Ch. 8 Flashcards
Valence electrons
The electrons important in chemical bonding. For main group elements, these electrons are in the outer most principal energy level.
Core electrons
Electrons in complete principal energy levels, and complete d and f sublevel’s, not including the valence electrons
Non-bonding atomic radius and van der waals radius 
The radius of an atom determined from the distance between atomic centres in an atomic solid.
Defined as one half the distance between the centre of adjacent, nonbonding Atoms in a Crystal
Bonding atomic radius/covalent radius
For nonmetals, 1/2 the distance between the two atoms bonded together. For metals, 1/2 the distance between the two Atoms next to each other in a crystal of the metal. 
Defined in nonmetals as one half the distance between two atoms bonded together, and in medals as one half the distance between two adjacent Atoms in a crystal of the metal
Atomic radius(radii)
Average bonding radius for a given element, determined from large numbers of compounds.
(Size can be defined as the 1/2 of the distance between Atoms)
-Represents the radius of an Atom when it is bonded to another Atom and is always smaller than the van der waals radius.
-Approximate bond length of any to covalently bonded atoms is simply the sum of their atomic radii.
-Atomic radii peak with each alkali metal.
-Core electrons are better at shielding.
General trends:
1. As we move down a column (or family) in the periodic table, atomic radius increases.
-New shells (n increasing), more shells of inner electrons that shield outer electrons = further away from nucleus therefor larger radius.
- As we move right across a period (or row) in the periodic table, atomic radius decreases.
-More, e-, p+, n0, but Zeff increases —> Net attraction to nucleus is greater.
-Gets smaller because more P+ = more attraction to the nucleus therefore must be smaller.
-To understand trends in radii, review the concept of effective nuclear charge Zeff = Z - S.
-Z = # of protons, S = shielding
-Higher Zeff more attraction it’s going to feel from the nucleus.
-as Zeff increases, radius decreases
Lanthanide contraction
The trend toward levelling off in size of the Atoms in the third and fourth transition rows due to ineffective shielding of the F sublevel electrons.
-Looking down a group in the d-block elements, we see a small but expected increase in size from the first transition metal row to the second, but the size of the elements in the third row is about the same as it is for those in the second row. This pattern is also different from that of the main group elements, especially when we consider that in any given column, the third transition row has 32 more electrons in the second row. The reason that the third row transition metals are not larger is as follows: 14 of the 32 electrons are in a large for 4f sublevel. These electrons are not effective at shielding the 5d and 6s electrons From the nuclear charge. Consequently, the outer electrons are held more tightly by the nucleus, offsetting the typical increase in size between the periods. 
-It’s not protected by simple Slater’s rules. This is because the shielding contributions for electrons in different types of orbitals are not truly equal. For example, a 4F electron should have a different shielding contribution than a 4d electron.
Ionization energy (IE)
The Energy required to remove an electron from an atom or ion and it’s gaseous state.
X(g) —> X + (g) + e-
•IE1 is The removal of the first electron.
•The electron comes from the outer shell (valence electron)
•If electron is closer to the nucleus the harder it will be to remove, so you’ll need a higher ionization energy.
Trends in IE1:
•Across a period, IE1 increases.
-Zeff Increases, more attraction, radius decreases, electrons held tighter to nucleus, therefore more energy required. 
•Down a group, IE1 decreases
-New shells when n increases, so electrons are further away from nucleus, therefore less energy required to remove electrons.
Anomalies in IE1 trends:
•There is a decrease in ionization energy between group 2 and Group 13 and between group 15 and group 16. 
-G2 & G13: Be 1s2 2s2, When removing from this, you are taking from 2s2. B 1s2 2s2 2p1, When removing from this you’re taking from 2p1. 2S2 (^2) is closer to the nucleus and therefore harder to remove then 2p.
-G15 & G16: N 1s2 2s2 2p3, This is a more stable configuration, because there’s less electron electron repulsion (Hunds rule). O 1s2 2s2 2p4
Electron affinity (EA)
The energy released when an electron is added to a neutral atom in the gas phase.
Or (means same thing)
Is the energy change associated with the addition of an electron to a neutral atom in the gas phase.
X(g) + e- —> X - (g)
•EA (exothermic) usually negative
•A negative value indicates the energy is released
•Not all Atoms have a negative electron affinity.
-Those that don’t will not bind to an extra electron
-Reverse process is exothermic
•No gas phase Adam will bind more than one extra electron
-Multiply charged anions only exist in condensed phases.
•Metals tend to lose electrons = forms cations.
•Nonmetals tend to gain electrons = forms anions.
Trends:
•EA Tends to increase across a period (more negative)
-Zeff Increases, electrons feel more attraction, closer to the nucleus.

•EA Tends to decrease down a group (less negative)
-Adding electrons = further away from the nucleus, not as stable.
Summary of trends of Zeff, radii, IE, and EA
Down Periodic table:
Zeff: constant
Radius: increases(larger) (adding e^-)
IE: decreases (need less energy to remove)
EA: less negative
Across periodic table (left to right):
Zeff: increases (b/c atoms get smaller, more attractive, more pts, less shielding)
Radius: decreases
IE: increases (b/c closer to nucleus therefore need more energy to remove)
EA: more negative
Isoelectronic series
Atoms or ions that have the same electron configuration.
-Smallest has the highest Zeff = Higher nuclear charge = higher the attraction = the smaller the atom.
Ionic radius(radii)
•Determined using interatomic distances in ionic compounds.
•adding an electron to an atom increases electron-electron repulsion.
-Anions are larger than neutral Atoms
-More electron electron repulsion (or shielding), adding new shells.
-making Zeff less = larger.
•Removing an electron reduces electron electron repulsion
-Cations are smaller than neutral Atoms.
-Outer most electrons feel greater Zeff as electrons removed.
-making Zeff greater = smaller.
- Zeff increases and ionic radius becomes smaller (when taking e-)
- Zeff decreases and ionic radius, becomes larger (when adding e-)
Z effective
The size of ions as measured by ionic radii varies in a svstematic manner. The size of the ion can be explained in part by effective nuclear charge, Zeff, which is the net nuclear charge felt by an electron.
The effective nuclear charge takes into account the actual nuclear charge and the shielding of this charge by inner electrons.
When an atom loses electrons, the resulting cation is smaller both because the remaining electrons experience a larger Zeff and because these electrons are usually in orbitals closer to the nucleus than the electrons that were lost. The more electrons that are lost, the smaller the ion becomes.
Similarly, when an atom gains electrons, the resulting anion is larger owing to both increased electron-electron repulsions and a reduction in Zeff. The more electrons that are gained, the larger the ion becomes.
Zeff increases across a period.
Zeff down a period is constant
