KEY - slhspapchem / SLHS Pre-AP Chemistry of Matter & Nuclear Chemistry STAAR-EOC Review 3 Activity A Properties of Matter & Nuclear Chemistry Station 1 Physical and Chemical Properties

SLHS Chemistry

Activity Booklet

STAAR-EOC

Chemistry

Review

KEY

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.

Physical Property

Chemical Property Extensive Intensive

1 Conductivity

2 Radioactivity

3 Volume

4 Number of Atoms

5 Boiling Point

6 State of Matter

7 Flammability

8 Color

9 Reactivity

10 Mass

11 Length

12 Freezing Point

13 Viscosity

14 pH

15 Calories per Gram


4 Chemistry

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

Physical properties are properties that can be measured or observed, in which no new substances are formed. Chemical properties are those that enable a substance to undergo a change in chemical composition.

4. How did you differentiate between an intensive and

extensive property?

Extensive properties depend on the amount of matter present in the substance. Intensive properties do not depend on the amount of matter present.


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.

Change Physical Chemical Chemical and Card Change Change Physical Change

A x

B x

C x

D x

E x

F x

G x

H x

I x

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

Card G illustrates a chemical change because the sodium bicarbonate and acetic acid react to form new substances: sodium acetate, water, and carbon dioxide. The production of a gas (CO2) is also evidence of a chemical change.


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

Card D is both physical and chemical change, because when carbon dioxide escapes from the solution, that is a separation of a mixture (physical), but at the same time, removal of CO2 prevents carbonic acid formation in the soda so it becomes less acidic, which is a chemical change.

Card E illustrates both a physical and a chemical change. When a dog chews a bone, the bone breaks into smaller pieces, which is a physical change. As the dog swallows and digests pieces of the bone, the digestion is a chemical change

Card F: Adding dye to hair is a physical change. Bleaching changes the hair’s original color, making it a chemical change. If curls are created with a curling iron or hair rollers, a physical change has occurred. If chemicals are applied to create curls, a chemical change has occurred.

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

chemical properties of matter?

Physical properties can be observed or measured and they do not cause a change in the composition of the substance. If a property causes a new substance to form, or if the property shows the ability of the substance to undergo a change that alters its composition, it is a chemical property.


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.


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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.

The cylinder of gas would show the greatest change in volume because the free space between the gas particles would allow the cylinder to be significantly compressed. The solid would show the least change in volume, as its particles are tightly packed together and would not easily compress.

3. What properties can scientists use to describe and classify

matter?

Matter can be described and classified according to its unique physical and chemical properties, its ability or inability to be separated from other substances, and its interactions with other substances.


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.

Characteristics of Radiation Cards Key

Which type of radiation presents the greatest risk? Why?

Gamma radiation presents the greatest risk because it has the greatest penetrating power.


10 Chemistry

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

Equations Strips. Insert the strips into the slots of the card. Move the strips until each equation is balanced.

Nuclear Equations Strips Answer Key

How do you know when a nuclear equation is balanced?

A nuclear equation is balanced when the sum of the mass numbers and the sum of the atomic numbers is equal on both sides of the equation.

3. Write balanced nuclear equations for the following:

Carbon-14 undergoes beta decay.

Answer:

Radium-236 undergoes alpha decay.

Answer:


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 in the table below.

Fission Both Fusion

Splitting of larger into smaller nuclei

Reactions release large amounts of energy

Smaller nuclei combine to form larger nucleus

Does not occur in nature often

Occurs in stars, like our sun

Produces highly radioactive particles

Produces few radioactive particles

Used to generate electricity in nuclear power plants

Takes large amounts of energy to start a reaction

Takes small amounts of energy to start a reaction


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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?

In the sun, lighter nuclei are fused into a larger nucleus through the process of nuclear fusion.

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

In nuclear power plants, large nuclei are split into smaller nuclei through the process of nuclear fission.

