Alberta Science 10
Unit A: Energy & Matter in Chemical Change
Ms. Terkper's Digital Classroom
Focusing Questions
"How has knowledge of the structure of matter led to other scientific advancements? How do elements combine? Can these combinations be classified and the products be predicted and quantified? Why do scientists classify chemical change, follow guidelines for nomenclature, and represent chemical change with equations?"
Program Outcomes
Outcome 1: Describe the basic particles that make up matter; evidence-based development of atomic model (Dalton, Thomson, Rutherford, Bohr); chemistry-based careers.
Outcome 2: Explain using periodic table how elements combine; IUPAC naming of ionic/molecular compounds and acids; WHMIS; solubility; molecular structure and properties.
Outcome 3: Identify and classify chemical changes; write balanced chemical equations; Lavoisier's law of conservation of mass; mole concept; Avogadro's number (6.02×1023).
Key Concepts
IUPAC Nomenclature
Atomic Models
Chemical Reactions
Conservation of Mass
Mole Concept
WHMIS Safety
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Atomic Theory
History of the Atomic Model
| Scientist | Year | Model | Evidence |
|---|---|---|---|
| Dalton | 1803 | Solid sphere ("billiard ball") | Law of definite proportions; Law of multiple proportions |
| Thomson | 1897 | "Plum pudding" — electrons embedded in positive mass | Discovery of electron via cathode ray tube |
| Rutherford | 1911 | Nuclear model — dense positive nucleus, electrons orbit far away | Gold foil experiment (alpha particle scattering) |
| Bohr | 1913 | Planetary model — electrons in fixed energy levels/shells | Hydrogen emission spectrum |
| Modern (Quantum) | 1926+ | Electron cloud/orbital model | Wave mechanics, Schrodinger equation |
Subatomic Particles
| Particle | Symbol | Charge | Location | Rel. Mass |
|---|---|---|---|---|
| Proton | p+ | +1 | Nucleus | 1 |
| Neutron | n0 | 0 | Nucleus | 1 |
| Electron | e- | -1 | Electron shells | 1/1836 |
Atomic Number, Mass Number, Isotopes
Atomic Number (Z): The number of protons in an atom. This defines the element. (e.g., Carbon's atomic number is 6, so every carbon atom has 6 protons).
Mass Number (A): The total number of protons + neutrons in the nucleus.
Isotopes: Atoms of the same element (same number of protons) but with a different number of neutrons (different mass number).
Example: Carbon Isotopes
- Carbon-12: 6 protons, 6 neutrons (Mass = 12)
- Carbon-14: 6 protons, 8 neutrons (Mass = 14)
Bohr Diagram Rules
- Nucleus contains protons and neutrons in the center.
- Electrons are placed in energy levels (shells) around the nucleus.
- Shell 1: Maximum 2 electrons
- Shell 2: Maximum 8 electrons
- Shell 3: Maximum 8 electrons
- Always fill the lowest energy level first before moving to the next.
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Periodic Table & Compounds
Periodic Table Groups and Trends
| Group / Category | Characteristics & Valence Electrons | Typical Ion Charge |
|---|---|---|
| Group 1: Alkali Metals | Highly reactive metals. 1 valence electron. | +1 |
| Group 2: Alkaline Earth Metals | Reactive metals. 2 valence electrons. | +2 |
| Groups 3–12: Transition Metals | Complex electron arrangements, multivalent (can form multiple ions). | Variable (e.g., +2, +3) |
| Group 17: Halogens | Highly reactive non-metals. 7 valence electrons. | −1 |
| Group 18: Noble Gases | Unreactive, stable non-metals. Full valence shell (8 electrons). | 0 |
Periods (rows): Indicate the number of occupied electron shells. Elements in period 3 have electrons in 3 shells.
Ion Formation
Metals lose electrons to form positive ions (cations).
Non-metals gain electrons to form negative ions (anions).
