Periodic Trends Flashcards
How is the periodic table blocks organized?
The blocks are named after the last electron that fills each respective atom in that section of the periodic chart.
An elements chemical properties depend on its _____/
Valence electrons.
Describe the trend on the periodic table, in relation to the effective nuclear charge.
As you move from left to right across a period in the periodic table, the effective nuclear charge increases.
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).
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.
Define the effective nuclear charge (Zeff).
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-.
How does valence shell affect periodic trends?
The valence shell affects the distance between between electrons and the nucleus.
What’s the periodic trend for atomic size?
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).
What’s the periodic trend for ionization energy?
The IA increases ( the energy required to remove the outermost e- from the atom in its gas phase, creating a cation).
What’s the periodic trend for electron affinity?
The EA increases ( the energetics associated with an atom gaining an e- in its gas phase).
What’s the periodic trend for electronegativity.
The EN increases ( tendency to hold shared e- with another within a BOND).
When doing MCAT questions on periodic trends, how should you handle F?
F tends to be on extremes.
What can cause inconsistency in the periodic trends we typically see?
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.
As a general rule, cations are _____ ( larger/smaller) than neutral atoms.
Smaller, because the loss of electrons allow the atom to compact more tightly.
Within a period, anions are _____( larger/smaller) than cations.
Larger. Cl- ion is larger than the Na+ cation.
How is atomic radius determined?
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.
The radius of an atom _____ (increases/decreases) as a family in the periodic table is ascended.
The radius of an atom decreases because the number of electronic shells decreases.
The radius of an atom ______ ( increases/decreases) as a period in the periodic table is scanned from left to right.
The radium decreases as a periodic table is scanned from L to R, because the effective nuclear charge increases.
Explain the sudden increase in atomic radii when one goes from Ne to Na.
The sudden increase is attributed to higher valence shell associated with the additional electron.
Describe the atomic radii for transition metals 21 - 30.
The radius of elements 21 - 30 stays roughly equal, since the e- are being added to the third quantum level (d-orbitals)
True or false. Transition metals in the same row have very similar atomic radii.
True.
The atomic radii trend (increase L to R and increases going up) is fairly consistent, except of helium and hydrogen. Why?
The explanation involves the shielding effect of the two neutrons in the helium nucleus and the e- repulsion experienced in the first quantum level.
The largest element would be found on the (lower/upper), (right, left) side of the periodic table.
The largest element would be found on the lower, left side of the periodic table.
Describe what ionization energy is.
Ionization is the process of losing an electron from the valence shell. When an atom is ionized, it becomes a cation.
What does the energy required to carry out ionization depend on?
The attraction of the electron to the nucleus, the distance from the nucleus, and the stability of its electronic configuration.
Within a row in the periodic table, the ionization energy (increases/decreases) as the atomic number increases
The IA increases as the atomic number increases.
What are exceptions to the IA trend?
Notable exceptions occur when there is half-filled stability of the energy level and when there is an s^2 shell.
Ionization energy (increases/decreases) as you ascend a column in the periodic table.
IA increases as you ascend a column in the periodic table, with the element higher up the column having the greater ionization energy.
Explain why the trend for IA for elements increases as you ascend the a column.
As the number of e- shell decreases, the proximity of a valence e- to the nucleus increases, and thus the attraction increases.
In terms of ionization energy, how do the transition metals from 21 - 30 behave?
As with atomic radius, elements 21 - 30 are roughly equal, because the e- are removed from the same 4s orbital.
What is the principal quantum number for Be? for Li?
Both are n = 2. The quantum number is the shell in which the electron resides.
The greater # of protons increases the attractive pull on the e-. Because the pull is greater, the effective nuclear charge is (greater, lesser).
The greater # of protons increases attractive pull and thus greater nuclear charge.
Define electron affinity.
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.
There’s a sudden increase in EA when one goes from Be to B. Explain why this is so.
The sudden increase in EA is due to the instability of one e- in the p-level.
There isn’t a trend for EA in transition metals. Is this true?
Yes it is!
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?
Exceptions occur when there is half filled stability of the energy level and when there is an s^2 shell.
What is electronic affinity analogous to? Oxidation potential, reduction potential, or electronegativity/
Reduction potential. EN deals with the sharing of e- in a bond. Reduction is the gain of an e-, so it’s the closest.
Define electronegativity.
EN is defined as the ability of an atom to attract toward itself the e- WITHIN a chemical bond.
The EN of an atom (increases/decreases) as the periodic table is ascended.
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.
The EN of an atom (increases/decreases) as the periodic table is traversed from left to right .
the EN increases, because the effective nuclear charge increases.
True or false. The trend in electronegativity is very clean, showing no exceptions.
TRUE. Finally.
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 _____.
Covelent when the EN is close. Ionic when the difference is greater than 2.0
How can we best determine the EN difference between bonded atoms? Looking at dipole moment, difference in EA, or difference in IA?
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.
The trend in EN increases with which of the following: IA, atomic number, number of VE?
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.
Describe properties of Alkali Metals (Group 1).
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.
All alkali metals react favorably with water to form ____.
All alkali metals react favorably with water to form metal hydroxide and hydrogen gas. The oxides vary from M2O to M2O2 to MO2.
Describe the properties of Alkaline Earth Metals (Group 2).
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.
All Alkaline earth metals, except beryllium, react favorably with water to form a ______ and _____.
All alkaline earth metals react favorably with water to form a metal hydroxide and hydrogen gas.
All Alkaline earth metals form oxides (MO) when oxidized by oxygen gas.
2M + O2 –> 2MO (s)
Alkali and alkaline earth metals can be oxidized into cations by halogens, nitrogen, and hydrogen. True?
This is true.
Describe the properties of Chalcogens (group 6).
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.
Describe the properties of Halogens (group 7).
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.
Describe the properties of Noble gases (group 8).
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.