Pure Substance/Mixtures & Thermochemical Equations STAAR-EOC

Review 13

Activity B Pure Substances/Mixtures & Thermochemical Equations Station 1 Classifying Pure Substance and Mixtures

1. Which bottles contain pure substances? Which contain

mixtures? Record your responses in the table below by placing an “x” in the appropriate column.

Bottle Name of Substance Pure Mixture Substance

A

Copper X

B

Aluminum X

C

Brass X

D

Copper sulfate solution

X

E

Magnesium sulfate X

F

Iron and Sulfur X


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

Copper and aluminum have atomic numbers listed, which makes them elements. All elements are pure substances.

Magnesium sulfate is a pure substance with a fixed chemical composition, MgSO4, and the elements of the compound cannot be separated by physical means.

Brass is a mixture made of the elements copper and zinc, and has a variable melting point.

Copper sulfate is mixed with water to form a solution. The copper sulfate crystals can be separated from the water.

The iron and sulfur are elements mixed together that can be separated using a magnet.

3. How are mixtures that are solutions identified?

Solutions are formed by dissolving one substance (solute) in another (solvent).

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

be done to separate the parts of the solutions?

Bottles C and D contain mixtures that are solutions.

The brass in Bottle C is a solution of copper and zinc. Using different properties of the two elements, such as melting point, could allow them to be separated.

The copper sulfate solution in Bottle D can be separated by evaporating the water, leaving copper sulfate crystals behind.


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

Pure substance

Made of one type of atom (element) B

Pure substance

Made of one type of atom (element) chemically combined in a specific ratio

C

Mixture

Three different substances mixed together D

Pure substance

Compound made of two elements combined in a specific ratio


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Heterogeneous and Homogeneous Mixtures

Examine the samples of mixtures below.

6. Which samples represent heterogeneous mixtures?

2, 3, and 4 7. Which samples should be classified as homogeneous

mixtures? What is the justification for this classification?

Samples 1, 5, and 6 are homogeneous, since the very small solute particles are evenly distributed throughout the mixtures.

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?

The iron and sulfur cannot be physically separated (only chemically). The iron sulfide is a pure substance (a compound).


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.

Energy Kinetic Potential

Thermal Energy

Chemical Energy

Motion of atoms or molecules

Energy stored in bonds

Measured as temperature

Calories in food

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

Answers may vary. Kinetic energy is energy of motion; potential energy is energy of position/stored energy.


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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)

Answer: ∆H = –393.5 kJ

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

burned?

∆H = 2 x (–393.5 kJ) = –787 kJ

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

The reaction is exothermic; releases thermal energy (heat).


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)

Answer: ∆H = –822.16 kJ

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?

∆H = +2808 kJ 7. Is this an exothermic or endothermic process?

The formation of glucose is an endothermic process. The sign of ∆H is positive.


20 Chemistry

8. The temperature of a 0.25 Kg block of an unknown metal

increases by 22.22°C after 5,000 Jof 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

Using Q = mCp∆T

Cp = Q/(m∆T) = 5,000J/[(0.25Kg)(22.22°C)] = 900 J/Kg•°C

The metal is aluminum.

9. How do energy changes occur during chemical reactions?

When a chemical change occurs, there is also a thermal energy change, which can be detected as a temperature change. If the products have less energy than the reactants, energy is lost. If the products have more energy than the reactants, energy is absorbed.


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.

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?

Because the temperature decreased, energy (heat) from the surroundings was absorbed into the system.

3. Was the change endothermic or exothermic?

Endothermic


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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?

Because the temperature increased, energy (heat) from the system was released into the surroundings.

5. Was the change endothermic or exothermic?

Exothermic


Review 23

6. Examine the graph.

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

150 KJ

Is this an exothermic or endothermic chemical reaction? Why?

The reaction is endothermic. The products have more energy than the reactants, so energy has been absorbed.

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.

