Periodicity Chemistry Worksheet
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.
Cs - It is the metal with the ↑ shielding & ↓ ENC
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 -2
Na +1
F -1
N -3
Ca +2
Ar 0
3. What is the most common oxidation number for calcium? Explain.
Ca+2 – Calcium has 2 valence electrons, it will lose 2 electrons
Name two more elements with that oxidation number and explain your choice.
Any of the two: Be, Mg, Se, Ba, Ra – Same valence electrons / same family
4. What element in period 3 is a metalloid? Silicon
5.
When element with atomic number 118 is discovered, what family will it be in? Noble Gas
Use the following word bank for questions 6 through 13
(Obviously, you will need to use several words more than once!  )
Alkali metals
Alkaline earth metals
Noble gases
Transition metals
Halogens
Noble gases
6. The __Alkali metals____________ have a single electron in the highest energy level.
7. The __Alkaline earth metals___ achieve the electron configurations of noble gases by losing
two electrons.
8. The __transition metals___ are metals that can hold up to 10 electrons in their sublevel
shape
9. The __halogens_____________ achieve the electron configuration of noble gases by gaining
one electron.
10. The __noble gases_________ have full s and p orbitals in the highest occupied energy
levels.
11. The __noble gases__________ are stable and un-reactive.
12. The __halogens________________ are highly reactive and readily form salts with metals.
13. The __Alkaline earth metals are metals that are more reactive than the transition elements
but less reactive than the alkali metals.
14. Predict the oxidation number based on the electron configuration shown.
1s2 2s2 2p6 3s2 _Mg +2________
1s2 2s2 2p6 3s1 ___Na +1___________
2
2
6
0
1s 2s 2p
__Ne _______
1s2 2s2 2p5
___F -1____________
2
2
1
+3
1s 2s 2p
___B ______
Periodicity Chemistry Worksheet - 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?
Na  K because Na has less shielding than K
3. Explain the difference in first ionization energy between lithium and beryllium.
Li  Be because Li has less ENC than Be
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.
Mg +2 has the noble gas configuration of [Ne]. The 3rd IE moves Mg+2 through a
noble gas configuration to Mg+3.
Spikes in IE energies will occur when moving through a STABLE noble gas
configuration.
What is the most likely ion for magnesium to become when it is ionized? Mg+2
5. Compare the first ionization energies for the noble gases.
Rn  Xe  Kr  Ar  Ne  He
6. Compare the first ionization energies for a noble gas with that of a halogen in the
same period.
Halogens  Noble gas due to filled orbital’s (especially octets) are more stable.
7. Where would the largest jump in ionization energies be for oxygen? (with the loss of
how many electrons?)
O+6 has a [He] noble gas configuration therefore the largest IE for Oxygen will
occur with the loss of the 7th electron
8. How can you tell from a list of ionization energies for an element where a kernel (non
valence) electron has been removed?
SPIKES in IE energies indicate that kernel electrons are being removed from
noble gas configurations.
Periodicity Chemistry Worksheet - page 3
C. Electronegativity and Electron Affinity
1. Arrange the following elements in order of increasing electronegativity.
a. gallium, aluminum, indium
In  Ga  Al
b. calcium, selenium, arsenic
Ca  As  Se
c. oxygen, fluorine, sulfur
SOF
d. phosphorus, oxygen, germanium
Ge  P  O
2. Will the electronegativity of barium be larger or smaller than that of strontium? Explain.
Ba  Sr
because Ba has greater shielding than Sr
3. Compare the electronegativity of tellurium to that of antimony. Explain your reasoning.
Te  Sb because Te has greater ENC than Sb
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?
IE – energy required to REMOVE electrons
Metal have LOW IE and EA
EA – energy change when ADDING electrons
Nonmetals have HIGH IE and EA*
* exception: Noble gases have the lowest EA values
6. What is the difference between electron affinity and electronegativity?
EA – energy change when ADDING electrons (backwards scale: positive to
negative
EN – the ability to attract (grab) electrons
7. If an element has a “large negative” electron affinity number where would it be located on the
periodic table?
Large negative EA number would be located in the top right (HALOGENS) section
of the periodic table exception NOBLE gases have very low positive EA values.
Periodicity Chemistry Worksheet - page 4
D. Definitions - match
Atomic radius
Decrease
Electron affinity
Electronegativity
1. _Ionization Energy
First ionization energy
Increase
Ionization energy
Metals
Noble gas configuration
Noble gases
Nonmetals
Metalloid
Shielding effect
_ is the energy required to remove an electron from an atom.
2. The energy change associated with the addition of electron is called _Electron Affinity
.
3. The energy needed to remove the most loosely held electron from a neutral atom is called
_First Ionization Energy
4. When they have a(n) _noble gas configuration, ions have a stable, filled outer electron
level.
5. Along with the increased distance of the outer electrons from the nucleus, the _Shielding
Effect__ 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) _Metals_______________________.
7. Ionization energies tend to _INCREASE_____________ across periods of the periodic table.
8. An element with a high ionization energy is classified as a (n) Nonmetals_or Noble Gases
9. The attraction an atom has for electrons is called __electronegativity_________________.
10. The distance from the nucleus to the outer most electron is known as _Atomic Radius___.
11. The _noble gases_ 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 __noble
gas configuration__.
E. Trend Chart
Ionization energy
Atomic Radius
Draw in the trends on the periodic table:
electronegativity
atomic radius
electron affinity
shielding effect
IE / EN* / EA*
Periodicity Chemistry Worksheet - 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?
