!!Name: _____________________________________ Block: ______________ Sc10 Veenstra!

!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!� of �1 20

Everything you can feel is made of atoms!Chemistry!

Our inquire into “Everything you can feel is made of atoms” will help us better understand:!• That the difference between atoms, ions, and molecules is caused by the difference in

their structure and components!• The classification of substances as acids, bases, and salts, based on their

characteristics, name, and formula!• The difference between organic and inorganic compounds!• Chemical reactions and the Law of conservation of mass !• How the rate of reaction is affected!• Radioactivity using modern atomic theory

Vocabulary!acids, alpha particle, atomic number, atoms, bases, beta particle, Bohr diagrams, bromothymol blue,catalyst, combustion, compounds, concentration, conservation of mass, covalent bonding, daughter isotope, decomposition, electron, fission, fusion, gamma radiation, half-life, indigo carmine, inorganic, ionic bonding, ions, isotope,Lewis diagrams, light, litmus paper, mass number, methyl orange, molecules, neutralization (acid-base), neutron, organic, parent isotope, phenolphthalein, polyatomic, proton,radioactive decay, salts, single and double replacement, surface area, symbolic equations, synthesis, valence electron      

Note:!If you lose this package it is your responsibility to print out a new copy from Ms. Veenstra’s

webpage: https://lveenstra.wordpress.com/science10/

Chapter 4 Learning Goal!!!!!!!!!!!!!!!!!!!!!!

!!!!!

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B D A E Reference

1. I can demonstrate knowledge of the three subatomic particles, their properties, and their location within the atom

Chapter 4.1

2. I can define and give examples of ionic bonding and covalent bonding

Chapter 4.1

3. I can with reference to elements 1 to 20 on the periodic table, draw and interpret Bohr models, including protons, neutrons, and electrons, of atoms (neutral), ions (charged), molecules - covalent bonding (e.g., O2, CH4

Chapter 4.1

4. I can draw and interpret Lewis diagrams showing single bonds for!simple ionic compounds and covalent molecules, and distinguish between lone pairs and bonding pairs of electrons in molecules

Chapter 4.1

5. I can use the periodic table and a list of ions (including polyatomic!ions)to name and write chemical formulae for common ionic compounds, using appropriate terminology

Chapter 4.2

6. I can convert names to formulae and formulae to names for!covalent compounds, using prefixes up to “deca”

Chapter 4.2

7. I can define and explain the law of conservation of mass Chapter 4.3

8. I can represent chemical reactions and the conservation of atoms using molecular models

Chapter 4.3

9. I can write and balance (using the lowest whole number coefficients) chemical equations from formulae, word equations, or descriptions of experiments

Chapter 4.3

Physics 11: Term 2 Report Card

Unit: Optics Knows: I can list the rules of ray drawing. I can list the characteristics of images. I can state the laws of reflection. I can distinguish between diffuse and specular reflection. I can label the parts of a ray drawing for plane mirrors. I can label the parts of a ray drawing for curved mirrors. I can describe different kinds of curved mirrors. I can define the index of refraction. I can describe Snell’s Law. I can label the parts of a ray drawing for refraction at a plane boundary. I can label the parts of a ray drawing for refraction by a lens. I can define critical angle. I can describe total internal reflection. I can give examples of common devices that reflect light. I can give examples of common devices that refract light.

LG = Mark = %

Do’s: I can draw labeled ray diagrams showing how images form in a plane mirror. I describe the characteristics of images formed in a plane mirror. I can experimentally distinguish between converging and diverging mirrors. I can experimentally find the focal length of a converging mirror. I can locate and describe the image formed by a curved mirror:

using a scaled ray diagram and using the mirror equation I can solve problems using:

The definition of index of refraction Snell’s Law

I can experimentally distinguish between converging and diverging lenses. I can experimentally find the focal length of a thin converging lens. I can locate and describe the image formed by a thin lens:

using a scaled ray diagram and using the mirror equation.