Answers:

2. What “clues” did you use to help arrange the periodic table?

Key of the table, size of element, electrons, ionization energy, electronegativity, color, etc.

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

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

Students’ answers will vary.


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?

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

Down in a group: As the elements go down in a group, electronegativity decreases because the distance between the nucleus and valence electrons in the elements’ atoms increases. More energy levels filled with electrons shield the nucleus.

Across in a period:

As elements move across a period, electronegativity increases as the nuclear charge increases.


26 Chemistry

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?

Ionization energy of Group 1 is low compared to that of Group 17.

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

The elements in Group 17 have larger effective nuclear charges than those in Group 1.

5. Positive ions (cations) tend to be smaller than their corresponding neutral atoms. What is a possible explanation for this?

Cations have fewer electrons than their corresponding neutral atoms, which increases the effective nuclear charge that draws the remaining electrons closer to the nucleus.

6. Negative ions (anions) tend to be larger than their corresponding neutral atoms. What could be an explanation for this?

Anions have more electrons, and repulsion forces between electrons push them further apart. When electrons outnumber protons, the nucleus cannot pull the electron cloud as tightly around itself.

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? Ionic radius decreases for positive and negative ions, although negative ions are larger than positive ions.

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

Ionic radius increases as energy levels are added.


Review 27

Station 3 Chemical Families Locate the Periodic Table Labels. Use the information in the

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.

(Answers shown below)

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

of elements in the Periodic Table of the Elements?

Classifying elements into groups with similar properties gives us information about their chemical reactivity.


28 Chemistry

5. Compare the reactivity of alkali metals and noble gases. What property of these groups of elements accounts for the differences in their reactivity?

The alkali metals are very reactive, losing their single valence electron and forming compounds easily. Noble gases are very stable as they have a complete outer energy level.

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

The alkali metals tend to lose their single valence electron, forming positive ions. Losing one electron allows the alkali metals to achieve a stable noble gas configuration. Gaining one electron allows the halogens to achieve a stable noble gas configuration.

The halogens tend to gain one electron, forming negative ions. The reactivity of the alkali metals increases as you move down the group because the outermost electron is more easily lost as it gets further away from the nucleus. The reactivity of the halogens decreases as you move down the group because electrons are less easily attracted by the nucleus due to increased shielding.

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

predict periodic trends in chemical families and periods?

Elements in the same group on the table have similar chemical properties. Elements in the same period on the table show trends in atomic and ionic radius, ionization energy, and electronegativity.


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 Chlorine dioxide HCl(aq) Hydrochloric acid NH4NO3 Ammonium nitrate Pb(NO3)2 Lead (II) nitrate Fe2O3 Iron (III) Oxide H2CO3 Carbonic acid C2Br6 Dicarbon hexabromide HNO3(aq) Nitric acid HCl(g) Hydrogen chloride KOH(aq) Potassium hydroxide


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

N2O3

Carbon tetrachloride

CCl4

Disulfur trifluoride

S2F3

2. Use the Covalent Compound Cards to determine the

chemical formulas for the following compounds:

a. A compound with carbon and fluorine:

CF4

b. A compound with silicon and hydrogen:

SiH4

c. A compound with sulfur and bromine:

SBr2


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.

Ionic Compound Name Ionic Compound Formula Copper (I) sulfate

Cu2SO4

Calcium carbonate

CaCO3

Copper (II) nitrite

Cu(NO2)2

Iron (III) oxide

Fe2O3

Ammonium phosphate

(NH4)3PO4

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

May include, but not limited to: PbI2, MgCl2, Be(OH)2, Cu(NO2)2, Fe(OH)3, NH4I, NH4OH

3. What features of compounds and chemical formulas do scientists use to determine the names for the compounds and formulas?

