polar species there is a symmetrical charge distribution

II. Instantaneous condition – interactions between electron clouds produce a displacement of electronic charge creating an instantaneous dipole, with charges + and −

III. Induced dipole – the instantaneous dipole on the left induces a charge separation on the right, resulting in a dipole-dipole interaction

+ −

+ −

+ −

ii) Dipole-Dipole Forces

• Dipole-dipole forces exist between ________ ____ __ molecules ONLY.

• They occur when the slightly positive side of one molecule is to the slightly negative side

of another molecule. This force of attraction is the intermolecular dipole-dipole interaction.

• These forces are than LDF.

• Note: Substances held together by dipole-dipole forces will also be held together in part by LDF. LDF will just

play a less significant role.

For example: HCI

iii) Hydrogen Bonding

• Hydrogen bonds are a special type of force. It is a very dipole-

dipole force that occurs between and another very electronegative atom usually ,

, or

For example: H 20

− +

− +

− +

+ −

+ −

+ −

3.Metallic Bonds

• These are special "bonds" created between two atoms in a sample of a metal.

• The metal nuclei and core electrons are held together by LDF and the electrons of each atom

leave their "original" atom and come together to create a "sea of electrons" found above the nuclei.

• The nuclei are arranged in rows and columns (like a lattice) and create very strong attractions; this can be seen in

the melting points of most metals.

• The strength of the metallic bond increases with an in valence electrons.

Follow-up Questions:

1.For each of the following:

i) Determine the type of intramolecular force.

ii) Draw a Lewis structure if is it ionic or covalent.

iii) If it is a molecular compound, determine if it has polar or non-polar molecules.

iv) Determine the type of intermolecular forces present in a sample of the compound.

a) H2CO

b) CS2 c) CHCI3

d) MgCI2

e) N2

f) Ni and Fe

g) SNF

h) HF

2.Which will have the highest boiling point:

a) HF, KCI and HBr? Explain why.

b) CCI4, CF4 or CBr4 ? Explain why.

c) HCI, Ar or F,? Explain why.

3.Which will have the highest melting point: AIN, AICI3 or Al2 03? Explain why.

Success Criteria

I can:

• Describe all type of intermolecular forces: metallic, ionic, London dispersion forces (LDF), dipole-dipole forces,

hydrogen bonds

• Given an example, identify the type of intermolecular forces in the substance

• Explain differences in physical properties of substances based on the IMF present

• Use data to justify the presence of specific IMF present in a sample of a pure substance

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Success Criteria - Unit 2 - Chemical Reactions

At the end of the unit, I am able to:

• Use appropriate terminology related to chemical reactions including, but not limited to: neutralization, precipitate, acidic, and basic

• Write balanced chemical equations to represent synthesis, decomposition, single displacement, double displacement, and combustion reactions

• Identify and predict the products of different types of synthesis and decomposition reactions, including reactions of metal oxides and non-metal oxides with water

• Identify and predict the products of single displacement reactions, using the metal activity series and non-metal (halogen) reactivity series

• Identify and predict the products of double displacement reactions (eg. the formation of precipitates, gases; neutralization)

• Use a solubility chart to predict the solubility of an ionic compound

• Explain the difference between a complete combustion reaction and an incomplete combustion reaction

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Chemical Reactions

• Understand the significance of reaction conditions in determining whether a complete or incomplete combustion reaction occurs

Synthesis and Decomposition Reactions A + B → and A →

Synthesis and decomposition reactions were studied in the SNC2D course and reviewed earlier this year. This course has

two additional types of synthesis reactions that are required learning:

A. Reaction of Metal Oxides and Water

When a metal oxide reacts with water, a _______________ solution results. This is because the product is a

____________ _____________ which are a type of ____________.

CaO (s) + H2O (l) →

Li2O (s) + H2O (l) →

B. Reactions of Non-Metal Oxides with Water

When a non-metal oxide reacts with water, an __________ is formed.

For example: SO3 (g) + H2O (l) → H2SO4 (aq) this is sulphuric acid – it is a very strong acid

Learning Goals:

• Use appropriate terminology related to chemical reactions including acidic, basic, binary, ternary, metal oxide, non-metal oxide, activity series, halogen series, precipitate, solubility, neutralization, complete and

incomplete combustion

• Identify synthesis, decomposition, single displacement, double displacement and combustion reactions

• Predict the products of different synthesis, decomposition, single displacement, double displacement and combustion reactions

• Write balanced chemical equations to represent synthesis, decomposition, single displacement, double displacement and combustion reactions

• Explain the chemical reactions that result in the formation of acids and bases from metal oxides and non-metal oxides

• Use the activity series of metals and the halogen series to determine whether a single displacement reaction will proceed or not

• Use a solubility chart to determine whether a precipitate is formed during a reaction

• Use criteria such as solubility, formation of gas and formation of water to predict whether a double displacement reaction occurs, and explain why these criteria are necessary

• Recognize stable and unstable products of double displacement reactions including water during neutralization reactions, and predict the stable products made when it decomposes if a product is unstable

• Plan and conduct an investigation to demonstrate the role solubility of the products plays in double displacement reactions

• Explain the difference between complete and incomplete combustion reactions, and know the products of each type

Single Displacement Reactions

One __________________ displaces or replaces an _______________ in a _______________________.

The Activity Series - a table of the reactivity of _________________ relative to each other. The most reactive metals are

at the _______________ of the activity series. The least reactive metals are found at the bottom.

____________________ is included in the series, even though it is not a ___________, because it forms a

________________ charged ion like a metal.

The activity series is based on two important generalizations:

1. One element can displace elements ____________ it from compounds in solution but cannot displace elements _____________ it.

2. The farther apart two elements are, the more likely it is that the displacement will occur _________________. To determine whether or not a reaction between an element and a compound will react, look at the relative reactivities

of the two metals. If the higher (more _______________) metal is the ___________________, the reaction proceeds. If

the higher metal is in the compound, ___________ reaction occurs.

A. Single Displacement Reactions and the Metal Activity A + BC →

There are 3 types of single displacement reactions that involve metals:

1. Reactions involving a Metal and an Ionic Compound

Examples: Mg(s) + CuSO4(aq) →

Pb(s) + Zn(NO3)2(aq) →

➢ The reaction between Mg and CuSO4 occurs because:

➢ The reaction between Pb and Zn(NO3)2 does not occur because:

2. Reactions involving a Metal and an Acid

For example: Li(s) + H2SO4(aq) →

➢ The reaction between Li and H2SO4 occurs because:

3. Reactions between a Metal and Water

For example: Na(s) + H2O(l) →

➢ The reaction between Na and H2O occurs because:

B. Single Displacement Reactions Involving Halogens A + BC →

The elements of the __________________ family have their own activity series.

Most reactive

The halogen activity series:

Least reactive

The single displacement pattern for halogens is slightly different from the pattern used for the metals. This is because a

__________________ ion (_________________) is being displaced rather than a positive ion (cation).

Examples: Cl2(g) + 2KI(aq) →

➢ This reaction occurs because:

I2(g) + CaBr2(aq) →

➢ This reaction does not occur because:

Combustion of Hydrocarbons

➢ Complete combustion of hydrocarbons: there is enough oxygen present to make H2O and CO2

➢ Incomplete combustion of hydrocarbons: there is not enough oxygen to make CO2 exclusively, so the products

are H2O and a mixture of CO and CO2.

For example: Write the balanced chemical equation for:

a) The incomplete combustion of C3H8 (assume no CO2 is made)

b) The complete combustion of C4H10

Other combustion reactions:

➢ Technically, a combustion reaction is any reaction where one of the reactants is oxygen.

➢ Most combustion reactions are exothermic.

For example:

• the combustion of magnesium is: 2Mg(s) + O2(g) → 2MgO(s)

heat and light are given off

Using a Solubility Table

This solubility table can also be found on page 173 or 665 of your textbook. It is always provided for tests and exams.

For each of the following compounds:

i) Use the table to find the solubility of each of the following compounds.

ii) Explain how you used the table to get your answer.

a) AgNO3

b) BaSO4

c) (NH4)2CO3

d) PbI2

e) Na2S

Double Displacement Reactions AB + CD →

Compounds __________________ their __________________ and _________________ pairs.

Double displacement reactions will only occur if:

1.

or

2.

or

3.

1. Reactions that produce a precipitate:

How do you know if a product is insoluble?

Examples: Do the following reactions proceed? If so, determine the products and balance the equation.

AgNO3(aq) + KCl(aq) →

➢ The reaction between AgNO3 and KCl happens because:

K2SO4(aq) + FeCl3 (aq) →

➢ The reaction between K2SO4 and FeCl3 does not happen because:

2. Double displacement reactions that produce a gas:

Sometimes a gaseous product is formed during a double displacement reaction. When it is formed, it escapes into the

air as bubbles of gas. A common gaseous product is H2S(g). It is formed when an acid reacts with a soluble sulfide such

as sodium sulfide.

Na2S(aq) + 2HCl(aq) →

Gases are also produced when an ________________ product of a double displacement reaction __________________.

This is a two-step process. The common unstable products to be watching for are: H2CO3 (carbonic acid) that

decomposes into H2O and CO2(g), and H2SO3 (sulfurous acid) that decomposes into H2O and SO2(g).

Example 1: carbonic acid produced, then decomposes

Overall: MgCO3 (aq) + H2SO4 (aq) →

Step 1: MgCO3 (aq) + H2SO4 (aq) →

Step 2:

Example 2: sulfurous acid produced, then decomposes

Overall: MgSO3 (aq) + HBr(aq) →

Step 1: MgSO3 (aq) + HBr(aq) →

Step 2:

3. Neutralization reaction:

An acid reacts with a base to form _______________ and __________________.

