staar-eoc · station 1 . physical and chemical properties. 1. locate the physical and chemical...

SLHS Chemistry

Activity Booklet

STAAR-EOC

Chemistry

Review

Table of Contents STAAR-EOC

2 Chemistry

Table of Contents

Activity A – Properties of Matter & Nuclear Chemistry 3

Activity B – Pure Substances/Mixtures &

Thermochemical Equations 13

Activity C – Chemical Families/Periodic Trends &

Naming Compounds/Chemical Formulas 24

Activity D – Atomic Structure and Electron Dot Formulas &

Molecular Geometry 32

Activity E – Moles and Gas Laws 39

Activity F – Balancing Equations & Types of Reactions 51

Activity G – Types of Solutions & Solubility Rules 59

Additional Notes - Things I need to remember 71 Credits 77

Properties of Matter & Nuclear Chemistry STAAR-EOC

Review 3

Activity A Properties of Matter & Nuclear Chemistry Station 1 Physical and Chemical Properties

1. Locate the Physical and Chemical Properties Cards. Sort the

cards into two groups—physical properties and chemical properties. Then, separate the Physical Properties Cards according to whether they represent extensive or intensive properties.

2. Record your work on the Physical and Chemical

Properties Table below, writing the title of each numbered card under the category it belongs to.

HINT: 4 Extensive, 7 Intensive, 4 Chemical.

Physical Property

Chemical Property Extensive Intensive

1

2

3

4

5

6

7

8

9

10

11

12

13

14

15


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3. How did you determine whether each card exhibited a physical property or a chemical property?

4. How did you differentiate between an intensive and

extensive property?


Review 5

Station 2 Physical and Chemical Change 1. Locate the Physical and Chemical Change Cards. Determine

whether the cards exhibit a physical or chemical change. Some observable changes may be the result of a physical and a chemical change. Sort the cards into three groups: Physical Change, Chemical Change, or Chemical and Physical Change.

2. Record your answers in the table below by placing an “x” in

the appropriate column. HINT: 3 Physical, 4 Chemical, 3 Both.

Change Physical Chemical Chemical and

Card Change Change Physical Change

A

B

C

D E F G

H

I

J 3. What evidence did you use to classify Card G?


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4. What evidence did you use to classify Cards D, E, and F?

5. How do the physical properties of matter differ from the

chemical properties of matter?


Review 7

Station 3 Solids, Liquids, and Gases 1. Locate the laminated Properties of Solids, Liquids, and

Gases Sheet and the Properties of Solids, Liquids, and Gases Cards. Remove the cards from the envelope and place them on the correct column of the laminated sheet. Record your placements below.

HINT: 5 Solid, 5 Liquid, 6 Gas


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2. If the same amount of pressure were applied to each of the

cylinders below, which state of matter would show the greatest change in volume? Which would show the least? Justify your response.

3. What properties can scientists use to describe and classify

matter?


Review 9

Station 4 Characteristics of Radiation 1. Locate the Characteristics of Radiation Cards. Sort the

cards into categories according to which type of radiation each card describes. Record your answers in the table below. HINT: 7 Alpha, 6 Beta, 6 Gamma

Which type of radiation presents the greatest risk? Why?


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Station 5 Nuclear Equations 1. Locate the Nuclear Equations Card and the Nuclear

Equations Strips. Examine the equations and strips and determine which strip contains the substance that will balance each equation. Insert the strip into the pocket behind the equation. Slide the strip until the substance that balances the equation shows in the window. Write your answers in the blank windows below.

Example:

How do you know when a nuclear equation is balanced?

3. Write balanced nuclear equations for the following:

Carbon-14 undergoes beta decay.

Radium-236 undergoes alpha decay.