Scientists use the compound or formula’s type of bonding (ionic or covalent), its location on the periodic table, and whether it is a main group or transition metal to name compounds and chemical 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. Find the block that represents carbon. Using the Energy

Levels Table, write the complete electron configuration for carbon:

C - 1s2 2s2 2p2

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. Atomic Number Element Electron Configuration

1 Hydrogen (H) 1s

2 Helium 1s2

3 Lithium 1s2 2s1

9 Fluorine 1s2 2s2 2p5

10 Neon 1s2 2s2 2p6

16 Sulfur 1s2 2s2 2p6 3s2 3p4

18 Argon 1s2 2s2 2p6 3s2 3p6

1 Hydrogen (H+) This ion has no electrons

11 Sodium ion 1s2 2s2 2p6

12 Magnesium ion 1s2 2s2 2p6

17 Chloride ion 1s2 2s2 2p6 3s2 3p6


Review 33

3. What is the noble gas configuration for

Calcium?

[Ar]4s2

Aluminum?

[Ne]3s22p1

4. Draw the Lewis Dot Structures for Elements 1–18 in the table below.


34 Chemistry

5. What correlation is there between the group numbers and the number of valence electrons you have drawn?

The group numbers, 1A–8A, are the same as the number of valence electrons.

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

electron configurations of atoms?

The periodic table reflects a repeating pattern of electron configurations in s, p, d, and f orbitals. Properties of atoms are largely determined by their electron configurations.


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.

The answers to Question 1 appear with the answers to Question 2.

2. Draw the electron dot formulas in the space provided on your

Student Pages. Notice the example of H2O has been provided. You may use a dashed line to represent a pair of bonding electrons between two atoms.


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?

It is important to know how many electrons are available for bonding so that each atom can gain a complete outer energy level.

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

BF3?

Be and B are exceptions to the octet rule. Be is stable with only four valence electrons, and B is stable with only six valence electrons.

5. Draw the possible Lewis Structures for the following ionic compounds:

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

The “x” represents electrons that were lost by the positive ions and gained by the negative ions.


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 Tetrahedral 2 B

BH3 Trigonal Planar 0 E

HCN Linear 1 A NH3 Tetrahedral 1 F CH4 Tetrahedral 0 D

SO2 Trigonal Planar 1 C

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

CH4 and H2O both have four areas of electron density. H2O has two lone pairs and two bonding pairs of electrons, while CH4 has four bonding pairs of electrons.


38 Chemistry

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

CH4 and NH3 both have four areas of electron density around the central atom.

How are they different?

Areas of electron density around CH4 are equally separated. The pairs of bonding electrons are closer together for NH3.

Why are they different?

The lone, nonbonding pair of electrons in NH3 has greater repulsive force than the bonding pairs.

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

different from the behavior of electrons in a covalent bond?

Ionic bonds result from the electrostatic attraction between positive and negative ions because electrons have transferred from metals to create positive ions and have been added to nonmetals to create negative ions. Covalent bonds occur when atoms share electrons to

complete their outer energy level. In addition, the

electrostatic attraction is due to the atoms having

gained or lost electrons (transferred).

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 L 1 week 7 days

1 kilometer 1,000 meters 1 microgram 1 x 10-6 grams

1 mole 6.02 x 1023 atoms

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

6.0 moles CO2 6.02 x 1023 molecules / 1 mole CO2

3.6 x 1024 molecules CO2

4.0 moles Fe 6.02 x 1023atoms / 1 mole Fe 2.4 x 1024 atoms Fe

2.5 moles NaCl 6.02 x 1023 ions Cl - / 1 mole NaCl

1.5 x 1024 ions Cl -


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. 2.5 mol CaCl2 x 6.02 x 1023 formula units CaCl2 x 2 Cl- ions 1 mol CaCl2 1 formula unit CaCl2 = 3.0 x 1024 Cl- ions How does your answer compare to the number of chloride ions in 2.5 moles of NaCl (see table on previous page)?

There are twice as many chloride ions in 2.5 moles of CaCl2 than in NaCl. Each formula unit of CaCl2 contains 2 chloride ions, but each formula unit of NaCl contains only 1 chloride ion.

