Question | Answer |
Noble Gases | All of them have completely filled electron shells. n2 np6
Since they have similar electronic structures, their chemical reactions are similar. |
Representative Elements | Are the elements in A groups on periodic chart.
These elements will have their “last” electron in an outer s or p orbital.
These elements have fairly regular variations in their properties. |
d-Transition Elements | Elements on periodic chart in B groups.
Each metal has d electrons.
ns (n-1)d configurations
These elements make the transition from metals to nonmetals.
Exhibit smaller variations from row-to-row than the representative elements |
f - transition metals | "inner transition metals"
Electrons being added to f orbitals, two shells below the valence shell!
Slight variations of properties
Outermost electrons have the greatest influence on the chemical properties of elements. (n quantum #) |
Outer most electrons | have the greatest influence on the properties of elements. Adding an e to an s or p orbital usually causes dramatic changes in the physical and chemical properties. Adding an e to a d or f orbital typically has a much smaller effect on properties. |
Outer most electrons | have the highest value of the principal quantum number, n. |
Atomic Radii | describe the relative sizes of atoms
increase within a column going from the top to the bottom of the periodic table
decrease within a row going from left to right on the periodic table. |
Description of atomic radii | shielding/screening effect
nuclear charge experienced by e- < actual nuclear charge
Inner e- block nuclear charge’s effect on outer e-
Moving across a period each element increases nuclear charge and e- go into same shell
Outer e- feel > nuclear charg |
Atomic radii increase going down a group because | e are being added to shells farther from the nucleus. Atomic radii decrease from left to right within a given period owing to increasing effective nuclear charge |
First ionization energy (IE1) | Atom(g)+Energy=Ion+(g)+e-
Measures how tightly e are bound to atoms.
Low ionization energies indicate easy removal of e and hence easy positive ion formation |
| it is always more difficult to remove a neg charged e from a pos charged ion than from the corresponding neutral atom |
Ionization Energy | IE1 decrease because the outermost e are farther from the nucleus.
1st IE1 generally increase from left to rt across the periodic table because the e are held more tightly by nucleus. |
Second ionization energy (IE2) | The amount of energy required to remove the second electron from a gaseous 1+ ion.
Symbolically:
ion+ + energy --> ion2+ + e- |
Periodic trends for Ionization Energy: | IE2 > IE1 always takes more energy to remove 2nd e- from an ion than from neutral atom
IE1 generally increases moving from Left to Right
EXCEPTIONS due to filled and half-filled subshells
IE1 generally decreases moving down
IE1 for Li > IE1 for Na, et |
Electron Affinity | amount of energy absorbed when atom(g) + e- + EA --> ion-(g)
Electron affinity is a measure of an atom’s ability to form negative ions.
Elements with very neg. EA gain electrons easily to form negative ions (anions) |
Sign conventions for electron affinity. | If electron affinity > 0 energy is absorbed.
If electron affinity < 0 energy is released. |
General periodic trend for electron affinity | values become more - from left to right
the values become more negative from bottom to top on the periodic chart.
EA is a precise and quantitative term, like ionization energy, but it is difficult to measure.
EA of anions always +Ionic Radii |
Ionic Radii | Cations (positive ions) are always smaller than their respective neutral atoms.
Negative ions (anions) are always larger than their respective neutral atoms.
Both cation and anion sizes increase going down a group. |
Within an isoelectronic series ionic radii | decrease with increasing atomic number because of increasing nuclear charge (more p+ to e-). |
isoelectronic species | have the same number of e. Comparing such species, the higher the nuclear charge (more protons), the smaller the radius. |
anion ionic radii | Anion (negative ions) radii decrease from left to right across a period.
Increasing electron numbers in highly charged ions cause the electrons to repel and increase the ionic radius. |
Electronegativity | a measure of the relative tendency of an atom to attract electrons to itself when chemically combined with another element |
For the representative elements, electronegativities usually... | increase from left to right across periods and decrease from top to bottom within groups |
High EN vs low EN | Elements with high EN (nonmetals) often gain E to form anions.
Elements with low EN (metals) often lose E to form cations. Cs and Fr
Large difference in EN --> ionic compounds |
Oxidation States | of an element in a simple binary (meaning 2) ionic compound is the number of e gained or lost by an atom of that element when it forms the compound.
the element with the pos oxidation state is written first, except NH3, and hydrocarbons CH4 H is +1 |
Electronegativity and oxidation states | The most EN atom in a compound is assigned a neg oxidation state, while the less EN atom lose e and have pos oxidation states. |
Guidelines for assigning oxidation numbers | Of any uncombined element = 0
Of an element in a simple ion = charge on the ion
Any compound sum of oxidation #s of all elements = 0
In a polyatomic ion, the sum of the oxidation numbers of the constituent elements = charge on the ion |
More Guidelines for assigning oxidation numbers | F oxidation number = –1 in its compounds.
H oxidation number = +1 unless combined with metals = -1
Oxygen usually has the oxidation number -2.
Exceptions: In peroxides O has oxidation number of –1.
In OF2 O has oxidation number of +2. |
| Periodic table helps assign oxidation numbers
IA metals have oxidation numbers of +1...+2...+3...-3...-2 |