Chemistry Final Exam Comprehensive Review

Day 1 – Date ________Teacher ________

Atomic Theory, Scientists, Basic Atomic Structure, Nuclear Chemistry2.01 Analyze the historical development of the current atomic theory.

Early contributions: Democritus and Dalton. The discovery of the electron: Thomson and Millikan. The discovery of the nucleus, proton and neutron: Rutherford and Chadwick. The Bohr model. The quantum mechanical model.

2.02 Examine the nature of atomic structure. Subatomic particles: protons, neutrons, and electrons. Mass number. Atomic number. Isotopes.

4.04 Analyze nuclear energy. Radioactivity: characteristics of alpha, beta and gamma radiation. Decay equations for alpha and beta emission. Half-life. Fission and fusion.

1) List the scientists who developed the atomic theory and their contributions:

2) Describe the experiment that was performed below:

3) Draw a sketch of what the Bohr model looks like:

4) What is meant by electron transition? Give examples:

5) Draw a sketch of the quantum mechanical model:

How is it different than the Bohr model?

6.) _________the smallest particle of an element that retains the properties of that element

_________a positively charged subatomic particle

_________a negatively charged subatomic particle

_________a subatomic particle with no charge

_________the central part of an atom, containing protons and neutrons

_________atoms with the same number of protons, but different numbers of neutrons in the nucleus of an atom

_________the total number of protons and neutrons in the nucleus of an atom

_________The number of protons in the nucleus of an element

_________the weighted average of the masses of the isotopes of an element

_________one-twelfth the mass of a carbon atom having six protons and six neutrons; “AMU”

7.) Complete the table below:

Element Symbol # Protons # Electrons # NeutronsAluminum ( ____ )Sodium ( ____ )Neon ( ____ )Potassium ( ____ )Sulfur ( _____ )

8.) What is an isotope and how can they be written?

9.) How many protons, neutrons, and electrons are there in 40K?

10.) How many protons, electrons, and neutrons are in Carbon-14.

11.) List the different types of radiation, their symbols, and their charges/number:

12.) Carbon -14 is produced when Nitrogen-14 captures an electron, write out the nuclear reaction for this process.

13.) What particle is needed to complete the following equation?N + ____ C + H

14.)

What is the definition for half-life?

How many half-lives pass if you only have 1/32 of parent remaining?

13) Compare fusion and fission:

Day 2 – Date ________Teacher ________

Periodic Trends, Electron Configuration, Bohr Model2.01 Analyze the historical development of the current atomic theory.

Early contributions: Democritus and Dalton. The discovery of the electron: Thomson and Millikan. The discovery of the nucleus, proton and neutron: Rutherford and Chadwick. The Bohr model. The quantum mechanical model.

3.01 Analyze periodic trends in chemical properties and use the periodic table to predict properties of elements.

Groups (families). Periods. Representative elements (main group) and transition elements. Electron configuration and energy levels. Ionization energy. Atomic and ionic radii. Electronegativity.

4.01 Analyze the Bohr model in terms of electron energies in the hydrogen atom. The spectrum of electromagnetic energy. Emission and absorption of electromagnetic energy as electrons change energy levels.

1.) Review Bohr model again……draw a sketch showing electron transition.

How is this related to atomic spectra?

What is EMR? Examples?

2.) Groups on the periodic table are ______________ Families on the periodic table are _________________

3.) List all the major groups of elements on the periodic table.

4.) What are electron energy levels?

5.) What is the electron configuration ofmercury _______________________________calcium __________________________________neon _________________________________fluorine __________________________________strontium ______________________________

iron _____________________________________

6.) Define and describe each of the following Periodic Trends -Electronegativity:

Ionization energy:

Atomic Radius (Radii):

7.) Sketch and explain IN DETAIL the 3 periodic trends on the periodic table:

8.) Give the longhand electron configuration for arsenic.

9.) The largest atoms are in the ___ corner of the table.

Day 3 – Date ________Teacher ________

Bonding, Nomenclature, Acid & Base Nomenclature2.03 Apply the language and symbols of chemistry.

Name compounds using the IUPAC conventions. Write formulas of simple compounds from their names.

2.06 Assess bonding in metals and ionic compounds as related to chemical and physical properties.

2.07 Assess covalent bonding in molecular compounds as related to molecular geometry and chemical and physical properties.

Molecular. Macromolecular. Hydrogen bonding and other intermolecular forces (dipole/dipole interaction, dispersion). VSEPR theory.

