safety matter and change
TRANSCRIPT
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Academic Chemistry Unit I
Safety &
Matter and Change
Name: __________________
Period: ______
Test Date: _______________
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Chemistry Calendar
Monday Tuesday Wednesday Thursday Friday
AUGUST 22 FIRST DAY OF SCHOOL Syllabus & Classroom Rules/Procedures
23 Lab Safety
24 Lab Safety
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26
29
30 31 SEPTEMBER 1 2
5 NO SCHOOL
6 7 8 9
12
13 14 15 16
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20 21 22 23
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27 28 29 30
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DAY ONE ACTIVITY: Classroom Rules & Expectations Classroom Rules:
1. Be respectful In your own words:
2. Be responsible In your own words:
3. Be prepared In your own words:
4. Be safe In your own words:
Classroom Procedures:
1. Bell to Bell‐ You will remain seated and in my classroom from the tardy bell to the end of class. Class begins when the bell rings and you should be working at that time. Class is dismissed when I dismiss you, not when the bell rings.
2. At the Bell‐ Class has begun and you should be in your seat working on the warm‐up. Anyone not present in the room when the bell rings will be marked tardy and will sign the tardy log. This is not the appropriate time to talk to friends, ask to use the bathroom, or sharpen pencils.
3. During Instruction & Independent Practice‐ When I’m talking, that is “my time.” You are to remain seated and quiet unless called upon. When working independently, you are to remain quiet and should be completing work alone.
4. Activities/Labs‐ This is “your time.” When working in groups, you may collaborate and talk about the assignment during this time.
5. Last few Minutes‐ You are expected to participate in any wrap‐up discussion and, when applicable, clean up lab stations or complete & turn in an exit ticket before leaving.
Consequences: 1. Verbal Warning 2. Teacher/Student private conversation 3. Parent contact 4. Written discipline referral
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DAY ONE ACTIVITY: Classroom Rules & Expectations In addition to following the classroom rules and procedures on the previous page, you will also be expected to follow all MCHS and Katy ISD rules, including:
1. Student IDs must be worn at all times. 2. Appropriate dress code must be adhered to. 3. Use of cell phones and electronic devices is not permitted. You will be fined and these devices will be
confiscated if they are being used during class. 4. All lab safety rules outlined in your lab safety contract must be followed in order to participate in lab
activities. Tutorial times: Student Resources:
Teacher Web
Edmodo.com
Notes/Other information:
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LAB SAFETY READING ACTIVITY: You will be given a lab safety scenario card to read through. While reading your scenario, make a list of your safety concerns in the space provided below: Add any additional safety concerns from posters of other groups: HOMEWORK:
1. Read through the rules on your lab safety contract and highlight or underline the rules that you believe are the most important for a chemistry lab.
2. You must also get your lab safety contract signed by a parent/guardian. 3. Lab safety QUIZ will be on .
WARM‐UP: 8/23/11 Identify 5 lab safety violations in the image to the left: 1. 2. 3. 4. 5.
SCENARIO:
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LAB SAFETY WARM‐UP: 8/24/11
1. Look around the room and write the location of the following: a. Goggle cabinet: b. Fire extinguisher: c. Fire blanket: d. Eye wash: e. Safety shower:
MSDS SHEET ACTIVITY: Read through the two MSDS sheets on pages 7‐9, highlighting or circling any vocabulary with which you’re unfamiliar, and then answer the questions below. 1) According to the MSDS, what should you do if you get yeast in your eyes?
2) How is potassium nitrate cleaned up following an accidental spill?
3) How is yeast cleaned up following a spill?
4) Is either of these chemicals flammable? If yes, which one(s)?
5) Does either chemical have an odor? If yes, which one(s)?
6) What are some of the hazardous products formed by the decomposition of yeast?
7) In the event of ingestion of potassium nitrate, what should you do according to MSDS?
8) Is sodium potassium nitrate soluble in water?
9) How should yeast be stored?
