chapter 7 periodic properties of the elements

Question 1

Which of the following has the largest atomic radius?
a. Na
b. Ca
c. Fr
d. Br
e. Bi

Answer 1

c. Fr
Atoms get larger down and to the left on the periodic table. Fr is in this corner.
‘You go to France, you get fat’

Question 2

Which of the following has the largest atomic radius?
a. Al
b. Si
c. P
d. S
e. Cl

Answer 2

a. Al

Question 3

Which of the following is the longest bond?
a. S-F
b. O-F
c. S-Cl
d. S-Br
e. S-I

Answer 3

e. S-I
The O-F bond is obviously smaller than the S-F bond so that one is out. The rest just go down the same column so Iodine is clearly the largest of these.

Question 4

Which of the following has the smallest ionic radius?
a. Li+
b. Na+
c. K+
d. Rb+
e. Cs+

Answer 4

a. Li+
This seems like a trick question because it’s so intuitive but it’s not.
Li has the smallest atomic radius by far and this will hold true as each of the possible answers loses only one electron

Question 5

Which of the following has an ionic radius that is smaller than its atomic radius?
a. Al
b. P
c. S
d. Cl
e. Br

Answer 5

a. Al
Despite being on the right-hand side of the table, Al has a +3 ionic charge. The rest of the possible answers have negative charges meaning they gain electrons causing the radius to be wider as these new electrons repel against the already present electrons.

Question 6

Which of the following has an ionic radius that is larger than its atomic radius?
a. Li
b. Na
c. Mg
d. Al
e. P

Answer 6

e. P
P is going to take on an electron making it larger all around. This electron will also push against the electrons already present.

Question 7

Which of the following has the smallest radius?
a. Se^2-
b. Br-
c. Kr
d. Rb+
e. Sr^2+

Answer 7

e. Sr2+
So basically, all of the answer choices have the same amount of electrons. But, since Sr has the greatest amount of protons, this will pull the most on those electrons causing this atom to be the smallest.

Question 8

Which of the following has the largest radius?
a. Te^2-
b. I-
c. Xe
d. Cs+
e. Ba^2+

Answer 8

a.
These all have the same amount of electrons. Te though, having the fewest amount of protons, has less pulling on it from the center. Less pull = larger radius.

Question 9

Which of the following has the highest ionization energy?
a. Na
b. Mg
c. Al
d. Si
e. P

Answer 9

e. P
Ionization energy trends the opposite of radii.
Ionization energy is the energy needed to be absorbed to discharge an electron.

Question 10

Which of the following explains why He has the highest ionization energy of all the noble gases?
a. Because it is the smallest.
b. Because its valence electrons experience the greatest effective nuclear charge.
c. Because its core electrons experience the greatest effective nuclear charge.
d. Because its valence electrons experience the smallest effective nuclear charge.
e. Because it violates the octet rule.

Answer 10

a.
These answers are idiotic but there is some takeaway.
1. He’s electrons are not considered ‘core’ electrons because there are no other electrons beyond them. They are, however, considered ‘valence’
2. The 2nd point its trying to make is… I’ll just copy it straight from the test solution comments…
Valence electrons of atoms within a group have very similar effective nuclear charges so it is better practice to compare the atomic size.
3. All that being said, I still don’t think that b. is incorrect

Question 11

Which of following has the highest ionization energy?
(careful: this is kind of a trick question)
a. Cs
b. Mg
c. Al
d. N
e. O

Answer 11

d. N
There are some exceptions to the general trend which will be relevant here. The element farthest up and to the right across the periodic table is oxygen, but nitrogen actually has a higher ionization energy than oxygen, one of the exceptions to the trend you’re supposed to remember.
This is why:
Nitrogen has a higher ionization energy than oxygen because the half-filled 2p orbital in nitrogen is more stable than the partially filled 2p orbital in oxygen, making it harder to remove an electron from nitrogen

Question 12

Which of the following has the highest ionization energy?
a. B
b. Ba
c. Be
d. Rb
e. Mg

Answer 12

c. Be
There are some exceptions to the general trend which will be relevant here. The element farthest up and to the right across the periodic table is oxygen, but nitrogen actually has a higher ionization energy than oxygen, one of the exceptions to the trend you’re supposed to remember.
Be has a nice 1s2 and 2s2 filled without any extra electrons just hanging out anywhere. It’s more solid than B because of this.

Question 13

Which of the following shows the equation for the 3rd ionization energy for Sc?
a. Sc3+ + e- → Sc2+
b. Sc3+ + e- → Sc2+
c. Sc2+ → Sc3+ + e-
d. Sc3+ → Sc2+ + e-
e. Sc → Sc3+ + 3e-

Answer 13

c.
This is easiest for me if I just write it out first without looking at the answer choices. Start with just Sc
Sc … Sc^1+ this would be the 1st energy
Sc^1+…Sc^2+ this would be the 2nd energy
Sc^2+…Sc^3+ this would be the 3rd energy
The only answer that starts with Sc2+ is c.

