Chemistry AChem. A.Final Exam Review

Objectives Covered

Chapter 1 Identify the location, relative mass, and charge for electrons, protons, and neutrons.

Chapter 2 Write and interpret the complete electron configuration of elements in the first four rows of the periodic

table. Describe the energy levels, sub levels, and orbitals of an atom. Determine the number of electrons that each energy level can hold. Determine the number of electrons that each sub-shell can hold Identify the orientation and shape of each orbital. Explain the Aufbau principle and place electrons accordingly.

Chapter 3 Write and interpret Noble Gas configurations for main group elements. Predict the general trend in atomic radius across a period and down a group and explain why these trends

occur. Predict the general trend in first ionization energy across a period and down a group and explain why

these trends occur. Predict the general trends in electronegativity across a period and down a group and explain why these

trends are present. Be able to interpret graphs that illustrate the three-periodic trends identified above (atomic radii, first

ionization energy and electronegativity). Predict oxidation states (charges) on main group elements. Predict bonding capacity for main group elements using their electron structure. (Explain why atoms form

ions and/or bond.) Using the ion reference chart, identify the various oxidation states on transition metals.

Chapter 4 Predict if the bonding between two atoms of different elements will be primarily ionic or covalent. Predict the formula for binary compounds of main group elements. Draw Lewis structures (dot diagrams) for simple compounds. Given the name, write the formula of a compound. Given a formula, name the compound. Using the ion reference chart, identify the charges and formulas of the polyatomic ions. Distinguish and explain the differences between an ionic, covalent and metallic bond. Identify and explain the properties of an ionic substance (melting point, boiling point, brittleness) Interpret and identify illustrations of ionic, covalent and metallic bonds.

Chapter 5 Predict the formula and name of covalent molecules.

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Predict the formula and name of acids. Draw Lewis structures (dot diagrams) for molecules. Determine the polarity of a bond using the eletronegativities of the atoms involved. Determine the polarity of a compound/molecule using the electronegativies and structure of the

compound/molecule. Predict the shape of individual molecules using VSEPR (valence shell electron pair repulsion.)

Chapter 6 Balance chemical equations applying the conservation of matter. Classify chemical reactions as single replacement, double replacement, synthesis, decomposition, and

combustion. Predict the products produced in a single replacement, double replacement, synthesis, decomposition,

and combustion reaction.

Resources Provided Periodic Table without group or period numbers Ion Reference chart VSEPR chart

Practice Review Questions

CHAPTER 1

1. Identify the charge, location and relative mass (amu) of each subatomic particle: proton, neutron, electrons

2. Describe the nature of the nucleus. (What particles are within it? Does it have a charge, if so what is it? How is the mass of the nucleus determined?)

3. Describe the general location of the electrons.

Review Questions Taken From Previous Chapter Reviews

CHAPTER 2

1. A. Write the electron configuration for each noble gas below.He: __________________________________________

Ne: __________________________________________

Kr: __________________________________________

B. What is special or unique about these three elements (noble gases)?

2. Determine the element each electron configuration represents. Then, determine the number of electrons in the highest energy level.

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A. 1ss2s2 _______ ______

B. 1s22s22p63s23p4 _______ ______

C. 1s22s22p63s23p64s23d104p1 _______ ______

D. 1s22s22p63s23p64s23d104p3 _______ ______

E. 1s22s22p63s23p64s23d104p65s1 _______ ______

3. Identify the name of each orbital(s) shown.

A. _________________ B. _____________

4. A. Circle one orbital in the Figure 1.

B. How many electrons can fit in each orbital? ______

C. How many total electrons can fit in these three orbitals? ____

D. In which sub-level would you expect to find these orbitals? _______

E. Circle one orbital in the Figure 2.

F. How many electrons can fit in 1 orbital? _____

G. How many total electrons can fit in these 5 orbitals?

H. In which sub-level would you expect to find these orbitals? _________

5. Complete the following chart:

Figure 1

Figure 2

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Sublevel Number of Orbitals within the sublevel

Total # of electronsthat can fit in the

sublevel

s

p

d

6. Complete the following chart:

Energy Level # of and Types of Sub-levels

Total # of electrons that can fit in the

energy level

1

2

3

4 - 7

Determine if the following statements are True or False. If false, re-write to make True.

