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Chemistry SOL Review Packet

CH.1 The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated produce observations and verifiable data.

a) designated laboratory techniques; b) safe use of chemicals and equipment; c) proper response to emergency situations; d) manipulation of multiple variables, using repeated trials; e) accurate recording, organization, and analysis of data through repeated trials; f) mathematical and procedural error analysis; g) mathematical manipulations (SI units, scientific notation, linear equations, graphing, ratio and proportion, significant digits, dimensional analysis); h) use of appropriate technology including computers, graphing calculators, and probeware, for gathering data and communicating results; and i) construction and defense of a scientific viewpoint (the nature of science).

In order to meet this standard, it is expected that students will

make connections between components of the nature of science and their investigations and the greater body of scientific knowledge and research.

demonstrate safe laboratory practices, procedures, and techniques.

demonstrate the following basic lab techniques: filtering, using chromatography, and lighting a gas burner.

understand Material Safety Data Sheet (MSDS) warnings, including handling chemicals, lethal dose (LD), hazards, disposal, and chemical spill cleanup.

identify the following basic lab equipment: beaker, Erlenmeyer flask, graduated cylinder, test tube, test tube rack, test tube holder, ring stand, wire gauze, clay triangle, crucible with lid, evaporating dish, watch glass, wash bottle, and dropping pipette.

make the following measurements, using the specified equipment:volume: graduated cylinder, volumetric flask, buretmass: triple beam and electronic balancestemperature: thermometer and/or temperature probepressure: barometer and/or pressure probe.

identify, locate, and know how to use laboratory safety equipment, including aprons, goggles, gloves, fire extinguishers, fire blanket, safety shower, eye wash, broken glass container, and fume hood.

design and perform controlled experiments to test predictions, including the

following key components: hypotheses, independent and dependent variables, constants, controls, and repeated trials.

predict outcome(s) when a variable is changed.

read measurements and record data, reporting the significant digits of the measuring equipment.

demonstrate precision (reproducibility) in measurement.

recognize accuracy in terms of closeness to the true value of a measurement.

determine the mean of a set of measurements.

use data collected to calculate percent error and discover and eliminate procedural errors.

use common SI prefixes and their values (milli-, centi-, kilo-) in measurements and calculations.

demonstrate the use of scientific notation, using the correct number of significant digits with powers of ten notation for the decimal place.

graph data utilizing the following:independent variable (horizontal axis), dependent variable (vertical axis)scale and units of a graph, regression line (best fit curve).

calculate mole ratios, percent composition, conversions, and average atomic mass.

perform calculations according to significant digits rules.

convert measurements using dimensional analysis.

use graphing calculators to solve chemistry problems.

read a measurement from a graduated scale, stating measured digits plus the estimated digit.

use appropriate technology for data collection and analysis, including probeware interfaced to a graphing calculator and/or computer and computer simulations.

summarize knowledge gained through gathering and appropriate processing of data in a report that documents background, objective(s), data collection, data analysis and conclusions.

explain the emergence of modern theories based on historical development. For example, students should be able to explain the origin of the atomic theory beginning with the Greek atomists and continuing through the most modern quantum models.


CH 1 Problems:

1. Review safety contract. Write precautions 27, 29, 33,39,40, and 42

2. Define: filtering, decanting, chromatography, and distillation. Indicate which can be used to separate heterogeneous mixtures and which can be used to separate homogeneous mixtures.

3. Identify laboratory safety equipment in the classroom and their locations. Be sure you understand how to use each one.

4. John performed an experiment in which he was concerned with the effects of changing temperature on the volume of a gas. In his experiment, he changed the temperature of a sealed balloon and found the volume. The following table shows the data he collected:

Trial # Temperature (°C)

Temperature (K)

Volume (mL)

Vol/Temp (mL/K)

1 -73.0 4032 -1.0 5463 0.0 5464 1.0 5455 10.0 6016 100.0 748

a) Complete the above table.b) If you wanted to graph this data which variable would go on

the x-axis? the y-axis?c) In which trial is it likely that experimental error has been

made? Explain.