For example: HNO3 (aq) + NaOH (aq) →

Success Criteria

I can:

• Use appropriate terminology related to chemical reactions including, but not limited to: neutralization, precipitate, acidic,

and basic

• Write balanced chemical equations to represent synthesis, decomposition, single displacement, double displacement, and

combustion reactions

• Identify and predict the products of different types of synthesis and decomposition reactions, including reactions of metal

oxides and non-metal oxides with water

• Identify and predict the products of single displacement reactions, using the metal activity series and non-metal (halogen)

reactivity series

• Identify and predict the products of double displacement reactions (eg. the formation of precipitates, gases; neutralization)

• Use a solubility chart to predict the solubility of an ionic compound

• Explain the difference between a complete combustion reaction and an incomplete combustion reaction

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Success Criteria – Unit 3 - Quantities in Chemical Reactions (Stoichiometry)

At the end of the unit, I am able to:

• Use laboratory equipment to measure quantities to the appropriate number of significant digits (estimating the last digit)

• Be able to identify the correct number of significant digits in a given measurement or numerical quantity

• Be able to give answers to calculations involving multiplication, division, addition and subtraction to the correct number of significant digits

• Use the factor-label method to carry out calculations involving multiplication or division

• Explain the concepts of mole and molar mass

• Determine the average atomic mass of an element

• Use Avogadro’s number and molar mass to convert between mass, moles and number of particles

• Determine percent composition, empirical and molecular formulas

• Apply the concepts of mole, molar mass and mole ratios in chemical reactions using the factor-label method

• Predict the amount or the mass of product/reactant expected in a chemical reaction

• Identify the limiting reactant and excess reactant in a chemical reaction and predict the amount or the mass of product expected

• Explain how the outcome of a chemical reaction can be affected by manipulating the limiting and excess reactants

• Calculate the amount of excess reactant remaining after a reaction

• Calculate percentage yield

• Analyze quantitative data from an experiment and determine percentage yield to compare the actual and theoretical yields

SCH3U

Significant Figures Rules:

• All non-zero digits DO count.

• Leading zeros (the zeros in _____________ of numbers) DON’T count.

• Captive zeros (_____________________) DO count.

• Trailing zeros DO count ONLY IF the number contains a __________________.

• Exact numbers (for example, 3 people) have _________________________ significant digits.

• For scientific notation, use the number of significant digits in the ______________________.

Examples:

567mm 0.0012g 101.3OC 706000s 210.0kg 5.6 x102m 56 apples

2502kJ 0.02102L

Number of sig figs

Significant Figures in Calculations:

1. When multiplying or dividing, round

the answer to the ___________ number of

___________________________________.

2. When adding or subtracting, round

the answer to the ___________ number of

___________________________________.

Examples:

a) 2.36 x 0.045 b) 123.25 – 4.108 c) 24.3 x (12.01 + 3.7)

Practice

A. Multiplying and Dividing:

a. When multiplying or dividing, multiple or divide, and then round the answer to the least number of significant

figures contained in the question.

b. When multiplying or dividing an uncertain value by an exact value, the answer has the same number of decimals

places as the number with uncertainty. Eg. 3 x 5.2 g = 15.6 g

c. Be careful with units.

Practice: Use your calculator to complete the required calculations. Give all answers to the correct number of

significant figures.

1. 22.4 g ÷ 1.2 g/mL [19mL]

2. 6.7 cm x 2.4 cm x 1.31 cm [21 cm3]

3. 600 x 4.521 g {600 is uncertain} [3000g]

4. 600 x 4.521 g {600 is exact} [2713g]

5. 600. X 4.521 g [2710g]

6. 19.3 g/mL x 4.5 mL [87g]

7. 4.123 cm x 2.5 mm [1.0cm2]

8. 1.2 x 10-4 x 3.02 x 10-23 [3.6 x 10-27]

9. 1.30 x 10-3 x 2.100 x 10-6 [2.73 x 10-9]

10. 260. x 10 {10 is uncertain} [3000]

11. 260. X 10. [2600]

B. Adding and Subtracting a. When adding or subtracting, add or subtract then round off the answer to the least number of decimal places

contained in the question.

b. When more that one unit is involved, convert all of the values to the largest unit used, and then add or subtract.

After, follow rule a. to round properly.

Practice: Use your calculator to complete the required calculations. Give all answers to the correct number of

significant figures.

12. 8.42 g + 3.6 g + 10.04 g [22.1g]

13. 760 km + 42.6 km [803km]

14. 6.54 mL + 10.1 mL + 4.63 mL [21.3mL]

15. 4.00 mm + 52.6 mm + 406.5 mm [463.1mm]

16. 942 m + 1.2 km [2.1km]

17. 852 mg + 1.76 g + 10.6 cg [2.72g]

18. 0.00022 L + 0.026 mL [0.00025mL]

19. 14.56 mL – 4.2 mL [10.4mL]

20. 1.26 kg – 9.5 g [1.25kg]

Success criteria

I can:

• Identify the correct number of significant figures for a given numerical value

• Be able to give answers to calculations involving adding and subtracting multiplication and division to the correct number of significant figures

SCH 3U

Factor-Label Method

Learning Goals:

• Use the units of the given and required quantities to construct calculation strategies (ie. Do not use known/memorized

formulas)

The factor label method, sometimes also called units analysis, is an approach that can be used to solve quantitative

problems without memorizing (or even knowing) the formula. Instead it requires that you analyze the UNITS in the

question.

In this method, the units for the given information must be converted into the units for the required answer. Units can

be treated like variables are in algebra.

Example: How many mL of milk if you have 2.5 L?

Note: Units can be treated like variables are in algebra.

Just as the variable b cancels out in this calculation:

So do the L in this one:

Example: How many mL of ethanol are present in 40.0g? (D = 0.80 g/mL)

Practice:

The most important thing is to show your work. The answer is less important at this stage in your learning!

a) If 1.00$US = 1.45$Can and 1.00$Can = 5.00pesos, how many $US can you get for 30.0pesos?

b) 1 crate holds 120 oranges and 1 truck hold 500 crates. How many trucks are required to transport 280 000

oranges?

c) How many seconds are in 3.5hours?

d) Light moves at 3.0 x 108m/s. How many hours does it take light to travel 1.0 x 1016km?

e) Light moves at 3.0 x 108m/s. How many km does is travel in one year?

Success criteria

I can:

• Use the factor-label method to carry out calculations involving multiplication or division

• Be able to give answers to calculations involving multiplication and division to the correct number of significant digits

SCH 3U

Average Atomic Mass

In nature, there is a mixture of the isotopes of a certain element. Look at the periodic table. Find the mass of chlorine.

Why is it not a whole number?

We can calculate the average atomic mass if you know the abundance of each isotope and its atomic mass using the

equation:

.%%

etcBof

abundance

Bisotopeof

massatomic

Aof

abundance

Aisotopeof

massatomic

mass

atomic

average

++=

For example: The relative abundance of 𝐶𝑙 1735 (also called Cl-35) is 75.76% and the relative abundance of 𝐶𝑙17

37 (Cl-37) is

24.24%. Calculate the average atomic mass of chlorine.

• Determine the average atomic mass, to the correct number of significant digits, of an element given the percentage abundance of each of its isotopes

• Give answers with the correct units

Learning Goals:

• Use appropriate terminology related to average atomic mass including isotopes, percentage abundance, molar mass • Be able to calculate an average atomic mass of an element based on the relative amounts (percentage abundance) of its

isotopes

Success Criteria

I can:

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Introduction to the Mole

Learning Goals:

• Use appropriate terminology related to the mole and molar mass including mole, Avogadro’s number, molar mass • To recognize the significance of a molar quantity and relate it to the number of particles and mass of particles in a given sample • To, through inquiry, establish the value of Avogadro’s number • Use the factor-label method (ie no formulas) to carry out calculations that convert between mole, mass and number of particles

(atoms, molecules) for a pure substance

Activity

substance Element (E) or

compound (C)?

Use the Periodic Table to determine the mass of one atom

(element) or one molecule (compound)

(u)

Use the scale to measure

the mass of 1 mole of the substance

(g)

What do you

notice?

Determine the number of atoms/molecules in the sample

given that

1u = 1.660 540 2 x 10-24 g

Zn

Al

NaCl

glucose C6H12O6

H2O

A mole is _________________________________ of anything.

Avogadro’s number is:____________________________

Written in standard form this is:

A mole of eggs is:

A mole of atoms is:

A mole of molecules is:

Molar mass is the mass of __________ mole of a substance.

To find molar mass, _________ the masses of the atoms together and give it the unit ___________.

For example: Find the molar mass of (NH4)2SO4:

• Explain the concepts of mole and molar mass

• Recognize a molar quantity

• Use Avogadro’s number and molar mass to convert between mass, moles and number of particles, giving answers to the correct number of significant digits with correct unit

Success Criteria I can:

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More on the Mole – Calculations

Calculations with the mole and Avogadro’s number

Example #1: How many molecules of CO2 are there in 4.56 moles of CO2?

Example #2: How many moles of water are 5.87 x 1022 molecules?

Example #3: How many atoms of carbon are there in 1.23 moles of C6H12O6?

Molar Mass Calculations

Example #1: What is the mass of 2.34 moles of carbon?

Example #2: How many moles of magnesium in 4.61 g of Mg?

Example #3: How many moles are 5.69 g of NaOH?

Example #4: How many grams are 3.5 moles of copper (II) chloride?

Example #5: How many atoms of lithium are in 1.00 g of Li?

Example #6: What is the mass of 3.45 x 1022 atoms of U?

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The Law of Definite Proportions and Percent Composition

Learning Goals:

• Use appropriate terminology related to chemical formulae including percent composition, empirical formula, empirical mass, molar mass, molecular formula

• Calculate percent composition

• Determine empirical and molecular formulae

The law of definite proportions states: the elements in a chemical compound are always present in the same

proportion by mass.

For example: CO ______________ and CO2 _______________ both contain only carbon and oxygen. They have very

different properties. Pure CO always contains 42.88% carbon and 57.12% oxygen by mass while pure

CO2 contains 27.29% carbon and 72.71% oxygen.

% composition = __mass of element____ x 100%

total mass of compound

Example: Determine the % composition of hydrogen and oxygen in hydrogen peroxide (H2O2).

Example: What is the percentage of water in copper (II) sulphate pentahydrate CuSO4. 5H2O?

Empirical Formula

The empirical formula is the _____________________ whole number _______________ of atoms in a substance.

Example: CH2, C2H4, and C4H8 all contain the same _____________ of hydrogen compared to carbon

The simplest whole number ratio of carbon to hydrogen in all these species is _________.

How do you determine the empirical formula?

1. Assume 100% is equal to __________ → change each of the substances percentages to a ________ in grams.

2. Determine the number of __________ of each → use ____________________

3. Express the moles in a chemical formula → put number of moles as __________________ to the right of the

element

4. Simplify to simplest ratio → divide each number of moles by the __________________________

Example: What is the empirical formula of a compound consisting of 80.0% carbon and 20.0% hydrogen?

Very important! You must make sure that you have a whole number ratio! If the ratio is between 0.2 and 0.8, then you

must multiply all numbers in the ratio so that all are whole numbers.

Example: What is the empirical formula of a compound containing 81.8% carbon and 18.2% hydrogen?

Hints:

Molecular Formula:

The molecular formula is just a _____________ of the empirical formula. It tells you the ____________ number of

atoms of each element in the molecule.

molecular formula = n x empirical formula where n = molar mass empirical mass Example: A molecule has an empirical formula of HO and a molar mass of 34.0g. What is the molecular formula?