When you’re finished, look at the supplies listed on pages 10 & 11 after the MSDS sheets to prepare for QUIZ.
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Laboratory Tools
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Laboratory Tools
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Matter & Change Part 1: Classifying Matter
NOTES #1 – Classifying Matter: Chemistry is the study of the of substances,
& the that substances undergo. Introduction: Western chemistry grew out of old alchemical ideas of Earth, Air, Fire, and Water. As more elements were discovered, ideas changed and by 1800 it was thought that there were two distinct types of matter: organic and inorganic. Most materials obtained from nature, organic and inorganic, are chemically complex, heterogeneous mixtures or composites.
A) A pure substance contains only one kind of matter. Examples of pure substances include:
1) ___________________________ cannot be separated into simpler substances by chemical techniques.
2) ___________________________ are substances that can be separated into simpler substances only by chemical means.
3) Classify each of the following as an element (E) or compound (C):
a. silver____________ c. baking soda __________ e. sugar____________
b. distilled water____________ d. oxygen____________ f. platinum____________
B) A mixture is a physical combination of 2 or more substances that can be separated by physical means.
4) ____________________________, also called solutions, have the same uniform appearance & composition throughout and may be gases, liquids, or solids.
5) ____________________________ do not have uniform composition & consist of visibly different substances or phases.
Types of Mixtures
Heterogeneous Homogeneous
Particle size: ___________________________________
_______________________________
Separation: Filtration or Centrifugation Distillation or Chromatography
Matter(solid, liquid, gas, plasma)
Pure Substances
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Write the following vocabulary words in the appropriate line with the correct definition. Brownian motion Emulsion Homogeneous mixture (solution) Tyndall effect Colloid Suspension Heterogeneous mixture Mixture
Term Definition
6)
physical blend of two or more substances whose compositions varies
7) mixture that is not uniform in composition; It can separate or settle upon standing because of its large particle size (Ex. oil & water).
8) heterogeneous mixture w/ particle size of 100nm or larger (Ex: sand, clay in water, tomato juice).
9) mixture that has the same uniform appearance and composition throughout; may exist as gas, liquid or solid, depending on state of solvent. It does not separate upon standing because of its small particle size (Ex. salt & water).
10) heterogeneous mixture w/ particle size of 1‐100nm that looks like a homogeneous mixture (Ex: Milk, fog, jello).
11) colloidal dispersions made of two liquids (Ex: soap & water, detergent & water)
12) scattering of light caused by colloid particles seen in a beam of light such as dust in air in a "shaft" of sunlight
13) chaotic movement caused by collisions of dispersed colloidal particles with water molecules, preventing the colloid from settling
14) Classify each of the following as a heterogeneous or homogeneous mixture.
a. blood ____________ d. motor oil ____________
b. chocolate‐chip ice cream ____________ e. alcohol in water ____________
c. brass (blend of Cu & Zn) ____________ f. beef stew ____________
ASSESSMENT: Label the following as an element, compound, heterogeneous mixture, or homogeneous mixture.
1. Flat soda pop
2. Salad dressing
3. Sugar
4. Soil
5. Aluminum foil
6. Black coffee
7. Sugar water
8. Pure water
9. Paint
10. Beer
11. Iron
12. Pure Air
13. Spaghetti Sauce
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Matter & Change Part 2: Properties of & Changes in Matter WARM‐UP: _________________________ (Date)
1. Name 3 ways in which I could physically change a piece of paper AND 3 ways in which I could chemically change a piece of paper:
PHYSICAL CHANGES CHEMICAL CHANGES
2. What is the difference between the two types of changes you listed above?
TEACHER DEMO:
3. Describe the properties of the paper before it is torn and set on fire.
4. Describe the properties of the paper after being torn and after being set on fire.
Notes #2 – Properties of & Changes in Matter
1. ________________________ is anything that takes up space and has mass. It is the stuff material things are made of.