LG = Mark = %

Strengths:

Unit: Kinematics Knows: I can define and relate the terms: clock reading, position and event. I can differentiate between a clock reading and a time interval. I can define and relate distance and average speed. I can define and relate displacement and average velocity. I can differentiate between scalars and vectors. I can define instantaneous velocity and instantaneous speed. I can define average acceleration.

LG = Mark = %

Do’s: I can solve problems involving: displacement, time interval, and average velocity. I can use position-time graphs to determine: • displacement & average velocity, distance travelled & average speed, instantaneous velocity I can construct position-time graphs based on data from various sources. I can construct velocity-time graphs based on data from various sources. I can use velocity-time graphs to determine: • instantaneous velocity, displacement, average velocity, acceleration

LG = Mark = %

Strengths:

!"#$%%$%#& '"(")*+$%#& ,--*.+)$/0"1& 23".+)456&Does not demonstrate a basic understanding of concepts.

Demonstrates a basic understanding of concepts.

Demonstrates a solid understanding of concepts.

Demonstrates a complete and deep understanding of concepts.

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Name: ___________________________

Student #: ______________ HR: _____

Date: _____________________

Teacher: Mr. J.Martens

Course: Physics 11

Block: _______

My estimate based on quizzes and assignments: !LG = Mark = %

Assign by Ms. V !LG = Mark = %

Test result: %

Chapter 4.1 Atomic Theory Lab 4.1: Building an Atom (PhET simulation)

PART I: ATOM SCREEN

1. Go to the website: phet.colorado.edu. Click on HTML5 simulations on top right of screen and choose the Build an Atom simulation (http://phet.colorado.edu/en/simulation/build-an-atom)

2. Explore the Build an Atom simulation with your group. As you explore, talk about what you find. List two things your group observed in the simulation.

a.

!b. !

!2. Click on the + sign for each of the boxes (element name, net charge and mass number) to view

changes as you change the number of particles in the atom.

3. What particle(s) are found in the centre of the atom?______________________________________________________

4. Play until you discover which particle(s) determine(s) the name of the element you build. _______________________

5. What is the name of the following atoms?

a. An atom with 3 protons and 4 neutrons: _____________

b. An atom with 2 protons and 4 neutrons: _____________

c. An atom with 4 protons and 4 neutrons: _____________

6. Play with the simulation to discover which particles affect the charge of an atom or ion. ________________________

7. Fill in the blanks below to show your results:

a. Neutral atoms have the same number of protons and electrons.

b. Positive ions have ________________________________ protons than electrons.

c. Negative ions have _______________________________ protons than electrons.

!!!!

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8. Develop a relationship (in the form of a single sentence or equation) that can predict the charge of the atom/ion based on the number and types of particle.

!!!9. Play with the simulation to discover what affects the mass number of your atom or ion.

______________________________

a. Whatisarulefordeterminingthemassnumberofanatomorion?!!10.Fillintheblanksbelowtoshowyourunderstandingofchargeandmass:

a. Protonshaveamassof___________amuandachargeof_____________.b. Neutronshaveamassof___________amuandachargeof_____________.

c. Electronshaveamassofnearly___________amuandachargeof_____________.

11. Practice applying your understanding by playing 1st and 2nd levels on the game screen.

!PART II: SYMBOL SCREEN

1. Using the Symbol readout box, figure out which particles affect each component of the atomic symbol and how the value of the numbers is determined.

3. Practice applying your understanding by playing the 3rd and 4th game levels. Play until you can get all the questions correct on the 4th level. Fill in the information here for your last screen of the 4th game level:

protons ____________

neutrons ___________

electrons ___________ !Finish Worksheet 4.1 Atomic Theory!