Practice Questions:

1. Calculate the empirical formula for the compounds whose composition is listed below: a) C 15.8%, S 84.2% b) S 47.5%, Cl 52.5% c) Ag 70.1%, N 9.1%, O 20.8% d) K 40.2%, Cr 26.9%, O 32.9% e) Na 29.1%, S 40.5%, O 30.4% f) Na 19.2%, H 0.8%, S 26.1%, O 53.3% g) K 14.0%, Al 9.7%, Si 30.2%, O 46.1%

2. Upon analysis a blue compound showed a copper content of 47.3%. Is it CuBr2, CuSO4, CuCl2 or Cu(NO3)2? 3. The simplest formula for glucose is CH2O and its molar mass is 360g. What is the molecular formula? 4. Determine the molecular formula for each compound given their % composition and molar mass (MM):

a) C 64.9%, H 13.5%, O 21.6%; MM = 74g b) C 52.2%, H13.0%, O 34.8%; MM = 46g c) Al 20.2%, Cl 79.8%; MM = 267g d) B 40.3%, N 52.2%, H 7.5%; MM = 80g

5. A compound was found to have: C 58.5%, H 4.1%, N 11.4% and O 26.0% and molar mass of approximately 125g. Determine: a) the molecular formula b) the exact molar mass of the compound Success Criteria I can:

• Determine the molecular formula given a molar mass and empirical formula • Determine percent composition of a pure substance given the chemical formula

• Determine the empirical formula of a pure substance given the percent composition of its constituent elements

SCH3U Caffe Latte

2003 Milk Calendar: Easy Cooking with Christine Cushing. Ontario: Daily Farmers of Ontario, 2002.

Use the recipe provided above to solve the following problems.

Helpful Hint: Use decimals when performing the calculations.

Part A: The Equation

Write a “chemical” equation to represent this recipe. The ingredients are the reactants and the Latte is the product. Be

sure to include coefficients indicating the amount of each ingredient needed and the amount of product formed.

Part B: Varying the recipe

When the quantity of one ingredient is changed in a recipe, all other ingredients must also be adjusted in order to yield

the same delicious product. In each of the following questions, however, it is only necessary to determining the amount

of ingredient or product asked for in the problem.

1. If 4 cups of milk are used, how many cups of ice cubes are required?

2. How many teaspoons of vanilla are needed if the amount of ice cream is increased to 4 ½ cups?

3. If you want to serve 8 delicious lattes on a hot summer day, how many cups of coffee do you need to make?

4. There are 1 ½ cups of coffee left in the pot from the morning brew. How many teaspoons of vanilla are required

to make the maximum number of lattes?

5. You have exactly 6 cups of milk that you need to use up before it expires. How many servings of lattes can you

make?

Recipe:

2 cups milk

½ cup ice cubes

1 ½ cups vanilla or coffee ice cream

¾ cup strong coffee

1 tsp vanilla

In blender, combine milk, ice cream, and desired flavouring; puree until smooth and frosty.

Serve in a chilled glass.

Preparation Time: 5 minutes

Yield: 2 servings

Part C: Converting for Comparison.

When the units in our equation do not match the units in the problem, conversions are necessary. To help you out:

250 mL = 1 cup

4 mL = 1 tsp

1 mole = 6.02 x 10 23

1. How many lattes can be made with 750mL of ice cream?

2. How much milk is required to combine with 15mL of vanilla?

Part D: Limiting Ingredient

There are times when the amount of product that can be produced is limited by the quantity of one ingredient. For

example, if you have an unlimited supply of milk, ice cream and coffee but only 1 tsp of vanilla left in the cupboard, you

can only make 2 serving of lattes.

1. There are 6 cups of milk in the refrigerator and 3 cups of coffee.

a) Identify the limiting ingredient. Provide mathematical proof.

b) Assuming an unlimited supply of ice, ice cream and vanilla, how many servings of lattes can be made with the

specified amount of milk and coffee?

c) How much of the excess ingredient remains?

SCH 3U

Stoichiometry – Quantities in Chemical Reactions

Learning goals:

➢ Develop a method for calculating, using the factor-label method (ie no formulas), amounts of either reactant or product in a

chemical reaction given the quantity of one other substance (either reactant or product) in a known chemical reaction

➢ Use appropriate terminology related to the mole and molar mass including mole, Avogadro’s number, molar mass, mole ratio

Stoichiometry is the field of chemistry that is concerned with the relative __________________ of reactants and

products in __________________________________.

Stoichiometry is based on the ______________________________________ as indicated by a

______________________ chemical equation.

Example 1: Given: Fe(s) + Cl2(g) → FeCl3(g)

➢ Ensure this equation is obeying the Law of Conservation of Mass

➢ If 4 Fe atoms react, how many FeCl3 are made? _____________________

➢ How many Fe atoms are needed to react with 30 Cl2 molecules? _______________________

Example 2: Given: Sr(s) + N2(g) → Sr3N2(s)

➢ Ensure this equation is obeying the Law of Conservation of Mass

➢ If 6 moles of Sr atoms react, how many moles of Sr3N2 are made? _____________________

➢ How many moles of Sr atoms are needed to react with 5 moles of N2 mol? _____________________

Mole ratio: The mole is the ratio of moles of atoms/molecules that react. This can be found be looking at the

coefficients of the balanced chemical equation.

Example: 2C2H6(g) + 7O2(g) → 4CO2(g) + 6H2O(l)

Some of the mole ratios that can be written for this reaction are:

Example 3: Given: C8H14(l) + 2H2(g) → C8H18(l)

➢ Are exactly 2 grams of H2 needed to react fully with exactly 1 g of C8H14? Explain.

The balanced chemical equation gives us information directly about ratios of __________________ and/or _________.

It does not give us direct information about _________________. To find information about masses, we must first find

the number of _____________ and then calculate the mass using the ___________________.

Example 4: Given: C8H14(l) + 2H2(g) → C8H18(l)

➢ How many grams of H2 are needed to react with 10.0 g of C8H14?

Stoichiometric Calculations:

Steps: 1. Write the balanced chemical equation

2. Look at the balanced chemical equation to find the mole ratio of the substances involved

3. If you are given mass or number of particles, convert to moles.

4. Use the mole ratio to convert from moles of one substance to moles of the other (the key step in stoich!)

5. If you are asked for mass or number of particles, do so.

Note: All stoichiometry involves steps 1, 2 and 4. Steps 3 and 5 are used only if necessary.

Example:

Ammonium sulfate, (NH4)2SO4, is used as a source of nitrogen in some fertilizers. It reacts with sodium

hydroxide to produce sodium sulfate, water and ammonia, NH3. What mass of sodium hydroxide required to

react completely with 15.4g of (NH4)2SO4?

Step 1: balanced chemical equation: (NH4)2SO4 + 2NaOH → Na2SO4 + 2H2O + 2NH3

Step 2: mole ratio from balanced chemical equation:

2 𝑚𝑜𝑙 NaOH

1 𝑚𝑜𝑙 (NH4)2SO4 or

1 𝑚𝑜𝑙 (NH4)2SO4

2 𝑚𝑜𝑙 NaOH

Step 3 – 5:

15.4 𝑔 (NH4)2SO4 x 1 𝑚𝑜𝑙 (NH4)2SO4

132.0𝑔 (NH4)2SO4 x

2 𝑚𝑜𝑙 NaOH

1 𝑚𝑜𝑙 (NH4)2SO4 x

40.0 𝑔 NaOH

1 𝑚𝑜𝑙 NaOH = 9.33 g NaOH

In-class Practice Problems:

1. Oxygen completely burns 1.56g of methane, CH4. How many moles of oxygen react?

2. Silver nitrate, 0.52g, reacts completely with sodium chloride. How many moles of silver chloride are

formed?

given Step 3:

convert to mol of given

Step 4: use mole ratio to

convert to moles of required

Step 5: convert to

mass of required

3. If 5.0 mol of copper (II) chloride react with sodium in the following reaction, how many atoms of

copper are formed?

CuCl2 + 2Na → Cu + 2NaCl

4. How many grams of oxygen gas are produced when 0.62 mol of mercury (II) oxide decomposes?

2HgO → 2Hg + O2

5. How many molecules of oxygen are needed to completely react with 3.2g of Fe to make Fe2O3?

6. What mass of glucose, C6H12O6, is needed to completely react with 25.0g of O2?

Success criteria:

I can:

• Apply the concepts of mole, molar mass, Avogadro’s number and mole ratios in chemical reactions using the factor-label method

• Predict the maximum amount or the mass of product/reactant expected in a chemical reaction to the correct number of significant digits and with the correct unit

• Clearly communicate all relevant steps in the calculation process

SCH3U

Stoichiometry Worksheet

Answer the following questions at the bottom of the sheet. Remember to do the following:

- balance your equations

- determine what is given and what is needed - convert to moles and use a mole ratio - convert back to the value asked for in the question

1) How many moles of iron(II) oxide are made from 2.4 moles of iron reacting with excess oxygen? 2) How many grams of copper are needed to react with excess chlorine gas to make 4.1g of copper(II) chloride? 3) How many grams of potassium iodide was reacted with lead(II) nitrate to produce 4.5g of lead(II) iodide? 4) How many grams of calcium chloride was reacted with excess sodium hydroxide solution to make 5.6moles of

calcium hydroxide? 5) Consider the combustion of carbon monoxide (CO) in oxygen gas:

2CO(g) + O2(g) → 2CO2(g)

a) Starting with 3.60moles of CO, calculate the number of moles of CO2 produced if there is enough oxygen gas to react with all of the CO.

b) How many grams of carbon dioxide is this? c) How many grams of oxygen are need to react?

6) When baking soda (sodium bicarbonate or sodium hydrogen carbonate NaHCO3) is heated, it releases carbon dioxide gas, which is responsible for the rising of cookies, doughnuts and bread.

a) write a balanced equation for the decomposition of the compound ( hint one of the products is Na2CO3) b) Calculate the mass of NaHCO3 required to produce 20.5g of CO2.

Answers:

1) 2.4 mol FeO 2) 1.9g 3) 3.2g KI 4) 6.2 x 102g 5a) 3.60 mol CO2 b) 158.4gCO2 c) 57.6 g O2 6b) 78.3g NaHCO3

SCH 3U

The Limiting Reactant

Learning Goals:

• Use appropriate terminology including limiting reactant, excess reactant, mole, mole ratio, product, molar mass, Avogadro’s number, theoretical yield

• To identify reactants that limit the amount of product made • Develop a method for calculating the maximum amount of product possible (theoretical yield) given a limiting reactant • Develop a method for calculating the amount of excess reactant remaining after a reaction is complete

A balanced chemical equation shows the moles ratios of the _________________ and the _________________. When

the amounts of the reactants are exactly enough to produce the amount of product called for by the mole ratio, then

the reactants are said to be in __________________________ amounts. This means that once the reaction is complete,

there are no reactants left.