Properties of Matter:
2. All substances have characteristics or properties used to identify or describe them.
3. ____________________________________: properties that do not change the chemical nature of matter & are readily observable. ‐ Examples: color, smell, size, melting point, boiling point, luster, solubility and density.
‐ Two types:
Intensive properties: those that do not depend on the size of the sample such as density, freezing point, color, melting point, reactivity, luster, malleability, and conductivity.
Extensive properties: those that do depend on the size of the sample such as length, volume, mass and weight.
4. __________________________________: properties that do change the chemical nature of matter &
are only observable during a chemical reaction. ‐ Examples: flammability, corrosion/oxidation, combustion, reactivity with water, toxicity, radioactivity, sensitivity to light, and pH.
5. Next to each answer, write the letters (P.P.) for physical property or (C.P.) for chemical property.
a. sweet odor______ d. melting point_____ g. flammability_____ j. fermentation_____
b. density______ e. solubility_____ h. hardness_____
c. freezing point_____ f. rusting _____ i. blue color_____
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Changes in Matter: 6. Matter undergo various changes all of the time, such as an increase in temperature or the combustion
of a piece of wood.
7. ______________________________: changes that do not result in the production of a new substance; only temperature, size or physical state is altered. ‐ Examples: melting, freezing, condensing, breaking, crushing, cutting, dissolving and bending.
8. ______________________________: new substances of different chemical composition are produced. ‐ Examples: any chemical reaction, digestion, respiration, photosynthesis, burning, & decomposition.
9. Next to the term, write the letters (P.C.) for physical change or (C.C.) for chemical change.
a. grinding_____ d. decomposing_____ g. breaking_____ j. baking_____
b. melting_____ e. dissolving_____ h. cutting_____
c. condensing_____ f. rusting_____ i. crushing_____
10. Chemical reactions occur when the original substances (called reactants) are consumed and new substances (called products) are formed. Some signs that a chemical reaction has occurred are observable changes in physical properties, such as color change, gas given off (bubbles), solid formed, heat given off, etc. The total mass of the reactants of a chemical reaction will always be equal to the total mass of the products. This is called .
ASSESSMENT: Directions: Label the following as Physical Properties or Chemical Properties.
1. Red Color
2. Density
3. Flammability
4. Solubility
5. Will react with acid to form hydrogen
6. Supports Combustion
7. Bitter Taste
8. Melting Point
9. Will react with water to form a gas
10. Will react with a base to form water
11. Hardness
12. Boiling point
13. Can neutralize a base
14. Luster
15. Odor
Directions: Label the following as Physical Changes or Chemical Changes.
16. Sodium Hydroxide dissolves in water
17. Hydrochloric acid reacts with sodium hydroxide to
produce a salt, water, and heat.
18. A pellet of sodium is sliced in two
19. Water is heated and change to steam
20. Potassium chlorate decomposes to potassium
chloride and oxygen gas
21. Iron rusts
22. Ice melts
23. Acid on limestone produces carbon dioxide gas
24. Milk Sours
25. Wood Rots
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Matter & Change Part 3: Density WARM‐UP: (Date)
RECAP & INTRODUCTION:
Yesterday we learned that extensive properties depend on the size of the sample. Examples of extensive properties are length, volume, mass and weight.
Intensive properties are just the opposite. They do NOT depend on size, like color, odor, temperature, etc. Density is the intensive property we will focus on today.
Density does NOT depend on size.
Density = mass / volume or D = m/V TEACHER DEMO ANALYSIS QUESTIONS:
1. When your teacher cut the piece of paper in half, what happened to the mass of the paper?
2. When your teacher cut the piece of paper in half, what happened to the volume of the paper?
3. If the original mass of the paper was 8.0 g and the original volume was 2.5 cm3, what was the original density of the paper? (Show your work)
4. Assuming the information given above is correct, what would the NEW density be for just one half of the original piece of paper? (Show your work)
5. When your teacher doubled the amount of water in the beaker, what happened to its mass?
6. When your teacher doubled the amount of water in the beaker, what happened to its volume
7. If the original mass of the water was 150.0 g and the original volume was 150 mL, what was the original density of the water? (Show your work)
8. Assuming the information given above is correct, what would the NEW density be for double the original amount of water? (Show your work)
GUIDED PRACTICE PROBLEMS: All problems must be set up and all work must be shown.
