Periodic Trends
Lecture 0802
Trends on the
Periodic Table
PERIODIC
TRENDS
Li
Na
K
Effective Nuclear Charge
Z*
The 2s electron PENETRATES the region occupied by the 1s electron.
2s electron experiences a higher positive charge than expected.
Effective
Nuclear Charge, Z*
Atom Z* e in Valence Orbitals
Li
Be
B
C est measured
1 +1.28
2 -------
Increase in
Z* across a
3 +2.58
period
4 +3.22
N
O
F
5 +3.85
6 +4.49
7 +5.13
General Periodic Trends
Atomic and ionic size
Ionization energy
Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger shells.
Electrons held less tightly.
Atomic Radius
Is taken as the covalent radius for non-metallic elements and as the metallic radius for metals
Atomic Radius
Covalent radius is one-half the distance between the nuclei of two identical atoms that are singly bonded to one another.
Chlorine Bond Length
Atomic Radius
Covalent radii for elements whose atoms do not bond to one another can be estimated by combining radii of those that do with the distances between unlike atoms in various molecules.
Atomic Radius
Metallic radius is one-half the closest internuclear distance in a metallic crystal.
Prediction!
200
150
100
50
0
1 3 5 7 9 11 13 15 17 19
Atomic Number
Atomic Size
Size goes UP on going down a group.
Because electrons are added further from the nucleus, there is less attraction.
Size goes DOWN on going across a period.
Atomic Size
Size decreases across a period owing to increase in Z*. Each added electron feels a greater and greater + charge.
Large
Small
100
150
200
250
Trends in Atomic Size
Radius (pm)
K
1st transition series
2nd period
Li
3rd period
Na
Kr
Ar
Ne
50
0
0
He
5 10 15 20
Atomic Number
25 30 35 40
Sizes of Transition Elements
3d subshell is inside the 4s subshell.
4s electrons feel a more or less constant Z*.
Sizes stay about the same and chemistries are similar!
General Periodic Trends
Atomic and ionic size
Ionization energy
Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger shells.
Electrons held less tightly.
Li,152 pm
3e and 3p
Ion Sizes form a cation?
Ion Sizes
Li,152 pm
3e and 3p
+
Li +, 78 pm
2e and 3 p
Forming a cation.
CATIONS are SMALLER than the atoms from which they come.
The proton/electron attraction has gone UP and so size
DECREASES.
Ion Sizes
Does the size go up or down when gaining an electron to form an anion?
Ion Sizes
F, 71 pm
9e and 9p
-
F- , 133 pm
10 e and 9 p
Forming an anion.
ANIONS are LARGER than the atoms from which they come.
The proton/electron attraction has gone DOWN and so size
INCREASES.
Trends in Ion Sizes
Trends in ion sizes are the same as atom sizes.
Redox Reactions
Why do metals lose electrons in their reactions?
Why does Mg form Mg 2+ ions and not Mg 3+ ?
Why do nonmetals take on electrons?
Ionization Energy
(General)
Is the energy required to remove the outermost electron from an atom or a positive ion in the ground state.
First Ionization Energy
Energy required to remove the first electron from a neutral atom in the gaseous state.
Ionization Energy
Mg (g) + 738 kJ Mg + (g) + e -
Prediction!
2500
2000
1500
1000
500
0
1 3 5 7 9 11 13 15 17 19
Atomic Number
2500
Trends in Ionization Energy
1st Ionization energy (kJ/mol)
He
Ne
2000
Ar
1500 Kr
1000
500
0
1 3
H Li
5 7 9 11 13 15 17 19 21 23 25 27 29 31 33 35
Na K
Atomic Number
Atomic Radii
Trends in Ionization Energy
IE increases across a period because Z* increases.
Metals lose electrons more easily than nonmetals.
Metals are good reducing agents.
Nonmetals lose electrons with difficulty.
Trends in Ionization Energy
IE decreases down a group
Because size increases.
Reducing ability generally increases down the periodic table.
Remember Li, Na, K
Second Ionization Energy
Energy needed to remove the outermost electron from a +1 ion.
Energy needed to remove the second electron from a neutral atom.
Ionization Energy
Mg (g) + 738 kJ Mg + (g) + e -
Mg + (g) + 1451 kJ Mg 2+ (g) + e -
Mg + has 12 protons and only 11 electrons.
Therefore, IE for Mg + > Mg.
8000
7000
6000
5000
4000
3000
2000
1000
0
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20
Atomic Number
Ionization Energy
Mg (g) + 735 kJ Mg + (g) + e -
Mg + (g) + 1451 kJ Mg 2+ (g) + e -
Mg 2+ (g) + 7733 kJ Mg 3+ (g) + e -
Energy cost is very high to dip into a shell of lower n.
This is why ox. no. = Group no.
General Periodic Trends
Atomic and ionic size
Ionization energy
Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger shells.
Electrons held less tightly.
Electron Affinity
A few elements GAIN electrons to form anions .
E.A. is the energy released or absorbed when an electron is added to the valence level of a gas-phase atom.
A(g) + e A (g) E.A. = ∆E
Prediction!
Trends in Electron Affinity
Electron Affinity of Oxygen
O atom [He]
+ electron
O
- ion [He]
EA = - 141 kJ
∆E is
EXO thermic because O has an affinity for an e .
Electron Affinity of Nitrogen
N atom [He]
N ion [He]
+ electron
EA = 0 kJ
∆E is zero for
N due to electronelectron repulsions.
Trends in Electron Affinity
See Figure 8.12 and
Appendix F
Affinity for electron increases across a period (EA becomes more negative).
Affinity decreases down a group (EA becomes less negative).
Atom EA
F -328 kJ
Cl -349 kJ
Br -325 kJ
I -295 kJ
General Periodic Trends
Atomic and ionic size
Ionization energy
Electron affinity
Higher effective nuclear charge
Electrons held more tightly
Larger shells.
Electrons held less tightly.