Position in symbol box Term to describe this information

Particle used to determine this

How the value is determined

a Element symbol protons # of p will identify the element

b Ion charge

c Atomic number

d Mass number

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cdab

Chapter 4.1 Bohr Models Notes!!Electron shells (orbitals, E levels) !1st shell holds only ______! ! ! ! ! Valence shell:! 2nd shell hold ______!3rd shell holds ______! ! ! ! ! ! Valence electrons:! 4th shell holds ______!!! ! ! ! ! ! ! ! ! ! Example: Sodium Atom! !Shortcomings of model?!!!!Advantages of model?!!!!Stable:!!Unstable:!!Atoms will tend to gain or lose e- in order to become stable!!ION! ! ! ! ! vs! ! ! ATOM! !!!!!!!!IMPORTANT NOTE:!Protons are locked in the nucleus so the positive charge can’t change...positive ions come from atoms that have lost e-, NOT from gaining protons!!!Example!Single e- in valence shell !Unstable !gain 7 or lose 1 e-?!!!!!!

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4.1 Bohr Model Compounds Notes!!Forming Compounds!Atoms wants to be stable (have a full valence shell). Atoms can become stable by!!1. !

- metals tend to ___________ and become ___ (called_______________)!

- non-metals tend to __________ and become ___ (called_______________)!

2. !!A. Ionic Compounds!Ionic bonds are formed between positive ions and negative ions.!• Generally, this is a metal (+) and a non-metal (-) ion.!• For example, lithium and oxygen form an ionic bond in the compound Li2O.!

!!!!!!!!Covalent Compounds!Covalent bonds are formed between two or more non-metals.!• Electrons are shared between atoms.!• For example, hydrogen and fluoride form a covalent bond in the compound HF.!!!!!!!!!!Ionic ! ! vs. ! ! ! Covalent! !!

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NOTE:!Metals are found

4.1 Lewis diagram Notes!Lewis Structures!Lewis diagrams illustrate chemical bonding by showing only an atom’s valence electrons and the chemical symbol.!• !

• !!Example!Draw Lewis structure for Ca, a chlorine ion and a beryllium ion.!!!!!

NOTE: Square brackets are placed around each ion and the charge is added outside the bracket.!

Ionic Compounds!Beryllium and chlorine can form an ionic compound:!

!!!Let’s try CaO:!!!!!!

Covalent Compounds!Lewis diagrams can also represent covalent bonds.!• The shared pairs of electrons are usually drawn as a straight line.!

Example:!

HF! ! ! H2O! ! Bonding e-:!! ! !! ! ! ! ! Lone pairs:!!!!

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4.2 Type I Ionic Compounds (metal + non-metal)!- contain a metal and a nonmetal !- the metal that is present only forms one type of cation (only one charge in the periodic table). !- Both the metal and the non-metal form ions, which is why it is called an ionic compound. !!

1) From the following list, cross out those compounds that do NOT belong in the category for Type I binary ionic compounds.!!

NaCl! FeCl2! CaCl2! TiO2! MgO! AlBr3! KCl! K2S! BeF2! Cu2O3!AgCl! Zn3N2!!Formula and name examples for Type I ionic compounds:!!KI = potassium iodide ! BaO = barium oxide ! ZnF2 = zinc fluoride !Na2S = sodium sulfide! Ag3N = silver nitride ! BeCl2 = beryllium chloride!!2) What type of element is always listed first (metal or nonmetal)? __________ second? _________!3) Is the name of the first element in the compound different from the element? (yes/no)!4) What is the common ending for all the names? ______________!5) In zinc fluoride, there are 2 fluoride atoms, are they indicated in the name? (yes/no)!6) What is the charge on the zinc ion? _______!7) What is the charge on the fluoride ion? _______!8) Why do you need one zinc ion and two fluoride ions for the formula for zinc fluoride?!9) Why do you need two sodium ions for every sulfide ion in sodium sulfide?!!10) As a team, determine the rules for naming type I ionic compound when given the formula.!!!!!11) As a team, determine the rules for writing the formula for a type I ionic compound when given the

name.!!!!12) Name each of the type I ionic compounds listed in question 1.!!! ! ! !________________________________ ____________________________________!! ! ! !________________________________ ____________________________________!! ! ! !________________________________ ____________________________________!! ! ! !________________________________ ____________________________________!! ! ! !________________________________ ____________________________________!!Finish Q. 1-11 on Worksheet 4.2 A + B!