Formula: D = m/V
9. A sample of gas has a mass of 0.0707 grams and has a volume of 48.9 milliliters. What is the density of this gas?
1. X = 5/3, solve for X.
2. 7 = Y/49, solve for Y.
3. 5.7 = P/0.3, solve for P.
4. 11.3 = 64/Q , solve for Q.
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10. A rock has a mass of 30.5 grams and, when placed in a beaker of water that has 20.0 ml in it, the volume increases to 22.9 ml. What is the density of the rock in grams per centimeter cubed?
11. Ammonium magnesium chromate has a density of 1.84 g/cm3. What is the mass of 6.96 cm3 of this substance?
12. Iron has a density of 7.87 g/cm3. What volume would 12.5 grams of iron occupy? INDEPENDENT PRACTICE PROBLEMS: All problems must be set up and all work must be shown.
1. If a piece of ivory had a mass of 10.22 kilograms, and took up 5347.5 cm3, what would the density be?
Answer:_________________
2. A weather balloon is inflated to a volume of 2200 Liters with 37.4 grams of helium. What is the density of helium in grams per liter?
Answer:_________________
3. Calcium chloride has a density of 2.15 g/cm3. What is the volume of 3.37 grams of this substance?
Answer:_________________
4. Barium perchlorate has a density of 2.74 g/cm3. What is the mass of 610 cm3 of this substance?
Answer:_________________
5. Cerium sulfate has a density of 3.17 g/cm3. What is the volume of 706 grams of this substance?
Answer:_________________
6. Bismuth phosphate has a density of 6.32 g/cm3. What is the mass of 86.0 cm3 of this substance?
Answer:_________________
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Matter & Change Part 3: Density of Sand vs. Gold Lab
Background: As you saw in the film, Indiana Jones tried to steal the gold idol without triggering the booby trap – and failed. He thought he could replace the gold idol on the pedestal with a bag of sand of an equal volume. Why didn’t the plan work? Purpose:
To calculate the density of sand.
To compare the densities of sand and gold, and then calculate the volume of sand Indiana Jones should’ve used in his plot.
Procedure:
1. Obtain a sample of sand. Determine its mass and record. (Remember to use weigh paper or a suitable container)
2. Measure the volume of your sand using a graduated cylinder. (Use precision in your measurement) 3. Record your results in the class data table and copy the remainder of the class data table. 4. Plot the class data on the following page (mass on the y‐axis and volume on the x‐axis). Calculate the
density of sand by drawing a best‐fit line and finding its slope. 5. Your graph should have labeled axes, a title, and a best‐fit line. Do NOT connect the dots. Be neat.
Group Data: Group # ________ Mass of sand sample = grams Volume of sand sample = milliliters Class Data:
Group # Mass (g) Volume (mL)
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2
3
4
5
6
7
8
9
10
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Analysis: (SHOW ALL WORK) 1. What is the slope of the line obtained from the
class data? Note: slope = y2‐y1 X2‐x1
2. The density of gold is 19.3 g/cm3. If the idol had a volume of 1,500 cm3, what was its mass in grams?
3. Indiana needed to swap the idol with a bag of
sand that had the same mass. What volume, in milliliters, of sand did he need to use?
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Matter & Change Part 3: Density of an Unknown Lab Problem: What is the density of a penny and what metal is it primarily made from? Research & Background Knowledge:
Metal Density (g/cm3)
water 1
aluminum 2.70
zinc 7.13
nickel 8.8
copper 8.96
silver 10.49
lead 11.36
mercury 13.55
gold 19.32
Hypothesis:
If I calculate the density of pennies, then I hypothesize that their density will be approximately __________ g/mL
because I believe pennies are made from _________________________ (metal(s) from chart).
