Download Science 10 Chem - Holy Trinity Academy

Science 10:
Matter and Energy in Chemical Change
Lesson 1:
Properties of Matter:
Matter: anything that contains mass and takes up space or volume.
 Physical Properties:
o Conditions that can be observed without changing the substance into a different
substance.
o Quantitative: measurable
o Qualitative: observable

Chemical Properties:
o Properties that describe how one substance reacts with others.
o Non-Reactive elements: rarely take part in chemical reactions
 Ex: inert elements, neon
o Reactive elements: take part in chemical reactions
 Ex: Radium, Barium, Hydrogen
Chemicals should be treated with respect:
 Environmental problems
 Heath problems (excessive use of alcohol and nicotine).
 DDT (food chains & malaria come back), mercury poisoning in fish, CFCs
WHMIS:
 Workplace Hazardous Materials Information System is a guideline for handling, storing,
and disposing of reactive materials.
 Each product must have labels in English and French, Material Safety Data Sheet
(MSDS), and workers need to complete an education program.
 Hand-out BLM 7.2
WHMIS Assignment:
Design your own form to record the following MSDS information:
 Chemical Name
 Chemical Formula
 Physical Properties
 Chemical Stability and Reactivity
 Potential Health Risks
 Handling and Storage
 Disposal
( http://www.hc-sc.gc.ca/hecs-sesc/whmis/ )
Lesson 2:
Classification of Matter:
Matter
Pure Substances
Mixtures
Elements Compounds
Mechanical Solutions
Mixtures
Matter: anything that occupies space (volume) and contains mass.
Mixture: a combination of 2 or more different pure substances where the properties can vary
depending on the quantities of the substances.
 Mechanical (Heterogeneous) Mixture: a mixture in which the different substances
are visible. e.g., soil
 Solution (Homogeneous): a mixture in which the different substances are not
visible. e.g., salt in water
Separating Mixtures:
 Flour and Salt: Separation Problem
 Provide each group of 2 with one nylon, two paper plates, breaker, water, filter paper,
salt, and flour.
o Ask the question: Which is the most efficient method of separation? (filter,
static, solubility)
o Review the lab report process
o -Provide procedure guidelines (filter, static, solubility)
o -Brainstorm a results table
Lesson 3:
Separating Mixtures: Lab
Lesson 4:
Pure Substance: made of only one kind of matter and has a unique set of properties (chemical
and physical). e.g., mercury (element) and sugar (compound).
 Element: a substance that cannot be broken down any further by a chemical reaction
into any simpler substance.
 pure substances that contain a single kind of atom
 Each element differs from the others because it has distinct physical and
chemical properties

e.g., helium, oxygen, carbon.
 Molecular elements: elements that naturally occur in combinations of 2-3 atoms. Ex:
H2, O2, N2, F2, Cl2, P4, S8
 Compound: when two or more elements are chemically combined together.
o They can’t be separated by ordinary physical means
o Fixed ratio of elements/never change
o e.g., water (H2O) and carbon dioxide (CO2).
Lesson 5:
The evolution of atomic models
 Dalton/Solid Sphere Model
o All matter is made up of small particles of matter called atoms.
o Atoms could not be divided, created, or destroyed.
o Atoms of an element are the exact same in mass and size. Atoms differ in size
between elements.
Compounds are formed when atoms combine in fixed proportions.
Chemical reactions change the way atoms are grouped but atoms are not
changed.
 Thomson/Plum Pudding Model
o Most of the atom consisted of one large positive charge and small negative
charges embedded that balances out the charges.
 Rutherford/Nuclear Model
o Atom contained a positive central core which made up all of the mass.
o Negative electrons were distributed around the core.
 Bohr Model
o Structural features of the atom.
o Energy exists in small units called quanta.
o Electrons have fixed amounts of energy and move in circular pathways at fixed
distances from the nucleus.
 de Broglie: Electron Cloud Model
o Electrons have distinct energy levels but the locations are not defined.
