Chemical Periodicity

Trends in the periodic tableTrends in the periodic table

Atomic Size

How do you measure the size of an atom?

The electron cloud doesn’t have a definite edge.

Can get around this by measuring covalent atomic radius.

Atomic Size

Atomic Radius = half the distance between two nuclei of a diatomic molecule.

}Radius

Atomic size is influenced by two factors:

Energy Level –more occupied levels = bigger atom

Charge on nucleus– More charge pulls electrons in closer

Group trends

As we go down a group electrons are added to higher energy levels so the atoms get bigger.

HLi

Na

K

Rb

Periodic Trends As you go across a period the radius

gets smaller. Same energy level, but protons pull

electrons closer to nucleus.

Na Mg Al Si P S Cl Ar

Trends in Atomic Radius

Questions:

Of the following elements, which has the largest atomic radius? Why?

a) Si, Mg, S

b) Al, Na, Cl

c) Li, Cs

Mg – same energy level, smallest nuclear charge

Na – same energy level, smallest nuclear charge

Cs – higher occupied energy levels

Ionic Size

Ionic SizeCations Positive ions - form by losing electrons. Metals form cations Cations of representative elements have noble gas

configuration. Smaller than the atom they come from because of

increased attraction by nucleus for fewer remaining electrons

+

Ionic sizeAnions Negative ions - form by gaining electrons. Nonmetals form anions. Anions of representative elements have noble gas

configuration. Larger than the atom they come from, because

nuclear attraction is less for an increased number of electrons.

-

Group trends

Ions get bigger as you go down (adding energy levels)

Li+1

Na+1

K+1

Rb+1

Cs+1

Periodic Trends

Across the period nuclear charge increases so both cations and anions get smaller from left to right.

Li+1

Be+2

B+3

C+4

N-3O-2 F-1

Questions:1. Of the following ions, which ones should have

the larger radius? Why?a) Na+ or Cs+

b) Br- or K+

2. The Mg2+ and Na+ ions have ten electrons surrounding the nucleus. Which ion would you expect to have the smaller radius? Why?

Cs+ It has more occupied energy levels

Br- Anions are larger than cations

Mg2+ Greater nuclear charge

Ionization Energy

Ionization Energy The amount of energy required to

completely remove an electron from a gaseous atom (how hard it is to pull an e- off an atom)

1st IE = removing 1 e-, 2nd IE=removing 2 e-

Na(g) Na+ + e-

Shielding

The electron on the outside energy level is shielded from the nucleus by the inner electrons

Group trends As you go down a group first IE

decreases because the electron is further away (more shielding)

Periodic trends

All the atoms in the same period have the same energy level (same shielding).

As you go from left to right, nuclear charge increases so IE generally increases.

Questions:1. Which element in the following sets has the

lowest ionization energy and why?a) B, C, F

b) K, Na, LiB – same energy level, smallest nuclear charge

K – electron farther away, more shielding

Electron Affinity

Electron Affinity The energy given off when an electron is

added to an atom how much an atom ‘wants’ an electron

F(g) + e - F -(g)

Electron AffinityGroup trends Generally decreases as we go down a group

because shielding increases

Periodic trends Increases from left to right as atoms become

smaller with greater nuclear charge

Questions:1. Of the following elements, which ones should

have the higher electron affinity? Why?a) Se or Te

b) Calcium or ChromiumSe – smaller atom

Chromium – greater nuclear charge

Electronegativity

Electronegativity

The tendency for an atom to attract electrons to itself when it is chemically combined (BONDED) with another element.

Big electronegativity means it pulls the electron towards itself.

Group Trends The further down a group the farther the

electron is away from the nucleus and the more electrons an atom has.

More willing to share = low electronegativity

So as you go down a group electronegativity decreases

Periodic Trends As we go from left to right across the table,

electronegativity increases, because nuclear charge is increasing and electrons are held in more strongly

Metals have low electronegativity Non-metals have high electronegativities

(they win the electron tug-of-war)

Questions:1. Which element would you expect to have the

highest electronegativity? Why?

2. Put the following elements in order of increasing electronegativity: Na, P, Cl

F smallest nonmetal

Na, P, Cl