Procedure:
1) Find the mass of the three different sets of DRY pennies using the scale. Record in data table. 2) Find the volume of the three different sets of pennies using the graduated cylinder. Add water and read
the volume of water in the graduated cylinder. Tilt the graduated cylinder, CAREFULLY add the pennies to it, and read the volume of the water after adding the pennies. Record in data table.
3) Dry off the pennies, clean up lab station, and return to your seat to calculate the volume of the pennies alone and the three density values.
Data:
# Pennies Mass (g) Volume of water (mL)
Volume of water and pennies (mL)
Volume of pennies (mL)
Density (g/mL)
Average Density:
Analyze: Compare your numbers with the chart of known densities in the research section above. Conclusion:
After measuring the mass and volume of the pennies, I found the average density of the pennies to be
______________g/mL which is closest to the density of ________________ _ . My hypothesis
was therefore .
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Matter & Change Part 4: States of Matter WARM‐UP: _________________________ Fill in information in the following chart for the cards you were given in your Ziploc bag:
SUBSTANCE FROM CARD:
Property 1) 2) 3)
Shape (Describe its shape, can it change?)
Volume (Describe its size, can it change?)
Compressibility (Can it be squeezed easily?)
Expansion on heating (does heat change size?)
Molecular Arrangement (ie, close together, far)
State of Matter (solid, liquid, gas, or plasma)
States of Matter:
1. The four states of matter are . 2. What do melting, freezing, boiling, and condensing have in common?
_____________________________________________________________________________________
____________________________________________________________________________________. 3. The Nature of Gases: Gas molecules are .
The hotter the gas becomes, the faster they move. When heated, a gas will either expand or its pressure will increase.
4. The Nature of Liquids: Liquid particles, like gases, are in motion & flow because they are free to slide past each other.
The weak attractive forces between the molecules are called , which prevent most from escaping the liquid phase.
Liquids & solids are denser than gases and can’t be compressed because the molecules are packed closely together.
Distinguish between a gas and a vapor:
______________________________________________________________________________
5. The Nature of Solids: Solids move through about fixed
points.
Ionic solids have high melting points because they are held together by stronger forces, which are harder to break.
Molecules of solids are held in fixed positions in a very regular lattice arrangement.
The hotter the solids and liquids become, the more they vibrate. This causes solids and liquids to expand slightly when heated.
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6. Based on your notes above, complete the following table:
Property Gas Liquid Solid
Shape indefinite
Volume indefinite
Compressibility almost incompressible
Expansion on heating great
Molecular Arrangement close, but not rigidly packed
Two Examples
Changes in State:
7. ANY change in state is a CHANGE. 8. Boiling, evaporation, & perspiration are cooling processes because the particles with the highest KE
(kinetic energy) escape first to the gas phase. The liquid particles left have a lower average KE, decreasing the temperature.
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Match the following terms with the correct definition: 9. ______ vaporization a) boiling point of liquid at pressure 101.3 kPa (1 atm) 10. ______ evaporation b) conversion of gas to liquid 11. ______ vapor pressure c) conversion of liquid to gas or vapor 12. ______ condensation d) temperature when vapor pressure of liquid is equal to 13. ______ boiling point external pressure 14. ______ normal boiling point e) conversion of liquid to gas without boiling 15. ______ dynamic equilibrium f) force produced above the liquid surface when liquids change 16. ______ melting point to gas strike (collide) with the walls of sealed container 17. ______ freezing point g) change from a solid to vapor without passing to liquid 18. ______ sublimation h) temperature at which liquid converts to solid
i) rate of evaporation (rate of condensation when closed container is saturated with vapors) j) temperature at which solid changes to liquid
19. Study the diagram above, and then explain what you think vapor pressure means in your own words:
20. Condensation forms on the warmer side of a container. One morning, you go to your car and find that
the windshield is “fogged.” When you touch it, you notice that the condensation (liquid) is on the inside of the car. What does that tell you about the temperature of the outside and inside of your car? Should you turn on the heater or the air conditioner to get rid of the condensation?