5. Sue found several samples of the same substance in a riverbed. She wanted to find the density of the substance to help identify it. She gathered the following data:

Mass (g) Volume (cm3)

Density (g/cm3)

Sample A 5.86 0.31Sample B 6.24 0.36Sample C 3.57 0.18Sample D 1.19 0.06Sample E 12.27 0.64Sample F 8.49 0.44

a) Find the mean (average) value for the density of the unknown substance.

b) According to the following density values, what is the substance Sue has found?

a. Silver – 10.49 g/mL d. Gold - 19.32 g/mLb. Titanium – 4.5 g/mL e. Lead – 11.3 g/mL c. Copper – 8.92 g/mL

c) Which sample is it likely that a procedural error has been made? Explain.

d) Discuss the precision and accuracy of Sue’s data.e) Convert the mass of Sample E to kilograms. Express the

answer in scientific notation.f) Convert the mass of Sample E to milligrams. Express the

answer in scientific notation.g) Calculate the % error of Sue’s data.


CH 2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of

a) average atomic mass, mass number, and atomic number;b) isotopes, half lives, and radioactive decay;c) mass and charge characteristics of subatomic particles;d) families or groups;e) periods;f) trends including atomic radii, electronegativity, shielding

effect, and ionization energy;g) electron configurations, valence electrons, and oxidation

numbers; h) chemical and physical properties; andi) historical and quantum models.


determine the atomic number, atomic mass, the number of protons, and the number of electrons of any atom of a particular element using a periodic table.

determine the number of neutrons in an isotope given its mass number.

perform calculations to determine the “weighted” average atomic mass.

perform calculations involving the half-life of a radioactive substance.

differentiate between alpha, beta, and gamma radiation with respect to penetrating power, shielding, and composition.

differentiate between the major atom components (proton, neutron and electron) in terms of location, size, and charge.

distinguish between a group and a period.

identify key groups, periods, and regions of elements on the periodic table.

identify and explain trends in the periodic table as they relate to ionization energy, electronegativity, shielding effect, and relative sizes.

compare an element’s reactivity to the reactivity of other elements in the table.

relate the position of an element on the periodic table to its electron configuration.

determine the number of valence electrons and possible oxidation numbers from an element’s electron configuration.

write the electron configuration for the first 20 elements of the periodic table.

distinguish between physical and chemical properties of metals and nonmetals.

differentiate between pure substances and mixtures and between homogeneous and heterogeneous mixtures.

identify key contributions of principal scientists including:atomos, initial idea of atom – Democritusfirst atomic theory of matter, solid sphere model – John Daltondiscovery of the electron using the cathode ray tube experiment,

plum pudding model – J. J. Thomsondiscovery of the nucleus using the gold foil experiment, nuclear

model – Ernest Rutherforddiscovery of charge of electron using the oil drop experiment –

Robert Millikanenergy levels, planetary model – Niels Bohrperiodic table arranged by atomic mass – Dmitri Mendeleevperiodic table arranged by atomic number – Henry Moseleyquantum nature of energy – Max Planckuncertainty principle, quantum mechanical model – Werner

Heisenbergwave theory, quantum mechanical model – Louis de Broglie.

differentiate between the historical and quantum models of the atom.

CH 2 Problems:


1. Electrons have little mass and a __________ charge. They are located in _____________, __________ the nucleus. Protons have a ___________ charge. Neutrons have a ____________ charge. Protons and neutrons are located in the center or ________________ of the atom and comprise most its mass.

2. An _____________ is an atom that has the same number of protons as another atom, but has a different number of __________. Some isotopes are radioactive; many are not.

3. The __________________ of an element is the same as the number of protons. In a neutral atom the number of _______ are always equal to the number of ______________. All atoms of the same element have the same number of _______.

4. The periodic table is arranged by increasing ______________.

5. Name the following elements:Group 2, Period 5Group 14, Period 2

6. Using the periodic table, determine the atomic #, mass #, # protons, # electrons, and # neutrons of the following:

Element Atomic #

Mass # # p+ # e- # n0

CICaMn

7. Determine the # protons, neutrons, electrons and mass # of the following:

137Ba+2 60Ni+3 31P-3

ProtonsNeutronsElectronsMass #

8. Write electron configurations for the following:a. chlorine

b. magnesiumc. oxygen

9. Define the periodic law.

10. Vertical columns are called _____________ and have similar properties because _________________________________. Horizontal rows are called ___________ and have predictable properties based on an increasing number of electrons in the outer orbitals.

11. Electronegativity ____________ from left to right and _________ from top to bottom within a group. Atomic radius _____________ from left to right and __________ from top to bottom within a group. Ionization energies generally __________ from left to right and ___________ from top to bottom within a group. Shielding effect is _____________ within a given period and ____________ within given groups from top to bottom.