For example: Given the caffe latte analogy. The balanced equation for the recipe is:

2 c milk + ½ c ice cubes + 1½ c ice cream + ¾ c coffee + 1 tsp vanilla → 2 caffe lattes

If you have exactly the amounts listed in the equation, then you can make exactly 2 caffe lattes with no ingredients left

over.

However, if you had 2 cups of milk, ½ cup of ice cubes, ¾ c of ice cream, ¾ c of coffee and 1 tsp of vanilla, then:

➢ How many lattes can you make?

➢ Which ingredient is limiting the number of lattes you can make?

➢ How much of the other ingredients are left over?

Determining the Limiting Reactant

Limiting reactant: the reactant that is ___________________ __________ ________. It determines how much

product is formed. Why?

Excess reactant: the reactant that is ____________ ____________. Once the limiting reactant is used up, no

more product is made regardless of how much excess reactant remains.

When you are given amounts of 2 or more reactants in a stoichiometric problem, you must identify the limiting

reactant. To do this, you must:

The reactant that would produce the __________________ amount of product is the limiting reactant.

Example: Lithium nitride reacts with water to form ammonia and lithium hydroxide according to the following

balanced chemical equation:

Li3N(s) + 3H2O(l) → NH3(g) + 3 LiOH(aq)

If 4.87 g of lithium nitride reacts with 5.80g of water, find the limiting reactant.

Solution:

➢ Determine the number of moles of one of the products 4.87g of Li3N can make:

4.87 𝑔 Li3N x 1 𝑚𝑜𝑙 Li3N

34.7 𝑔 Li3N x

1 𝑚𝑜𝑙 NH3

1 𝑚𝑜𝑙 Li3N = 0.140mol NH3

➢ Determine the number of moles of the same product that 5.80g of H2O can make:

5.80 𝑔 H2O x 1 𝑚𝑜𝑙 H2O

18.0 𝑔 H2O x

1 𝑚𝑜𝑙 NH3

3 𝑚𝑜𝑙 H2O = 0.107mol NH3

➢ Since _________________ produces fewer moles of product, it is the limiting reactant.

Example Problem:

White phosphorus consists of a molecule made up of four phosphorus atoms. It burns in pure oxygen to

product tetraphosphorus decaoxide according to the following equation:

P4 + 5O2 → P4O10

A 1.00g piece of solid white phosphorus is burned in a flask filled with 5.00g of oxygen gas.

a) Determine the limiting reactant.

b) How would you use the results from part a) to determine the mass of tetraphosphorus decaoxide.

c) Determine the mass of P4O10 produced in this reaction.

Steps for completing stoichiometric questions that involve limiting reactants*

1. Ensure that you have a balanced chemical equation.

2. For each reactant, use the method you learned in the last section to determine the number of ______________

of one of the ____________________.

3. The reactant that produces the ____________________ number of moles of product is the _________________

reactant.

4. Use the moles of product made by the __________________ reactant to complete the question. You may be

asked to find moles, mass or number of atoms/molecules.

*To identify a question with a limiting reactant, see if you are given quantities of 2 or more reactants; if so,

proceed as indicated. If not, then it is a standard stoichiometry problem (last note).

More examples:

1. The combustion of propane proceeds by the following reaction: C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)

If 100.0g of propane is reacted with 50.0g of oxygen:

a) Which reactant is limiting and which is excess?

b) How many grams of water will be formed?

2. Given the following reaction of iron with water: 3Fe + 4H2O → Fe3O4 + 4H2

How many grams of Fe3O4 is produced if 25.0g of Fe is reacted with 30.0g of water?

3. If 20.0g of H2 reacts with 100.0g of O2 according to the reaction: 2H2(g) + O2(g) → 2H2O(l)

Which reactant is limiting, which is excess and how grams of product is formed?

4. If 58.4g of FeCl2, 14.0g of KNO3 and 40.0g of HCl are mixed and allowed to react according to the following

equation:

3FeCl2 + KNO3 + 4HCl → 3FeCl3 + NO + 2H2O + KCl

Which reactant is limiting and how many grams of FeCl3 is produced?

Success criteria

I can:

• Identify the limiting reactant and excess reactant in a chemical reaction and predict the amount or the mass of product expected

• Explain how the outcome of a chemical reaction can be affected by manipulating the limiting and excess reactants

• Calculate the amount of excess reactant remaining after a reaction

• Clearly communicate all relevant steps in the calculation process

SCH 3U Extra Stoichiometry Practice including Limiting Reactants

1) Chlorine is used by textile manufacturers to bleach cloth. Excess chlorine is destroyed by its reaction with sodium thiosulfate, Na2S2O3:

Na2S2O3(aq) + 4Cl2(g) + 5H2O(aq) → 2NaHSO4(aq) + 8HCl(aq) a. How many moles of Na2S2O3 are needed to react with 0.12mol of Cl2? b. How many moles of HCl can form from 0.12mol of Cl2? c. How many moles of H2O are required for the reaction of 0.12mol of Cl2? d. How many moles of H2O react if 0.24mol HCl is formed?

2) The incandescent white of a fireworks display is caused by the reaction of phosphorous with O2 to give P4O10.

a. Write the balanced chemical equation for the reaction. b. How many grams of O2 are needed to combine with 6.85g of P? c. How many grams of P4O10 can be made from 8.00g of O2? d. How many grams of P are needed to make 7.46g P4O10?

3) In dilute nitric acid, HNO3, copper metal dissolves according to the following equation:

3Cu(s) + 8HNO3(aq) → 3Cu(NO3)2(aq) + 2NO(g) + 4H2O(aq) How many grams of HNO3 are needed to react with 11.45g of Cu?

4) The reaction of powdered aluminum and iron(II)oxide, 2Al(s) + Fe2O3(s) → Al2O3(s) + 2Fe(l)

produces so much heat the iron that forms is molten. Because of this, railroads use the reaction to provide molten steel to weld steel rails together when laying track. Suppose that in one batch of reactants 4.20mol Al was mixed with 1.75mol Fe2O3.

a. Which reactant, if either, was the limiting reactant? b. Calculate the mass of iron (in grams) that can be formed from this mixture of reactants.

5) Silver nitrate, AgNO3, reacts with iron(III) chloride, FeCl3, to give silver chloride, AgCl, and iron(III) nitrate,

Fe(NO3)3. A solution containing 18.0g AgNO3 was mixed with a solution containing 32.4g FeCl3. How many grams of which reactant remains after the reaction is over?

Answers:

1) a. 0.030mol Na2S2O3 b. 0.24mol HCl c. 0.15mol H2O d. 0.15mol H2O

2) a. 4P + 5O2 → P4O10 b. 8.85g O2 c. 14.2g P4O10 d. 3.26g P

3) 30.31g HNO3

4) a. limiting reactant is Fe2O3 b. 195g Fe is formed 5) 26.7g of FeCl3 are left over

SCH3U

Determining the Amount of Excess Reactant

Sometimes with limiting reactant questions you are asked to determine the amount of excess reactant that remains. A

good way to deal with this is to construct an i-c-f (initial-change-final) chart to help you keep track of the quantities

involved.

Example: 25.0g of cesium carbonate is reacted with 20.0g of zinc fluoride. Which reactant is in excess and how

many grams of this reactant remains once the reaction is complete?

Example: 1.25g of silver nitrate is reacted with 2.00g of lithium chloride. Which reactant is in excess and how

many grams of this reactant is left at the end of the reaction?

Learning goals:

• Use appropriate terminology including limiting reactant, excess reactant, mole, mole ratio, product, molar mass, Avogadro’s number, theoretical yield

• Develop a method for calculating the amount of excess reactant remaining after a reaction is complete

Another example: 10.0g of NaCl is mixed with 5.0g of Pb(NO3)2. Which reactant is in excess and what mass of it

remains?

Practice:

1. Suppose 316.0 g aluminum sulfide reacts with 493.0 g of water. What mass of the excess reactant remains?

[265.5 g]

2. In this reaction: CaCO3 + 2HCl ---> CaCl2 + CO2 + H2O

6.088 g CaCO3 reacted with 2.852 g HCl. What mass of CaCO3 remains unreacted? [2.174 g]

3. What weight of each substance is present after 0.4500 g of P4O10 and 1.5000 g of PCl5 are reacted completely?

P4O10 + 6PCl5 ---> 10POCl3 [1.8408g of POCl3; 0.1092 g of P4O10; 0g of PCl5]

• Identify the limiting reactant and excess reactant in a chemical reaction

• Calculate the amount of excess reactant remaining after a reaction

• Clearly communicate all relevant steps in the calculation process

Success Criteria I can:

SCH 3U

Percentage Yield

Formula for Percentage Yield:

Example: Under very controlled and careful conditions, a chemist mixes 3.00g of Na metal with 5.00g of Cl2 gas. If 6.80g

of product is formed, what is the percentage yield?

Possible reasons for not achieving the theoretical yield.

• Reaction may stop short of completion so that reactants remain unreacted. • There may be competing reactions that give other products and therefore reduce the yield of the desired one. • In the process of separation and purification, some of the product is invariably lost. • There may be impurities in the reactants that do not form the desired products

Learning goals:

• Use appropriate terminology including percentage yield, theoretical yield, actual yield, error • Use the percentage yield formula

• Identify procedural errors that affect the actual yield

Success Criteria I can:

• Analyze quantitative data from an experiment and determine percentage yield to compare the actual and theoretical yields

• Calculate percentage yield

• Give reasons for loss of product

SCH3U

Success Criteria - Unit 4 - Gases

At the end of this unit, I can:

• Use appropriate terminology related to gases and atmospheric chemistry including but not limited to: standard pressure, molar volume, and ideal gas

• Use the kinetic molecular theory to explain the properties and behavior of gases in terms of types and degrees of molecular motion

• List the gases that make up the earth’s atmosphere as well as how atmospheric pressure is created.