Common Ions
Cations (+)
Na+, K+, Ca2+, Mg2+, Al3+
Fe2+, Fe3+, Cu+, Cu2+
NH4+ (Ammonium)
Fe2+, Fe3+, Cu+, Cu2+
NH4+ (Ammonium)
Anions (−)
Cl-, O2−, N3−, S2−
OH−, NO3−
SO42−, CO32−, PO43−
OH−, NO3−
SO42−, CO32−, PO43−
Formula Writing
Cross-Multiply Method
To write the formula for an ionic compound, balance the total positive and negative charges to zero.
1. Calcium chloride: Ca2+ Cl− → CaCl2
2. Aluminum oxide: Al3+ O2− → Al2O3
3. Iron(III) sulfate: Fe3+ SO42− → Fe2(SO4)3
4. Magnesium nitride: Mg2+ N3− → Mg3N2
5. Sodium carbonate: Na+ CO32− → Na2CO3
2. Aluminum oxide: Al3+ O2− → Al2O3
3. Iron(III) sulfate: Fe3+ SO42− → Fe2(SO4)3
4. Magnesium nitride: Mg2+ N3− → Mg3N2
5. Sodium carbonate: Na+ CO32− → Na2CO3
IUPAC Naming Rules
Ionic Compounds (Metal + Non-metal)
Name = metal name + non-metal name (change ending to -ide).
If the metal has a variable charge, use Roman numerals in parentheses.
NaCl = sodium chloride
FeCl2 = iron(II) chloride
FeCl3 = iron(III) chloride
CaCO3 = calcium carbonate
Al2(SO4)3 = aluminum sulfate
FeCl2 = iron(II) chloride
FeCl3 = iron(III) chloride
CaCO3 = calcium carbonate
Al2(SO4)3 = aluminum sulfate
Molecular Compounds (Non-metal + Non-metal)
Use Greek prefixes: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca.
First element gets a prefix ONLY if more than one. Second element always gets a prefix and ends in -ide.
CO = carbon monoxide
CO2 = carbon dioxide
N2O4 = dinitrogen tetroxide
PCl5 = phosphorus pentachloride
CO2 = carbon dioxide
N2O4 = dinitrogen tetroxide
PCl5 = phosphorus pentachloride
Common Acids
Binary (no oxygen): hydro- + root + -ic acid.
Oxyacids (with oxygen): polyatomic ion name determines ending (−ate → −ic, −ite → −ous).
HCl = hydrochloric acid
H2S = hydrosulfuric acid
H2SO4 = sulfuric acid
HNO3 = nitric acid
H3PO4 = phosphoric acid
CH3COOH = ethanoic (acetic) acid
H2S = hydrosulfuric acid
H2SO4 = sulfuric acid
HNO3 = nitric acid
H3PO4 = phosphoric acid
CH3COOH = ethanoic (acetic) acid
Properties Comparison
| Property | Ionic | Molecular |
|---|---|---|
| Elements | Metal + Non-metal | Non-metals only |
| Bonding Type | Transfer of electrons | Sharing of electrons |
| State at RT | Solid crystal lattice | Solid, liquid, or gas |
| Melting Point | High | Relatively low |
| Conductivity | Conducts in solution/liquid | Does not conduct |
| Solubility | High solubility in water | Variable solubility |
| Examples | NaCl, CaCO3 | H2O, CO2, C6H12O6 |
Solubility Quick Rules
- • All Group 1 and NH4+ compounds are soluble.
- • All nitrates (NO3−) are always soluble.
- • Most chlorides are soluble (except AgCl, PbCl2).
- • Most sulfates are soluble (except BaSO4, PbSO4).
- • Most carbonates are generally insoluble (except Group 1/NH4+).