Review 11

Station 6 Fission and Fusion

1. Locate the Fission–Fusion Venn Diagram and the Fission–Fusion Cards. Place the cards in the appropriate spaces on the diagram. Record the results from the Venn Diagram in the table below. HINT: 6 Fission, 1 Both, 5 Fusion

Fission Both Fusion


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2. Nuclear reactions change the structure of nuclei and involve

a tremendous amount of energy. The sun is fueled by nuclear reactions. Nuclear power plants are also fueled by nuclear reactions.

A. What changes occur in the nuclei of atoms involved in

the reactions that fuel the sun?

B. What changes occur in the nuclei of atoms involved in the reactions that fuel nuclear power plants?

Pure Substance/Mixtures & Thermochemical Equations STAAR-EOC

Review 13

Activity B Pure Substances/Mixtures & Thermochemical Equations

Station 1 Classifying Pure Substance and Mixtures

Locate the six plastic bottles containing different substances at this station at this station. Examine the substances in each of the bottles without removing the lids. You may use the hand lens and magnet to make observations. Rotate the sleeve attached to the bottle to reveal information about the substance inside the bottle. Use this to determine whether the substance is a pure substance or a mixture.

1. Which bottles contain pure substances? Which contain mixtures? Record your responses in the table below by placing an “x” in the appropriate column. HINT: 3 of each.

Bottle Name of Substance Pure Mixture Substance

A

B

C

D

E

F


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2. What properties helped you determine which bottles contained pure substances or mixtures?

3. How are mixtures that are solutions identified?

4. Which bottles contain mixtures that are solutions? What can

be done to separate the parts of the solutions?


Review 15

5. Classify the following diagrams as pure substances or mixtures and justify your choices in the Classifying Substances Table below.

A B C D Classifying Substances Table

Pure Substance Justification

or Mixture

A

B

C

D


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Heterogeneous and Homogeneous Mixtures

Examine the samples of mixtures below.

6. Which samples represent heterogeneous mixtures?

7. Which samples should be classified as homogeneous

mixtures? What is the justification for this classification? 8. When iron and sulfur are heated together in a test tube over

high heat until they are completely reacted, they form a new substance called iron sulfide. Do you think the iron and sulfur can be returned to their original condition? Is the iron sulfide a pure substance or a mixture?


Review 17

Station 2 Energy Cards and Thermochemical Data Table Use the Energy Cards to complete the table below. Record each type of energy in its proper category. HINT: 3 in each category.

Form of Energy Kinetic Potential

1. Based on your choices above, come up with definitions of kinetic energy and potential energy, in your own words.


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Thermochemical Data Table Locate the Thermochemical Data Table and use it to answer the questions below. 2. Calculate the enthalpy change for the following reaction: C(s) + O2(g) CO2(g)

3. What would the enthalpy change be if 2 moles of carbon were

burned?

4. What does the sign of ΔH tell you about the combustion of carbon?


Review 19

5. Calculate the enthalpy change for the production of 1 mole of iron oxide (rust) according to the following equation:

4Fe(s) + 3O2(g) 2Fe2O3(s)

6. The combustion of 1 mole of glucose in cellular respiration releases 2,808 kJ of energy according to the following equation:

C6H12O6(s) + 6O2(g) 6CO2(g) + H2O(l) ΔH = –2,808 kJ

What is the enthalpy of formation of 1 mole of glucose?

7. Is this an exothermic or endothermic process?


20 Chemistry

8. The temperature of a 0.25 Kg block of an unknown metal increases by 22.22°C after 5,000 J of heat energy is added. What is the unknown metal?

Use the Thermochemical Data Table and the information about aluminum, copper, and zinc metals in the table below to determine your answer. You must show your calculations.

Metal Specific Heat Aluminum 0.900 J/g•°C

Copper 0.385 J/g•°C Zinc 0.388 J/g•°C

9. How do energy changes occur during chemical reactions?


Review 21

Station 3: Endothermic and Exothermic Changes 1. Use the graduated cylinder to measure 10 mL of distilled

water. Pour the water into one of the test tubes. Use the thermometer to measure the temperature in degrees Celsius, and record it in Data Table A.