• Convert between the different methods of measuring gas pressure – mmHg, atm, torr and converting between degrees Celcius and Kelvin scale

• Explain the quantitative relationships between pressure and volume, pressure and temperature, volume and temperature, moles and volume

• Be able to state and describe all the gas laws by name (Boyle’s law, Charles’ law, Gay-Lussac’s law, the combined gas law, Dalton’s law of partial pressures, and the ideal gas law) and as a mathematical expression as well as, given a graph, identify the law represented

• Solve quantitative problems by performing calculations based on Boyle’s law, Charles’ law, Gay-Lussac’s law, the combined gas law, Dalton’s law of partial pressures, and the ideal gas law

• Explain Avogadro’s hypothesis and how his contribution to the gas laws has increased our understanding of the chemical reactions of gases

• Use stoichiometry to solve the problems related to chemical reactions involving gases (ex. Problems involving moles, number of atoms, number of molecules, mass and volume)

SCH 3U

Gas Laws Simulation Activity

1. Go to: http://phet.colorado.edu/en/simulation/gas-properties and run now!

2. By adjusting the settings investigate and answer the following questions:

➢ As the number of gas particles increases, what happens to the pressure?

➢ If you keep the temperature and the number of particles the same, what happens to the pressure as the

volume decreases? Increases?

➢ If you keep the volume and the number of particles the same, what happens to the pressure as the

temperature increases? Decreases?

➢ If you wish to increase the pressure but keep the temperature and volume the same, what can one do?

➢ As you change the temperature, what happens to the speed of the particles?

➢ How does the speed of the heavier particles compare to the speed of the lighter particles?

SCH3U

Introduction to Gases

Learning goals:

• Understand that the three states of matter have different properties, such as the ability to flow or be

compressed, based on the forces that exist between the atoms and molecules

• Use Kinetic Molecular Theory to describe the behavior of gases in terms of the type of molecular motion

• Identify the components of the Earth’s atmosphere and understand how atmospheric pressure is calculated

• Explain pressure of a gas on a molecular level and convert between different units of pressure

Review: Properties of gases

Some common, basic physical (observable) properties (circle correct one):

Definite/indefinite volume

Definite/indefinite shape

Highly compressible/incompressible

Process of solid to liquid to gas is exothermic/endothermic

External pressure increase/decrease causes the volume to increase

Temperature increase/decrease causes the volume to decrease

Density is usually highest/lowest of the three states of matter

Gases also have a range of other physical properties: may be clear or colourful ____________________, odorless or

smelly _________________, have higher or lower densities ___________________, be gases or vapors

_______________________.

Gases have a wide variety of chemical properties: may be necessary for life or poisonous ____________________,

reactive or inert ___________________.

Conceptual properties of gases:

The distances between gas molecules is ______________________, speed and kinetic energy content is the

___________________ of all states of matter. Pressure of a gas is cause by

___________________________________________________________. Attraction between the molecules is

____________________________________________ and intermolecular forces are ___________________________.

Despite the wide differences in chemical properties of gases, they more or less obey the same set of ________________

laws → the gas laws. The ______________________________________ theory (KMT) accounts for the nonpolar and

low molecular weight gases (eg. H2, N2, O2, He, Ar, CO, CO2, CH4, N2O, NO2) which come close to ideal gas behavior at

______________ pressure and ______________ temperature. Deviations from ideal behavior increase as pressure

increases and temperature decreases.

The kinetic molecular theory accounts for the properties of an ideal gas in terms of a set of assumptions about the

nature of gases:

1. A gas consists of an extremely large number of very tiny particles that are in constant, ________________

motion. This is referred to as ____________________ motion.

2. The gas particles occupy a relatively _________________ volume, negligible to the volume of their container.

3. The gas molecules have ___________________ attractive forces for one another because of their

_______________________ distances apart.

4. Collisions of the gas particles with each other and with the walls of the container are perfectly ______________.

5. The average kinetic energy of the gas molecules is proportional to __________________ temperature.

Our atmosphere:

The atmosphere surrounding Earth contains:

N2 ____________, O2 _____________,

Ar _____________, CO2____________,

as well as many other gases in very small

quantities.

STP and SATP:

Standard Temperature and Pressure (STP) is: ______________________ and ______________________

Standard Ambient Temperature and Pressure (SATP) is: ______________________ and ______________________

Atmospheric pressure:

Pressure = Force/area Units: Newtons/m2

Atmospheric pressure is the ______________________________ exerted by air on all objects.

• 1 standard atmospheric pressure at sea level (1atm) is the force exerted by:

• 140km column of atmosphere or

• 760mm column of mercury (1mmHg = 1torr) or

• 10.3m column of water

on an area of 1m2.

• Therefore:

1atm = 760 mmHg = 760 torr = 101.3kPa

Example: Convert 1.3atm to mmHg, pKa and torr

.

Gas laws:

To describe the state or condition of a gas, it is important to specify four variables:

• ___________________ (kPa)

• ___________________ (K)

• ___________________ (L)

• ___________________ (mol)

The relationship between these four variables is described in a set of physical laws… THE GAS LAWS.

The gas laws that we will study: Prediction of relationship

1. Boyle’s Law [P,V change; T,mol constant]

2. Charles’ Law [V,T change; P,mol constant]

3. Gay-Lussac’s Law [P,T change; V,mol constant]

Also:

4. The combined gas law [P,V,T change; mol constant]

5. Dalton’s Law of Partial Pressures [relates P to the partial P of each gas in a sample]

6. Avogadro’s Principle [V,mol change; P,T constant]

7. Ideal Gas Law (equation) [P,V,T,mol change]

Success Criteria

I can:

• Describe the physical properties of gases and explain how the attraction between gas molecules contributes to

these properties

• Use KMT to explain the properties and behavior of gases in terms of types and degrees of molecular motion

• List the gases that make up the Earth’s atmosphere

• Describe how atmospheric pressure is created and names the various units for describing atmospheric pressure

• Convert between the different units of measuring gas pressure – kPa, mmHg, torr, atm

SCH3U

Charles Law: Inquiry

Learning goals: • Understand that absolute zero on the Kelvin scale is the theoretical temperature at which entities have to kinetic energy so a

gas exerts no pressure and has no volume

• Understand that the volume of a gas is directly proportional to its temperature in kelvin, as long as the pressure and amount of gas remain constant

In the study of the behavior of gases, there are four variables that are interrelated: number of molecules of gas, volume,

pressure and temperature. In order to study the relationship between any two of these variables, the other two must

be held constant.

Academic task:

To graph and interpret experimental data measuring the relationship between the volume and temperature of a fixed

mass of gas at constant pressure.

Materials:

One lab sheet, laptop with excel or google sheets for graphing, calculator

Directions:

1. Your teacher will assign one experimental data set to you. Record the data set assigned.

2. Using the data set assigned to you, fill in the following table:

T(C) V(mL) V/T (mL/C)

3. Using the data set assigned to you, plot a graph of volume vs. temperature, with the horizontal axis running from –350oC to 100oC, and the vertical axis running from 0 to 30 mL. Extrapolate the curve to zero volume and obtain a value for the x-intercept, identified as To.

To = ___________________

0C 20C 40C 60C 80C

DATA

SETS

A V(mL) 16.0 17.2 17.9 19.4 20.1

B V(mL) 20.6 22.5 24.1 25.2 27.1

C V(mL) 22.5 24.2 25.7 27.0 28.5

D V(mL) 19.0 19.9 21.8 22.7 24.3

E V(mL) 13.0 13.6 14.9 15.6 16.9

F V(mL) 23.5 25.2 26.7 28.0 29.5

G V(mL) 15.0 16.2 16.9 18.4 19.1

H V(mL) 20.0 20.9 22.8 23.7 25.3

4. Sketch your graph here. Make sure you label your axes, include units and indicate the x-intercept at To.

5. Using the data set assigned to you, the result you obtained for To and the equation: Tx = T – To to set up and fill in the following chart:

T(C) Tx(C) V(mL) V/Tx (mL/?)

6. Answer the following questions: a. Explain the significance of extrapolating the temperature to zero volume.

b. What is your x-intercept? _______________ What is the generally accepted value for this intercept? ___________________

What is it called? _______________________________

Calculate the experimental error, given that: experimental error = |experimental value – accepted value| x 100% accepted value

c. Compare the results obtained for the volume/temperature ratios for the two different tables. What does this suggest with respect to any future calculations?

Kelvin Scale Both the Fahrenheit and Celsius scales are ______________. They are based on the ______________ and

________________ temperatures of water. The Kelvin scale is ______________. The zero corresponds to the

______________ temperature _________________.

The formula for calculating the temperature in Kelvin (K) is:

Example: Liquid oxygen is used in rockets. Its boiling point is -183oC. What is this temperature in K?

Charles’ Law Charles’ Law states that:

This can be given the formula: This relationship has the following graph:

Example: Anesthetic gas is normally given to a patient when the room temperature is 20.0oC and the patient’s

body temperature is 37.0oC. What would this temperature change do to 1600.0 mL of gas if the pressure

and mass stay constant?

Success criteria:

I am able to:

• Calculate the temperature in Kelvin when given the temperature in degrees Celsius

• Explain the quantitative relationship between the volume and temperature of a gas

• State and describe Charles’ law by name and as a mathematical expression as well as, given a graph, identify the law represented

• Sketch a graph that accurately depicts the relationship between the temperature and volume of a gas

• Solve quantitative problems by performing calculations based on Charles’ law

SCH3U

Boyle’s Law, Gay-Lussac’s Law and Combined Gas Law

Learning goals:

• Explain, in words, the quantitative relationships between pressure and volume, pressure and temperature, volume and temperature, moles and volume of a gas

• State and describe Boyle’s law, Gay-Lussac’s law and the combined gas law by name and as a mathematical expression as well as, given a graph, identify the law represented

• Sketch a graph that accurately depicts the relationship between the pressure and volume of a gas and the pressure and temperature of a gas

• Solve quantitative problems by performing calculations based on Boyle’s law, Gay-Lussac’s law and the combined gas law

Boyle’s Law Inquiry

Consider the following data from an investigation in which a fixed amount of gas is trapped in a J-tube and the volume of

the gas is altered and the pressure recorded:

Compression of a fixed sample of air (temperature and moles kept constant)

Volume (cm3) Pressure (kPa) PV (cm3kPa) 1/P (kPa-1)

50.0 66.7

31.0 107.5

21.0 158.6

15.0 222.3

Sketch graphs for the following and describe your observations:

Pressure 1/P PV

Volume Volume Pressure

Comments:

Boyle’s Law states that the ________________ of a fixed ___________ of gas at constant _________________ is

____________________ proportional to the _________________ of the gas (i.e. As the pressure increases, the volume

________________ by a reciprocal factor).

Formula:

Example: If a given mass of gas occupies a volume of 8.4L at a pressure of 101 kPa, what is its volume at a

pressure of 180 kPa at the same temperature?

Gay-Lussac's Law Gay-Lussac's Law states that: Equation: Graph: Example: A cylinder of the highly flammable fuel acetylene should be stored at 22oC and 62 kPa. By how

much would the pressure change if the temperature rose to 36oC?

