SCPS Chemistry Worksheet � Periodicity A. Periodic table 1. Which are metals? Circle your answers: C, Na, F, Cs, Ba, Ni Which metal in the list above has the most metallic character? Explain. 2. Write the charge that each of the following atoms will have when it has a complete set of valence electrons forming an ion. O Na F N Ca Ar 3. What is the most common oxidation number for calcium? Explain. 4. Name two more elements with that oxidation number and explain your choice. 5. What element in period 3 is a metalloid? 6. When element with atomic number 118 is discovered, what family will it be in? 7. Make an argument for placing hydrogen in the halogen family rather than the alkali metals.
Alkali metals Alkaline earth metals
Transition metals Halogens
Noble gases 8. The _____________________________ have a single electron in the highest energy level. 9. The ___________ achieve the electron configurations of noble gases by losing two electrons. 10. The ______________________ vary in the number of electrons in the highest energy level 11. The __________ achieve the electron configuration of noble gases by gaining one electron. 12. The ___________________ have full s and p orbitals in the highest occupied energy levels. 13. The ________________________ are stable and un-reactive 14. The __________________________ are highly reactive and readily form salts with metals. 15. The _____________ are metals that are more reactive than the transition elements but less reactive than the alkali metals. 16. Predict the oxidation number based on the electron configuration shown.
1s2 2s2 2p6 3s2
1s2 2s2 2p6 3s1
1s2 2s2 2p6
1s2 2s2 2p5
1s2 2s2 2p1
SCPS Chemistry Worksheet � Periodicity - page 2 B. Ionization Energy 1. Choose the element with the greatest first ionization energy:
Carbon or aluminum
Calcium or strontium
Helium or lithium
Chlorine or argon
Chlorine or fluorine
Sulfur or chlorine
2. Which has the larger ionization energy � sodium or potassium? Why? 3. Explain the difference in first ionization energy between lithium and beryllium. 4. The first and second ionization energies of magnesium are both relatively low, but the third ionization energy requirement jumps to five times the previous level. Explain. What is the most likely ion for magnesium to become when it is ionized? 5. Compare the first ionization energies for the noble gases. 6. Compare the first ionization energies for a noble gas with that of a halogen in the same period. Support your comparison with an orbital diagram. 7. Where would the largest jump in ionization energies be for oxygen? (with the loss of how many electrons?) 8. How can you tell from a list of ionization energies for an element where a kernel (non valence) electron has been removed?
SCPS Chemistry Worksheet � Periodicity - page 3 C. Electronegativity and Electron Affinity 1. Arrange the following elements in order of increasing electronegativity. a. gallium, aluminum, indium
b. calcium, selenium, arsenic
c. oxygen, fluroine, sulfur
d. phosphorus, oxygen, germanium
2. Will the electronegativity of barium be larger or smaller than that of strontium? Explain. 3. Compare the electronegativity of tellurium to that of antimony. Explain your reasoning. 4. The family within any period with the greatest negative electron affinity is usually the ___. a. alkali metals b. transition metal c. halogens d. noble gases 5. Contrast ionization energy and electron affinity. In general, what can you say about these values for metals and non-metals? 6. What is the difference between electron affinity and electronegativity. 7. Why is it difficult to determine electron affinities for metals?
SCPS Chemistry Worksheet � Periodicity - page 4 D. Definitions - match
Atomic radius Decrease
Electron affinity electronegativity
First ionization energy Increase
Ionization energy Metals
Noble gas configuration
Noble gases Nonmetals Semimetal
Shielding effect
1. _______ _________________ is the energy required to remove an electron from an atom.
2. The attraction of an atom for an additional electron is called _______________________.
3. The energy needed to remove the most loosely held electron from a neutral atom is called
____________________.
4. When they have a(n) ___________________, ions have a stable, filled outer electron level.
5. Along with the increased distance of the outer electrons from the nucleus, the ______ _____ of the inner electrons causes ionization energy to decrease going down a column of the periodic table.
6. A low ionization energy is characteristic of a(n) ___________________________________.
7. Ionization energies tend to _______________________ across periods of the periodic table.
8. An element with a high ionization energy is classified as a (n) ________________________.
9. The attraction an atom has for shared electrons is called ____________________________.
10. The distance from the nucleus to the outer most electron is known as _________________.
11. The ______________________ do not have measured electronegativites since they do not commonly form compounds.
12. The electron arrangement with a complete outermost s and p sublevel is known as
_____________.
E. Trend Chart Draw in the trends on the periodic table: Ionization energy electronegativity atomic radius electron affinity shielding effect
SCPS Chemistry Worksheet � Periodicity - page 5 F. Atomic Radius 1. Circle the atom in each pair with the larger atomic radius?
