Unit 2 Review KEY

Unit 2 Review: Atomic Structure, Nuclear chemistry, Quantum Theory, Periodic Table
Basic Atomic Structure
1. Complete the following chart.

2. Atomic mass is a decimal. Why? It is a weighted average of all naturally occurring isotopes.
3. Define Isotope. The same element, different number of neutrons, therefore a different mass.
4. Positively charged ions are formed when atom _loses_ (lose, gain) electrons.
5. Calculate the atomic mass of the following sample of Silicon. 92.21 % 28Si, 4.70% 29Si, and 3.09% 30Si. Answer to 2 decimal places, remember units.
0.9221 x 28 = 25.8188 amu
0.0470 x 29 = 1.363 amu
0.0309 x 30 = 0.927 amu 28.1088 amu = 28.11 amu
6. When at atom gains two electrons, it becomes an ion with a charge of _-2___.
Nuclear Chemistry
Complete the following nuclear decay reactions. Type of Decay (α, β, γ)
7. __gamma____ 99m43Tc  9943Tc + __00γ____
8. __beta______ 24795Am  0-1e + _24796Cm____
9. __alpha________ 17593Np  42He + _17191Pa___
10. What is the difference between fission and fusion? Which releases more energy?
Fission is a big nucleus breaking apart. Fusion is small nuclei combining, this releases more energy.
11. What is the difference between nuclear reactions and chemical reactions? (hint: which parts of the atom are involved)
In nuclear reactions, the nucleus changes, so the identity of the atoms involved changes. In a chemical reaction, the distribution of electrons changes, but the identity of all atoms involved stays the same.
The Bohr Model and Quantum Theory
12. The probability region through which an electron may move is a(n) __orbital________.
13. At a p sublevel, there are (how many) _3__ orbitals. At an s sublevel, there are __1_ orbitals. At a d sublevel there are _5__ orbitals.
14. The maximum number of electrons possible to any p sublevel is ___6________.
15. What was Bohr’s major contribution to our understanding of atomic structure?
Bohr originated the idea of different energy levels (rings, shells) in which electrons could be found.
16. If an electron has absorbed energy and has shifted to a higher energy level, the electron is said to be in an Excited state .
17. When all the electrons in an atom are in the lowest available energy levels, the atom is in the ground state .
18. Spectral lines of elements are caused by (1) electrons turning in their orbitals, (2) electrons jumping to higher energy levels, (3) the vibration of the nucleus, (4) electrons falling to lower energy levels.
19. Write the full electron configurations (not shortcut) for Sulfur (S) and Manganese (Mn). Circle the valence electrons.
S: 1s22s22p63s23p4 Mn: 1s22s22p63s23p64s23d5
20. Write the shortcut electron configuration for Strontium (Sr) and Oxygen (O). Circle the valence electrons.
Sr: [Kr]5s2 O: [He]2s22p4
Organization of the Periodic Table
21. Elements may react to form ions that have electron configurations like those of the noble gases.
22. Which element is in group 15 and period 2? _P____
23. Which elements are halogens? What charge of ions will they make? Why (what is happening)?
F, Cl, Br, I (sometimes At is included)  they will make (-1) ions because they each will gain an electron
24. Which elements are alkaline earth metals? What charge of ions will they make? Why (what is happening)?
Be, Mg, Ca, Sr, Ba, Ra  they will make (+2) ions because they will each lose two electrons
25. List at least three elements that are metals. List three characteristics of metals.
Anything below and to the left of the stairstep line. Malleable, ductile, good conductors of heat and electricity, tend to lose electrons when forming ions.
26. List at least three elements that are nonmetals. List three characteristics of nonmetals.
Anything above and to the right of the stairstep line. Poor conductors of heat and electricity, tend to gain electrons when forming ions, brittle, many are gases at room temperature.
27. Which elements are the metalloids? B, Si, Ge, As, Sb, Te (sometimes Po and At are included)
28. Describe electronegativity. The ability of an atom to attract electrons in a chemical bond (how tightly at atom is able to hold on to electrons).
29. What happens to the atomic radius as you move left to right across one period? Why? Radius decreases because there are more protons and electrons pulling on one another (higher electronegativity) without gaining any more energy levels.
30. Which element has the highest electronegativity on the whole table? F
31. An atom is chemically stable when all of the orbitals in the outermost energy level are filled.
32. Which group of elements has NO electronegativity and VERY HIGH ionization energy? Why? Noble Gases, no electronegativity because they don’t bond (see definition of electronegativity), high ionization energy because they are so stable by themselves it is very very hard to remove an electron.
33. Is an oxygen ion (oxide) bigger or smaller than an oxygen atom? Why? Oxide is bigger because it has gained two electrons.
34. Is a potassium ion bigger or smaller than a potassium atom? Why? Potassium ion is bigger because it has lost an electron.
35. Label the regions of the periodic table with group and period numbers as well as group names. Also label the s, p, d, and f orbital blocks.
Group 1 – alkali metals, Group 2 – alkaline earth metals, d-block – transition metals, group 17 – halogens, group 18 – noble gases, metalloids – see above. Groups 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18