WHMIS 2015 Safety Symbols
Flammable
Corrosive
Toxic
Oxidizing
Explosive
Compressed Gas
Environmental Hazard
Biohazardous
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Chemical Reactions & Equations
Evidence of Chemical Change
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Energy change: Release of light, heat, or sound (e.g., burning wood).
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Formation of a gas: Bubbling (e.g., mixing baking soda and vinegar).
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Formation of a precipitate: A solid forms when mixing two solutions (e.g., mixing lead nitrate and potassium iodide).
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Colour change: Unexpected color appears (e.g., rusting iron).
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Temperature change: Solution gets hotter (hand warmers) or colder (cold packs).
Exothermic vs Endothermic
| Exothermic | Endothermic | |
|---|---|---|
| Energy flow | Released to surroundings | Absorbed from surroundings |
| Temperature change | Surroundings get warmer | Surroundings get cooler |
| Examples | Combustion, cellular respiration, neutralization, hand warmers | Photosynthesis, cold packs, baking soda + citric acid |
5 Types of Chemical Reactions
1. Synthesis (Formation)
A + B → AB
• 2H2 + O2 → 2H2O
• N2 + 3H2 → 2NH3
• 2Na + Cl2 → 2NaCl
• N2 + 3H2 → 2NH3
• 2Na + Cl2 → 2NaCl
2. Decomposition
AB → A + B
• 2H2O → 2H2 + O2
• 2H2O2 → 2H2O + O2
• CaCO3 → CaO + CO2
• 2H2O2 → 2H2O + O2
• CaCO3 → CaO + CO2
3. Single Replacement
A + BC → AC + B
Metal replaces metal, or nonmetal replaces nonmetal.
• Zn + 2HCl → ZnCl2 + H2
• 2Na + 2H2O → 2NaOH + H2
• Mg + CuSO4 → MgSO4 + Cu
• 2Na + 2H2O → 2NaOH + H2
• Mg + CuSO4 → MgSO4 + Cu
4. Double Replacement
AB + CD → AD + CB
• NaCl + AgNO3 → AgCl + NaNO3
• BaCl2 + Na2SO4 → BaSO4 + 2NaCl
• HCl + NaOH → NaCl + H2O
• BaCl2 + Na2SO4 → BaSO4 + 2NaCl
• HCl + NaOH → NaCl + H2O
5. Hydrocarbon Combustion
CxHy + O2 → CO2 + H2O
• Complete: CH4 + 2O2 → CO2 + 2H2O
• Complete: C3H8 + 5O2 → 3CO2 + 4H2O
Incomplete combustion (less O2) produces CO or C (soot) instead of CO2.
• Complete: C3H8 + 5O2 → 3CO2 + 4H2O
Incomplete combustion (less O2) produces CO or C (soot) instead of CO2.
Balancing Chemical Equations
Step-by-Step Guide
- Rules: Coefficients go in front; subscripts CANNOT change; same number of atoms on both sides; conservation of mass.
- Step 1: Write the unbalanced equation.
- Step 2: Count atoms of each element on both sides.
- Step 3: Add coefficients to balance. Tip: Balance metals first, then non-metals, then hydrogen, then oxygen last.
- Step 4: Recount to verify.
Worked Examples
Ex 1: __H2 + __O2 → __H2O
Answer: 2H2 + O2 → 2H2O
Answer: 2H2 + O2 → 2H2O
Ex 2: __CH4 + __O2 → __CO2 + __H2O
Answer: CH4 + 2O2 → CO2 + 2H2O
Answer: CH4 + 2O2 → CO2 + 2H2O
Ex 3: __Fe + __O2 → __Fe2O3
Answer: 4Fe + 3O2 → 2Fe2O3
Answer: 4Fe + 3O2 → 2Fe2O3
Ex 4: __Al + __HCl → __AlCl3 + __H2
Answer: 2Al + 6HCl → 2AlCl3 + 3H2
Answer: 2Al + 6HCl → 2AlCl3 + 3H2
Ex 5: __C3H8 + __O2 → __CO2 + __H2O
Answer: C3H8 + 5O2 → 3CO2 + 4H2O
Answer: C3H8 + 5O2 → 3CO2 + 4H2O
Law of Conservation of Mass
Lavoisier's law states that matter cannot be created or destroyed in a chemical reaction. The total mass of reactants equals the total mass of products.