All these theories and models combined to form the Modern Atomic Theory
o
o
Modern Atomic Theory
o tiny nucleus containing (+) protons & neutral neutrons
o (-) electrons occupying large space
o nucleus mass almost = the total mass of atom
o atom is electrically neural b/c # protons = # electrons
Lesson 6:
Develop a Theory Activity:
1. Design a mystery box (objects that can move around on the inside and a few taped to the
inside) that will allow for simple tests or experiments to take place like probing or shaking.
2. Make a drawing of your box based on inferences you think a person could make out
about it.
3. Exchange boxes with someone and perform tests, record information in the below chart.
Test Conducted
Observations
Inferences
4. Draw a model of mystery box and state your inferences as a theory.
Lesson 7:
Atomic Structure Defines the Characteristics of an Element
 Three sub-atomic particles which have consistent masses and charges
 Elements differ in the number of protons
Protons:
 Positive charges
 1 Atomic Mass Unit (AMU) = 1.67 x 10 (-24)
 Atomic Number is the number of protons
Neutrons:






Neutral Charge
1 AMU mass
Neutrons vary which is the reason for variation in atomic mass.
Protons don’t change therefore the atomic number remains the same as well as the
element.
Isotopes are elements that have the same atomic number but different atomic
masses. Ex. Carbon 14 vs Carbon 16
Periodic table’s atomic mass is the average mass for all isotopes.
Atomic Masses:
Carbon =12.01 AMU, 6 AMU proton and 6.01 AMU neutron
Neon = 20 AMU
Argon = 36 AMU
Electrons:
 Negative charges
 Mass is so very small that it doesn’t affect the total mass of the atom
o -electrons are located in energy levels and the level number depends on # of electrons
o -electrons are lazy (want to stay close to the nucleus/less energy need)
o *electrons gain energy (heated) they may jump into next level
o *once electron emits this energy, falls back down
o -electrons move b/w energy levels by losing or gaining energy (usually that is a
specific amount of energy)
o -Electrons can’t exist in-between levels
 -Atoms have nucleus and electron energy levels.
o 1st level - 2 electrons
o 2nd level - 8 electrons
o 3rd level - 8 electrons
th
o 4 level – 18 electrons
 -Period tells you # of energy levels
 -Family tells you # of valence electrons (electrons outer layer)
o -Atoms tend towards stability
o groups tend to react together = give/take to become stable
o group/column 18 = full outer layer
o group/column 1 = 1 extra electron
o group/column 17 = 7 extra electrons
Draw out an example figure (Chlorine) label parts
7e
Cl 17 (mass 35)
8e
2e
p 17
n 17
Assignment: Finish filling in element worksheet
Element
Name
Element
Symbol
Atomic
Number
Atomic
Mass
Number
of
Protons
Average
Number
of
Neutrons
Number
of
Electrons
(1st
Level)
Number
of
Electrons
(2nd
Level)
Number
of
Electrons
(3rd
Level)
Number
of
Electrons
(4th
Level)
1 to 20
Lesson 8:
The Periodic Table:
 Provides information about the properties of elements and its structure.
 Vertical Groups (families 1-18)
o Families have similar chemical properties/different intensity
o Group 1: Alkali metals (Very reactive metals)
o Group 17: Halogens (Most reactive non-metal)
o Metal Reactivity increases top to bottom (groups 1 - 2 only)
o Non-metal Reactivity decrease top to bottom (17-18 only)
o Group 18: Noble gases (unreactive)
 Periods or Horizontal Rows (7)
o reactive properties of elements gradually or periodically change from left to right
o two series at the bottom: convince, same properties

Staircase line: line that separates metals (L) from non-metals (R)
o Properties of Metals (75% of elements on earth)
 Ductility: the ability to draw into a thin, flexible metal thread or wire.
 Malleability: the ability to be hammered into shape without breaking.
 strength, hardness, durability, ductility, and malleability, lustrous,
reactive
 good conductors of heat and electricity
 Vitamins-metals are important for living organisms (potassium -plants &
animals)
 Alloys: mixtures or solutions of metals
o Properties of Non-Metals
 lack the properties of metals
 Nitrogen, phosphorus, oxygen are examples...
o Metalloids: the elements B, Si, Ge, As, Sb, Te, & Po exhibit both metallic and
non-metallic properties.