.

12. Name groups 1, 2, 3-12, 17, and 18 on the periodic table.

13. Name the metalloids.

14. Which of the following ahs the largest atomic radius? Ca Ga Li

15. Which of the following has the lowest ionization energy? Na K Rb

16. Which of the following has the greatest electronegativity? Li N F

17. States the Pauli Exclusion Principle.


18. State Hund’s Rule.

19. Identify the groups that represent the following orbitals: s, p, d, f

20. When an atom gains an electron, it becomes ________ charged and is called a(n) _______. When an atom loses and electron, it becomes _______ charged and is called a(n) ________.

21. Define physical and chemical properties.

22. Write the formulas for the seven (7) diatomic molecules.

23. Classify the following as element (E), compound (C), heterogeneous mixture (het), or homogeneous mixture (hom)

a. Waterb. Carbonc. Salt waterd. Aire. Tap waterf. Salad dressing

24. Explain the periodic trend of reactivity.

25. Briefly describe the contributions to the modern atomic theory of the following scientists:

a. Democritus

b. Dalton

c. JJ Thomson

d. Lord Rutherford

e. Millikan

f. Bohr

g. Mendeleev

h. Plank

i. Heisenberg

j. de Broglie


CH 3 The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. Key concepts includea) nomenclature;b) balancing chemical equations;c) writing chemical formulas;d) bonding types;e) reaction types; and f) reaction rates, kinetics, and equilibrium.


name binary covalent/molecular compounds.

name binary ionic compounds (using the Roman numeral system where appropriate).

predict, draw, and name molecular shapes (bent, linear, trigonal planar, tetrahedral, and trigonal pyramidal).

transform word equations into chemical equations and balance chemical equations.

write the chemical formulas for certain common substances, such as ammonia, water, carbon monoxide, carbon dioxide, sulfur dioxide, and carbon tetrafluoride.

use polyatomic ions for naming and writing formulas of ionic compounds, including carbonate, sulfate, nitrate, hydroxide, phosphate, and ammonium.

draw Lewis dot diagrams to represent valence electrons in elements and draw Lewis dot structures to show covalent bonding.

use valence shell electron pair repulsion (VSEPR) model to draw and name molecular shapes (bent, linear, trigonal planar, tetrahedral, and trigonal pyramidal).

recognize polar molecules and non-polar molecules.

classify types of chemical reactions as synthesis, decomposition, single replacement, double replacement, neutralization, and/or combustion.

recognize that there is a natural tendency for systems to move in a direction of randomness (entropy).

recognize equations for redox reactions and neutralization reactions.

distinguish between an endothermic and exothermic process.

interpret reaction rate diagrams.

identify and explain the effect the following factors have on the rate of a chemical reaction: catalyst, temperature, concentration, size of particles.

distinguish between irreversible reactions and those at equilibrium.

predict the shift in equilibrium when a system is subjected to a stress (Le Chatelier’s Principle) and identify the factors that can cause a shift in equilibrium (temperature, pressure, and concentration.)


CH 3 Problems:

1. State the law of definite proportions.

2. Define empirical formulas, molecular formulas, and structural formulas.

3. What is the difference between an ionic bond and a covalent bond in terms of electrons.

4. Define ionization energy.

5. Define electronegativity.

6. Define polarity.

7. Name the following compounds:a. PbCl2

b. Fe2O3

c. P2O5

d. Cu2CO3

e. BCl3

f. Au(NO3)3

g. Pb(C2H3O2)4

8. Write chemical formulas for the following:a. Water

b. Carbon monoxide

c. Calcium carbonate

d. Lithium nitride

e. Carbon dioxide

f. Iron(III) sulfate

g. Sulfur dioxide

h. Carbon tetrachloride

i. Methane

j. Ammonia

9. Draw the lewis dot diagram for the following:S Br Ca

10.Draw examples of molecules with the following molecular shapes:Linear Trigonal Planar

Bent Tetrahedral

Trigonal Pyramidal

11.Reaction Type ______Write general form a. Synthesisb. Decompositionc. Single replacementd. Double replacemente. Combustion


12.Define oxidation-reduction reaction, oxidation, reduction, oxidizing agent, reducing agent.

13.Define exothermic and endothermic reactions.

14.State Le Chatelier’s Principle. Explain what stresses affect reactions.