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

63.564 g/mole

Ar

39.948 g/mole

NaOH

39.997 g/mole

NaHCO3

84.006 g/mole


Review 41

Pb(NO3)2

331.208 g/mole

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.

Answers for Questions 2–4 will vary depending on the masses of the samples provided.

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?

3.34 x 1024 molecules of H2O

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

6.68 x 1024 atoms of H (2 atoms H per molecule H2O)

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

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

The mole is a way to count atoms, ions, or molecules from weight. One mole of any substance contains 6.02 x 1023 particles.


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.

The balloon will not inflate because the bottle is already filled with air.


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?

As the balloon inflated, air was pushed out of the bottle through the small hole in the side. Keeping the hole covered once the balloon is inflated keeps air from rushing back into the bottle, pushing on the balloon, and forcing air out through the top.

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

The balloon deflates as air rushes in the bottle through the hole.


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.

The balloon in the warm water bath is inflated. The balloon in the ice bath is not inflated.

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.

The balloon that was inflated is now deflated, and the


46 Chemistry

balloon that was deflated is now inflated.

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

As the temperature of the gas (i.e., the air inside the balloon) increases as a result of the warm water bath, its volume increases. As the temperature of the gas decreases due to the cold water bath, its volume decreases.

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?

V1/T1 = V2/T2

(2.5 L)/(303 K) = x/(283 K)

x = 2.3 L


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?

The balloon decreases in size as the air added to the flask increases pressure on the outside of the balloon.

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

changes in the balloon in Steps 10-12 illustrate Boyle’s Law.

As the pressure inside the flask and on the balloon is increased, the volume of the balloon decreases. Letting the added air out of the flask reduces the pressure on


48 Chemistry

the outside of the balloon and the volume increases again.

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?

P1V1 = P2V2

(1.0 atm)(25 L) = (0.85 atm)x x = 29 L

Graph B shows a decrease in pressure as temperature increases, but pressure actually increases as temperature increases, per Charles’ Law.


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

P1V1 = P2V2

B

V1/T1 = V2/T2

C

PV = nRT

D

PT = P1 + P2 + P3…

E

V1/n1 = V2/n2

F

P1V1 = P2V2

G

P1V1/n1T1 = P2V2/n2T2

H

PV = nRT

I

V1/T1 = V2/T2

J

P1V1 = P2V2

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

selected will solve the problem.

Answers will vary.


50 Chemistry

Gas Stoichiometry

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

O2 and H2.

O2 + 2H2 2H2O

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

(100. g H2O) × (1 mole O2 / 31.998 g) = 3.13 moles O2

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

O2 + 2H2 2H2O

(3.13 moles O2) × (2 moles H2O / 1 mole O2) × (22.4 L / 1 mole) = 140. L O2

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

related?

As gas pressure increases at a constant temperature, the volume of the gas decreases. As the temperature of a gas increases at a constant pressure, the volume of the gas increases.

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?

The mass of the reactants and the mass of the products are equal even though the atoms are arranged differently.

2Cu + O2 2CuO 2HBr + Mg H2 + MgBr2

CH4 + 2O2 CO2 + 2H2O

2NaI + Pb(NO3)2 2NaNO3 + PbI2

AlCl3 + 3NH4OH 3NH4Cl + Al(OH)3

2NaOH + H2SO4 Na2SO4 + 2H2O


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.

The Reactants bag has more mass than the Products bag.

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.

Mg + 2 HCl MgCl2 + H2

There is too much HCl in the Reactants bag. 3. From your observations of the contents of the bags, which

reactant is limiting, and why?

The magnesium is limiting because there is more hydrochloric acid than needed to react all the magnesium.


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?

The mass of Setup #1 would be greater because the balloon would trap the hydrogen gas produced; in Setup #2, the hydrogen gas would escape.

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?

According to the balanced equation, 2 moles of HCl are needed for every 1 mole of Mg. There are 0.10 moles of HCl and 0.010 moles of Mg available.

