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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