Obtain the ammonium nitrate, NH4NO3, sample from your teacher. Use the balance to carefully measure 5 grams of ammonium nitrate, NH4NO3, and add it to the test tube containing the water. Stir the mixture with a glass rod to dissolve all the crystals. Place the test tube in the test tube rack. Carefully insert a thermometer into the test tube. Observe for two minutes then measure the temperature of the solution and record it in Data Table A.

Data Table A

Mixture Initial Temp (distilled water) Final Temp

(after 2 min)

NH4NO3

2. How did the energy change as the ammonium nitrate

dissolved in the water? 3. Was the change endothermic or exothermic?

SAFETY: Safety goggles MUST be worn for this activity.


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Pour 10 mL of distilled water into the second test tube. Measure the temperature of the distilled water and record it in Data Table B. Add 5 grams of calcium chloride, CaCl2, to the test tube containing the distilled water. Stir the mixture with a glass rod to dissolve all of the substance. Place the test tube in the test tube rack. Carefully insert a thermometer into the test tube. Observe for two minutes then measure the temperature of the solution and record it in Data Table B.

Data Table B

Mixture Initial Temp (distilled water) Final Temp

(after 2 min)

CaCl2

4. How did the energy change as the calcium chloride dissolved

in the water?

5. Was the change endothermic or exothermic?


Review 23

6. Examine the graph.

What is the change in enthalpy for the chemical reaction pictured in the graph?

Is this an exothermic or endothermic chemical reaction? Why?

SAFETY: Properly dispose of all chemicals in the sink. Rinse and dry all glassware.

Chemical Families/Periodic Trends & Naming Compounds/Chemical Formulas STAAR-EOC

24 Chemistry

Activity C Chemical Families/Periodic Trends & Naming Compounds/Chemical Formulas Station 1 Periodic Table of Elements Cards 1. Arrange the Periodic Table Cards in a logical order on the

metal sheet, creating a Periodic Table of the Elements 2. What “clues” did you use to help arrange the periodic table?

3. In what way did your method compare and contrast with how

Russian chemist, Dimitri Mendeleev, arrange the first periodic table of elements?


Review 25

Station 2 Periodic Trends 1. Using the information on the cards, place the Periodic

Trend Arrows around the periodic table, showing the direction of increase in each trend.

2. Electronegativity is the ability of an atom to attract

electrons. Based on the information on the Periodic Table Cards, which element has the greatest electronegativity?

3. Explain the trend in electronegativity as elements go . . .

Down in a group:

Across in a period:

4. Ionization energy is the energy required to remove an

electron from an atom. Based on the Periodic Table of the Elements, how do the ionization energies of Group 1 compare to the ionization energies of Group 17?

What causes the differences in ionization energies for these two groups?


26 Chemistry

5. Positive ions (cations) tend to be smaller than their

corresponding neutral atoms. What is a possible explanation for this?

6. Negative ions (anions) tend to be larger than their

corresponding neutral atoms. What could be an explanation for this?

7. The following table shows the ionic radius for elements in

Periods 2 and 3:

Period Li+ Be2+ B3+ C4+ N3- O2- F-

2 60 31 20 15 171 140 136

Na+ Mg2+ Al3+ Si4+ P3- S2- Cl-

3 95 65 50 41 212 184 181

What is the general trend that occurs across each period?

8. What do you predict the trend will be for ionic radius down a group?


Review 27

Station 3 Chemical Families Locate the Periodic Table Labels.

1. Place the Periodic Table Labels for alkali metals, alkaline

earth metals, transition metals, halogens, and noble gases on the large Periodic Table of the Elements.

2. Why is it important to know the properties of different groups

of elements in the Periodic Table of the Elements?

3. Compare the reactivity of alkali metals and noble gases.

What property of these groups of elements accounts for the differences in their reactivity?