(0.25 g Mg) x (1 mol Mg/24.305 g) = 0.010 mol Mg

Since you have 10 times as much HCl as Mg, the Mg is limiting.


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?

0.25 g Mg x (1 mol Mg/24.305 g) x (1 mol Mg/1 mol H2) x (2.016 g/1 mol H2) = 0.021 g H2

(0.018/0.021) x 100 = 86%

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

produced from, a chemical reaction be calculated?

According to the law of conservation of mass, a chemical equation is balanced when there are equal numbers of each type of atom involved on both sides of the equation.

The coefficients of a balanced equation give information about the proportional amounts of each substance involved in the reaction; they allow you to determine the amounts of reactants used and the amounts of products produced.


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.

Acid-Base Precipitation Oxidation/ Reduction

C B A

G F D

H J E

L K I

P M N

Q R O

T V S

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

are the variations?

Oxidation-reduction reactions have the most different variations. These reactions occur when two substances combine to form a single product, when compounds decompose, when more active elements replace less active elements, and during combustion.


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

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

Type of Reaction

Oxidation-reduction

Supporting Evidence

Hydrogen gains electrons and magnesium loses electrons.

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

Balanced Equation

2NaOH(aq) + Cu(NO3)2(aq) Cu(OH)2(s) +2 NaNO3(aq)

Type of Reaction

Precipitation

Supporting Evidence

Copper II hydroxide is an insoluble product.


58 Chemistry

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

Balanced Equation

HC2H3O2 (aq) + NaOH (aq) NaC2H3O2 (aq) + H2O (l)

Type of Reaction

Acid–base

Supporting Evidence

An acid and a base are combined to produce a salt and water.

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?

Electron transfer is occurring, so this is an oxidation-reduction reaction. 2. What characteristics can be used to identify different types

of aqueous chemical reactions?

Precipitation reactions occur when an insoluble product forms; the reaction of an acid and a base can result in formation of water and a salt or gas production; oxidation- reduction reactions always involve the transfer of electrons

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 20 – 25oC

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 answers will vary

Mass of NaCl used answers will vary


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.

Depending on the temperature of the water, students should be able to dissolve between 35 and 38 grams of NaCl in the water. This would be a reasonable amount according to the Solubility Graph. Students should recognize that the line on the graph represents a saturated solution, and that anything below that line indicates an unsaturated solution.

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

needed to make a saturated solution at the same temperature?

Between 90 and 95 grams. 6. Stir another spoonful of NaCl into the solution. Observe the

results and describe what happens.

Additional NaCl settles to the bottom of the solution and cannot be dissolved.


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

Saturated

95 g of KNO3 at 60oC

Unsaturated

20 g of NH3 at 55oC

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

45 g

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.

Saturated Unsaturated Supersaturated


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.

The positively charged side of the polar water molecules are attracted to the negatively charged chloride ions and the negatively charged side of the water molecules are attracted to the positively charged sodium ions. The ions are separated from each other in solution.


Review 63

2. Which diagram below illustrates an accurate

representation of NaCl when it is completely dissolved in water? Justify your answer.

Incorrect Correct—mostly*

Students should recognize that in soluble ionic compounds, ions are separated and completely surrounded by water molecules. 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. *Swap the sodium and chloride ions’ positions so that the partially negative oxygen attracts to the positive sodium ion, and the partially positive hydrogens attract to the negative chloride ion.

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

Many biochemical reactions in cells take place in aqueous solutions.

The hydrogen bonding in water molecules gives water the properties of cohesion and surface tension, which allows capillary action to transport water and other substances in plants.


64 Chemistry

4. Examine the Solutions Cards. Classify the solutions

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

Electrolyte Nonelectrolyte

A

B

D

C

E

F

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?

Both solutions are clear and colorless. The salt solution has dissolved ions and will conduct electricity. The sugar solution has no dissolved ions and will not conduct electricity.


