Honors Chemistry
Name ____________________________
Chapter 5: Periodic trends Review Worksheet
Date _____/_____/_____
Period _____
1. Refer to a table of atomic and ionic radii, if necessary.
Decreases
a. Within a period, does the size of the atom increase or decrease with
increasing atomic number?
Increases
b. Within a family, does the size of the atom increase or decrease with
increasing atomic number?
Smaller
c. When metallic atoms lose electrons do they form smaller or larger ions?
Larger
d. When nonmetallic atoms gain electrons do they form smaller or larger
ions?
2. From each of the following pairs, circle the larger particle.
a. Na
Li
c. F
F-1
e. K
K+1
b. Br
I
d. Cs
Ba
f. Ne
Ar
3. Predict the oxidation number (charge), given the electron configuration for the following neutral atoms.
+2
A. 1s22s22p63s2
–1
D. 1s22s22p5
+1
B. 1s22s22p63s1
+3
E. 1s22s22p1
0
C. 1s22s22p6
4. Write the correct empirical formula for each of the following combinations of elements from question 3.
(If no reaction, write none.) After writing the empirical formula, circle the element that has the higher
electron negativity.
BD
a. B and D
AD2
c. A and D
ED3
b. E and D
None
d. A and C
5. Complete each of the following statements.
Electron Affinity
a. The __________ of an element is determined experimentally by
bombardment of atoms in the vapor phase with a stream of electrons.
First Ionization Energy b. The energy necessary to completely remove the most loosely held
electron from an atom is called the ________ ________ ________.
Metal
c. A low ionization energy is characteristic of a __________.
Increase
d. As a trend on the periodic table, ionization energies __________
across rows or periods.
Decrease
e. As a trend on the periodic table, ionization energies __________
down columns or groups.
Nonmetal
f. An element with a high ionization energy is classified as a _______.
Effective Nuclear Charge
Metal
High
Electronegativity
g. The attraction an atom has for its outer electrons is called ______.
h. An element having a low electron affinity is classified as a ______.
i. Nonmetals have __________ electron affinities.
j. As the number of valence electrons increase, the __________
increases and the pull on the bonded electrons in a chemical bond
increases.
6. Factors affecting ionization energy include nuclear charge, the shielding effect, the atomic radius, and
the electron arrangement in a sublevel. Use the appropriate factors to explain the peaks and valleys
(the circled points) in the graph between lithium and neon.
Helium and Neon have a high first IE because the atoms have a small atomic radius (due to the fact
that they have a strong nuclear charge – attraction from the protons to electrons), small amount of
shielding (there is a small number of inner electrons to shield the valence electrons from the
attraction of the protons), and a full octet (full s and p sublevels – which gives the atoms greater
stability).
Beryllium has a slightly elevated first IE because the atom has a small atomic radius, strong nuclear
charge, small amount of shielding, and a full sublevel (2s2 – higher stability in a filled sublevel which
means that it will take more energy to remove an electron).
Nitrogen has a slightly elevated first IE because the atom has a small atomic radius, strong nuclear
charge, small amount of shielding, and a half filled sublevel (2p3 – higher stability in a half filled
sublevel which means that it will take more energy to remove an electron).
Even though Lithium has a half filled sublevel (2s1), its first IE is lower than would be expected. The
lower IE is due to the fact that if it loses one electron it will end up with a full duet.
Boron also has a lower first IE than would be expected because it has an ending electron
configuration of 2s22p1. Having that one electron in the p sublevel gives it instability (like adding
weight to one blade of a fan). This instability is what leads Boron to having a lower first IE than
would be expected.
Oxygen also has a lower first IE than would be expected because it has an ending electron
configuration of 2s22p4. Having four electrons in the p sublevel also gives it instability (like adding
weight to one blade of a fan). This instability is what leads Oxygen to having a lower first IE than
would be expected.