Calculation Example:
If 10g of H2 reacts completely with 80g of O2, what mass of water forms?
10g + 80g = 90g (Product)
If 10g of H2 reacts completely with 80g of O2, what mass of water forms?
10g + 80g = 90g (Product)
Word Equations → Chemical Equations
"magnesium reacts with oxygen to form magnesium oxide"
2Mg + O2 → 2MgO
"hydrogen peroxide decomposes to water and oxygen gas"
2H2O2 → 2H2O + O2
"sodium reacts with water to form sodium hydroxide and hydrogen gas"
2Na + 2H2O → 2NaOH + H2
"methane burns in oxygen to form carbon dioxide and water"
CH4 + 2O2 → CO2 + 2H2O
"aluminum reacts with iron(III) oxide to form aluminum oxide and iron"
2Al + Fe2O3 → Al2O3 + 2Fe
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The Mole Concept
What is a Mole?
1 mole = 6.02 × 1023 particles
This is Avogadro's number. It is the SI unit for amount of substance. Just like a "dozen" means 12 of something, a "mole" means 6.02 × 1023 of something.
Molar Mass (M)
The mass of 1 mole of a substance in grams. It equals the atomic mass from the periodic table, measured in g/mol.
1 mol of C = 12.01 g
1 mol of H2O = 18.02 g (2×1.01 + 16.00)
1 mol of NaCl = 58.44 g (22.99 + 35.45)
The Calculation Triangle
n = m / M
m = n × M
M = m / n
n = moles (mol) • m = mass (grams) • M = molar mass (g/mol)
5 Worked Mole Problems
1. How many moles is 36g of water?
n = m / M = 36 / 18.02 = 2.00 mol
2. What is the mass of 3.5 mol of NaCl?
m = n × M = 3.5 × 58.44 = 204.5g
3. How many molecules are in 2.5 mol of CO2?
2.5 × 6.02×1023 = 1.505×1024 molecules
4. How many moles of Fe2O3 contains 4.5×1024 formula units?
n = 4.5×1024 / 6.02×1023 = 7.47 mol
5. What is the molar mass of Ca3(PO4)2?
3(40.08) + 2(30.97 + 4×16.00)
= 120.24 + 2(94.97) = 310.18 g/mol
= 120.24 + 2(94.97) = 310.18 g/mol
Moles and Balanced Equations
Introduction to stoichiometry. The coefficients in a balanced equation represent the mole ratio between reactants and products.
2H2 + O2 → 2H2O
The "2:1:2" ratio means: 2 moles of H2 reacts exactly with 1 mole of O2 to produce 2 moles of H2O.
Simple Ratio Problem:
If you start with 4 moles of O2, how many moles of H2O can you make?
Answer: Since the ratio of O2 to H2O is 1:2, you multiply by 2. You can make 8 moles of H2O.
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Interactive Practice & Quizzes
Equation Balancer
Adjust the coefficients to satisfy the Law of Conservation of Mass.
Interactive Equation Balancer
Equation 1 of 8
Atom Tracker
Reactants
Element
Products
IUPAC Naming
Test your knowledge of ionic, molecular, and acid nomenclature.
IUPAC Naming Practice
Score: 0 / 0
Name this compound
Reaction Types
Identify the five major types of chemical reactions.
Classify the Reaction
Score: 0
Question 1 of 10
Mole Calculations
Practice converting between mass, moles, and number of particles.
Mole Calculation Practice
Problem 1 / 8Vocabulary Flashcards
Review key terms for the unit exam. Click a card to flip it.
Click to flip
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