 Poor conductors of heat, can conduct electricity
Examining Metals and Non-Metals:
Problem: What are the properties of metals and non-metals?
Variables:
Hypothesis:
Materials: Samples, magnet, conductivity apparatus, hammer and block, and hydrochloric acid.
Procedure:
1. Examine each sample and determine its state, colour, ductility and magnetism.
2. Using the conductivity apparatus to determine the conductivity of the element.
3. Place a few drops of hydrochloric acid on the element and record the observations.
4. Clean up.
Observations:
Substance
State
Colour
Ductile
Magnetic Conduction
Acid
Aluminum
Carbon
Copper
Iron
Magnesium
Sulphur
Zinc
Analysis:
Conclusion:
Errors:
Lesson 9:
Metal lab
Lesson 10:
Chemical Solitaire: Science 9
Part B:
Lesson 1:
Compounds: Combination of 2 or more elements.
 Two forms of compounds: ionic and molecular
Ionic Compounds
Bonds are created by the transfer of electrons
Molecular Compounds
Bonds are created by the sharing of electron
High melting point
Distinct crystal shape
Formed from metallic and non-metallic elements
Forms ions in solution
Conducts electricity when dissolved in water
Solid at room temperature
Low melting point
Not always form crystals
Formed from only non-metallic elements
Does not form ions in solution
Does not conduct electricity when dissolved in
water
Solid, liquid, or gas at room temperature.
Activity: Ionic or Covalent: Track those electrons (Page 33 in Sciencefocus 10)
Lesson 2:
Atoms of elements (neutral): the number of protons and electrons are the same.
 # of electrons can change b/c they are free to move
 atoms want to have a full outer energy level
o -gain (non-metals) or lose
o -share electrons
Ions: Atoms or groups of atoms that have lost/gained electrons
o -Atoms lose an electron = + charge (more protons than electrons)
o -Atoms gain an electron = - charge (more electrons than protons)
o -Atoms can lose or gained more than one electron = shown by number with the
charge
-Worksheet page 3-39, 40 & Go over the answers
Types of Ions:
o -Monoatomic ions: simple ions
 ions formed from atoms of an element
 examples: F- or Fe 2+
o -Polyatomic ions: complex

ions made of a groups of atoms acting as an unit
 as an unit they will lose or gain electrons
 charge is shared
 -examples: NH4+
 -Show alchem periodic table
Chemistry Bingo with Ions
Lesson 3:
Ionic Bonds:
o -Ions bond by transferring electrons to form ionic compounds
o –Example:
Na +
Cl -
Transfer and Attraction
o
o
-Force that holds 2 opposing charged ions together (+ -)
o -Ions with different charges (strongly attracted)
o -Ions with similar charges (strongly repelled)
Characteristics:
o -involve the transfer of electrons
o -involve a change in energy (add/remove electrons)
o -form between metallic and non-metallic elements
o -solid at room temperature
o
o -conducts electricity
Lattice - How ions arrange themselves to become stable
Writing an ionic compound:
-compound does not carry a charge
-sum of the charges must be zero
-example:
Al 3+
Al
Cl -
Al 3+
Cl
Al
AlCl3
O2-
O
Al2O3
-Hand-out work-sheet
Lesson 4:
Naming Binary (2 elements) Ionic Compound
#1. Name includes both elements in compound
#2. Metal element appears first
#3. Non-metal element appears last and has its ending
changed to -ide
#4. Name of compound doesn’t mention the number of ions
#5. Name is not capitalized
example:
Na +
+
Cl-

NaCl
Sodium Chloride
Ionic Compound’s Chemical Formulas
#1. Symbol for metallic ion is first
#2. Symbol for non-metallic ion is second
#3. Subscripts indicate the ratio of ions in compound
Example:
Na + ClCa 2+ Cl-
NaCl ratio 1 to 1
CaCl2 ratio 1 to 2
Remind them: ionic compounds must be neutral (total # of + = # of -)
Naming Ionic compounds with 2 or more charges.
-Usually involves transitional metals.