Question 14

Which of the following has the highest 2nd ionization energy?
a. Li
b. Be
c. B
d. C
e. N

Answer 14

a. Li
The 2nd electron removed from Li would be a core electron and it takes a ton of energy to remove core electrons. So even though it isn’t the right-most on the table, it would still take the most energy to remove this electron.

Question 15

Which of the following shows the equation for the electorn affinity for sulfur?
a. S- + e- → S
b. S- → S + e-
c. S → S- + e-
d. S + 2e- → S2-
e. S + e- → S-

Answer 15

e.
The question doesn’t really specifiy so I assumed it wants the equation for gaining just one electron.

Question 16

Which of the following has the highest (most negative) electron affinity?
a. C
b. N
c. O
d. F
e. Ne

Answer 16

d. F
Noble gases have a positive electron affinity since it would REQUIRE energy to stick an electron to a noble gas (since they don’t really want one).
The electron affinity trend (other than noble gases) is up and to the right.
Also, it would seem the F would be more negative than Cl but this isn’t true because F is smaller and the incoming electron would be closer to the nucleus which would try to repel it.

Question 17

Which of the following has the highest (most negative) electron affinity?
a. Na
b. Mg
c. S
d. P
e. Cl

Answer 17

e. Cl
Just remember that Cl has the most negative electron affinity than any other element. It is easier for an electron to move to a Cl than anything else.

Question 18

Which of the following is the most electronegative?
a. N
b. O
c. F
d. Cl
e. Br

Answer 18

c. F
Electronegativity, finally, is the one where I don’t think there are any exceptions to the rule (other than to discount noble gases). It’s just up and to the right so Fluorine is king.

Question 19

Which of the following is the most electronegative?
a. Sb
b. Cl
c. Te
d. Br
e. I

Answer 19

b.
Up and to the right. No exceptions with electronegativity except no noble gases

Question 20

Which of the following is the most polar bond?
a. P-I
b. P-S
c. P-Cl
d. P-F
e. P-Br

Answer 20

d. P-F
Although not the farthest apart physically on the periodic table, the greatest difference in electronegativity is between these two elements making it the most polar.

Question 21

Which of the following would combine with an alkaline earth metal in a 1:1 ratio and an alkali metal in a 1:2 ratio?
a. helium
b. sulfur
c. carbon
d. hydrogen
e. nitrogen

Answer 21

b. sulfur
sulfur want to gain 2 electrons
alkaline earth metals want to lose 2 electrons and alkali metals want to lose 1

Question 22

Which of the following would most likely by expected to be brightly colored?
a. N2
b. CO2
c. NaHCO3
d. Na2Cr2O7
e. KCl

Answer 22

d.
This one is a little tricky. If you know that compounds containing transition metals are brightly colored, then d. is the only answer containing a transition metal.

Question 23

Which of the following pairs of elements would be expected to have the most similar chemical properties?
a. Mg and Cs
b. Na and I
c. F and Cl
d. N and O
e. Sc and Zn

Answer 23

c. F and Cl
elements in the same column (group) tend to have similar properties

Question 24

Which of the following is most similar in reactivity to chlorine?
a. sodium
b. oxygen
c. bromine
d. sulfur
e. calcium

Answer 24

c. Br
Br is in the same column (group) and has the same amount of valence electrons

Question 25

Which of the following would be expected to be most reactive with water? a. Cl b. He c. Rb d. Be e. Sr

Answer 25

c. Rb alkali metals are the most reactive with water (due to their low ionization energy) followed by alkaline earth metals

Question 26

Which of the following is most reactive with water? a. aluminum b. oxygen c. copper d. potassium e. calcium

Answer 26

d. potassium The alkali metals are highly reactive with water (they are oxidized while water is reduced) with the alkaline earth metals being next most reactive. The only alkali metal listed is potassium and it is therefore the most reactive with water.

Question 27

Which of the following would react most violently with water? a. lithium b. sodium c. potassium d. rubidium e. cesium

Answer 27

e. The alkali metals are highly reactive with water (they are oxidized while water is reduced) often accompanied by the production of a flame and/or an explosive reaction. They tend to react more violently as you go down the group as the atomic radius increases and ionization energy decreases. Therefore cesium will react the most violently with water.

Question 28

Which of the following is most likely to be chemically inert? a. lithium b. fluorine c. argon d. rubidium e. strontium

Answer 28

c. The noble gases are well known for being chemically inert and therefore argon is the correct answer. The larger noble gases are known to participate in a couple of reactions but it definitely isn’t common.

Question 29

Atomic radium generally increases ________ a group and to the _______ across a period. a. up, left b. down, right c. up, right d. down, left e. There is no clear trent in atomic radius

Answer 29

d.