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4.2 Type II Ionic Compounds (transition metal +non-metal)!- contain a metal and a non-metal !- the metal that is present here can form more than one type of cation (blank in the ! periodic table above)!- both the metal and the nonmetal form ions, and it is still called an ionic compound. !!

1) From the following list, cross out those compounds that do NOT belong in the category for Type II ionic compounds.!!

! AlP! FeCl2! Ag2O! VBr5! CoS! SnF2! K3N! SrF2! CuBr! AuCl3! ZnO! HgS!!Formula and name examples for Type II ionic compounds:!!Fe2O3 = iron(III) oxide! FeO = iron(II) oxide!CuS = copper(II) sulfide ! CuCl = copper(I) chloride!MnO2 = manganese(IV) oxide ! MnCl2 = manganese(II) chloride !!2) What type of element is always listed first (metal or nonmetal)? ____________ second?

___________!3) Is the name of the first element in the compound different from the element? (yes/no)!4) What is the common ending for the nonmetal portion of the names? ______________!5) In the compound FeO, what is the charge on iron? _______!6) In the compound Fe2O3, what is the charge on iron? ________!7) What does the Roman number after the metal name represent? !!8) As a team, determine the rules for naming type II ionic compound when given the

formula.!!!!11) As a team, determine the rules for writing the formula for a type II ionic compound when

given the name.!!!12) Name each of the type II ionic compounds listed in Question 1 of Type II section.!!! ! ! !_____________________________ _________________________________!! ! ! !_____________________________ _________________________________!! ! ! !_____________________________ _________________________________!! ! ! !_____________________________ _________________________________!Finish Q. 12-22 on Worksheet 4.2 A + B

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4.2 Compounds Containing Polyatomic Ions!Polyatomic ions are ions that as a group have a set charge. Polyatomic ions are usually

recognized in a formula by the grouping of more than one nonmetal elements after a metal. Your book has a table listing polyatomic ions. Use p. 5 in the data booklet to fill in the following table with the appropriate names/formulas of the polyatomic ions. !!

!!Use your knowledge of Type I and Type II metals as well as the appropriate polyatomic name/formula to fill in the following table.!!

!!Check your work:!Are the compound formulas you filled into the table above neutral in charge?!!Do all type II metals in the table above have their charge indicated by either a Roman numeral?!!Are all type I metals listed without a Roman numeral?!!Finish Q. 23-48 on Worksheet 4.2 A + B

Name Formula Name Formula

ammonium chlorite

nitrite

cyanide carbonate

! NO3–

! CH3COO–

!CrO42–

!OH– !SO3

2–

Name Cation Anion Formula

aluminium hydroxide

iron(II) nitrate

calcium sulfite

ammonium nitrate

KCN

Cu(NO

! !

!

!OH–

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4.2 Covalent compounds (non-metal and non-metal)!! Compounds that do not contain metals have covalent bonds instead of ionic bonds. A covalent bond is formed by sharing one or more pairs of electrons. The pair of electrons is shared by both atoms. For example, in forming H2, each hydrogen atom contributes one electron to the single bond.!!1) From the following list, cross out those compounds that do NOT belong in the category

for binary compounds containing only nonmetals or metalloids.!!CCl4! AlCl3! CO! SeF6! SiO2! SrI2! P4O10!TiO2! SeO3! IrCl! ZrO2! N2O5!!