7. Energy level 5 is closer to the nucleus than energy level 4.

8. The energy of each sublevel within an energy level is the same.

9. The energy of each oribital within a sublevel in the same.

10. Every electron in sublevel 3p has the same energy.

11. Every electron in energy level 3 has the same energy.

12. The lowest energy sublevel is d.

13. Each d sublevel has 10 oribtals.

14. An electron's position and speed can be determined at the same time.

15. Electrons behave like waves and particles.

16. Atomic orbitals are regions in space where electrons are likely to be found.

17. In its ground state, hydrogen's electron will be found in the 3rd energy level.

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18. All elements have the same ground state electron configuration.

19. In most instances, 3d is filled before 4p because 3d has less energy.

20. All orbitals within a sublevel must be filled with 2 electrons.

21. Electrons are stationary and do not move or spin.

22. Electrons orbit the nucleus within orbitals (energy levels), like the planets orbit the sun.

23. Orbitals are the same as energy levels.

Determine the electron configuration24. calcium

25. phosphorus

26. strontium

27. arsenic

28. zinc

Complete the chartAtom # of

Valence Electrons

Ion each will form (if no ion, write no ion)

29.) Sr

30.) He

31.) N

32.) O

33.) C

34.) Ne

35.) Li

36.) Al

37.) Br

CHAPTER 3

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Determine the element represented by the following notations. Then determine if the element identified is a metal, non-metal or metalloid.

1. [Ne]3s23p5 _____________ __________________________________

2. [Ar]4s23d104p4 _____________ __________________________________

3. [Ar]4s2 _____________ __________________________________

Determine the noble gas configuration for the following elements. Then determine if the element identified is a metal, non-metal or metalloid.

4. S __________________________ _______________________________

5. Ti ___________________________ ________________________________

6. Se ____________________________ _________________________________

Answer the following questions:

7. An element with an atomic number of 7. What does its noble gas configuration start with?

8. An element is located in period 4. What does its noble gas configuration start with?

9. An atom with a noble gas configuration: [Ne]3s2, belongs to what group?

10. An atom with a noble gas configuration: [Kr]5s24d105p4 belongs to what group?

Based on periodic trends, answer the following questions for these atoms: Li, Be, Mg, Na

20. Which element has the highest ionization energy? _____________21. Which element has the lowest electronegativity? ______________22. Which element has the largest radius? _________

Based on periodic trends, answer the following questions for these atoms: P, S, Cl, F

23. Which element has the highest ionization energy? _______________24. Which element has the lowest electronegativity? ________________25. Which element has the largest radius? _________________

Based on periodic trends, answer the following questions for these atoms: Li, Be, N, O

26. Which element has the highest ionization energy? _______________27. Which element has the lowest electronegativity? ________________28. Which element has the largest radius? _________________

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Identify which of two graphs depicts the trend of electronegativity and which depicts ionization energy. Then explain how you were able to determine the identity of each.

Determine if each statement if True or False. Make the False statements True.35. Ionization energy is the energy required to remove the most tightly held electron from the inner

energy level of that atom.

36. The 2nd, 3rd and 4th ionization energies require less energy and separate an electron from a particle that has increasingly greater net positive charge.

37. With increasing atomic number within a group, ionization energy decreases, because each successive member of a group has 1 less energy level of electrons.

39. With increasing atomic number within a group, ionization energy decreases, because there are fewer energy levels that shield the outer electrons from the attractive force of the nucleus.

40. Electronegativity is the measure of the energy required to remove an electron from an atom.

41. Electronegativity is the measure of the ability of an element to attract electrons to itself when bonding with another substance.

42. An increasing number of protons across a period in addition to an increase in the energy levels, results in a smaller radii.