1.) Name the following chemical compounds:

a) PbBr2 __________________________________

b) NH3 __________________________________

c) P4 __________________________________

d) CaS __________________________________

e) H2SO4 __________________________________

f) V(CO3)2 __________________________________

g) P2O5 __________________________________

2.) Write the chemical formulas and identify the type of bonding involved in the following compounds:

a) ammonium nitrate __________________________________

b) fluorine __________________________________

c) boron trichloride __________________________________

d) iron (III) phosphate __________________________________

e) nitric acid __________________________________

f) potassium carbonate __________________________________

g) dinitrogen tetrachloride __________________________________

3.) Write the formula and name the type of bonding that occurs in each of the following compounds:

a) ammonium sulfate __________________________________

b) boron trichloride __________________________________

c) hydrogen fluoride __________________________________

d) dihydrogen monoxide ________________________________

4.) Compare and contrast: (i.e. how are they formed?, how do their physical properties differ?, what are their representative particles called?, etc.)

Ionic Compounds Molecular Compounds

5.) Draw the Lewis Structures for each of the following molecular compounds, and describe the shape of each molecule using VSEPR theory:

water carbon dioxide methane

Find 5 other examples of molecular compounds you can draw the Lewis structures for:

6.) What are the main 5 geometric shapes associated with the VSEPR theory? Draw each of them.

7.) Classify the following as chemical or physical changes, and explain why:rusting of iron

digestion of meat

boiling water

8.) Write formulas for the compounds in:

magnesium fluoride

dinitrogen pentoxide

sodium sulfate

phosphoric acid

9.) Name the compounds:

KNO3 HBr SO3 FeCl3

10.) Draw the Lewis diagram & specify the molecular geometry & polarity:

Day 4 – Date ________Teacher ________

Chemical Reactions, Reaction Types, Predicting Products5.01 Evaluate various types of chemical reactions.

Analyze reactions by types: single replacement, double replacement (including acid-base neutralization), decomposition, synthesis, and combustion including simple hydrocarbons.

AsH3 BF3

Predict products.

5.02 Evaluate the Law of Conservation of Matter. Write and balance formulas and equations. Write net ionic equations.

5.03 Identify and predict the indicators of chemical change. Formation of a precipitate. Evolution of a gas. Color changes. Absorption or release of heat.

5.06 Assess the factors that affect the rates of chemical reactions. The nature of the reactants. Temperature. Concentration. Surface area. Catalyst.

1.) Identify the 5 types of chemical reactions and give an example for each:Reaction Type Example

2.) What is a hydrocarbon?

Examples:

3.) For each of the following reactions, identify the reaction type, predict the products and balance the equation. Include physical states. Word equations must first be converted to formulas.

____ Li(s) + ____ N2(g)

Reaction Type: _______________________________

____ Mg(s) + ____ CrCl3(aq)

____ CoBr2(s)

What does the delta mean?

____ C4H8(g) + ____ O2(g)

Aqueous solutions of potassium bromide and silver nitrate react to form a white precipitate.

Solid nickel is added to an aqueous solution of iron(II) sulfate.

4.) Write the reaction and balance:

Hydrogen and nitrogen react together to produce ammonia gas (note that the reaction is a reversible one - ammonia also breaks up to form hydrogen and nitrogen):

Reaction Type: _______________________________

Reaction Type: _______________________________

Reaction Type: _______________________________

Reaction Type: _______________________________

Reaction Type: _______________________________

5.) Write the net ionic equation for the following double displacement reaction:

H PO (aq) Ca(OH) (aq) Ca (PO ) (aq) H O(l)

6.) Consider the following double displacement reaction: copper(II) oxide reacts with sulfuric acid to produce copper(II) sulfate and water.

I. Write a balanced chemical equation that describes this reaction.

II. Write a complete and net ionic equation for this reaction.

7.) Predict the precipitate that forms when aqueous solutions of silver nitrate and potassium chloride react to form products in a double-replacement reaction. Include the complete and net ionic equations for this chemical equation as well.

8.) What are some observations that can be made to indicate that a CHEMICAL change has occurred? DESCRIBE

8) What does “reaction rate” mean?

What factors can affect the rate at which a reaction occurs? LIST AND DESCRIBE EACH.

9.) Draw the energy diagrams for an exothermic and endothermic reaction. Label the important part of the energy diagram that we learned in class….

Day 5 – Date ________Teacher ________

The Mole, Stoichiometry, Percent Composition, Empirical/Molecular Formula3.02 Apply the mole concept, Avogadro's number and conversion factors to chemical calculations.

Particles to moles. Mass to moles. Volume of a gas to moles. Molarity of solutions. Empirical and molecular formula. Percent composition.