28 Chemistry

4. Compare the reactivity of alkali metals and halogens. What accounts for their differences?

5. How can the Periodic Table of the Elements be used to

predict periodic trends in chemical families and periods?


Review 29

Station 4 Naming Compounds 1. Locate the How Compounds Are Named Flowchart and use

it to name the compounds in the table below.

Compound Name

ClO2

HCl(aq)

NH4NO3

Pb(NO3)2

Fe2O3

H2CO3

C2Br6

HNO3(aq)

HCl(g)

KOH(aq)


30 Chemistry

Station 5 Chemical Formulas (Covalent Bonds) 1. Using the How Compounds Are Named Flowchart to

complete the table by writing the chemical formulas for the following covalent compounds.

Example: Tetrasulfur dinitride – S4N2

Covalent Compound Name Covalent Compound Formula

Dinitrogen trioxide

Carbon tetrachloride

Disulfur trifluoride

2. Use the Covalent Compound Cards to determine the

chemical formulas for the following compounds:

a. A compound with carbon and fluorine:

b. A compound with silicon and hydrogen:

c. A compound with sulfur and bromine:


Review 31

Station 6 Chemical Formulas (Ionic Bonds) 1. Locate the Anion and Cation Cards. Match the cards and

arrange them on the baking sheet so that they represent the neutral compounds listed in the table. Complete the table by giving the correct chemical formulas for the compounds.

Example: Copper (I) phosphate – Cu3PO4

Ionic Compound Name

Ionic Compound Formula

Copper (I) sulfate

Calcium carbonate

Coppr (II) nitrate

Iron (III) oxide

Ammonium phosphate

2. What are three other ionic compounds you can make using

these cards?

3. What features of compounds and chemical formulas do

scientists use to determine the names for the compounds and formulas?

Atomic Structure & Electron Dot Formulas/Molecular Geometry STAAR-EOC

32 Chemistry

Activity D Atomic Structure & Electron Dot Formulas and Molecular Geometry Station 1 Atomic Structure

1. Look at the Periodic Table. Find the block that represents

carbon. Using the Energy Levels Table, write the complete electron configuration for carbon:

Locate the Atom Board, which is a model that represents the energy levels around the nucleus of an atom. The marbles represent the subatomic particles:

• Blue – protons • Green – neutrons • Yellow - electrons

Locate the envelope containing the Element / Ion Isotope Cards, as shown below.


Review 33

2. Using the Atom Board, build models of the atoms and ions shown on the Element / Ion Isotope Cards and record the electron configurations in the following table. The Energy Levels Table may also be useful.

Atomic Number Element Electron Configuration

1 H

2

3

9

10

16

18

1 H+

11

12

17

Electron configurations can also be written as noble gas electron configurations. The noble gas before Silicon is Neon, so its noble gas configuration is [Ne]3s23p2. 3. What is the noble gas configuration for

Calcium?

Aluminum?


34 Chemistry

Valence electrons are represented for each element as Lewis Dot Structures. Carbon’s electron configuration is 1s22s22p2. 4. Draw the Lewis Dot Structures for Elements 1–18 in the table below.

5. What correlation is there between the group numbers and

the number of valence electrons you have drawn?

6. What is the relationship between the Periodic Table and the

electron configurations of atoms?


Review 35

Station 2 Electron Dot Formulas

1. Locate the Electron Dot Atom Cards. Arrange these cards to show the electron dot formulas for the covalent molecules listed in the table.

For example, locate two hydrogen cards and one oxygen card. Slide the cards together to make the H2O electron dot configuration shown below.

2. After you have constructed the physical models,

draw the electron dot formulas in the space provided below. Notice the example of H2O has been provided. You may use a dashed line to represent a pair of bonding electrons between two atoms. Chemical Formula Electron Dot Formula

H2O

BeH2

CH4

BF3

NH3

CO2


36 Chemistry

3. When drawing Lewis Structures, why is it important to know the number of valence electrons for each atom involved in covalent bonding?