_____________________________________________________________________________________
21. Why do heated liquids evaporate faster?
_____________________________________________________________________________________
22. Why is gasoline more volatile than water? Relate to increased vapor pressure.
_____________________________________________________________________________________
23. Is it possible to boil liquid without increasing temperature?
_____________________________________________________________________________________
24. Why does water boil at lower temperature in Denver than Houston?
_____________________________________________________________________________________
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Academic Chemistry Unit 1 Test Conceptual Review Prentice Hall Chemistry Textbook – Chapter 2 (pp. 28 – 49)
TEST on
1) Lab safety General lab safety rules
Procedure for mixing acids & water – Add acid to water (hint: think A&W root beer)
2) Structures of matter Pure substances – Elements and compounds
Mixtures – Homogenous and heterogeneous
3) Properties of matter Physical properties – Do not change chemical nature of matter Two types: Intensive (doesn't depend on size) or extensive (depends on size)
Chemical properties – Changes chemical nature of matter thru a chemical reaction
4) Changes in matter Physical changes – do not result in production of a new substance
Chemical changes – results in the production of a new substance
5) Law of conservation of mass The total mass of the reactants equals the total mass of the products
6) States of matter Properties of solids, liquids, gases, and plasmas Shape, volume, molecular arrangement, expansion upon heating, etc
Changes between states of matter
7) Density Definition
Calculate density, given m and v. (D = m/v)
Calculate volume, given D and m
Calculate mass, given D and v
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Prentice Hall Chemistry Textbook – Chapter 2 (pp. 28 – 49) TEST on
Lab safety 1. When diluting acids, should you add water to acid or acid to water?
2. What would be the proper safety procedure when a lab produces too many fumes? 3. Why do you waft chemical vapors rather than inhale them directly?
Structures of matter 4. What is a compound?
5. What is matter?
6. What is the difference between a homogenous and heterogeneous mixture?
7. Which type of mixture is also known as a solution?
Properties of matter 8. What are physical properties? Give 5 examples.
9. What are chemical properties? Give 3 examples.
10. How will the density of 100mL of water change when half is poured out?
11. What is the difference between intensive and extensive properties? Give examples of each.
12. Identify as physical or chemical properties: A. luster B. color C. solubility D. oxidation E. ability to rot/spoil F. flammability G. boiling point H. freezing point I. composition (what elements make it)
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Changes in matter 13. What is a physical change?
14. What is a chemical change?
15. Identify as physical or chemical changes:
A. changing color of leaves B. spoiling of milk C. ability to dissolve sugar in water D. rusting of an old bike E. reaction with acid F. polishing silver G. water boiling to form steam
Law of conservation of mass 16. If in the following reaction you end up with 30 grams of iron III oxide and had 8 grams of oxygen
before the reaction, what mass of iron did you begin with? 4Fe + 3O2 ‐‐‐> 2Fe2O3
17. What does the law of conservation of mass say?
States of matter 18. Which state(s) of matter have a definite volume?
19. Which state(s) of matter have a definite shape?
20. What happens to the movement of the molecules and the volume of a gas when it is heated?
21. When water evaporates to a gas, how does the change in volume affect the density?
Density 22. If the mass of a substance is 12.5 g and the volume is 7.3mL, what is its density?
23. If the mass of a substance is 450 g and its density is 3.59 g/cm3, what is the volume?
24. If the volume of a substance is 7.2cm3 and its density is 17.25 kg/cm3, what is its mass? 25. If these three substances (from # 21, 22, & 23) were liquids and they were all poured into the same
beaker, then comparing their densities, which substance would sink to the bottom, which would be in the middle, and which would float on top?