1. 2.3 Periodic trends

2. Why trends?
Atomic structure and Coulombic attractions are responsible for the patterns in the following periodic properties
Specifically :
The attraction between protons in nucleus and valence electrons (nuclear charge)
The repulsion between electrons and other electrons in the electron cloud (shielding effect)

3. Nuclear charge
Nuclear charge – the net positive charge experienced by an electron from the attractive force of the proton in nucleus
Nuclear charge increases across a period
The number of protons increases but the energy level and number of core (inner) electrons stays the same

4. Nuclear charge
PRACTICE: From each of the following pairs of atoms, select the atom that feels a GREATER force from the nucleus.
Na, Cl
Se, Ca
Cl – more nuclear charge
Se-more nuclear charge

5. Shielding Effect / energy levels
DEFINITION: Inner e- shield the valence e- from the attractive force of the nucleus. Caused by inner e- and valence e- repelling one another Li K
v.e–
TOUGHER
TO REMOVE
EASIER

6. Shielding Effect GROUP TREND: TOP to BOTTOM INCREASES
New energy levels filled with electrons that shield the nucleus
PERIODIC TREND: LEFT to RIGHT
DOES NOT CHANGE
No change in the number of energy levels

7. Trends in Atomic Size
Atomic Radius-one half the distance between the nuclei of two atoms of the same element when the atoms are joined

8. Trends in Atomic Size
Group Trend: Atomic radii of atoms increase as you move from TOP to BOTTOM within a group.
As you move from top to bottom on the PT, the principal QN (n) increases. As n increases, the size of the electron cloud increases.

9. Trends in Atomic Size
Periodic Trend: Atomic radii of atoms decrease slightly as you move from LEFT to RIGHT on the periodic table.
As you move from left to right on the PT, n remains constant, but the positive charge on the nucleus increases by one proton for each element in the period. This results in the electron cloud being pulled in a little tighter.
We call the positive charge that an electron experiences from the nucleus the nuclear charge.

10. Trends in Atomic Size
SUMMARY OF TREND: The size of atomic radii increase top to bottom within a group and decrease left to right across a period.

13. PRACTICE: From each of the following pairs of atoms, select the atom that is larger in radius.
B, C
N, P
Cl, O
Ca, Sc
Ar, Ne
B - less nuclear charge
P - larger “n”
Cl - larger “n”
Ca - less nuclear charge
Ar - larger “n”

14. Trends in Ionic Size
Metals: Metallic ions (cations), on the left and in the center of the table, are formed by loss of e-. When an electron is lost, electron-electron repulsion becomes less and the positively charged nucleus is now attracting fewer electrons. Metallic ions are therefore smaller than the atoms from which they are formed.

15. Trends in Ionic Size Nonmetals: Nonmetallic ions
(anions), located on the right
side of the table, are formed
by gain of e-. When an electron
is gained, electron-electron
repulsion is greater resulting in
anionic radius larger than the
atomic radius. Nonmetallic ions
are therefore larger than the
atoms from which they are
formed.

16. Trends in Ionic Size Group Trend: Ions increase in size as
you move from TOP to BOTTOM
within a group.
Due to increasing value of n.

17. Trends in Ionic Size
Periodic Trend: Ions decrease in size as you move from LEFT to RIGHT amongst metals, and LEFT to RIGHT amongst non-metals on the periodic table.
In an isoelectronic series, ions have the same number of electrons, however there is increasing nuclear charge.

18. Trends in Ionic Size
SUMMARY OF TRENDS: Metallic ions are smaller than their respective atoms, whereas nonmetallic ions are larger than their respective atoms. The size of ions increase as you move from top to bottom within a group and decrease as you move from left to right (among metals and among non-metals) across a period.

19. PRACTICE: From each of the following pairs of atoms, select the atom that is larger in radius.
S or S2-
Ca or Ca2+
S2-
Non-metals: anions are larger than atoms due to greater e- repulsion Ca
Metals: cations are smaller than atoms due to stronger pull on less e-

20. Predicting Oxidation States
When an atom forms an ion, the stability of an octet of electrons enables you to predict the number of electrons to be gained or lost. You can then predict the charge on the ion formed.
When a single atom takes on a charge, it is called a monatomic ion.
The charge on a monatomic ion is known as the oxidation state (oxidation number) of the atom.
Oxidation state (number): the apparent charge on an atom if the electrons in a compound are assigned according to established rules

21. Predicting Oxidation States1+ 2+ 3+ 3- 2- 1- or 4+
Tend to have more than one
oxidation state

22. Trends in Ionization Energy
Ionization energy-the energy required to remove the highest energy e- from a neutral atom
First ionization energy is that energy required to remove first electron
Second ionization energy is that energy required to remove second electron, etc.
It requires more energy to remove each successive electron.