3.03 Calculate quantitative relationships in chemical reactions (stoichiometry). Moles of each species in a reaction. Mass of each species in a reaction. Volumes of gaseous species in a reaction.

1.) Write the “double mole map” here:

2.) What are examples of representative particles forMolecular compounds:

Ionic compounds:

Individual elements:

3.) How many grams of methane are there in 1.23 x 1024 molecules?

How many liters of carbon dioxide are there in 45 grams?

Convert 5.6678 moles of arsenic to grams:

89.3 liters of water vapor is equivalent to ________________ molecules of water vapor.

How many moles of ammonium sulfate do I have if my sample weighs 496 grams?

Calculate the number of moles in 13.5 g magnesium nitrate (Mg(NO3)2)

4.) Write the formula for molarity:

How is molarity useful / important?

5.) Solve the following molarity problems. Show all work and include correct units!

Find the molarity of a solution in which 58 g of NaCl are dissolved in 2.5 L of solution.

How many grams of KMnO4 should be used to prepare 2.00 L of a 0.500M solution?

What volume of 0.25M solution can be made from 5.0 g of KCl?

Find the molarity of a 450 mL solution containing 13.7 g of ZnSO4.

How many grams of CuCl2 are required to make 75 mL of a 0.20M solution?

6.) A saline solution contains 0.90 g of NaCl per 100.0 mL of solution. What is its molarity?

7.) What is the definition of empirical formula?

8.) Identify each of the following as an empirical formula or a molecular formula:

C2H4 __________________ NaCl __________________H2O __________________ C6H12O6 __________________H2O2 __________________ N2O4 __________________NH4 __________________ H2S __________________

9.) Determine the empirical formula of a compound that contains 36.5% sodium, 25.4% sulfur, and 38.1% oxygen.

10.)  An organic compound has an empirical formula of CH and a molecular mass of 78 g.  What is the molecular formula?

11.) Determine the percent composition of Fe3(PO4)2

12.) A 350.0 gram sample of iron ore has been tested to contain 97.7 grams of iron. What is the percentage of iron in the ore?

What percentage of the iron ore IS NOT iron?

13.) What is the empirical formula of a compound that is 52.5% phosphorus and 47.5% oxygen?

14.) Determine the empirical formula of the following compound:10.0% C 0.80% H 89.1% Cl

15.) How many grams of carbon dioxide will be made when 100 grams of methane burn in an excess of oxygen?

16.) Write the empirical formula for a compound that is 26.5% K, 35.4% Cr, 38.1% O.

17.) Review this again!! Write the formula and name the type of bonding that occurs in each of the following compounds:

a) ammonium sulfate __________________________________

b) boron trichloride __________________________________

c) hydrogen fluoride __________________________________

d) dihydrogen monoxide ________________________________

18.) Find the percentage composition of iron(III) oxide. Formula: _______________

19.) Find the mass percentage of water in copper(II) sulfate pentahydrate.

20.) How many grams of iron can be obtained from a 268-g sample of iron(III) oxide?

21.) A compound used to test for the presence of ozone in the stratosphere contains 96.2% thallium and 3.77% oxygen. What is its empirical formula?

22.) The molecular mass of benzene, an important industrial solvent and know carcinogen, is 78.0 g/mol and its empirical formula is CH. What is the molecular formula of benzene?

23.) Ascorbic acid, or vitamin C, has a percent composition of 40.9% C, 4.58% H, and 54.5% O. Its molecular mass is 176.1 g/mol. Find its empirical and molecular formulas. (HINT: Multiply by 2, 3, or 4 to get whole number subscripts.)

24.) How many magnesium sulfate molecules are in 25.0 g?

25.) Find the molarity of a 750 mL solution containing 346 g of potassium nitrate.

26.) Calculate the number of grams required to make a 50.0 mL solution of 6.0M NaOH.

27.) Find the % composition of copper(II) chloride.

Day 6 – Date ________Teacher ________

Gas Laws, Basic Matter2.04 Identify substances using their physical properties:

Melting points. Boiling points. Density. Solubility.

2.05 Analyze the basic assumptions of kinetic molecular theory and its applications: Ideal Gas Equation. Combined Gas Law. Dalton's Law of Partial Pressures.

2.08 Assess the dynamics of physical equilibrium. Interpret phase diagrams. Factors that affect phase changes.

1.) Identify the gas laws that explain these situations. Specify the variables involved and whether they have a direct/inverse relationship.

A balloon pops after floating high into the atmosphere.

A balloon pops in a hot car on a summer day.

Do not store aerosol cans at temperatures above 120F. “Danger of explosion.”

2.) Identify the gas law and solve the problem.A jar is tightly sealed at 22C and 772 torr. What is the pressure inside the jar after it has been heated to 178C?