The Combined Gas Law Combining Charles, Boyle’s and Gay-Lussac’s Laws we get the Combined Gas Law: Example: A sample of gas at SATP in a 3.5L vessel is heated to 45oC and 5.0L. What is the final pressure of

the gas?

Success criteria: I can: • Explain, in words, the quantitative relationships between pressure and volume, pressure and temperature, volume and

temperature, moles and volume of a gas

• State and describe Boyle’s law, Gay-Lussac’s law and the combined gas law by name and as a mathematical expression as well as, given a graph, identify the law represented

• Sketch a graph that accurately depicts the relationship between the pressure and volume of a gas and the pressure and temperature of a gas

• Solve quantitative problems by performing calculations based on Boyle’s law, Gay-Lussac’s law and the combined gas law

SCH3U

Dalton’s Law of Partial Pressures

Learning goals:

• Understand that Dalton’s Law of Partial Pressures state’s that the total pressure of a mixture of gases is equal to the sum of

the partial pressures of the individual gases.

Dalton’s Law states that:

Total pressure of a mixture of gasses A, B, C…

To determine the partial pressure of each gas in a mixture, one must first consider the mole fraction () of that gas in

the mixture.

Mole fraction of gas A: 𝑋𝐴 = 𝑚𝑜𝑙 𝐴

𝑡𝑜𝑡𝑎𝑙 𝑚𝑜𝑙

To calculate the partial pressure of that gas use the following formula:

Pressure of gas A = A x total pressure

Example: A vessel contains 2.50 mol of O2, 0.50 mol of N2 and 1.00 mol of CO2. The total pressure of the gas

mixture is 200.0kPa. Find the partial pressure exerted by O2 in the mixture.

Application of Dalton’s Law of Partial Pressures:

Many gases prepared in school laboratories are collected by a

_________________________ displacement of water. This method is used for gases

that are not soluble in water. Some water vapor is inevitably mixed with the

collected gas.

Total pressure exerted by wet gas =

The partial pressure of water vapor at different

temperatures is known and can be found on a

table:

Example: 400.0mL of carbon dioxide gas is

collected over water at 27oC and at 100.0kPa.

a) Find the pressure exerted by dry CO2.

b) Calculate the volume of dry CO2 collected at STP.

Success criteria:

I can:

• State and describe Dalton’s law of partial pressures by name and as a mathematical expression

• Solve quantitative problems by performing calculations based on Dalton’s law of partial pressures

• Perform an inquiry based lab experiment to collect gas over water and use their data to calculate the molar volume of a gas

SCH3U

Avogadro’s Law and Ideal Gas Law

Learning goals:

• Understand that Avogadro’s Law explains that the volume of a gas is directly proportional to the amount of the gas when temperature and pressure of the gas remain constant

• Understand that the molar volume of a gas in the volume that one mole of a gas occupies at a specific temperature and pressure • Understand that ideal gases do not condense to a liquid when cooled, have no volume and do not attract each other • Understand that the ideal gas law is a combination of Charles’ Law, Boyle’s Law and Avogadro’s Law that describes the

behaviour of ideal gases

Avogadro’s law (also called Avogadro’s hypothesis) states:

This means that one mole of a gas occupies the same _________________ as one mole of a different gas, at the same

temperature and pressure.

The volume of one mole of a gas is called its _________________________ and is expressed in _____________.

The molar volume of any gas at STP is _________________.

Example 1: What is the volume of 3.2 mol of O2 at STP?

Example 2: CH4 reacts with oxygen in the following reaction:

CH4 + 2O2 → CO2 + 2H2O

If 4.5L of O2 reacts, what volume of CO2 is produced (assuming the pressure and temperature remain

constant)?

Ideal Gas Law

Because of Avogadro’s law, we are able to relate the quantity (moles) of a gas to its volume and from there also to

temperature and pressure for ideal gases.

Characteristics of “ideal gases” are:

1. composition:

2. volume:

3. attractive forces:

4. movement:

The Ideal gas law allows us to able to relate the quantity (moles) of a gas to its volume and from there also to

temperature and pressure using the following equation:

PV = nRT where R = gas constant

=

=

Example 1: a) A gas cylinder with a capacity of 105L contains helium at a pressure of 6.70 x 103 kPa and a

temperature of 27oC. Calculate the mass of helium gas in the cylinder.

b) Find the density of this gas.

Example 2: Find the volume of 0.48mol of a gas at 32oC and 99.6kPa.

Example 3: Find the molar mass of a gas if 8.0g of it occupies 5.8L at 18oC and 101.3kPa.

Example 4: Benzene consists of 92.24% of C and 7.76% of H. When a sample with a mass of 15.62g was placed in a

container with a volume of 3.78L and heated to 110oC, the benzene vaporized and created a pressure of

168.4kPa. What is the molecular formula of benzene?

hints: 1. Find moles; use moles you calculated and mass given to find the molar mass of the gas

2. Find the empirical formula of benzene

3. Use answers to 1. and 2. To find the molecular formula

Extra Practice problems:

i) As geologists study the area where an ancient marsh was located, they discover an unknown gas seeping

from the ground. They collect a sample of the gas and take it to a lab for analysis. Lab technicians find that

the gas is made up of 80.0% carbon and 20.0% hydrogen. They also find that a 4.60g sample occupies a

volume of 2.50L at 1.5atm and 25.0oC. What is the molecular formula of this gas? [C2H6]

ii) A gaseous compound contains 93.1% carbon and 7.69% hydrogen by mass. 4.35g of the gas occupies 4.16L

at 22.0oC and 738 torr. Determine the molecular formula of the gas. [C2H2]

Success criteria: I can: • Explain Avogadro’s hypothesis and how his contribution to the gas laws has increased our understanding of the chemical

reactions of gases

• State and describe the ideal gas law by name and as a mathematical expression

• Solve quantitative problems by performing calculations based on Avogadro’s hypothesis and the ideal gas law

SCH 3U

Gas Stoichiometry

Learning goals:

• Use stoichiometry to solve problems related to chemical reactions involving gases • Understand how to use quantities in chemical reactions to solve problems involving all of the gas laws

Example 1: What volume of carbon dioxide (at STP) is produced when 25.0 g of propane, C3H8, is burned in an

excess of oxygen?

Example 2: What mass of magnesium is needed to react with an excess of hydrochloric acid to produce 6.5L of

hydrogen at STP?

Example 3: How many litres of hydrogen gas, measured at 23oC and 103kPa, can be obtained by the reaction of

75.0g of aluminum with 12.0mol of sulfuric acid?

2Al(s) + 3H2SO4(aq) → Al2(SO4)3(aq) + 3H2(g)

Example 4: What mass of aluminum is required to produce 36.3L of hydrogen, collected over water at 22oC and

103.0kPa, according to the following equation?

2Al(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2(g)

Success criteria:

I can:

• Write out balanced chemical equations for reactions involving gases

• Use mole ratios to solve problems involving moles, numbers of atoms, numbers of molecules, mass and volume using appropriate IUPAC symbols

• Calculate quantities of reactants and/or products in chemical reactions involving gases and use these quantities to solve problems involving the various gas laws learned in this unit

SCH3U Unit 5 - Solutions and Solubility

You can use this chart to keep track of your progress in this unit. As you complete the work, indicate the date in the spaces provided. Answers sheets will be provided, that you should check and then you can move onto the next topic.

Please note that this unit is designed for you to complete the work individually as indicated. Do not just copy the answer sheets. You will understand the material far better by following all of the instructions on the handouts.

All homework should be complete and corrected before moving to the next topic. All handouts as well as videos and animations will also be posted to our course’s weebly.

Formative quizzes are to be written at certain stages of your progress as indicated in the chart. Please contact your teacher who will forward a copy to you when you are ready. These are to be written at home and submitted either through pdf, photograph or screen-shot.

On-line quizzes will be held at the end of each day on the topics indicated by the completion date.

Topic Completion date

time to complete Unit Component

1

approx.

2.5 hours

Solutions and Solubility – Important Review from Unit 1

1 p. 375 #1-3

1 Solutions and Solubility

1 p. 381 #1,5,8,9

1 Read Tutorial 1 on p. 383-384 and p. p384 #1

1 p. 389 #2-10,14

1 p. 397 #1,4,6,10 & p. 394 #1-3

1 Formative quiz on Topic 1

2

approx. 2.5 hours

Reaction of Ions in Solution

2 p. 427 #1 & p. 428 #1-6

2 BLM 9-5: Answer 1-5 on a separate page

2 Qualitative Analysis

2 Read Tutorial on p. 438 and do p. 439 #1a,b

2 Formative quiz on Topic 2

3

approx. 2.5 hours

Amount concentration

3 p. 400 #1-4

3 Dilutions

3 p. 405 #1-4

3 Concentration of Ions in Solution

3 p. 448 #1-3

3 Concentration of Consumer Products

3 p. 411 #2-7

3 Formative quiz on Topic 3

4 approx. 2 hours

Stoichiometry of Solutions

4 p. 447 #1,2 & p. 449 #1-9

4 Formative quiz on topic 4

5

approx. 4 hours

Acids and Bases

5 P. 469 #2-7

5 P. 475 #4-7,9-11

5 p. 484 #1-3 & 485 #9-13

5 Lab – Virtual Titration

5 Formative quiz on Topic 5

Review Morning

Test summative

SCH3U

Success Criteria: Solutions and Solubility

By the end of the unit students should be able to…

Solubility and solutions :

• use appropriate terminology related to aqueous solutions and solubility including but not limited to:

concentration, solubility, precipitates, ionization, dissociation, pH, dilute solute and solvent

• describe what is a solution and all the components that make up a solution ( solute, solvent)

• describe the properties of water (ex. Polarity, hydrogen bonding) and explain why these properties make water

a good solvent

• list and explain all the different types of solutions

• explain why and how things dissolve

• explain the formation of solutions that are produced by dissolving ionic and molecular compounds (ex. Salt and

oxygen) in water and for solutions that are produced by dissolving non-polar solutes in non-polar solvents (ex.