4. What is unique about the Lewis Structures for BeH2 and

BF3?

5. Lewis structures for ionic compounds can be represented in

several ways. For example, the Lewis Structure for LiF can be represented as:

Draw the possible Lewis Structures for the following ionic compounds:

Chemical Formula Lewis Structure

NaCl

MgI2

6. Why is an “x” used to represent some of the electrons?


Review 37

Station 3 Molecular Geometry 1. Locate the molecular geometry models and the VSEPR

Chart. Use these to determine the molecular structure for each of the following molecules. Record the letter of the model that matches each component in the table below.

Molecule Electron

Geometry Number of Lone

Pairs of Electrons

Model Letter

H2O

BH3

HCN

NH3

CH4

SO2

3. Compare the electron geometries of CH4 and H2O. How are they similar and how are they different?


38 Chemistry

3. How are the molecular geometries of CH4 and NH3 similar?

How are they different?

Why are they different?

4. How is the behavior of electrons in an ionic bond

different from the behavior of electrons in a covalent bond?

Moles & Gas Laws STAAR-EOC

Review 39

Activity E Moles & Gas Laws Station 1 Units of Measure and Units of Conversions 1. Locate the Units of Measure Cards in the envelope.

Arrange the cards into sets so that you create correct conversion factors. Record your results below. Example: 1 L = 1000 mL.

Given Quantity and Unit = Value Unit

1 L = 1000 mL

2. Locate the Unit Conversion Cards and the Unit Conversion

Calculations Sheet. Arrange the cards to solve problems with dimensional analysis. Record your results below.

Given (in moles) x Conversion Factor =

Number of Atoms, Ions, or Molecules


40 Chemistry

Extension question: Consider that you have 2.5 moles of calcium chloride, CaCl2. How many chloride ions are in that sample? Set this up using dimensional analysis and solve below. How does your answer compare to the number of chloride ions in 2.5 moles of NaCl (see table on previous page)?

3. The periodic Table shows the grams per mole of every element. For example, the mass of carbon is 12.011 grams per mole. The mass of one mole of the compound HCl can be determined by adding the mass of H and the mass of Cl to get 36.461 grams per mole. Complete the following table using the information on the Periodic Table.

Substance Molar Mass (Grams per Mole)

Cu

Ar

NaOH

NaHCO3

Pb(NO3)2


Review 41

Station 2 Molar Mass

1. If the mass of a substance is known, then the number of

moles it contains can be determined using the substance’s molar mass from the Periodic Table.

Example: 10.0 grams Mg x 1 mole Mg = 4.11 x 10-1 moles Mg

24.305 grams Use the electronic balance to determine the mass of the three samples at the station: copper metal sample, 2 scoops baking soda, and aluminum can. Be sure to subtract the mass of the beaker when weighing the baking soda. Determine how many moles of each substance you have. Record your data below.

Sample Mass (grams) Moles

Cu

NaHCO3

Aluminum soda can

2. How many atoms of Cu are in the sample you measured?


42 Chemistry

3. The compound sodium hydrogen carbonate (NaHCO3) is formed from Na+ ions and HCO3– ions. How many ions of HCO3– are in the sample you measured?

4. Locate the aluminum soda can at the station. How many

atoms of Al are in the soda can? 5. If 1.0 mL of H2O has a mass of 1.0 g, how many molecules

are in 100 mL of H2O?

Extension: How many atoms of hydrogen are in the water sample above?

6. What is the unit used to convert grams of a substance to

atoms, ions, or molecules? What numerical value does this unit represent?


Review 43

Station 3 Balloon in a Bottle

1. Place a balloon into the opening of a plastic water bottle. Stretch the balloon opening over the mouth of the plastic water bottle so that the balloon is suspended inside the bottle.

Locate the small hole on the side of the water bottle and cover it with your finger. Keeping the hole covered, place your mouth over the opening and try to inflate the balloon by blowing into the opening. Explain what happens.