Formula and name examples for covalent compounds:!CO2 = carbon dioxide ! ! ! H2O = dihydrogen monoxide! !P4S10 = tetraphosphorus decasulfide! ! NO = nitrogen monoxide!IF5 = iodine pentafluoride ! ! ! BF3 = boron trifluoride!!2) Which element is listed first in the name?!3) Is the name of the first element in the compound different from the element? (yes/no)!4) What is the common ending for all the names? ______________!5) What do the prefixes (di-, mono-, penta-, tri-) in the names above mean?!6) Is the prefix mono- used when there is only one atom of the first element? (yes/no) !7) Is the prefix mono- used when there is one atom of the second element? (yes/no)!8) As a team, determine the rules for naming covalent compound when given the formula.!!!!!!!!!!!!!!!8) Name each of the covalent compounds listed above.!!! ! ! !_____________________________ _________________________________!! ! ! !_____________________________ _________________________________!! ! ! !_____________________________ _________________________________!! ! ! !_____________________________ _________________________________!Finish Worksheet 4.2 C + D Covalent Compounds

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!NOMENCLATURE REVIEW

1. Type 1 ionic compounds: metal (only one possible charge) + nonmetal!!! Name: metal written first!and “-ide” ending! ! example: !! ! ! !! Formula:! *balance charges! ! ! example:!! ! ! ! ! !! ! ! !!2. Type 2 ionic compounds: transition metal (two or more possible charges) + non-metal!! ! ! ! !! * include roman numerals in the name! ! example: ! !!!3. Poly atomic ions: includes ions with 2 or more elements!! ! ! ! ! ! ! ! ! ! ! example:!! *have different endings, such as “ite” and “ate”!! *exceptions: Cyanide (CN-), hydroxide (OH-), bisulfide (HS-)!!! ! !4. Hydrogen:!

• Some areas of common confusion concerning hydrogen are listed below:!!! ! HBr ! hydrogen bromide! ! ! NaOH sodium hydroxide!!! ! NaH ! sodium hydride!

!5. Covalent Compounds: 2 nonmetals together!! !! *no balancing of charges!! *use prefixes (what you see is what you get)!! *no roman numerals!!! example: ! !!! ! ! ! ! !6. Elements:! ! !! !! *Diatomic gases: I2 Br2 Cl2 F2 O2 N2 H2!

pneumonic: "I Bring Clay For Our New House ! (or “7" plus one)!!! * Monoatomic elements: All other elements (Na, K, Fe, etc.)

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Chapter 4.3 Chemical Equations Notes!

!4.3 Law of conservation of mass notes!

!!!!

4.3 Law of conservation of atoms notes!!!!

Chemical Equation:

Word equation:!

Symbolic equation: Skeleton Equation!!!!! reactants products!!Balanced Equation!!!!!coefficient subscript!

States of Matter: !

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Chapter 4.3 Lab!The Mass of Reactants and Products!

!Question: In a chemical reaction, will the reactants weigh the same as the products?!

Background: !

In this experiment, we investigate two chemical reactions. !

1. The reaction between calcium chloride and sodium carbonate:!

! calcium chloride + sodium carbonate → calcium carbonate + sodium chloride!

! CaCl2 + Na2CO3 → CaCO3 + NaCl!

2. The reaction of vinegar (the chemical name for vinegar is acetic acid) and baking soda (the chemical name for baking soda is sodium bicarbonate):!

! baking soda + vinegar → carbon dioxide + water + sodium acetate!

! NaHCO3 +HC2H3O2 → CO2 + H2O + NaC2H3O2 !

We will use this experiment to find out if the reactants weigh the same as the products.!

Pre-lab questions: !

1. How do you know that a chemical reaction have occurred?!

!2. What do you think happens to the mass in a chemical reaction?!

!Materials: !

2 beakers!! ! ! ! ! ! sodium carbonate (Na2CO3)!

