Periodic Trends

Question 1

How is the periodic table blocks organized?

Answer 1

The blocks are named after the last electron that fills each respective atom in that section of the periodic chart.

Question 2

An elements chemical properties depend on its _____/

Answer 2

Valence electrons.

Question 3

Describe the trend on the periodic table, in relation to the effective nuclear charge.

Answer 3

As you move from left to right across a period in the periodic table, the effective nuclear charge increases.

Question 4

As you _____ (ascend/descend) a family in the periodic table, the valence shell increases, so the distance between the atom’s valence e- and its nucleus _______ (increases/decreases).

Answer 4

As you descend a family in the periodic table, the valence shell increases, so the distance between the atom’s valence e- and its nucleus in increases.

Question 5

Define the effective nuclear charge (Zeff).

Answer 5

The net charge exerted upon the outermost e- (valence e-). It can be approximated by the equation: Zeff = Z - S, where Z is the atomic number and S is the number of shielding (core) e-.

Question 6

How does valence shell affect periodic trends?

Answer 6

The valence shell affects the distance between between electrons and the nucleus.

Question 7

What’s the periodic trend for atomic size?

Answer 7

The atomic size decreases ( the radius of the atom is defined by the distance from the center of nucleus to the exterior of the valence e- cloud).

Question 8

What’s the periodic trend for ionization energy?

Answer 8

The IA increases ( the energy required to remove the outermost e- from the atom in its gas phase, creating a cation).

Question 9

What’s the periodic trend for electron affinity?

Answer 9

The EA increases ( the energetics associated with an atom gaining an e- in its gas phase).

Question 10

What’s the periodic trend for electronegativity.

Answer 10

The EN increases ( tendency to hold shared e- with another within a BOND).

Question 11

When doing MCAT questions on periodic trends, how should you handle F?

Answer 11

F tends to be on extremes.

Question 12

What can cause inconsistency in the periodic trends we typically see?

Answer 12

Half-filled and filled-shell stability. For example, N has a larger IA than O because upon ionization, N loses its half-filled p-shell. O gains half filled stability upon being ionized.

Question 13

As a general rule, cations are _____ ( larger/smaller) than neutral atoms.

Answer 13

Smaller, because the loss of electrons allow the atom to compact more tightly.

Question 14

Within a period, anions are _____( larger/smaller) than cations.

Answer 14

Larger. Cl- ion is larger than the Na+ cation.

Question 15

How is atomic radius determined?

Answer 15

By dividing bond distances b/t like atoms in half. Because of overlapping of e- clouds, this method doesn’t generate a true atomic r, but rather a covalent bonding r.

Question 16

The radius of an atom _____ (increases/decreases) as a family in the periodic table is ascended.

Answer 16

The radius of an atom decreases because the number of electronic shells decreases.

Question 17

The radius of an atom ______ ( increases/decreases) as a period in the periodic table is scanned from left to right.

Answer 17

The radium decreases as a periodic table is scanned from L to R, because the effective nuclear charge increases.

Question 18

Explain the sudden increase in atomic radii when one goes from Ne to Na.

Answer 18

The sudden increase is attributed to higher valence shell associated with the additional electron.

Question 19

Describe the atomic radii for transition metals 21 - 30.

Answer 19

The radius of elements 21 - 30 stays roughly equal, since the e- are being added to the third quantum level (d-orbitals)

Question 20

True or false. Transition metals in the same row have very similar atomic radii.

Answer 20

True.

Question 21

The atomic radii trend (increase L to R and increases going up) is fairly consistent, except of helium and hydrogen. Why?

Answer 21

The explanation involves the shielding effect of the two neutrons in the helium nucleus and the e- repulsion experienced in the first quantum level.

Question 22

The largest element would be found on the (lower/upper), (right, left) side of the periodic table.

Answer 22

The largest element would be found on the lower, left side of the periodic table.

Question 23

Describe what ionization energy is.

Answer 23

Ionization is the process of losing an electron from the valence shell. When an atom is ionized, it becomes a cation.

Question 24

What does the energy required to carry out ionization depend on?

Answer 24

The attraction of the electron to the nucleus, the distance from the nucleus, and the stability of its electronic configuration.

Question 25

Within a row in the periodic table, the ionization energy (increases/decreases) as the atomic number increases

Answer 25

The IA increases as the atomic number increases.

Question 26

What are exceptions to the IA trend?

Answer 26

Notable exceptions occur when there is half-filled stability of the energy level and when there is an s^2 shell.

Question 27

Ionization energy (increases/decreases) as you ascend a column in the periodic table.

Answer 27

IA increases as you ascend a column in the periodic table, with the element higher up the column having the greater ionization energy.

Question 28

Explain why the trend for IA for elements increases as you ascend the a column.

Answer 28

As the number of e- shell decreases, the proximity of a valence e- to the nucleus increases, and thus the attraction increases.

Question 29

In terms of ionization energy, how do the transition metals from 21 - 30 behave?

Answer 29

As with atomic radius, elements 21 - 30 are roughly equal, because the e- are removed from the same 4s orbital.