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)

These two substances produced an insoluble precipitate. The KI and Pb(NO3)2 ionized in solution allowing PbI2 to form. PbI2 is insoluble.

Soluble Insoluble

Salts of K+, Na+, NH4+ Most metal hydroxides

Salts of NO3- Most sulfides

NaOH MgCO3

MgSO4 PbCl2

Ba(NO3)2 Fe2S3


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

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

NaCl( ) +LiOH( ) no reaction NaCl(aq)+ LiOH(aq) no reaction

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


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?

Based on the reactions above, hydroxide ions are insoluble with metals except the alkali metals.

Test Tube Precipitate

B Fe(OH)3(s)

C Co(OH)2(s)

D None

E Ni(OH)2(s)


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 AgI(s)

G AgCl(s)

H AgBr(s)


Review 69

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.”

Based on the information above, the halogens are soluble with the alkali metals, but not with silver.

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?

Compounds of lead and mercury

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)

BaSO4(s) + CuCl2(aq)

KBr(aq) + CuSO4(aq)

No reaction

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

CaCO3(s)+ NaNO3(aq)

Na2CrO4(aq) + CuSO4(aq)

CuCrO4(s)+ Na2SO4(aq)


70 Chemistry

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.

Answers will vary. 12. How can you predict whether a precipitate will form

when aqueous solutions of two ionic compounds are mixed?

Determine the ions present in the reactants, and consider possible combinations of ions. Use solubility rules to determine whether any of the combinations is insoluble.

Additional Notes STAAR-EOC

Review 71

I need to remember…

Physical and Chemical Changes: • Physical properties are characteristics of

matter that can be observed or measured that do not affect the chemical composition of the substance.

• Physical changes do not produce new substances, but they may cause a change in energy or state of matter.

• Chemical properties are characteristics of matter that enable it to undergo a change in composition.

• Chemical changes cause new substances to form.

• Intensive physical properties do not depend on the amount of matter present.

• Extensive physical properties depend on the amount of matter present.

• Solids have a definite shape and volume with particles fixed in place. Solids have little free space between particles and are not easily compressible.

• Liquids have a definite volume, but take on the shape of their container. Particles move past each other easily, allowing liquids to flow. There is little free space between the particles of a liquid, and they are not easily compressible.

• Gases assume both the shape and the volume of their container. Gas particles are far apart from each other, moving freely at high speeds. Gases are compressible.


72 Chemistry


Pure Substances and Mixtures: • A pure substance cannot be separated into

two or more substances by physical or mechanical means.

• A pure substance has a uniform composition, and constant properties, throughout the entire sample.

• A pure substance has a constant chemical composition.

• Mixtures can be separated into two or more substances by physical or mechanical means.

• Mixtures display properties of the substances making up the mixture.

• The composition of a mixture can vary. • Homogeneous mixtures have components that

are equally distributed throughout. • Heterogeneous mixtures have components that

are not uniformly distributed. Chemical Families and Periodic Trends:

• Alkali metals are the most reactive metal family, with one valence electron that is easily lost, forming ions with a +1 charge.

• Alkaline earth metals are reactive and form ions with a +2 charge.

• Transition metals are typical metals that can have multiple oxidation states.

• Halogens tend to gain electrons, forming –1 ions.

• Noble gases have extremely low reactivity because they have full outer energy levels.

• A greater effective nuclear charge increases the attraction of the nucleus and pulls the electron cloud closer to the nucleus, resulting in a smaller atomic radius.


Review 73


Chemical Families and Periodic Trends Cont.: • An increased number of energy levels

increases the distance over which the nucleus must pull and reduces the attraction for electrons.

• Full energy levels provide shielding between the nucleus and valence electrons.

Atomic Structure:

• Electrons fill the lowest available energy states available before filling higher states.

• The most important electrons for each element are the outermost electrons.