Grease in vegetable oil)

• list and explain the factors that affect solubility and the rate of dissolving

• explain the effects of changes in temperature and pressure on the solubility of solids, liquids and gases

• read and interpret solubility curves

• plan an investigation to analyse qualitative and quantitative properties of solutions

Concentration:

• calculate molarity of compounds and ions

• use the formula C1V1=C2V2

• prepare solutions of a given concentration by dissolving a solid solute in a solvent and making dilutions using

C1V1=C2V2

Stoichiometry in Solutions:

• write balanced total ionic equations and net ionic equations

• Qualitative analysis of solutions based on solubility of compounds

• Use stoichiometry to solve problems involving solutions and solubility

Acids and Bases:

• Explain the Arrhenius theory of acids and bases

• Identify acids, bases and their conjugates in an equation

• Explain the difference between a strong and weak acid and between a strong and weak bases in terms of degree

of ionization

• Calculate pH, pOH and concentrations of acids and bases from pH and pOH

• Use stoichiometry to solve problems involving acids and bases

• Determine the concentration of an acid or a base in a solution (ex. The concentration of acetic acid in vinegar),

using the acid-base titration technique

SCH3U

Solutions and Solubility - Important Review from Unit 1

Use your notes from Term 1, your textbook (p.102-118), the internet or your own ideas to complete the following:

1. Intermolecular Forces:

❖ Intermolecular forces act ________________ molecules.

Some important types of intermolecular forces that we will be talking about in this unit on solutions are listed below.

Describe each of the following:

1. Dipole-dipole forces:

2. Hydrogen bonding:

3. LDF:

4. Ion-dipole interactions (this is new – use page 383 to fill this in)

2. The Structure of Water:

Summarize the structure of water, its properties and the significance of hydrogen bonding. Draw at least one diagram

as part of your description.

✓ Do p. 375 #1-3

A good video you may want to watch on the topic of polar molecules, water etc. is:

https://www.youtube.com/watch?v=PVL24HAesnc

SCH3U

Solutions and Solubility

Part 1 – Definitions for Common Terms and Ideas

You may use your textbook (p.374-397), the glossary, the internet, your prior knowledge and/or you own ideas to

complete the following chart:

Term Definition

Solution

Heterogeneous

Homogeneous

Solvent

Solute

Aqueous solution

Miscible

Immiscible

Alloy

Solubility

Concentration

Concentrated

solution

Dilute solution

Saturated

solution

Unsaturated

solution

Supersaturated

solution

Part 2 – Types of Solutions

You may use your textbook (p.379), the internet and/or you own ideas to complete the following chart:

Type of Solution (at room temperature)

Original State of Solute

State of Solvent Example

Gas in gas

Gas in liquid

Liquid or solid in gas

Liquid in liquid

Solid in liquid (amalgam)

Solid in gas

Solid in liquid

Solid in solid (alloy) *at time of mixing *at time of mixing

✓ Complete page 381 #1,5,8,9

Part 3 - The Dissolving Process

When a solute dissolves in a solvent the individual particles of the solute separate from the other particles of the solute and move between the spaces of the solvent particles. The solvent particles must collide with the solute particles and forces of attraction between solute and solvent particles "hold" the solute particles in the spaces. See the diagram below.

There are 3 steps to the dissolving process:

1. The solvent particles must move apart to make room for solute particles. This process requires energy

to overcome forces of attraction between solvent particles. 2. The solute particles must separate from the other solute particles. This process also requires energy to

overcome the forces of attraction between the solute particles. 3. When the solute particles move between the solvent particles new forces of attraction between solute

and solvent take are established. This process releases energy.

Part 4 – Factors that Affect the Rate of Dissolving

The rate of dissolving is: how _________________ a solute dissolves in a solvent.

The rate of dissolving depends on several factors:

• the temperature

• the amount of mixing/agitation

• the size of the solute particles (eg. powder vs. lump of sugar)

The rate of dissolving is greater at higher temperatures. Why do you think this is?

The rate of dissolving is greater with mixing and agitation. Why do you think this is?

The rate of dissolving is greater with smaller solute particles? Why do you think this is?

Part 5 – Solubility

Solubility is: how _____________________ solute dissolves in a solvent.

Watch the video: https://www.youtube.com/watch?v=0cPFx0wFuVs

Compare and contrast dissolving ionic and covalent compounds in the following chart:

Compare (commonalities) Contrast (differences)

A. Solubility of Ionic Compounds: During the hydration process, the water molecules surround the ions forming new ion-dipole interactions to replace the hydrogen bonds and ionic bonding. Not all ionic compounds dissolve in water. The following activity will investigate an important factor in determining whether an ionic compound is soluble. Activity: Look at the list of ionic compounds in the following table. Without looking at a solubility chart, predict if

the compound will dissolve in water [Hint: consider the charges on the cation and anion.] Explain your prediction. Then check your prediction with a solubility table. Write a new explanation if your prediction doesn’t match your observation. Summarize your findings.

Predict: Will this dissolve in water? (Y or N)

Explain your prediction: Observation: Explain your observation (if different from prediction):

KCl

NaNO3

MgO

CaCO3

Summary:

B. How Ionic Compounds Dissolve in Water: Using page 383 as a reference, draw a diagram to show the ionic compound NaCl dissolving in water: The type of intermolecular force exists between the separated ions of NaCl and water are ________________________. The process in which ions are surrounded by water molecules is called ____________________. Dissociation is: An ionic compound dissolving in water can be represented by a dissociation equation. An example is:

Na2SO4 (s) → 2 Na+ (aq) + SO42- (aq)

But how are water molecules able to dissolve ionic compounds? Ionic compounds are held together very strong interactions while ion-dipole interactions are generally much less strong (the energies need to break ionic bonds are in the range of 400-4000kJ/mol while ion-dipole interactions are in the range of 40-600kJ/mol). Give a possible explanation why water is able to break ionic bonds and dissolve so many ionic compounds:

✓ Read Tutorial 1 on p. 383-384 and do p. 384 #1

C. Solubility of Molecular Compounds

Activity: Fill in the empty cells in each of the following tables. Then see if you can find the rule for determining

whether a solute is going to dissolve well or poorly in a given solvent.

What do you think is the rule is for determining if a molecular compound is soluble in a given solvent?

This rule for molecular compounds is commonly summarized as:

____________ DISSOLVES ______________

This means that polar solvents can dissolve ______________ solutes while non-polar solvents can dissolve ________________ solutes.

Table 1 - Solvents

Lewis Dot Diagram Polar or Non-Polar?

Water H2O

Hexane C6H14

Table 2 - Solutes

Lewis Dot Diagram Polar or Non-Polar?

Solubility in Water

Solubility in Hexane

CH3OH

good poor

CH4

poor good

Br2

poor good

H2CO

good poor

D. How Molecular Compounds Dissolve in Water:

Using page 384 as a reference, draw a diagram to show the polar compound ethanol dissolving in water:

The most significant type of intermolecular force between ethanol and water is ______________________________.

E. Surfactants – Soaps and Detergents

Watch the video: http://www.learnnext.com/CBSE/Class-10/Science/Chemistry/Carbon-and-its-Compounds/Soaps-

and-Detergents/L-921.htm#container

Summarize how these molecules clean grease from an object. You may also wish to refer to p. 386 -387 of your

textbook.

✓ Do p. 389 #2-10,14

F. Other Factors that Affect Solubility

Temperature:

• The solubility of most solids in liquids increases with increasing temperature. Why do you think this is?

• The solubility of liquids in liquids is not affected by temperature changes.

• The solubility of gases decreases with increasing temperature. Why do you think this is?

Molecule size:

• Larger molecules are often less soluble than smaller molecules as t is more difficult it is for solvent molecules to

surround bigger molecules.

Pressure:

• Changes in pressure have no effect on the solubility of solids and liquids.

• Changes in pressure will affect the solubility of a gas in a liquid solvent. Why do you think this is?

G. Solubility Curves

These are used to determine the solubility of a solute at various

temperatures.

1. Read tutorial on p. 394

2. Answer the following using the curve shown:

Solubility of: KClO3 at 40oC: ___________________

NH3 at 20oC: ___________________

Temperature when each of the following have a solubility of 50g/100g H2O:

KNO3: _____________

KClO3: _____________

✓ Do p. 397 #1,4,6,10 and p. 394 #1-3

H. Summary: Why are some substances soluble while others are not?

When a solute dissolves in a solvent:

• The intermolecular forces between solute particles must break.

• The intermolecular forces between solvent particles must break.

• New intermolecular forces between solute and solvent particles form.

For this to happen:

• Each new interaction must be energetically more favorable (stronger) than the old interaction

and/or

• The number of new interactions is much greater.

Example: Explain why NaCl and H2CO are each soluble in water while AlN and CH4 are each not.

SCH3U

Reactions of Ions in Solution

A. Double Displacement Reactions

Recall from term 1 that for a double displacement reaction to occur, one of the following must be observed:

➢ the formation of a __________________ OR

➢ the formation of a __________________ OR

➢ the formation of a __________________

To predict the formation of a precipitate one must:

For example: silver nitrate + sodium hydroxide

If you consider what is actually changing during this reaction one can look at the states of the reactant and products:

This equation that includes all the ions and compounds is called the __________________________ equation.

However, _______________ and ______________ do not react. These are called ___________________ ions since they

are present in the solution but not actively involved in the reaction.

So the overall reaction is really just:

This equation that only includes the ions involved in the reaction and the product is referred to as the

____________________________ equation.

B. Steps for writing a net ionic equation:

1. Include only ions and compounds that have reacted. Do not include _______________________ ions.

2. Write the soluble ionic compounds as ________________. For example, write Na+(aq) and Cl- (aq) instead of

NaCl(aq).

3. Write insoluble ionic compounds as _________________, not ions. For example, since barium hydroxide is

insoluble you would write it as Ba(OH)2 (s), not Ba2+ and OH-.

4. Since covalent compounds do not produce ions in aqueous solution, write their full molecular formulas.

5. Write strong acids in their ionic form. You need to memorize these six strong acids:

➢ Hydrochloric acid ______________________

➢ Hydrobromic aicd ______________________

➢ Hydroiodic acid ______________________

➢ Sulphuric acid ______________________

➢ Nitric acid ______________________

➢ Perchloric acid ______________________

6. Always check that the net ionic equation is balanced for charges as well as for atoms.

Example: Sodium carbonate and lead (II) nitrate react. Write and balance the formula, total ionic and net ionic

equations for this reaction. Identify the spectator ions.

C. Single Displacement Reactions

Recall from term 1, for a single displacement reaction to occur:

For example: Cu(s) + AgNO3(aq) → because:

Write and balance the formula, total ionic and net ionic equations for this reaction. Identify the spectator ions.