44 Chemistry

2. Remove your finger from the hole and try again to inflate

the balloon by blowing into the opening of the plastic water bottle. This time, the balloon will inflate. Place your finger over the hole as soon as the balloon is inflated. Why does the balloon stay inflated even though the top of the balloon is open?

3. Remove your finger from the hole. Explain what you observe.


Review 45

Station 4 Charles’ Law

1. Locate the flask with the balloon apparatus in the hot water bath and the flask with the balloon apparatus in the cold water bath.

Describe the difference in the two balloons.

2. Switch the two flasks so that the flask from the warm water bath is in the cold water and the flask from the cold water is in the warm water. Wait several minutes, then describe what has happened to the balloons.


46 Chemistry

3. Explain how the changes in the two balloons in Steps 1–3 illustrate Charles’ Law.

4. Leave one flask in the ice bath and one flask in the warm water bath for the next group.

5. Application: A balloon has a volume of 2.5 L on a sunny day, when the temperature is 30oC. If the air temperature drops to 10oC overnight, what will the volume of the balloon be?


Review 47

Station 5 Boyle’s Law 1. Locate the 1,000-mL flask with the balloon inside. Do not

remove the stopper or pump from the flask.

2. Carefully pump the handle of the balloon pump several

times until it becomes difficult. What happens to the balloon? How can you explain this?

3. Pull the pump out of the stopper. Explain how the

changes in the balloon in Steps 1-2 illustrate Boyle’s Law.


48 Chemistry

4. Application: A balloon is filled with 25 L of air in Corpus Christi, where the pressure is 1.0 atm at sea level. What will the volume of the balloon be in mountains of Denver, where the pressure is 0.85 atm?


Review 49

Station 6 Other Gas Laws

1. Locate the Gas Law Problem Cards and the STAAR

Chemistry Reference Materials. Determine which formula is needed to solve each of the problems. Record your answers below.

Card Formula

A

B

C

D

E

F

G

H

I

J

2. Select any one of the cards and justify why the formula you

selected will solve the problem.


50 Chemistry

Gas Stoichiometry

3. Write a balanced equation for the production of water from O2 and H2.

4. How many moles are in 100. g of O2?

5. How many liters of water vapor can be made from 3.13 moles of oxygen gas with excess hydrogen at STP?

6. How are pressure, volume, and temperature of gases

related?

Balancing Equations & Types of Reactions STAAR-EOC

Review 51

Activity F Balancing Equations & Types of Reactions Station 1 Balancing Equations

1. Locate the Balancing Equations Sheet. Use the dry erase marker to predict the products of the equations.

2. Locate the double pan balance and the bag containing the

Element/Ion Cubes (Unifix cubes). Use the cubes to help you determine the coefficients needed to balance the equations as shown below.

Write the balanced equations on the laminated Balancing Equations Sheet, using a dry erase marker. Record your answers below. The pan balance is VERY important. The scale will balance only when the equation is correctly balanced.


52 Chemistry

3. What is the relationship between the mass of the reactants and the mass of the products?


Review 53

Station 2 Conservation of Mass 1. Locate the double pan balance, the bags labeled Reactants

and Products, and the Masses of Substances Sheet. Do not open the bags or change the arrangement of the cubes. Place the Reactants bag on the left side of the balance and the Products bag on the right. Record your observations below.

2. Carefully examine the contents of the two bags and consult

the Masses of Substances Sheet. According to your observations, what is the reason for the discrepancy in mass between the two bags? HINT: It may be helpful to write a balanced equation for this chemical reaction.

3. From your observations of the contents of the bags, which

reactant is limiting, and why?


54 Chemistry

4. The following investigation was conducted using two

different procedures.

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)

Magnesium ribbon is added to 1.0 Molar hydrochloric acid in a flask In the first setup, a balloon covers the opening of the flask. In the second no balloon is used.

How would you expect the final mass of the products to differ in the two flasks after the reaction is complete? How do you account for the difference?