23. Trends in Ionization Energy
When all valence electrons have been removed, the ionization energy increases enormously.

24. Trends in Ionization Energy
Group Trend: Ionization energy of atoms decrease as you move from TOP to BOTTOM within a group.
As you move from top to bottom on the PT, the distance of outermost electrons from the nucleus increases, as does the shielding effect (inner electrons block the attraction of the nucleus for outer electrons) due to increase in n.
Less energy is required to remove first e-.

25. Trends in Ionization Energy
Periodic Trend: Ionization energy of atoms increases as you move from LEFT to RIGHT across the periodic table.
As you move from left to right on the PT, shielding effect remains constant, but nuclear charge increases.
More energy is required to remove first e-.

26. Trends in Ionization Energy
SUMMARY OF TREND: The ionization energy of atoms decrease top to bottom within a group and increase left to right across a period.

27. Trends in First Ionization Energy
Two deviations from this trend exist.

28. Trends in First Ionization Energy
The first occurs between Groups IIA and IIIA.
A IIIA element (ns2 np1) has a smaller IE than the preceding IIA element (ns2).
Electron more easily removed
from p-orbital rather than
s-orbital
e- n farther from nucleus
Small amount of
repulsion by s e-

29. Trends in First Ionization Energy
The second occurs between Groups VA and VIA.
A VIA element (ns2 np4) has a smaller IE than the preceding VA element (ns2 np3).
e- more easily removed
from a doubly occupied np
orbital than a singly occupied
np orbital.
Repulsion from other
e- in orbital helps in its
removal.

30. Factors Affecting Ionization Energy
Nuclear Charge- the larger the nuclear charge, the higher the IE
Shielding Effect- the greater the shielding effect, the lower the IE
Radius- the greater the distance between the nucleus and the outer electrons of an atom, the lower the ionization energy
Sublevel- an e- from a full or half-full sublevel requires more energy to be removed

31. PRACTICE: From each of the following pairs of atoms, select the atom that has a lower first ionization energy.
Be, Mg
Al, P
Be, B
Mg-greater shielding, larger radius
Al-smaller nuclear charge
B-less energy required than removing e- from full 2s

32. Trends in Electronegativity
Electronegativity-the measure of the ability of atom to attract electrons when the atom is in a compound
Can be used to determine the type of bond that will form during a reaction
Can be represented in a table, however the noble gases do not appear because they are not very reactive
The values are expressed in units called Paulings

33. The least electronegative elements are Cesium and Francium at 0.7
The most electronegative element is Fluorine at 4.0
Atoms that have little tendency to gain e- (metals) have a low electronegativity
Atoms that tend to gain e- easily (nonmetals) have a high electronegativity

34. Trends in Electronegativity
Influenced by the same factors that affect ionization energy
Group Trend: Electronegativity values decrease as you move from top to bottom within a group.
Periodic Trend: For representative elements, the values tend to increase from left to right across a period. Electronegativity values among the transition elements are not regular.

36. PRACTICE: Arrange the following elements in order of increasing attraction for electrons in a bond.
S, F, In, Se
Fr, Ga, Ge, P, Zn
In < Sb < Se < F
Fr < Zn < Ga < Ge < P

37. Trends in Reactions of Metals & Nonmetals
Group Trend: Metals become more reactive as we move from TOP to BOTTOM on the PT. Nonmetals become more reactive as we move from BOTTOM to TOP on the PT. As we move “down” the PT, each successive element has one greater EL filled with e-. This increases the shielding of the nucleus and therefore e- are more easily “lost” (i.e., less pull from nucleus) from larger atom metals. A “smaller” atom (fewer EL’s) will have a greater attraction for e- due to less shielding and therefore e- are more easily “gained” by smaller atom nonmetals.