300.0 mL of gas has a pressure 75.0 kPa. When the volume is decreased to 125.0 mL, what is its pressure?

50.0 L of gas has a temperature of 75C. What is the temp in Celsius when the volume changes to 110 L?

What is the volume of a container that holds 48.0 g of helium at a pressure of 4.0 atm and temperature of 52C?

3.) A gas occupies 325 L at 25C and 98.0 kPa. What is its volume at 70.0 kPa and 15C?

4.) Define: STP Temperature

Kelvin Volume

Air Pressure

5.) What is the meaning of DENSITY?

What is the formula for density?

What are the units for density?

How is density important?

6.) Limestone has a density of 2.72 g/cm3. What is the mass of 24.9 cm3 of limestone?

7.) Helium has a density of 0.017 g/L. What is the volume of a weather balloon that contains 38 g of helium?

8.) A certain metal has a mass of 50 grams and a volume of 18.5 cm³. What is the density of the metal? Identify the metal.

51.15 J of heat was applied to a metal that has a mass of 5 grams. The metal increased in temperature by 10 degrees C. Identify the metal.

9.)

What is normal boiling?

What is normal melting?

What is the triple point?

What substance is this a phase diagram for?

10.) How many moles of chloroform gas, CHCl3, are required to fill a 253-mL flask at 100.0C and 940 torr?

11.) You want the pressure inside a bottle to be 75.0 kPa at 23C. At what temperature in Celsius should you seal the bottle when the pressure is 1.12 atm?

12.) A Marshmallow Peep® has a volume of about 45.0 cm3 at 101 kPa. What pressure is required to increase its size to 150.0 cm3 assuming no air escapes from the Peep®.

13.) A bottle containing hydrogen, nitrogen, and oxygen gases has a total pressure of 80 kPa. If the partial pressure of the hydrogen gas and the nitrogen gas in the container equals 10 kPa each, what is the partial pressure of the oxygen gas?

Formula:Day 7 – Date ________

Teacher ________

Solutions, Acids/Bases5.04 Identify the physical and chemical behaviors of acids and bases.

General properties of acids and bases.

Concentration and dilution of acids and bases. Ionization and the degree of dissociation (strengths) of acids and bases. Indicators. Acid-base titration. pH and pOH.

1.) List the general properties of acids and bases for each category below:ACIDS BASES

Taste

pH range

Feel

What ion does it produce when dissolved in water?

Electrolytes it can produce:

Arrhenius Theory states:

Brønstead-Lowry Theory states:

2.) Refresh on the pH / pOH formulas. Write down each formula and when you would use it:

A student prepares a 0.80 M solution of potassium hydroxide.

Find the:

a) concentration of hydronium ions

b) pH

c) the concentration of the hydroxide ions.

3). State whether the following are acids or bases:

Have a sour taste.

React with metals.

Release H+ ions in water

Feel slippery

Turn blue litmus paper red.

Release OH- ions in water

4.) Define acids and bases according to the Arrhenius theory & Brønsted-Lowry theory:

5.) Identify each substance as acid, base, conjugate acid, or conjugate base.H2S + H2O HS – + H3O+

6. Give the conjugate acids of: NH3 and Br –.

7. Give the conjugate bases of: H3O+ and HSO4–.

8. Find the pH and pOH of 0.75M HCl.

9. Find the molarity of a KOH solution with a pH of 9.5

10. If 43.5 mL of 0.15 M HBr is required to neutralize 25.0 mL of NaOH, what is the molarity of the NaOH?

VOCAB: hydronium ion neutralization reaction

hydrogen ion hydroxide ion

titration equivalence point

strong/weak acid/base

Day 8 – Date ________Teacher ________

Reviewing Reference Tables, Thermochemistry ___________________________________________________________________

4.02 Analyze the law of conservation of energy, energy transformation, and various forms of energy involved in chemical and physical processes.

Differentiate between heat and temperature. Analyze heating and cooling curves. Calorimetry, heat of fusion and heat of vaporization calculations. Endothermic and exothermic processes including interpretation of potential energy. Diagrams (energy vs. reaction pathway), enthalpy and activation energy.1.) How many joules are required to heat 250 grams of liquid water from 5° to 95° C?

2.) How many joules are required to melt 90 grams of water?

3.) How many joules are required to boil 150 grams of water?

4.) How many joules are required to heat 100 grams of water from 25°C to 105°C?

5.) How many joules are given off when 120 grams of water are cooled from 55°C to -15°C?