#1. Rules are the same as ionic compounds with 1 charge
#2. Use a roman numeral to show which one is being used.
Example:
Cl - + Cu+
Cl - + Cu2+
 Copper (I) Chloride (CuCl)
 Copper (II) Chloride (CuCl2)
Worksheet
Lesson 5:
Polyatomic Ionic Compound Chemical Formulas
#1. Surround the polyatomic ion with parentheses
#2. Use subscript to indicate the required number of units
examples:
Mg 2+
OH-
Al3+
Mg (OH)2
SO4 2-
Al2(SO4)3
Naming Polyatomic Ionic Compound
o Same as binary ionic compounds
-Refer to “Common Polyatomic Ion Chart” Handout P. 466
example:
beryllium carbonate
Be2+ CO3 2Be2(CO3)2
BeCO3
Worksheet (B14- Alchem 10)
Lesson 6:
Ionic Solubility:
o Some ions react in solution (usually water) & produce solid ionic compounds (precipitate)
o example:
Ag+ (aq) + Cl- (aq)  AgCl(s)
o
To see if an ionic compound will be solid look at “Solubility of Ionic Compounds in Water”
in Appendix D
o (-) ion on top
o (+) ion on bottom - highly solubility = compound dissolve
o low solubility = compound form solid
* Practice problems on page 278
Lesson 7:
Remind them that atoms want to fill their outer energy layer
Covalent/Molecular Compounds
o -non-metal and non-metal elements
o -atoms share electrons (from outer layer)
o -no ions are formed (because electrons don’t move from one atom to another)
example:
H2
HCl
1P
1P
share 2 e
1P
17P
share 2 e
Covalent Bonds:
o -atoms are “glued” together
o -weaker than ionic bonds
o -liquids, solids, & gases at room temperature
Naming Binary Molecular Compounds:
#1. Write first element’s entire name
#2. Change the second element’s ending to -ide
#3. Use prefix to indicate the number of each type of atom in the
formula
mono = 1
di = 2
tri = 3
tetra = 4
penta = 5
hexa = 6
Hepta = 7
octa = 8
Nana = 9
deca = 10
#4. Write the name in lower case letters.
Examples:
CO
carbon monoxide or monocarbon monoxide
Writing Formula for Binary Molecular Compounds:
#1. Write symbol for elements in the same order
#2. Use subscripts to indicate the number of each type of atom
Examples:
Sulfur dioxide SO2
Hydrogen chloride HCl
Problems (Alchem 20, F12; Alchem 10, B5, & student worksheet)
Complex Molecular Compounds:
 H2O (l) – Water
 H2O2 (l) – Hydrogen peroxide
 CH4 (g) – Methane
 CH3COOH (aq) – vinegar
 C6H12O6 (s) – glucose
 C12H22O11 (s) – sucrose
 NH3 (g) – ammonia
 C2H5OH (l) – ethanol
 CH3OH (l) – methanol
 O3 (g) – ozone
Lesson 8:
Molecular Model Lab (Alchem 10, B6 & B7)
-Groups of two to four per kit
Lesson 9:
Ionic or Molecular? Page 58 to 59 (Sciencefocus 10)
Lesson 10:
pH is an other way to classify matter,
o -acids, bases, and neutral substances.
o -Show pH scale
o Properties of bases and acids
o Acid Properties:
 Soluble in water, conducts electricity, turns litmus paper red, turns
bromthymol blue yellow, turns phenolphalein colorless, reacts with
zinc to produce gas, tastes sour, seen in any physical state.
o Base Properties:

o
Feels slippery, turns phenolphalein pink, turns litmus paper blue,
conducts electricity, doesn’t react with zinc, tastes bitter, usually
seen as a solid.
Bases neutralize acids
Naming and Writing Formulas for Acids:
o Two systems:
IUPAC
#1. Follow the guidelines for naming ionic compounds.
#2. Insert the word aqueous in front of the compound.
Example:
HCl
aqueous hydrogen chloride
* Hint - Any ionic compound that has aqueous and hydrogen in front of the negative ion is an
ACID.