5. If you repeated the above experiment using 0.10 moles of

HCl and 0.25 grams of Mg, which reactant would be the limiting one?


Review 55

6. The above experiment is performed using 0.25 grams of

magnesium ribbon with excess hydrochloric acid. If o.018 grams of hydrogen gas are collected, what is the percent yield?

7. How can the amounts of chemicals needed for, or

produced from, a chemical reaction be calculated?


56 Chemistry

Station 3 Types of Reactions 1. Locate the Reaction Cards. Sort the cards into the following

three reaction categories. Record the letter of each Reaction Card in the appropriate column on the table below. HINT: One card will be used twice. 8 Acid-Base, 7 Precipitation, 8 Redox.

Acid-Base Precipitation Oxidation/ Reduction

2. Which reaction type has the most different variations? What

are the variations?


Review 57

Scenarios 1. Consider the following four scenarios. Using your knowledge

of chemical reactions, write the balanced chemical equation for the reaction, determine what type of reaction occurs, and provide evidence to support your choice.

Scenario 1 – Several drops of hydrochloric acid are added to a strip of magnesium metal, producing a gas. Record your answers in the table below.

Balanced Equation

Type of

Reaction

Supporting Evidence

Scenario 2 – Sodium hydroxide is mixed with copper II nitrate in a test tube. Record your answers in the table below.

Balanced Equation

Type of Reaction

Supporting Evidence


58 Chemistry

Scenario 3 – Vinegar, HC2H3O2, is mixed with sodium hydroxide, NaOH. Record your answers in the table below.

Balanced Equation

Type of Reaction

Supporting Evidence

Scenario 4 – Electricity is produced in a battery according to the following equations:

What type of reaction occurs in the battery? What evidence supports your choice?

2. What characteristics can be used to identify different types

of aqueous chemical reactions?

Types of Solutions & Solubility Rules STAAR-EOC

Review 59

Activity G Types of Solutions & Solubility Rules Station 1 Types of Solutions 1. Pour 100 mL of distilled water into the 250-mL beaker. Use

the thermometer to measure the temperature of the water and record it below. Following proper laboratory procedures, measure out 100 grams of NaCl onto the weighing paper. Measure out 30 grams on another piece of weighing paper from the 100 grams on the weighing paper and add it to the distilled water.

Water temperature

2. Using a plastic spoon or scoopula, slowly add additional

NaCl from the remaining 70 g to the solution while constantly stirring with the stirring rod. Do not use the spoon to stir the solution.

Continue adding NaCl until no more will dissolve—a few undissolved grains will be left at the bottom of the beaker.

3. Determine the mass of the NaCl remaining on the

weighing paper from the original 100 grams and the mass of the NaCl used, and record them below.

Mass of NaCl remaining

Mass of NaCl used


60 Chemistry

4. Examine the Solubility Graph. Is the amount of NaCl

previously dissolved in the distilled water a reasonable amount to make a saturated solution? Justify your answer using information from the Solubility Graph.

5. If NaNO3 were used instead of NaCl, how much would be

needed to make a saturated solution at the same temperature?

6. Stir another spoonful of NaCl into the solution. Observe the

results and describe what happens.


Review 61

7. Using information from the Solubility Graph, classify the following solutions as saturated or unsaturated.

Amount of Solute and

Temperature Type of Solution

25 g of KClO3 at 50oC

95 g of KNO3 at 60oC

20 g of NH3 at 55oC

8. How much NH3(g) can be dissolved in 100 g H2O at 25oC?

9. Examine the test tubes below and determine which one

contains the unsaturated solution, which one contains the saturated solution, and which one contains the supersaturated solution. Write your responses in the blanks.


62 Chemistry

Station 2 Water, Electrolytes and Nonelectrolytes

1. Locate the Water Molecule Kit. Remove the models of chlorine, sodium, and water.

Bring the sodium and chlorine models together and allow them to attach, making NaCl. Next, observe the Water Molecule Models and examine how they attach to the NaCl molecule.