• Electrons are negatively charged particles that occupy energy levels surrounding the positively charged nucleus of an atom.

• Electrons repel each other. • Lewis Dot Structures show valence electrons. • The Periodic Table is organized by increasing

atomic number. • The atomic number tells us how many protons

an atom has. • The atomic number for neutral atoms is also the

number of its electrons. • Electron configurations show the arrangement

of the electrons around the nucleus. Nuclear Equations and Radioactive Decay:

• In balanced nuclear equations, the sum of the mass numbers is balanced on both sides of the equation.

• In balanced nuclear equations, the sum of the atomic numbers is balanced on both sides of the equation.


74 Chemistry


Naming Compounds and Chemical Formulas: • Use prefixes to name covalent compounds. • For ionic compounds, name the metal or

cation first, then name the nonmetal or anion, changing the end of the names of single element anions to –ide.

• The name of an acid depends on the type of anion it has.

• Chemical formulas are based on neutral compounds.

Electron Dot Formulas and Molecular Geometry:

• Ionic bonds are formed by the attraction between oppositely charged ions.

• Covalent bonds are formed when atoms share electrons.

• Molecular geometry is predicted by assuming that atoms in molecules will be positioned to minimize electron repulsion.

• Nonbonding pairs of electrons take up more space and have greater repulsion than bonding pairs of electrons.

Moles:

• The mass of one mole of any element or compound can be determined from the Periodic Table.

• A mole of any substance contains 6.02 x 1023 atoms, ions, or molecules.

• Moles are used to convert mass to atoms, ions, or molecules.


Review 75


Balancing Equations: • Balanced chemical equations have equal

numbers of participating atoms on either side of the equation.

• Mass is not created or destroyed in reactions. • Balanced chemical equations show

proportional amounts of reactants and products and allow for the determination of the amounts of reactants used and the amounts of products produced.

Gas Laws:

• Pressure and volume of a fixed mass of gas are inversely related.

• The pressure of a fixed volume of gas is directly related to its Kelvin temperature.

• The volume of a gas at constant pressure is directly related to its Kelvin temperature.

• The volume of one mole of any gas at standard temperature and pressure is 22.4 L.

Thermochemical Equations:

• Endothermic reactions absorb energy (heat). • Exothermic reactions release energy (heat). • Thermal energy is a measure of the motion of a

substance’s atoms or molecules. • Chemical energy is stored in bonds. • Thermochemical equations can be used

to classify reactions as endothermic or exothermic.


76 Chemistry


Types of Solutions: • Unsaturated solutions can hold more

solute at a given temperature without becoming supersaturated.

• Saturated solutions cannot hold more solute at a given temperature without becoming supersaturated.

• Supersaturated solutions hold more solute than is normally possible at a given temperature.

• Supersaturated solutions are unstable, and the solute will precipitate if disturbed or if more solute is added.

• Electrolytes are solutions with dissolved ions that can conduct electricity.

• Nonelectrolytes have no dissolved ions and cannot conduct electricity.

• The polar nature of water molecules allows water to be an effective solvent.

• Many important biochemical reactions occur in solution.

Solubility Rules:

• When two aqueous solutions are combined, a chemical reaction that results in the formation of an insoluble product called a precipitate may occur.

• A substance that is insoluble will not dissolve in a solvent.


Review 77

Types of Reactions: • Acid–base reactions occur when an acid is

mixed with a base and the products are salt and water.

• Precipitation reactions occur when an insoluble product is formed after two aqueous salt solutions are mixed together.

• Oxidation-reduction reactions occur when electrons are transferred during the reaction. One substance loses electrons and is oxidized, while one substance gains electrons and is reduced.

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

University of Texas at Austin, STAAR Chemistry Assessments


KEY - slhspapchem / SLHS Pre-AP Chemistry of Matter & Nuclear Chemistry STAAR-EOC Review 3 Activity A Properties of Matter & Nuclear Chemistry Station 1 Physical and Chemical Properties

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