7. Underline the atom in each of the following pairs that has the lower first ionization energy.
a. Li
Na
c. Cs
Ba
e. F
Ne
b. Kr
Rb
d. Cl
Br
f. S
Cl
Honors Chemistry
Name
Chapter 5: Periodic Trends Review WS
Date _____/_____/_____
Period _____
1. What is atomic radius? What are the trends for atomic radius?
AR is half the distance between two identically bonded atoms or, in other words,
the distance from the nucleus of an atom to the outer regions of the electron cloud.
AR increases down each group (because of additional energy levels and shielding)
and decreases, from left to right, across the period (because the larger protons have
a greater attraction for the electron cloud and pull it in closer to the nucleus).
2. List the following atoms in order of increasing atomic radius: N, Au, Al
N, Al, Au
3. List the following atoms in order of decreasing atomic radius: Cl, K, Cu
K, Cu, Cl
4. What is the shielding effect?
SE is the reduction of the nuclear attraction that the valence electrons have for the
protons in the nucleus. This reduction is due to the inner electrons blocking out or
getting between the valence electrons and the positive charge in the nucleus.
5. How are shielding effect and the size of the atomic radius related?
As the number of energy levels are filled, the inner electrons block out or shield
the attraction the valence electrons have for the protons in the nucleus. That
reduction in attraction causes the atomic radius to get larger.
6. What is ionic radius? What are the trends for ionic radius?
IR is the change in atomic radius due to the loss or gain of electrons. For each
element, its the radius of the most commonly found ion of the element.
The trend for going across the PT, from left to right, is that the IR radius decreases
across the metals and then increases greatly at the nonmetals and then continues
to decrease.
The trend for going down the PT is that the IR radius increases because of the atomic
radius trend.
7. How are neutral atoms converted into cations? How are neutral atoms converted in anions?
Atoms that have 4 or less valence electrons tend to lose electrons to obtain a full
octet. It is the loss of these valence electrons that allows the atom to gain a full
octet and gain greater stability. It is the loss of an electron or electrons that forms
a cation.
Atoms that have 4 or more valence electrons tend to gain electrons to obtain a full
octet. It is the gaining of electrons that allows the atom to gain a full octet and
gain greater stability. It is the gaining of an electron or electrons forms an anion.
8. Metals usually form what type of ions? Nonmetals usually form what type of ions?
Metals form cations and nonmetals usually form anions.
9. When an atom becomes an anion, what happens to its radius?
When an atom forms an anion the radius increases due to the additional electron
being added to the atom. This means that there are more electrons sharing the
attraction for the positively charged protons in the nucleus (less attraction – larger
radius).
10. When an atom becomes a cation, what happens to its radius?
When an atom forms a cation the radius decreases due to the loss of an electron
or electrons. This means that there are less electrons sharing the attraction for the
positively charged protons in the nucleus (more attraction – smaller radius).
11. For each of the following pairs, circle the atom or ion that would have the larger radius.
a. S or O
c. Na+1 or K+1
e. S-2 or O-2
b. Ca or Ca+2
d. Na or K
f. F or F-1
12. For each of the following pairs, identify the smaller ion.
a. K+1 or Ca+2
c. C+4 or C-4
e. O-2 or F-1
b. F-1 or Cl-1
d. S-2 or F-1
f. Fe+2 or Fe+3
13. What is ionization energy? What are the trends for ionization energy? What are the
exceptions for ionization energy? What are the units for ionization energy?
The energy needed to remove a valence electron from a neutral atom.
IE increases as you go across the PT, from left to right, and decreases as you go
down a group on the PT.
The exceptions for IE occur at the full and half full sublevels. The IE increases at
these points because the atoms have greater stability.
The units for IE are kJ / mole and have a positive sign. The positive sign indicates
that the energy is being absorbed by the atom.
14. What is the meaning behind the first, second, and third ionization energies for a particular
atom?
The first IE is the energy needed to remove one electron from a neutral atom.
The second IE is the energy needed to remove an electron from a cation that has
a positive one charge.
The third IE is the energy needed to remove an electron from a cation that has a
positive two charge.
15. Why does each successive ionization require more energy than the one before?
As an electron is lost by an atom the remaining electrons are pulled in closer to the
nucleus. The fact that the remaining electrons are pulled in tighter to the nucleus
means that it is going to take more energy to remove an additional electron. This
means that the second and third ionizations energies are going to be higher than the
first.