Explain the process through which substances such as NaCl can dissolve in water.


Review 63

2. Which diagram below more accurately depicts NaCl when it is completely dissolved in water? Justify your answer.

Diagram B still contains an error having to do with the orientation of the water molecules. Thinking about the partial charges on the water molecule, redraw Diagram B so that it is completely correct.

3. Most living organisms are 60–90% water. Why is the polar nature of the water molecule important to the survival of living organisms?


64 Chemistry

4. Examine the Solutions Cards. Classify the solutions

represented on the cards as electrolytes and nonelectrolytes. Record your answers below.

Electrolyte Nonelectrolyte

5. Two solutions are prepared, one by dissolving 20 g of NaCl (salt) in 100 mL of distilled water at room temperature, and the other by dissolving 20 g of C12H22O11 (sugar) in 100 mL of water at room temperature. How are these solutions similar? How are they different?


Review 65

Station 3 Solubility Rules 1. Locate the Soluble and Insoluble Cards. Using the

STAAR Chemistry Reference Materials, sort the cards according to whether they are soluble or insoluble.

2. Locate test tube A at this station. Two colorless aqueous

solutions were mixed together in test tube A. According to the reaction below, what is the source of the yellow colored substance and how did it form?

2KI(aq) + Pb(NO3)2(aq) PbI2(s) + 2KNO3(aq)

Soluble Insoluble


66 Chemistry

3. Observe test tubes B through E. Study the solutions used to make them and the products that were formed. Using the STAAR Chemistry Reference Materials, label each reactant and product as aqueous (aq) or solid (s).

Test tube Solution and Products

FeCl3( ) +NaOH( ) Fe(OH)3( ) + 3NaCl( )

CoCl2( )+2NaOH( ) Co(OH)2( ) +2NaCl( )

NaCl( ) + LiOH( ) no reaction

NiCl2( ) + 2KOH( ) Ni(OH)2( ) + 2KCl( )


Review 67

4. What precipitates were formed by the solutions in Test Tubes

B through E?

6. Based on the data above, what generalization can you make about the solubility of the hydroxide ion (OH-) with metals?

Test Tube Precipitate

B

C

D

E


68 Chemistry

6. Examine test tubes F through H. Study the table below

that shows the solutions used to make them and the products that were formed:

Test tube Solution and Products

AgNO3(aq) + KI(aq) AgI(s) + KNO3(aq)

AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)

AgNO3(aq) + LiBr(aq) AgBr(s) + LiNO3(aq)

7. What precipitates were formed by the solutions in Test

Tubes F through H?

Test Tube Precipitate

F

G

H


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8. Based on the chemical equations shown in the table,

what generalizations can you make about the solubility of halide compounds? “Halide” means “containing a halogen.”

9. Locate the STAAR Chemistry Reference Materials at this

station. Review the solubility of common ionic compounds in water section. Which other metals form insoluble compounds with halogens?

10. Continue using the STAAR Chemistry Reference Materials to complete the following equations. If no precipitate forms, write “no reaction.”

Reactants Products

CuSO4(aq) + BaCl2(aq)

KBr(aq) + CuSO4(aq)

Na2CO3(aq) + Ca(NO3)2(aq)

Na2CrO4(aq) + CuSO4(aq)


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11. Look again at the Soluble and Insoluble Cards. Choose

any two cards showing soluble substances that, when mixed together, will form a precipitate. Write a balanced equation below.

12. How can you predict whether a precipitate will form

when aqueous solutions of two ionic compounds are mixed?

Additional Notes STAAR-EOC

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I need to remember…


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Credits

Materials & Activities are provided by: Charles A. Dana Center at the University of

Texas at Austin, STAAR Chemistry Assessments

staar-eoc · station 1 . physical and chemical properties. 1. locate the physical and chemical...

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