Periodic Trends Homework
1. Rank the following elements from smallest to largest: N, Sb, Ar. N, Ar, Sb
2. Given any two elements within a group, is the element with the larger atomic number likely to have a larger or smaller
radius? larger, the larger the atomic number in a group means more core electrons which increase the number of
energy levels which causes the electron shielding to increase and therefore less attraction between the nucleus and
the valence electrons resulting in a larger atomic radius
3. Explain why is it harder to remove an inner shell electron than a valence electron from an atom. An inner shell
electron will be closer to the nucleus. The greater attraction from being closer to the nucleus means that more energy
will be required to remove the electron.
4. An element forms a negative ion when ionized. On what side of the periodic table is the element located? right
(nonmetals form anions)
5. Of the elements magnesium, calcium, and barium, which forms the ion with the largest radius? The smallest? What
periodic trend explains this? Barium will have the largest radius while magnesium will have the smallest. These
elements are within a group. As you proceed down a group each element gains an energy level. Having additional
energy levels means more core electrons and more electron shielding, making the valence electrons farther from the
nucleus.
6. Define an ion: an ion is a charged particle; it becomes charged when electrons are gained (anion, negative) or lost
(cation, positive)
7. How does the ionic radius of a nonmetal compare with its atomic radius? Explain why the change in radius occurs.
Nonmetals form anions, meaning they gain electrons. Since the number of protons remains constant but the number
of electrons increases, the increase in electron repulsion causes the ion to have a larger radius.
8. Explain why atomic radii decrease as you move left-to-right across a period. The number of energy levels and
electron shielding remain the same but number of protons increases, increasing the effective nuclear charge. Since
the valence electrons feel a greater attraction to the nucleus, the radius decreases.
9. Arrange the elements oxygen, sulfur, tellurium, and selenium in order of increasing atomic radii. Is your order an
example of a group trend or a period trend? O, S, Te, group trend
10. Arrange the elements in order of increasing atomic radii and explain: Mg, N, P, Sr.
N, P, Mg, Sr
Nitrogen will be the smallest since it is the only atom with just two energy levels. Phosphorus and magnesium both have
three energy levels, but phosphorus has more protons and a greater effective nuclear charge, so it will be smaller
than magnesium. Strontium will be the largest with five energy levels.
11. Which has a higher ionization energy: K or Ca? Explain. Calcium. As you go from left to right across the periodic
table the effective nuclear charge increases (because there are more protons), but the core electrons and
shielding remain the same. Calcium’s valence electrons are pulled in closer to the nucleus than Potassium’s
valence electrons. It requires more energy to remove an electron that is closer to the nucleus.
12. Which has a higher ionization energy: P or S? Explain. Phosphorous. Even though sulfur is a smaller atom, it is
easier to remove an electron from phosphorous because the electron being removed is the P 4 electron (the first
electron that is paired in the P sublevel) and removing this electron alleviates repulsion due to pairing.
1)
Why does atomic radius decrease as you move across the periodic table?
As you move from left to right on the periodic table the number energy
levels (shells) and electron shielding stays the same and the effective
nuclear charge increases (because the # protons increase) which pulls
the valence electrons closer to the nucleus, resulting in a smaller
radius.
2)
How are the trends of atomic radius and ionization energy related? Explain in both groups and periods.
These trends are inverses of each other. As the size of the atom
increases, the ionization energy (energy required to remove the
electron) decreases because the electron being removed is farther from
the nucleus in a larger atom since there are more energy levels. Less
energy is required to remove an electron further from the nucleus. [AR
increases down a group while IE decreases; AR decreases across a
period while IE increases.]
3)
What’s the difference between electronegativity and ionization energy?
Electronegativity is related to the pull of a nucleus on an electron being
shared in a bond. Ionization energy is the energy required to remove an
electron from an unbonded atom. Electronegativity is a ranking and has
no units. Ionization energy is an amount of energy and has units such
as kilojoules.
4)
What does it mean to be isoelectronic? Identify the largest ion and smallest ion of the group: Se 2-, Br1-, Kr, Rb1+,
Sr2+. Explain why.
Isoelectronic means a group of atoms/ions having the same number of
electrons. All of the atoms in this group have 36 electrons; each
atom/ion has the same number of Energy levels (core electrons) and
the same amount of electron shielding. What they don’t have in
common (and what accounts for the difference in size) is the number of
protons. The greater the number of protons the greater the effective
nuclear charge (ENC). The atom/ion with the greater ENC will be
smallest (Sr2+) while the atom/ion with the lowest ENC will be largest
(Se2-).
5)
What does the octet rule have to do with the periodic trends?
A trend like ionization energy is very closely related to the octet rule.
The halogens are the smallest of atoms and the closest to the having a
full octet. Since gaining an electron would give these atoms a full octet,
the amount of energy required to remove an electron (IE) would be very
high. When determining the charge an atom will have (as in number 4),
the number of electrons gained or lost is the number required to have a
full octet of electrons.
6)
SKIP