15.Write balanced chemical equations for the following reactions and classify the reaction.

a. Phosphorus reacts with oxygen to product

diphosphorus pentoxide.

b. Iron(II) oxide reacts with aluminum to produce

aluminum oxide and iron.

c. Copper(II) hydroxide and potassium sulfate are

produced when potassium hydroxide reacts with

copper(II) sulfate.

16. In the following reactinos, tell which element is oxidized and which is reduced and identify the oxidizing and reducing agents:

a. 2HgO 2Hg + O2

b. S + O2 SO2

c. Cu + 2AgNO3 Cu(NO3)2 + 2Agd. Mg + Cl2 2MgCl2

17. Identify the following as an endothermic reaction or an exothermic reaction:

a. C3H8 + 5O2 3CO2 + 4H2O + 2043kJb. C + H2O + 113kJ CO + H2


CH 4 The student will investigate and understand that chemical quantities are based on molar relationships. Key concepts includea) Avogadro’s principle and molar volume;b) stoichiometric relationships;c) solution concentrations; andd) acid/base theory; strong electrolytes, weak

electrolytes, and nonelectrolytes; dissociation and ionization; pH and pOH; and the titration process.


perform conversions between mass, volume, particles, and moles of a substance.

perform stoichiometric calculations involving the following relationships:- mole-mole;- mass-mass;- mole-mass;- mass-volume;- mole-volume; - volume-volume;- mole-particle;- mass-particle; and- volume-particle.

identify the limiting reactant (reagent) in a reaction.

calculate percent yield of a reaction.

perform calculations involving the molarity of a solution, including dilutions.

interpret solubility curves.

differentiate between the defining characteristics of the Arrhenius theory of acids and bases and the Bronsted-Lowry theory of acids and bases.

identify common examples of acids and bases, including vinegar and ammonia.

compare and contrast the differences between strong, weak, and non-electrolytes.

relate the hydronium ion concentration to the pH scale. perform titrations in a laboratory setting using indicators.

CH 5 The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. Key concepts include

a) pressure, temperature, and volume;b) partial pressure and gas laws;c) vapor pressure;d) phase changes;e) molar heats of fusion and vaporization;f) specific heat capacity; andg) colligative properties.


explain the behavior of gases and the relationship between pressure and volume (Boyle’s Law), and volume and temperature (Charles’ Law).

solve problems and interpret graphs involving the gas laws.

identify how hydrogen bonding in water plays an important role in many physical, chemical, and biological phenomena.

interpret vapor pressure graphs.

graph and interpret a heating curve (temperature vs. time).

interpret a phase diagram of water.

calculate energy changes, using molar heat of fusion and molar heat of vaporization.

calculate energy changes, using specific heat capacity.examine the polarity of various solutes and solvents in solution formation.


CH 4 Problems

1. How many liters of carbon dioxide will be produced if 1.47 mol of oxygen is used in the following reaction?

C3H8 + 5O2 3CO2 + 4H2O + 2043kJ

2. Methane (CH4) reacts with oxygen to produce carbon dioxide and water. If 5.28L of oxygen is used in the reaction at STP, what mass of water will be produced?

3. Calcium carbonate and hydrochloric acid react to produce calcium chloride and carbonic acid. How much calcium chloride will be produced if 7.89g of hydrochloric acid are used?

4. How many moles of oxygen are produced from the decomposition of 19.55 moles of water?

5. How many atoms of iron are required to react with 8.2x1024 molecules of oxygen according to the following equation:

4Fe + 3O2 2Fe2O3

6. How many liters of N2 are needed to produce 6.84L of NH3 in the synthesis of ammonia?

7. Aluminum reacts with sulfuric acid to produce aluminum sulfate and hydrogen. How many grams of aluminum sulfate will be produced from 0.58 mol of sulfuric acid?

8. List the postulates of the kinetic-molecular theory.

9. Write equations for Boyle’s Law, Charles’s Law, combined gas law, and Dalton’ Law.

10.The pressure and volume of a sample of gas at a constant temperature are _________ proportional to each other. (____________ Law) At constant pressure, the volume of a fixed amount of gas is _________ proportional to its temperature. (____________ Law)

11. If 75.3g of aluminum and 110.1g of oxygen react according to the following equation, which is the limiting reactant?

4Al + 3O2 2Al2O3

12.125g of Al2O3 is recovered from the above reaction, What is the percent yield?