Question 30

The trend in atomic radii across a period can be 'best' explained by which of the following? hint: Zeff = z - s z = # of protons s = # of shielding electrons shielding electrons = any electron that's not a valence electron a. electron repulsion b. effective nuclear charge c. principal quantum number of valence orbitals d. shielding effect e. nuclear mass

Answer 30

b. effective nuclear charge I really don't like this question for many reasons. The wording itself is terrible. 'Across' a period to me can mean moving left OR right. But here, we're talking about just left to right (the radii get smaller). Also, a lot of these answers to hand-in-hand or one causes another so... The book's official answer is: Atomic radius decreases moving across the periodic table as a result of increasing effective nuclear charge as proton number increases. Increasing effective nuclear charge attracts electron orbitals more strongly which causes a general decrease in atomic radius.

Question 31

Which of the following has the smallest atomic radius? a. Sb b. Se c. Br d. I e. Bi

Answer 31

c. Br

Question 32

Which of the following species has the largest radius? a. Al^3+ b. K+ c. Cl- d. S^2- e. Ca^2+

Answer 32

d. S^2+ The positive ions will have small radii from getting rid of electrons. Between Cl and S, Cl has the greater Zeff which means greater nuclear pull which means smaller so... It says they're all part of an isometric series. My logic is what's stated above and I think it holds water. The book's explanation just says 'the most negative' is the largest.

Question 33

Which of the following correctly describes ionization energy? a. the degree of negative charge on an ion b. the number of valence electrons in an atom c. the tendency of an atom to attract electron density d. the total number of electrons lost of gained by an atom e. the amount of energy required to remove an electron

Answer 33

e. the amount of energy required to remove an electron (it will be the most loosely bound, or outermost, electron)

Question 34

Going down a group in the periodic table, ionization energy ___________ and electronegativity __________ . a. increases, increases b. increases, decreases c. decreases, increases d. decreases, decreases e. remains constant, increases

Answer 34

d. Ionization energy is the energy required to remove an electron. Let's take halogens for example. Cl wants to gain electrons not lose them so it will already have a high ionization energy. It will be easier to pluck an electron from this same column is the atom is bigger and the valence electron is farther away from the pull of the nucleus. Electronegativity is the atom's ability to attract an electron. Cl can do this easily because it wants electrons but also because its a smallish atom and the positive nucleus will help pull electrons towards it. Therefore, it has a higher electronegativity compared to Iodine for example where the large size of the atom and the distance from the nucleus hinders the attraction the its ability to easily attract an electron. So the electronegativity decreases as the go down and the atoms get larger.

Question 35

Of the following elements, which has the highest first ionization energy? a. He b. H c. F d. Li e. Fr

Answer 35

a. He While hydrogen does have a pretty high ionization energy, He's is higher. It doesn't say why but I can only suppose that it's the attraction of the 2 proton nucleus versus hydrogen having only one.

Question 36

Which of the following has the highest first ionization energy? a. F b. K c. O d. Ba e. Ne

Answer 36

e. Ne It is hard to pry an electron away from a noble gas

Question 37

Which of the following shows the 2nd ionization energy of Fe? a. Fe+ + e- = Fe b. Fe2+ + 2e- = Fe c. Fe2+ = Fe3+ + e- d. Fe+ = Fe2+ + e- e. Fe = Fe2+

Answer 37

d. You're going to start out with a + ion because one ion has already been removed. The a. option leaves you right back with the original element which wouldn't make any sense.

Question 38

Which of the following is NOT an exception to the general trend for ionization energy? a. N b. Be c. P d. Zn e. Kr

Answer 38

e. Kr

Question 39

Which of the following correctly defines electron affinity of an atom? a. The energy released when a valence electron is removed from an atom or ion b. The energy required to remove a valence electron from an atom or ion c. The energy released when an electron is added to a neutral atom in the gaseous phase d. The energy released when an electron atom is added to a metal atom in the solid phase e. The energy absorbed by an electron while moving to a higher energy level

Answer 39

c. electron

Question 40

Which of the following groups in the periodic table have the highest (most negative) electron affinities? a. noble gases b. alkali metals c. halogens d. transition metals e. alkaline earth metals

Answer 40

c. halogens

Question 41

Which of the following has the lowest (least negative) electron affinity? a. S b. O c. C d. N e. B

Answer 41

d. N nitrogen has 7 electrons. the 2p shell has 3 electrons in it which makes it perfectly halfway full. because of this stability, nitrogen doesn't necessarily want to gain any more electrons than it already has. think of all then nitrogen in the air and how it's probably good that it's not very reactive.

Question 42

Which of the following has the lowest electronegativity? a. Si b. Cl c. S d. Rb e. Cu

Answer 42

d. Rb it's a big atom that wants to lose that one valence electron anyway

Question 43

sodium is classified as a...? a. noble gas b. alkali metal c. halogen d. transition metal e. alkaline earth metal

Answer 43

b. alkali metal

Question 44

krypton is classified as a...? a. noble gas b. alkali metal c. halogen d. transition metal e. alkaline earth metal

Answer 44

a. noble gas

Question 45

Which of the following forms an acidic oxide? a. nitrogen b. magnesium c. lithium d. calcium e. iron

Answer 45

a. nitrogen I guessed Fe and I was wrong. But there is a nugget of information in this answer that's definitely worth remembering. "Non-metals tend to form acidic oxides while metals form basic oxides"