✓ Practice: do p. 428 #1-6 and handout BLM 9-5

SCH3U

Qualitative Analysis

Define qualitative analysis:

Part A – Qualitative analysis involving differences in solubility

By carefully adding a series of anion to cation mixtures, it is possible

to determine which ions are present in an unknown mixture. For

example, you have a mixture you expect may contain lead and/or

barium ions. You could follow the scheme outlined in the chart on

the right:

Part B – Qualitative analysis involving flame colour

When salts containing certain cations are heated in a flame, they produce a distinctive colour. Watch the following

video and complete the chart:

https://www.youtube.com/watch?v=NEUbBAGw14k

cation flame colour

Li+

Na+

K+

Ca2+

Sr2+

Ba2+

Cu3+

✓ Read Tutorial 1 on p. 438 and do p.439 #1a,b\

Thought Lab Activity - Qualitative Analysis Design

Read p. 452: Investigation 9.3.1 Cation Qualitative Analysis

➢ You are designing for an unknown mixture that contains 1,2 or 3 of the possible ions

Testable Question: Which of the following ions are present in the unknown cation solution; Ba2+(aq), Ag+(aq),

and/or Zn2+(aq)?

Hypothesis:

In the space to the below, construct a flow chart

describing the qualitative analysis you plan to carry out.

Experimental Design:

Plan a sequence of steps, based on your flow chart and

using the materials listed on p.452, to determine

whether the solution you have been given contains

Ba2+(aq), Ag+(aq), and/or Zn2+(aq). Show your teacher

your design

SCH3U

Concentrations of Solutions – Molarity

Amount concentration:

Amount concentration is also called molarity. It is the most commonly used measure of concentration used by

chemists. The formula for molarity is:

𝑐 =𝑛

𝑉 where n is the number of moles of solute (in mol) and v is the volume of solution (in L)

The unit for molarity is mol/L which can be simplified to M. For example, 12.0mol/L can be written as 12.0M.

Example 1: What is the amount concentration of potassium chloride if 2.1g of potassium chloride are dissolved in

500.0mL of water?

Example 2: How many grams of barium chloride are needed to make a 1.6M solution with a volume of 300.0mL?

✓ Do p. 400 #1-4

SCH3U

Dilutions

Dilution:

During dilution, the number of moles of solute stays the same. The following equation describes the relationship

between the concentrations and volumes of the original and diluted solutions:

𝑐1𝑉1 = 𝑐2𝑉2

Why does this work? Watch video: https://www.youtube.com/watch?v=c9uJ42_7DEI

Explain in your own words:

We will now use the animation at the site: https://phet.colorado.edu/sims/html/molarity/latest/molarity_en.html

Use copper sulfate as the solute and click on “Show Values”.

1. If you add more solute and keep the volume the same, what happens to the concentration?

2. If you keep the amount of solute the same and add increase the volume (ie dilue it), what happens to the

concentration?

Example calculation: Freda needs to make 2.0L of 0.10M sulfuric acid using a stock solution that is 18M. How much

of the stock solution does she need to make the new solution with the correct concentration?

✓ Do p. 405 Practice #1-4

SCH3U

Concentration of Ions in Solutions

Imagine we have a 1.60M solution of Na2SO4(aq). Since the sodium sulfate is dissolved, the ions have separated and are

surrounded by water molecules. We will calculate the concentration of the separated ions in this solution.

Examples: Calculate the concentration (in mol/L) of chloride ions in each solution:

a) 13.2g of sdium chloride in 100.0mL of solution [ans. 2.26mol/L]

b) 38.4g of magnesium chloride in 125.0mL of solution [ans. 6.44mol/L]

c) after mixing the two solutions together [ans. 4.58mol/L]

SCH 3U

Concentrations and Consumer Products

Read p. 406-407 of your textbook.

Use Table 2 on page 411 to complete the following chart:

name abbreviation equation note: application

percentage volume/volume

use same unit for both volumes

percentage weight/volume

use g for weight(mass) and mL for volume

percentage weight/weight

use same unit for both weights (masses)

parts per million

use same unit for both weights (masses)

parts per billion

use same unit for both weights (masses)

parts per trillion

use same unit for both weights (masses)

Try these example questions. Show all work. The answers are given; if you do not get the correct answer, check the

screencast solution posted to clement.

1. An intravenous solution for a patient was prepared by dissolving 17.5 g of glucose in distilled water to make 350

mL of solution. Find the percent weight/volume (%W/V) concentration of the solution. [5.0%]

2. Health Canada’s guideline for maximum mercury content in commercial fish is 0.5 ppm. When a 1.6 kg salmon

was tested, it was found to contain 0.6 mg of mercury. Would this salmon be safe to eat? [yes; this salmon has

0.4 ppm mercury]

✓ Do p. 411 #2-7 note: for water 1.00 mL = 1.00 g

SCH3U

Stoichiometry of Solutions

A. Finding the Minimum Volume to Precipitate

Example: What is the minimum volume of 0.50M magnesium chloride required to precipitate all the silver ions in

80.0mL of 0.60M silver nitrate?

B. Finding the Mass of a Precipitated Compound

Example 1: 30.00mL of 0.0075M aqueous sodium sulfide is added to 67.4mL of 0.15M mercury (II) nitrate. What

mass of mercury (II) sulfide precipitates?

Example 2: Silver chromate, Ag2CrO4, is insoluble. It forms a brick-red precipitate. Calculate the mass of solver

chromate that forms when 75.0mL of 0.200M silver nitrate reacts with 25.0mL of 0.250M sodium

chromate.

✓ Do p. 447 #1,2 p. 448 #1-3 p. 449 #1-9

SCH3U

Acids and Bases

You may use your textbook (Chapter 10), the glossary, the internet, your prior knowledge and/or you own ideas to

complete the following handout.

Part 1 – Properties of Acids and Bases

Property Acids Bases

Taste

Electrical conductivity in solution

Feel of solution

Reactions with acid-base

indicators

Reaction with litmus paper

Reaction with active metals

Reaction with carbonate

compounds

Reaction with carbon dioxide

Part 2 - Naming Acids and Bases

Explain the conventions used to represent a general formula for an acid.

Write the names of the following acids:

HF __________________________________ HI __________________________________

HCl __________________________________ H2S __________________________________

HBr __________________________________ HCN __________________________________

Review naming oxyacids using the learning tip and the tutorial on pg. 467. Summarize here:

Write the chemical formula or name for each of the following acids:

a) hydrofluoric acid b) sulfurous acid c) hypochlorous acid d) perbromic acid e) HFO f) HIO4

✓ Do p. 467 #2-7

Part 3 – Neutralization Reactions

Distinguish between dissociation and ionization:

The ionization equation of HCl is:

The dissociation equation of NaOH is:

Overall equation for the neutralization of HCl with NaOH:

Total ionic equation for the neutralization of HCl with NaOH:

Net ionic equation for the neutralization of HCl with NaOH:

Part 4 - Acid/ Base Theories

A. Arrhenius Explain and use reactions to show what the definition of an acid and base are according to Svante Arrhenius.

B. Bronstead-Lowry

The Arrhenius theory of acids and bases works well for acids that contain a H+ ion and bases that contain OH- ions. But

we know of many compounds that can act as a base and not contain OH-ions. For example, ammonia NH3 is a common

base, but there is no OH- in the chemical formula. We must now refine our definition of acids and bases. This leads to

the Bronstead-Lowry definition of acids and bases.

Bronstead – Lowry acid: a compound from which a proton (H+ ion) can be removed.

Bronstead – Lowry base: a compound that can remove a proton (H+ ion) from an acid.

Like an Arrhenius acid, a Bronstead-Lowry acid must contain H in its formula. This means that all Arrhenius acids are

also Bronstead-Lowry acids. However, any negative ion (not just OH- ) can be a Bronstead-Lowry base.

According to this theory, there is only one requirement for an acid-base reaction. One substance must be the proton

donor or provider and the other substance must receive that proton (proton acceptor). In other words, an acid base

reaction involves the transfer of a proton.

For example:

HCl(aq) + H2O(l) → H3O+(aq) + Cl-

(aq)

HCl is acting as an acid because it is donating its proton (H+) to water.

H2O is acting as a base because it is accepting the proton from the HCl.

Identify the acid and base in the following reaction:

NH3(aq) + H2O(l) → NH4+ (aq) + OH-

(aq)

Water is the _______________________ because _____________________________________.

Ammonia is the ____________________ because_____________________________________.

Conjugate Acid-Base Pairs:

Two molecules or ions that are related by the transfer of a proton are called a conjugate acid-base pair. (Conjugate

means linked together.) The conjugate base of an acid is the particle that remains when a proton is removed from the

acid. The conjugate acid of a base is the particle that results when the base receives the proton from the acid.

For example:

HBr(aq) + H2O(l) → H3O+ (aq) + Br - (aq)

HBr is the acid because it has a proton to donate. The base, H2O is the molecule that accepts the proton from HBr. On

the right side of the arrow, the particle that has one less proton than the acid on the left hand side of the equation is the

conjugate base (Br - in this case). The particle on the right side that has one proton more than the base on the left side

is the conjugate acid of the base (H3O+).

HBr(aq) + H2O(l) → H3O+ (aq) + Br - (aq)

Practice: For the following reactions, identify the Bronstead-Lowry acid, base, conjugate-acid and conjugate-base.

a) HCN(aq) + H2O(l) → H3O+(aq) + CN-

(aq)

b) CH3COO-(aq) + H2O(l) → CH3COOH(aq) + OH-

(aq)

c) HCOOH(aq) + CN-(aq) → HCOO- (aq) + HCN(aq)

d) H2PO4- (aq) + OH-

(aq) → HPO42-

(aq) + H2O(l)

✓ Do p. 475 #4-7,9-11

Part 5 - Strong/Weak Acids and Bases

Definition Examples and uses

Strong acid

Weak acid

Strong base

Weak base

Part 6 - pH Calculations

Define pH:

pH can be calculated using the formula:

If you are given the pH, you can also calculate the concentration of your H3O+ by using the formula:

Calculate the pH of each solution given the hydronium ion concentration:

a) [H3O+] = 0.00365mol/L b) [H3O+] = 6.28 x 10-6 mol/L c) [H3O+] = 9.27 x 10 -12 mol/L

Calculate the hydronium ion concentration given the following pH values:

a) pH = 1.3 b) pH = 4.5 c) pH = 9.6 d) pH = 14

pH = - log ([H3O+])

[H3O+] = 10-pH

Part 7 - Acid-Base Titrations

Definition

Titration

Equivalence point

End point

Titration Stoichiometry

Example: 13.84 mL of hydrochloric acid, HCl(aq), just neutralizes 25.00 mL of a 0.1000 mol/L solution of sodium

hydroxide, NaOH(aq). What is the concentration of the hydrochloric acid?

Example: What volume of 0.250 mol/L sulfuric acid, H2SO4(aq), is needed to react completely with 37.2 mL of

0.650 mol/L potassium hydroxide, KOH(aq)?

✓ Do p. 484 #1-3 & p. 485 #9-13