13.A gas occupies a volume of 60.0cm3 at 36°C. What will the same gas occupy at standard temperature if the pressure remains constant?


14.A mixture of gas has a pressure of 725mmHg. The mixture contains hydrogen, nitrogen, and argon. The pressure of the hydrogen and argon are 378mmHg and 64mmHg, respectively. What is the pressure of the nitrogen in kPa?

15. If some oxygen gas at 101 kPa and 25°C is allowed to expand from 5.0dm3 to 10.0dm3 without changing the temperature, what pressure will the oxygen gas exert?

16.How many moles can be obtained in a 1.44L flask at 32°C and 93.5kPa?

17.What is the molarity of 2.50x102mL of solution containing 9.46g cesium bromide?

18.What is the molarity of 5.23g of iron(II) nitrate dissolved in 100.0cm3 of solution?

19.What volume of 12M HCl is neeed to prepare 500mL of a 2M HCl solution?

20.List five properties of acids and five properties of bases.

21.Name the following and indicate whether it is an acid or base:

a. KOHb. HClc. H2SO4

d. NaOHe. HNO3

f. H2CO3

22.What is an acid/base titration?

23.Explain the difference between a strong and a weak electrolyte.


CH 5 Problems

Use the following heating curve for water to answer the questions.

1. At which point on the heat curve is heat of fusion represented? Heat of vaporization?

2. At which point(s) on the curve is there no increase in average kinetic energy?

3. Label solid, liquid, gas, melting point, boiling point, freezing, condensation, melting, and vaporization.

4. Define sublimation and deposition.

5. A 21.0g sample of water is cooled from 34.0°C to 28.0°C. How many joules of heat were removed from the water?

6. An 18.7g sample of platinum metal increases in temperature by 2.3°C when 5.7J of heat are added. What is the specific heat of platinum?

7. Define heat of fusion and heat of vaporization.

8. Polar substance dissolve ___________ substances and nonpolar substances dissolve _____________ substances.

9. Label the following on the above phase diagram for water: solid, liquid, gas, melting and boiling points at 1atm, critical point, and triple point.

10.Draw and label arrows that represent freezing, vaporization, and sublimation.

AB

C

D

E

Heat Added


Heats of vaporization and heats of fusion of some substancesSubstance Heat of

vaporizationHeat of fusion

Water 40.7 kJ/mol 6.00 kJ/molMethane 9.2 kJ/mol 0.84 kJ/molMercury 59.4 kJ/mol 2.33 kJ/mol11.How much energy is required to vaporize 22.5g of liquid

mercury?

12.Draw the following molecules and determine if they are polar or non-polar molecules.

a. H2O

b. CH4

c. CO2

d. NH3

e. Cl2

13.Distinguish between inter- and intra-molecular forces

14.List the intramolecular forces in order of decreasing strength.

15.List the intermolecular forces in order of decreasing strength.


CH.6 The student will investigate and understand how basic chemical properties relate to organic chemistry and biochemistry. Key concepts includea) unique properties of carbon that allow multi-carbon

compounds; andb) uses in pharmaceuticals and genetics,

petrochemicals, plastics, and food.


describe how saturation affects shape and reactivity of carbon compounds.

draw Lewis dot structures, identify geometries, and describe polarities of the following molecules: CH4, C2H6, C2H4, C2H2, CH3CH2OH, CH2O, C6H6, CH3COOH.

recognize that organic compounds play a role in natural and synthetic pharmaceuticals.

recognize that nucleic acids and proteins are important natural polymers.

recognize that plastics formed from petrochemicals are organic compounds that consist of long chains of carbons.

conduct a lab that exemplifies the versatility and importance of organic compounds (e.g., aspirin, an ester, a polymer).

CH 6 Problems

1. Which of the following polymers are naturally occurring?a. Nylonb. Proteinc. Polyesterd. Teflone. DNA

2. Which statement describes how plastics differ from nucleic acids?

a. Plastics are synthetic polymers but nucleic acids are natural

b. Plastics are formed from repeated subunits, but nucleic acids are not

c. Plastics are formed from organic compounds, but nucleic acids are not

d. Plastics are polymers, but nucleic acids are monomers

3. Draw Lewis structures for the following and determine molecular polarity:

a. Methane

b. Ethane

c. Ethene

d. Ethyne

e. Ethanol

f. Formaldehyde

g. Benzene

h. Acetic Acid

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