periodic properties of the elements chapter 7. effective nuclear charge orbitals of the same energy...
TRANSCRIPT
Periodic Properties of the Elements
Chapter 7
Effective Nuclear Charge
Orbitals of the same energy are said to be degenerate.
Effective nuclear charge is the charge experienced by an electron on a many-electron atom.
The effective nuclear charge is not the same as the charge on the nucleus because of the effect of the inner electrons.
Effective Nuclear Charge
Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nucleus.
The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons.
As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases.
As the distance from the nucleus increases, S increases and Zeff decreases.
Energies of Orbitals
The result of the Effective nuclear charge on the electronic configuration is a shift of the orbital ordering for large (n = 4 or more) electronic systems.
Energies of Orbitals
Sizes of Atoms
Electron Shells in AtomsConsider a simple diatomic molecule.The distance between the two nuclei is called the bond distance.If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom.
Atomic Radii•As a consequence of the ordering in the periodic table, properties of elements vary periodically.•Atomic size varies consistently through the periodic table.•As we move down a group, the atoms become larger.•As we move across a period, atoms become smaller.
There are two factors at work:•principal quantum number, n, and•the effective nuclear charge, Zeff.
Atomic Radii decreasing across a period is due to:
shielding or screening effect inner electrons [He] or [Ne], etc. block the nuclear charge for 2 or 10 or
__ electrons consequently the outer electrons feel
a stronger effective nuclear charge Li [He] shields effective charge is +1 Be [He] shields effective charge is +2
Atomic Radii
Atomic Radii
All radii are in angstroms.
Ionic Radii
Cations (+ ions) are always smaller than their neutral atoms
Li Be1.52 1.11
Li+ Be2+0.6 0.31
Na Mg Al1.86 1.6 1.43
Na+ Mg2+ Al3+0.95 0.65 0.5
Ionic Radii Anions (- ions) are always larger
than their neutral atomsN O F
0.7 0.66 0.64
N3-
O2-
F-
1.71 1.4 1.36
Ionization Energy minimum amount of energy required to
remove the most loosely held electron from an isolated gaseous atom measure of an element’s ability to form
positive ions
Mg(g) + 738kJ/mol Mg+ + e-
Atom(g) + energy ion+(g)
+ e-
first ionization energyfirst ionization energy
Ionization Energy second ionization energysecond ionization energy
energy required to remove a 2nd electron
ion+ + energy ion2+ + e-
Mg+ + 1451 kJ/mol Mg2+ + e-
• can have 3rd, 4th, etc. ionization energies
Ionization Energy generally increases as you go
across a period important exceptions at Be & Mg, N & P
generally decreases as you go down a group
Ionization Energy (kJ/mol) vs Atomic Number
H1
He2
B5
C6
N7
O8
F9
Ne10
Mg12 Al
13
Si14
P15
S16
Cl17
Ar18
Ca20
Be4
Li3
Na11 K
19
0
500
1000
1500
2000
2500
0 5 10 15 20
Ionization Energy First 4 ionization Energies (kJ/mol) -
Period 3
IA IIA IIIA IVANa Mg Al Si
1stIE 496 738 578 7862ndIE 4562 1451 1817 15773rdIE 6912 7733 2745 32324thIE 9540 10,550 11,580 4356
Ionization Energy
these energies are exactly why these ions form
Na becomes Na+
Mg becomes Mg2+
Al becomes Al3+
Si does not form simple ions
Ionization Energy
Electronegativity
The attraction of an atom to electrons
This is not a measurable property BUT it is very useful in helping to predict bonding (attraction of electrons)
Electronegativity