Question 30

What is the principal quantum number for Be? for Li?

Answer 30

Both are n = 2. The quantum number is the shell in which the electron resides.

Question 31

The greater # of protons increases the attractive pull on the e-. Because the pull is greater, the effective nuclear charge is (greater, lesser).

Answer 31

The greater # of protons increases attractive pull and thus greater nuclear charge.

Question 32

Define electron affinity.

Answer 32

EA measures the tendency of an element to gain an electron. It is a measure of the E absorbed or released when an e- is added into the valence shell.

Question 33

There's a sudden increase in EA when one goes from Be to B. Explain why this is so.

Answer 33

The sudden increase in EA is due to the instability of one e- in the p-level.

Question 34

There isn't a trend for EA in transition metals. Is this true?

Answer 34

Yes it is!

Question 35

In general, an element releases more E upon gaining an e- as you move from left to right in the periodic table. Where are exceptions?

Answer 35

Exceptions occur when there is half filled stability of the energy level and when there is an s^2 shell.

Question 36

What is electronic affinity analogous to? Oxidation potential, reduction potential, or electronegativity/

Answer 36

Reduction potential. EN deals with the sharing of e- in a bond. Reduction is the gain of an e-, so it's the closest.

Question 37

Define electronegativity.

Answer 37

EN is defined as the ability of an atom to attract toward itself the e- WITHIN a chemical bond.

Question 38

The EN of an atom (increases/decreases) as the periodic table is ascended.

Answer 38

The EN of an atom increases as the periodic table is ascended because the # of e- shells decreases, causng the attraction to the nucleus to increase.

Question 39

The EN of an atom (increases/decreases) as the periodic table is traversed from left to right .

Answer 39

the EN increases, because the effective nuclear charge increases.

Question 40

True or false. The trend in electronegativity is very clean, showing no exceptions.

Answer 40

TRUE. Finally.

Question 41

When the EN values of two atoms within a bond are close, the bond is _____. When the EN difference exceeds 2.0, then the bond is _____.

Answer 41

Covelent when the EN is close. Ionic when the difference is greater than 2.0

Question 42

How can we best determine the EN difference between bonded atoms? Looking at dipole moment, difference in EA, or difference in IA?

Answer 42

Since EN measures the tendency to share an e-, and the dipole moment represents the degree of sharing b/t two atoms in a bond, this is the aswer. EN can be estimated knowing both EA and IA, but not just one of them.

Question 43

The trend in EN increases with which of the following: IA, atomic number, number of VE?

Answer 43

EN follows a clear trend, so not IA, which is erratic. Atomic number is wrong because when a new shell is formed, the EN drops. Increasing VE doesn't affect shells, so that's the answer.

Question 44

Describe properties of Alkali Metals (Group 1).

Answer 44

Their valence shell is ns^1. As neutral elements they are strong reducing agents because they readily lose an e- to become 1+ cation (which gives them an octet). Their reactivity increases as you descend the column, because it's easier to lose an s e- from a further out shell.

Question 45

All alkali metals react favorably with water to form ____.

Answer 45

All alkali metals react favorably with water to form metal hydroxide and hydrogen gas. The oxides vary from M2O to M2O2 to MO2.

Question 46

Describe the properties of Alkaline Earth Metals (Group 2).

Answer 46

Their valence shell is ns^2. As neutral elements, they are strong reducing agents, because they readily lose 2 e- to become a 2+ cation with a filled octet. They alkali metals they readily react with compound or elements that has a high EA.

Question 47

All Alkaline earth metals, except beryllium, react favorably with water to form a ______ and _____.

Answer 47

All alkaline earth metals react favorably with water to form a metal hydroxide and hydrogen gas.

Question 48

All Alkaline earth metals form oxides (MO) when oxidized by oxygen gas.

Answer 48

2M + O2 --> 2MO (s)

Question 49

Alkali and alkaline earth metals can be oxidized into cations by halogens, nitrogen, and hydrogen. True?

Answer 49

This is true.

Question 50

Describe the properties of Chalcogens (group 6).

Answer 50

Chalcogens are metalloids and nonmetals. Included are O, S, Se, etc. Their valence shell that is ns2np4. They form covalent bonds w/ nonmetals. When neutral, they're oxidizing agents, bc they gain 2 e- to become a -2 anion with a filled octet. Their reactivity decreases as you descend a column, bc their EA isn't as great.

Question 51

Describe the properties of Halogens (group 7).

Answer 51

Halogens are nonmetals. Their valence shell is ns2np5. They form covalent molecules w/ nonmetals and ionic compounds w/ metals. When neutral, they are strong oxidizing agents, bc they readily gain an e- to become a -1 anion w/ a filled octet. Their reactivity decreases as you descend a column, bc the EA is not as great.

Question 52

Describe the properties of Noble gases (group 8).

Answer 52

Noble gaeses are non-metals. Included are He, Ne, Xe, etc. Their valence shell is ns2np6. For the most part, they form no bonds and exist as monatomic atoms. When exposed to F gas, Xe and Kr form molecular compounds.