43. When going across a period, electronegativity decreases due to the need to complete the valence shell.

44. The noble gases have the lowest ionization energy and lowest electronegativity.

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CHAPTER 4For each description, write the formula of the compound.

1. Element X from group 1A and element Y from group 5A.

2. Element X from group 2A and element Y from group 7A.

3. Iron (III) and an element Y from group 6A.

4. Ammounium and element Y from 5A.

5. Manganese (III) and element Y from group 7A.

6. Element X from group 2A and sulfite.

7. Element X from group 3A and chlorite.

8. A transition metal X with a charge of +2 and nitrate.

9. A transition metal X with a charge of +4 and element Y from group 7A.

10. A transition metal X with a charge of +1 and chromate.

Determine the best answer.

11. Which of the following choices is an example of a polyatomic iona. CO2

b. Mg+2

c. MnO4

d. OH-1

12. Which of the following contain a polyatomic ion?a. NO2

b. NaClc. NH4Cld. FeO

13. True or False: Monatomic cations end with–ite? (If false, fix to make true.)

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14. An ionic bond …a. occurs when electrons are shared between atoms.b. occurs when electrons are shared between ions.c. is the force that holds oppositely charged ions together.d. is the force that causes ions to form when electrons are transferred.

For each substance determine the substance contains ionic or covalent bond(s)

15. H2O ___________________

16. HCl ___________________

17. MnF3 ___________________

18. Na2O ___________________

19. BeSO4 ____________________

20. O2 ____________________

21. NH3 ____________________

22. NH4Fl ____________________

23. H2 ____________________

24- 26 For each picture, determine the type of bond shown. Then explain why.

For Ionic Compound(s) determine if the property listed is high or low. Then explain why

High or Low

27. Melting Point ________________

28. Boiling Point ________________

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29. conductivity as a solid _______________

30. conductivity in solution _______________

CHAPTER 5Part 1A. For each substance determine if it represents an ionic bond, covalent bond or acid. B. Then name or write the formula of the substance.

1. Fe2O3 7. ammonium phosphate 13. Cu2CO3

A. A. A.

Explanation

Explanation

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B. B. B.

2. nitrous acid 8. hydrobromic acid 14. N2OA. A. A.

B. B. B.

3. potassium sulfite 9. sulfuric acid 15. Sn(SO4)2

A. A. A.

B. B. B.

4. PCl5 10. MgCrO4 16. nitric acidA. A. A.

B. B. B.

5. HI 11. carbon tetrachloride 17. silver nitrateA. A. A.

B. B. B.

6. nitrogen trihydride 12. hydrochloric acid 18. H3NA. A. A.

B. B. B.

Choose the best answer

19. All of the following have bent molecular shapes EXCEPT ______________.a. H2Beb. H2Sc. H2Od.H2Se

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20. Which of the following is not a polar molecule?a. H2Sb. CCl4

c. H2Od. HCl

True/False - If False correct to make true.

22. Boron is stable with 3 bonds.

23. The VSEPR theory predicts that lone pairs of electrons have no influence on the bond angles in molecules.

24. The attractive force that holds a molecule together is called an ionic bond.

25. If two atoms combine with very similar electronegativities the bond formed will be more polar in nature.

26. A bond between two atoms may be polar even if the molecule to which they belong is non-polar.

For questions 34 - 38, use the electronegativity chart on page 263 in your textbook.A. Determine if the bonds involved within each molecule are polar.B. Determine if the molecule itself is polar. You will need to draw the structural model of the molecule.

34. HCl A. B.

35. H2O A. B.

36. CCl4 A. B.

37. CF4 A. B.

38. NH3 A. B.

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39. Study the structural model of boron triflouride. Determine if the molecule is polar or non-polar? Explain.

For each moleculeA. Draw the Lewis Dot DiagramsB. Draw the structural model be sure to show the lone pairs on the central atom.C. Determine the molecular shape and the bond angles

40. H2Se 42. CCl4 44. BCl3

A. A. A.

B. B. B.

C. C. C.

41. N2 43. NH3 45. SiO2

A. A. A.

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B. B. B.

C. C. C.

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