6.) If I burn 0.315 moles of hexane (C6H14) in a bomb calorimeter containing 5.65 liters of water, what’s the molar heat of combustion of hexane if the water temperature rises 55.4°C?

7.) Study the energy diagram below:

Is this reaction exothermic or endothermic?

How can you tell?

Draw a line indicating how a catalyst would affect the forward reaction.

The activation energy for the forward reaction is _____________________________

The activation energy for the reverse reaction is _____________________________

The heat of the reaction for the forward reaction is ___________________________

The change in enthalpy for the forward reaction is _____________________________

8.) Study the heating curve below:

What is happening to the average kinetic energy of the molecules in the sample during section 2?

As a substance goes through section (2), what happens to the distance between the particles?

What is the name of the process happening during section (4)?

What would be the name of the process happening during section (4) if time were going the other way?

What is the melting point of this substance?

At what temperature would this sample finish boiling?

When a given quantity of water is heated at a constant rate, the phase change from liquid to gas takes longer than the phase change from solid to liquid because

a.     The heat of vaporization is greater than the heat of fusionb.     The heat of fusion is greater than the heat of vaporizationc.     The average kinetic energy of the molecules is greater in steam than in waterd.     Ice absorbs energy more rapidly than water does

The temperature at which a substance in the liquid state freezes is the same as the temperature at which the substance

a.     Meltsb.     Sublimesc.     Boilsd.     Condenses

Is this curve showing an endothermic or an exothermic process?

34oC is equal to ____________K.

–128oC is equal to ____________K.

How many kJ are in 750 J?

If 150 grams of water is heated from 20oC to 30oC, the number of joules of heat energy absorbed is…

If a 2.0 g sample of water at 5.0oC absorbs 21.8 J of heat energy, the temperature of the sample will be raised by…

The temperature of 50.0 grams of water was raised to 50oC by the addition of 4180 J of heat energy. What was the initial temperature of the water?

A sample of water is heated from 10.oC to 15oC by the addition of 130 joules of heat. What is the mass of the water?

Day 9 – Date ________Teacher ________

Redox Reactions, Practice Questions5.05 Analyze oxidation/reduction reactions with regard to the transfer of electrons.

Assign oxidation numbers to elements in REDOX reactions Identify the elements oxidized and reduced. Write simple half reactions. Assess the practical applications of oxidation and reduction reactions.1.) Determine the oxidation states of each of the following compounds:

a) ammonium nitrate __________________________________

b) fluorine __________________________________

c) boron trichloride __________________________________

d) iron (III) phosphate __________________________________

e) nitric acid __________________________________

f) potassium carbonate __________________________________

g) dinitrogen tetrachloride __________________________________

2.) In each of the following reactions, determine what was oxidized and what was reduced:

Ca + H2O → CaO + H2

Element Oxidized ___________________________

Element Reduced ___________________________

Oxidizing Agent ___________________________

Reducing Agent ___________________________

2 H2 + O2 → 2 H2OElement Oxidized ___________________________

Element Reduced ___________________________

Oxidizing Agent ___________________________

Reducing Agent ___________________________

Cu + 2 AgNO3 → 2 Ag + Cu(NO3)2

Element Oxidized ___________________________

Element Reduced ___________________________

Oxidizing Agent ___________________________

Reducing Agent ___________________________

2 AgCl + Co → CoCl2 + 2AgElement Oxidized ___________________________

Element Reduced ___________________________

Oxidizing Agent ___________________________

Reducing Agent ___________________________

Mg + Ca(OH)2 → Ca + Mg(OH)2

Element Oxidized ___________________________

Element Reduced ___________________________

Oxidizing Agent ___________________________

Reducing Agent ___________________________

Cu + Zn+2 → Zn + Cu+2

Element Oxidized ___________________________

Element Reduced ___________________________

Oxidizing Agent ___________________________

Reducing Agent ___________________________

3.) Give the oxidation state of each element in the following compounds:

Zn

Pb(NO3)2

Zn(NO3)2

HI

HIO4

H2SO4

CO

CO2

CO3-2

NH3

H2O

H2O2

NO3-

O2

N2

CH4

Other Important Things to Know/Remember:

What are the diatomic molecules?

What does STP mean? What values do they include?

What are the two gas laws equations you should know?

What are the formulas for methane, hydrochloric acid, sulfuric acid, and water vapor?

What are the charges on Al, Zn, and Ag ions, respectively?

What does a combustion equation ALWAYS look like?

What unit is molar mass given in?

What is the molar mass of water, carbon dioxide, oxygen gas, nitrogen gas, and hydrogen gas?

How do the solubility rules work? When do you use them?

What is the activity series for metals and halogens? When do you use them? What do they tell you?