Li or K
Ca or Ni
Ga or B
O or C
Cl or Br
Be or Ba
Si or S
Fe or Au
2. Chlorine selenium, and bromine are located near each other on the periodic table. Which of
these elements is the smallest atom and which has the highest ionization energy? 3. Which of the following atoms is smallest: nitrogen, phosphorus, or arsenic? Which of these atoms has the most negative electron affinity? 4. Which of the following is the largest: a potassium atom, a potassium ion with a charge of 1+ or a rubidium atom? 5. Which of the following is the largest: a chlorine atom, a chlorine ion with a charge of 1- or a bromine atom? 6. Which of the following is the smallest: a lithium atom, a lithium ion with a charge of 1+ or a sodium atom? 7. Use the atomic theory to explain why within a family such as the halogens, the ionic radius
increases as the atomic number increases. 8. In terms of electron configuration and shielding, why is the atomic radius of sodium smaller than that of potassium? 9. In terms of electron configurations and shielding, Why do atoms get smaller as you move across a period?
SCPS Chemistry Worksheet � Periodicity - page 6 G. Concept Mastery Questions 1. The shielding effect increases with increasing atomic number within a ___. a. period b. group c. both d. neither 2. In any ___, the number of electrons between the nucleus and the outer energy level is the same. a. period b. group c. both d. neither 3. Within a ____, the nucleus has a stronger ability to pull on the outermost (valence) electrons in elements of high atomic number. a. period b. group c. both d. neither 4. In a ____, electron affinity values become more negative as atomic number increases. a. period b. group c. both d. neither 5. The halogens are considered a ____. a. period b. group c. both d. neither 6. Which atom has the greater nuclear charge? a. Na b. Al c. P d. Ar 7. Which atom demonstrates the greatest shielding effect? a. Na b. Al c. P d. Ar 8. The atoms Na, Al, P, and Ar all have the same a. shielding b. size atomic radius c. number of valence electrons d. number of kernel electrons 9. Which element on the periodic table has a. lowest ionization energy b. highest second ionization energy c. highest electronegativity d. highest ionization energy e. largest atomic radius 10. Explain the relationship between the relative size of an ion to its atom and the charge on the
ion. 11. Explain why noble gases are inert and do not form ions. 12. Will the shielding effect be more noticeable in metals or non metals? Explain your answer.
SCPS Chemistry Worksheet � Periodicity - page 7 13. Why do elements in the same family generally have similar properties? Choose one as an example to support your reasoning. 14. Arrange each of the following in order of increasing ionization energy and explain your
reasoning: Calcium, iron, copper, bromine and krypton. 16. I am an element. I have a high electron affinity, (highly negative value), and my atomic number is X. The element with atomic number X-1 has a lower ionization energy and a lower electron affinity. The element with atomic number N+1 has a higher ionization energy and basically no electron affinity (positive value). I am toxic in my elemental state, but I am very commonly found in my nontoxic ionic form. Within my group, I have the second highest ionization energy. Who am I? (support each step of your reasoning). 17. Write a chemical equation that shows the process or events in the formation of an anion. 18. What do transition metals have in common with respect to their electron configurations?
Ar
NaIE
10 15 20 atomic number
15. Factors affecting ionization energy include nuclear charge, the shielding effect, the atomic radius and the electron arrangements in a sublevel. Use the appropriate factors to explain the overall trend indicated by the dark line and the exceptions to it
SCPS Chemistry Worksheet � Periodicity - page 8 19. Consider the table of the first four ionization energies for an element we will call A.
1st 2nd 3rd 4th Ionization energy in kJ/mol 578 1817 2745 11580 a. In which group does A appear on the periodic table? b. What is the most likely oxidation number for element A? c. What is the minimum number of electrons that A must have? d. Write the valence electron sublevel configuration for this element. (sublevel and number of
electrons in them) 20. Can anions of two different elements have the same valence electron arrangement? If so, give examples and discuss. If no, explain why not. 21. When an atom loses an electron to become an ion, what happens to its electric charge? To its size? Write a chemical equation that shows the process or events in the formation of this ion. What energy value is associated with this process?