16. Which of these elements has the highest first ionization energy: Sn, As, or S?
S has a slightly higher first ionization energy than As, and a much larger first
ionization energy than Sn because it has a smaller atomic radius which makes it
more difficult for an electron to be removed.
17. What is electronegativity? What are the trends for electronegativity?
Electronegativity is the attraction that an atom in a chemical bond has for the
bonding electrons. The atom in the bond that has a higher EN has a greater
attraction for the bonded electrons and pulls the bonded electrons closer to itself.
EN increases as you from left to right across the PT and decreases as you go
Down a group on the PT.
18. List the following atoms in order of increasing electronegativity: O, Al, Ca
Ca, Al, O
19. List the following atoms in order of decreasing electronegativity : Cl, K, Cu
Cl, Cu, K
20. What is electron affinity? What are the trends for electron affinity? What are the
exceptions to the trends for electron affinity? What are the units for electron affinity?
EA is the energy change that occurs when a neutral atom gains an electron. The
trend for EA is based upon the loss of energy when an electron is gained by a
neutral atom.
The trends for EA are that it increases as you go from left to right across the PT and
decreases as you go down a group on the PT.
The exceptions for EA occur when the atom as a full sublevel and half filled sublevel
(these atoms have a slightly greater stability and do to this stability do not want to
gain an electron as much as the atoms before and after it in the trends).
EA has the units of kJ/ mole and has a negative sign. The negative sign indicates
that the energy is being released by the atom
21. List the following atoms in order of increasing electron affinity: F, S, Cl
S, F, Cl
22. List the following in order of decreasing electron affinity: Mg, K, Ca
K, Mg and Ca
23. Why does chlorine have a higher electron affinity than fluorine?
Cl has a higher EA due to the fact that its radius is a little larger than F and the
electrons are spread out a little farther. Since F electrons are packed so tightly it is
harder to fit an additional electron in the atom.
24. What is melting point? What are the trends for melting point?
The temperature at which a substance goes from a solid state to a liquid state.
MP decreases as you go across the periodic table and increases as you go down
a group on the PT.
25. What is thermal conductivity? What are the trends for thermal conductivity?
TC is how easily heat energy passes through a substance.
TC decreases as you go across the periodic table and increases as you go down
a group on the PT.
26. What is metallic character? What are the trends for metallic character?
MC are the chemical and physical properties that are used to characterize an
element as a metal. For example:
chemical properties – metals lose electrons to gain stability and a full octet
physical properties – ductile, malleable, conductor of heat and electricity, and
lustrous
27. What are the trends for the number of valence electrons?
The number of valence electrons increases as you go from left to right across the
PT (excluding the transition metals).
The number of valence electrons stay the same as you go down a group.
28. What is an isoelectronic series? Give an example.
An isoelectronic series is a group of ions that has the same number of electrons
and the same electron configuration.
C-4, N-3, O-2, F-1, Na+1, Mg+2, Al+3, Si+4
All of these have the same number of electrons as Neon (10 electrons).
29. What is effective nuclear charge?
ENC is the attraction that the outer electron (valence electrons) have for the
protons in the nucleus. It is the nuclear charge (the total attraction that the
valence electrons should have) minus the shielding effect (blocking of some of the
attraction) that occurs.
30. How would the difference in the above properties / trends affect the reactivity of a metal
versus a nonmetal? Include how the number of energy levels and the number of valence
electrons would affect reactivity.
As you go down a group on the PT the AR increases due to the electrons being
distributed in the different energy levels. An increase in AR means that there is a
larger number of inner electrons that can block out part of the attraction the valence
electrons have for the protons in the nucleus.
This would affect the reactivity of metals and nonmetals differently.
Metals would be more reactive with an increase in AR and shielding because they
lose electrons when they react.
Nonmetals would be less reactive as the AR and shielding increases because
nonmetals gain electrons when they react. The increase in AR and shielding would
reduce the attraction that nonmetals have for gaining electrons.