Classical System
Ionic Name
#1. hydrogen
#2. hydrogen
#3. hydrogen
Example:
ide
ate
ite
hydro
Acid Name
ic acid
ic acid
ous acid
HCl hydrogen chloride  hydrochloric acid
HClO3 hydrogen chlorate  chloric acid
HNO2 hydrogen nitrate  nitrous acid
Talk about naming bases - same as ionic compounds with aqueous in the front.
o Practice problems (B22 Alchem 10) & answer.
Lesson 11:
Objectives:
The students will be able to:
 Identify the properties of Acids and Bases
Identifying Properties of Acids and Bases:
Problem: What are the Properties of Acids and Bases?
Variables:
Hypothesis:
Materials: Hydrochloric acid, sodium hydroxide, household substances, distilled water, 50ml
beakers, litmus paper, conductivity apparatus, bromthymol blue, phenolphthalein, and zinc.
Procedure:
1. Place each substance in a 50ml beaker and test with litmus paper, pH paper, conductivity
apparatus, zinc, and the chemical indicators.
2. In between each substance wash out the beaker with distilled water.
3. Clean up area.
Observations:
Substance
Red
Blue
pH
Conductivity
Brom
Pheno
Zinc
Litmus
Litmus
paper
Apparatus
Blue
Vinegar
Lemon
Juice
Vitamin C
Asprin
Oven
Cleaner
Soap
Baking
Soda
Ammonia
Analysis: Page 295 5 to 7
Conclusion:
Errors:
Lesson 12:
Structure of Water:
o Properties of water are determined by the attractive forces between water molecules
caused by the positive (protons) and negative forces (electrons) that are contained in
every atom.
o Hydrogen bonds are extremely strong.
o Polar Molecule: ~ electrically neutral (protons = electrons)
Hydrogen
+
-
Oxygen
Positive charges spend more time at the hydrogen atoms
Negative charges spend more time at the oxygen atom.
Properties of Water:
o High surface tension: demonstrate
o Cohesion: forces of attraction between molecules of the same substance.
o Adhesion: forces of attraction between molecules of different substances.
o Surface tension: the tendency for molecules to be pulled from the surface to the
interior of the liquid.
o
Addition of soap or detergents (surfactants) will lower the surface tension of
the liquid in which they are dissolved.
o
Water has a concave meniscus and shows capillary action:
o Capillary action is the force that draws water up from the roots to the leaves in
tall trees.
o
Large specific heat capacity:
o
o
Specific Heat Capacity: the amount of heat it takes to raise the temperature of a
specific mass of a substance by one degree Celsius.
o Since it takes a lot of heat to raise the temperature of water, organisms’
internal temperature is regulated and remains constant.
Density of ice is less than the density of liquid water:
o Water expands as it freezes and contracts as it is heated.
D=m/v
~ m = Dv
0 C ~ 100 g water sample will have expanded and increased its volume
therefore its density will have dropped.
4 C ~ 100 g water sample will have contracted and decreased its volume
therefore its density will have increased.
m=m
Dv=Dv
o
Enables organisms to survive winter conditions under the ice.
o
Can exist in more than one state at the same time:
o
High melting and boiling points: requires a lot of energy to break the bonds
0 C
100 C
Part C
Lesson 1:
Physical Change: substance changes shape, texture, state, and/or size but retains similar
properties as before the change. Reversible.
Chemical Change: causes one or more new substances, with new properties, to be formed and
may be difficult or impossible to reverse..
Evidence:
 Heat or light energy is produced or absorbed. When gasoline burns in a car
engine and heat is released.
 Change in colour. Bleach on a denim jacket
 Change in odour. Striking a match
 Formation of a solid or gas (precipitate or bubbles). Vinegar and baking soda
produces bubbles.
o Reactants = the form of matter that go into a reaction
o Products = the forms of matter that come out of a reaction
Lab Activity: Mystery Solid Laboratory
Lesson 2
Energy change always accompanies chemical reactions
 Exothermic reactions - release of excess energy (out of)
 more energy stored in the reactant bonds than needed to form the product bonds
DEMONSTRATION
Place the pellets of NaOH into the water
 Endothermic reactions - absorb energy from surroundings (into)
 more energy is need to form the products than can be provide by breaking
reactant bonds.