Atomic Radius: Se  Br  Cl Chlorine has ↓ shielding & greatest ENC
due to Cl smallest radius it also has the highest IE
3. Which of the following atoms is smallest: nitrogen, phosphorus, or arsenic? Which of these atoms has
the most negative electron affinity?
Atomic Radius: Ar  P  N
….Nitrogen has the least amount of shielding
due to N small radius, it also has the largest negative EN value.
4. Which of the following is the largest: a potassium atom, a potassium ion with a charge of 1+ or a
rubidium atom?
Atomic Radius: K+1  K  Rb
5. Which of the following is the largest: a chlorine atom, a chlorine ion with a charge of 1- or a bromine
atom?
Atomic Radius: Cl  Cl-1  Br
Shielding predominates (rules)
6. Which of the following is the smallest: a lithium atom, a lithium ion with a charge of 1+ or a sodium
atom?
Atomic Radius Na  Li  Li
+1
7. Explain why within a family such as the halogens, the ionic radius increases as the atomic number
increases.
Ionic radius increases within a family/group due to the INCREASE shielding (distance).
8. In terms of electron configuration and shielding, why is the atomic radius of sodium smaller than that
of potassium?
Na & K have the same VALENCE ELECTRON configuration (“ns1”),
however K has greater shielding than Na
9.
In terms of electron configurations and shielding, why do atoms get smaller as you move across a
period?
Shielding remains fairly constant across the period,
however the ENC increases due to increase # of valence electrons in the electron configuration.
Periodicity Chemistry Worksheet - 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. number of valence electrons
9. Which element on the periodic table has
a. lowest ionization energy
b. size atomic radius
c. number of kernel electrons
__Fr ____
b. highest second ionization energy __Li +1____
c. highest electronegativity
__F
____
d. highest ionization energy
__He
e. largest atomic radius
__Fr ____
___
10. Explain the relationship between the relative size of an cation and anion (ionic radius) to its atom
(atomic radius).
Cation (positively charged ion)  Atomic Radius  Anion (negatively charged ion)
11. Explain why noble gases are inert and do not form ions.
Noble gases have an OCTET (FILLED s & p orbitals)
12. Will the shielding effect be more noticeable in metals or non metals? Explain your answer.
Nonmetals have more valence electrons (partial shield) and ↑ ENC than metals.
Periodicity Chemistry Worksheet - page 7
13. Why do elements in the same family generally have similar properties? Choose one as an example to
support your reasoning.
Families/groups have similar properties because they have the same # of valence electrons.
Alkali metals all have ONE valence electron “ns1” and react with water.
14. Arrange each of the following in order of increasing ionization energy and explain your reasoning:
Calcium, iron, copper, bromine and krypton.
Ca  Fe  Cu  Br  Kr
ENC increases across a period creating a stronger magnet
which required increasing IE to pull electrons off.
15. Factors affecting ionization energy include effective nuclear charge, the shielding effect, the atomic
radius and the electron arrangement in a sublevel. Use the appropriate factors to explain the overall
trend indicated by the dark line and the exceptions to it.
1) Effective Nuclear energy: Increase, Decrease or Constant
2) Shielding: Increase, Decrease or Constant
3) Atomic Radius: Increase, Decrease or Constant
4) Explain exceptions to the overall trend based
on electron configuration.
_filled and ½ filled orbital are more stable
due to an equal distribution of electrons along all axis of orientation.
16. What element am I?. (Brief periodic table location description for each clue)
Clue #1)
I have a high electron affinity, (highly negative value), and my atomic number is X.
HALOGEN FAMILY
Clue #2)
The element with atomic number X-1 has a lower ionization energy and a lower electron affinity.
OXYGEN FAMILY
Clue #3)
The element with atomic number X+1 has a higher ionization energy and basically no electron
affinity (positive value).
NOBLE GAS FAMILY
d) Within my group, I have the second highest ionization energy.
What element am I? CHLORINE
17. What do transition metals have in common with respect to their electron configurations?
Transition metal all have FILLED “s” blocks – ns2nd1-10
Periodicity Chemistry Worksheet - page 8
18. Consider the table of the first four ionization energies for an element we will call A.
1st
576
Ionization Energy
(kJ/mole)
2nd
1817
a. In which group does A appear on the periodic table?
3rd
2745
4th
11580
GROUP 13: Boron family
b. What is the most likely oxidation number for element A? +3
c. What is the minimum number of electrons that A must have?
FIVE electrons; Boron is the smallest element in the group with FIVE electrons.
d. Write the valence electron sublevel configuration for this element.
(Hint: use “n” as the energy level)
ns2np1
n = row number, energy level
19. Can anions of two different elements have the same valence electron arrangement? If so, give
examples and discuss. If no, explain why not.
Anions of two different elements CAN have the same valence electron arrangement,
if they are from the SAME GROUP/FAMILY
Example Sulfur and Oxygen both have an “ns2np4 ” valence electron arrangement.
20. The first ionization energy of beryllium is 9.322 eV, the second ionization energy is 18.211 eV, and
the third ionization energy is 153.893 eV. Explain why the third ionization energy of beryllium so
much higher than the first two.
The 3rd IE will higher due to Be+2 has the [He] noble gas configuration.
Removing KERNEL electrons from a noble gas configuration requires a substantial
amount of energy.