Balance! ! ! ! ! ! ! Baking Soda (NaHCO3)!

calcium chloride!(CaCl2)! ! ! ! Vinegar (CH3COOH)! !

Procedure/Data/Observations:!

1. Put on safety goggles.!

Part 1!

2. Add approximately 10 mL of calcium chloride to one beaker and 10 mL of sodium carbonate to another beaker.!

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3. Record the mass of the two beaker. Mass of reactants: _________________!

4. Add the content of one beaker to the other beaker. Record your observations:!! (Change in colour, temperature, formation of bubbles or similar)!

!!

5. Record the mass of the two beakers again. Mass of products: ____________________!

6. Rinse the beakers.!

Part 2!

7. Add approximately 10 mL of vinegar to one beaker and 1 tsp of baking soda to another beaker.!

8. Record the mass of the two beaker. Mass of reactants: _________________!

9. Add the vinegar to the baking soda. Record your observations:!! (Change in colour, temperature, formation of bubbles or similar)!

!!

10.Record the mass of the two beakers again. Mass of products: ____________________!

11.Rinse the beakers.!

Analysis:!

1. How did you know that reactions occurred?!

!2. In part 1 did the mass of the reactants equal the mass of the products?!

!3. What is the law of conservation of mass? How does it relate to this experiment?!

!!4. In part 2, did the mass of the reactants equal the mass of the products?!

!5. In part 2, was mass destroyed?!

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4.3 Conservation of Atoms Activity!1. Pick up a bag of blocks from Ms. Veenstra. !2. Record the colours symbolizing the following atoms (as written on the note in the bag): !! Na =! ! ! H =! ! ! C =! ! ! O =!!3. Make block models of the two reactants in the chemical reaction represented below. !! ! NaHCO3 +HC2H3O2 → CO2 + H2O + NaC2H3O2 !

4. Make a simplified drawing of the reactants in the table below (include colour)!!

5. Countthenumberofdifferentatomsintheallofthereactantsandaddittothebo:omrightofthetable.!

5. Rearrange the blocks to represent the products in the chemical reaction. Make a simplified drawing in the table.!

6. Count the number of different atoms in the all of the products and add it to the bottom right of the table.!

!Questions: !

1. What does a block represent?!

2. What does a collection of blocks represent?!

3. What can you say about the number of blocks before and after the ‘reaction’?!

!4. What is the law of conservation of atoms and how is it related to this activity?!

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Number of the different atoms in the reactants!!Na = H = C = O =!

Number of the different atoms in the products!!Na = H = C = O =!

� ��

�� �

!!!!!insert balancing equation activity

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!!!!insert balancing activity

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4.3 Balancing Notes!Counting atoms!

Mg3(PO4)2! ! ! !

2Ca(NO2)2

!Balancing Basics!

• If a chemical reaction is _____________, then mass, atoms and charge are _________

• The number of each atom is the same on both sides of the arrow.

• Balance by placing ______________________ in front of each compound.

Tips for Balancing!

• Balance _____________________ first and ______________________ last

• Balance hydrogen and oxygen ______________

• You can often treat polyatomic ions, such as SO32-, as a unit

• If an equation is balanced by using half a molecule (1/2O2), you must double all the

coefficients to get whole numbers.

• Examples:

____ AlBr3 + ____ K2SO4 → ____ KBr + ____ Al2(SO4)3

!!!!

C4H10 + O2 → CO2 + H2O !!!

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4.3 Tips for changing word equations to symbolic equations!• Memorize the following:!

! methane =!

! ammonia =!

! water =!

• The following elements are diatomic (they come in pairs when NOT in a compound)!

! H2, N2, O2, F2, Cl2, Br2, I2 ________________________________________!

• All other elements are not diatomic and no subscripted is used. Eg. Pb, Na, V!

!Example: ! !

iron(III) oxide + hydrogen → water + iron!

!!!!

sodium phosphate + calcium hydroxide → sodium hydroxide + calcium phosphate

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