DEMONSTRATION

add ammonium chloride to water
The release/absorption of energy doesn’t affect the mass and must be conserved.
Demonstration: Endothermic or Exothermic? Page 91 (Sciencefocus 10)
How do we represent chemical reactions?
 chemical equations which are recipes for chemical changes that tells you
what to put in and what you get out
Law of Conservation of Mass (Lavoisier)
 matter cannot be created nor destroyed only changed
Example:
40 kg (wood) -------> 40 kg (ashes)
Therefore, the total number of atoms on the reactant side must equal the total number of atoms
on the product side. Equations must be balanced.
Three Rules for Balancing Equations:
#1. Write the correct chemical formulas for both reactants and
products.
-reactants left/products right & arrow in between
#2. Balance each atom, one at a time, using whole number
coefficients
-start at the atom with the greatest number
#3. Leave hydrogen and oxygen atoms to balance last.
Example:
HCl + NaOH  H2O + NaCl
C6H12O6 + O2 
C
H
O
6
12
8
C6H12O6 + 6 O2 
C
H
O
H2O +
CO2
1
1
3
6 H2O + 6 CO2
6
12
18
6
12
18
Hand-out practice sheets and extra help sheet (3-76 to 3-81)
Lesson 3
Types of Chemical Reactions:
#1. Formation Reactions - simple elements combine to form compounds.
#2. Decomposition Reactions - reactions that produce products simpler than the reactants
#3. Combustion/Oxidation Reactions - reaction of a substance with oxygen (burning).
 Always exothermic and have carbon dioxide and water as products
#4. Replacement Reactions - exchanges occur between the reactants to produce products that
are neither more nor less complex than reactants
(I) single replacements: 2 metals switch places
(II) double replacements: when 2 single replacements occur in the same reaction
Examples:
1. O2 + 2H2  2 H20
2. HgO  2Hg + O2
3. CH4 + 3O2  CO2 + 4H2O
4. Cl2 + 2NaI  I2 + 2NaCl
5. Na2CO3 + Ca(OH)2  2NaOH
+ CaCO3
Worksheets -post answers on the board
Lesson 4
Chemical Tests:
 Hydrogen: a burning splint is lowered into a jar. (+) Pop
 Oxygen: a glowing splint is lowered into a jar. (+) splint will ignite
 Carbon Dioxide: limewater is added to a jar. (+) limewater will become cloudy
 Water: cobalt (ii) chloride paper will be added to the solution. (+) paper turns pink
Assignment: Putting it together: Classifying Chemical Reactions (Page 110 to 113 Sciencefocus
10)
Lesson 5:
Moles and Molar Mass:
 n=m/M
n = moles or the number of atoms of that element
rd
~ 1 mol = 6.02 x 10 to the 23 power (Avogradro’s Number)
m = mass or the amount of a substance in grams
M = molar mass or the mass of a mole of a compound
Molar Mass:
 Use the atomic mass given in the periodic table and multiply it by the number of
moles for each element in the compound.
 Add up all the values.
Example:
Molar mass of Cu(ClO3)2:
1 Cu = 1 x 63.55 = 63.55
2 Cl = 2 x 35.45 = 70.90
6 O = 6 x 16.00 = 96.00
230.45 g/mol
Molar Mass of CCl4? answer: 153.81 g/mol
Calculating the Number of Moles from Mass:
 n = m/M
Example: Determine the number of moles of magnesium oxide in 8.06g of the compound.
MgO 1 Mg = 1 x 24.31 = 24.31
n = m/M
1 O = 1 x 16.00 = 16.00
= 8.06 g/ 40.31 g/mol
40.31 g/mol
= 0.200 mol
Calculating Mass from the Number of Moles:
 m = nM
Example: Determine the mass of 0.25 mol of copper (ii) sulfate.
CuSO4
1 Cu = 1 x 63.55 = 63.55
m = nM
1 S = 1 x 32.00 = 32.00
= (0.25 mol) x (249.71g/mol)
4 O = 4 x 16.00 = 64.00
= 62 g
249.71 g/mol