I II III

Unit 5 AP Chemistry

Periodic Table Trends

Periodic Law

When elements are arranged in order of

increasing atomic #, elements with similar

properties appear at regular intervals.

0

50

100

150

200

250

0 5 10 15 20

Ato

mic

Ra

diu

s (p

m)

Atomic Number

ALL Periodic Table Trends

Influenced by three factors:

1. Energy Level Higher energy levels are further away

from the nucleus.

2. Charge on nucleus (# protons) More charge pulls electrons in closer. (+ and – attract each other)

3. Shielding effect(blocking effect)

What do they influence?

Energy levels and Shielding

have an effect on the GROUP

Nuclear charge has an effect

on a PERIOD

Atomic Radius size of atom

© 1998 LOGAL

Atomic Radius

Atomic Radius Average distance in an atom between

the nucleus and the outermost electron

#1. Atomic Size - Group trends

Increases going down a group). . .

each atom has another energy level

so the atoms get

bigger.

HLi

Na

K

Rb

#1. Atomic Size - Period Trends

Going from left to right across a period, the size gets smaller.

Electrons are in the same energy level.But, there is more nuclear charge.Outermost electrons are pulled closer.

Na Mg Al Si P S ArCl

1

2

3

4 5

6

7

Atomic Radius Increases to the LEFT and DOWN

Atomic Radius

#2. Trends in Ionization Energy

Ionization energy is the amount of energy required to completely remove an electron (from a gaseous atom).

M + energy M+1 + e-

Removing one electron makes a +1 ion.The energy required to remove only the

first electron is called the first ionization energy.

Ionization Energy

The second ionization energy is the energy required to remove the second electron. Always greater than first IE.

The third IE is the energy required to remove a third electron. Greater than 1st or 2nd IE.

What factors determine IE

The greater the nuclear charge, the greater IE.

Greater distance from nucleus decreases IE

Filled and half-filled orbitals have lower energy, so achieving them is easier, lower IE.

Shielding effect

Shielding

The electron on the outermost energy level has to look through all the other energy levels to see the nucleus.

Second electron has same shielding, if it is in the same period

Ionization Energy - Group trends

As you go down a group, the first IE decreases because...The electron is further away from the attraction of the nucleus, and

There is more shielding.

Ionization Energy - Period trends

All the atoms in the same period have the same energy level.

Same shielding.But, increasing nuclear chargeSo IE generally increases from left

to right.Exceptions at full and 1/2 full orbitals.

Trend is opposite of atomic radius.Why?

In small atoms, e- are close to the nucleus where the attraction is stronger

Ionization Energy Trends

1

2

3

4 5

6

7

First Ionization Energy Increases UP and to the RIGHT

E. Ionization Energy

First Ionization Energy

0

500

1000

1500

2000

2500

0 5 10 15 20Atomic Number

1s

t Io

niz

ati

on

En

erg

y (k

J)

E. Ionization Energy

KNaLi

Ar

NeHe

Firs

t Ion

izat

ion

ener

gy

Atomic number

HeHe has a greater IE He has a greater IE than H.than H.

Both elements have Both elements have the same shielding the same shielding since electrons are since electrons are only in the first level only in the first level

But He has a greater But He has a greater nuclear chargenuclear charge

H

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He Li has lower IE than H

more shielding further away These outweigh

the greater nuclear charge

Li

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Be has higher IE

than Li

same shielding

greater nuclear

charge Li

Be

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He B has lower IE than Be

same shielding greater nuclear

charge By removing an

electron we make filled s -sublevelLi

Be

B

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

Oxygen breaks Oxygen breaks the pattern, the pattern, because because removing an removing an electron leaves it electron leaves it with a 1/2 filled p with a 1/2 filled p orbitalorbital

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Ne has a lower Ne has a lower IE than HeIE than He

Both are full,Both are full,

Ne has more Ne has more shieldingshielding

Greater Greater distancedistance

Firs

t Ion

izat

ion

ener

gy

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Na has a lower IE than Li

Both are s1

Na has more shielding

Greater distance

Na

Successive Ionization Energies

Mg 1st I.E. 736 kJ

2nd I.E. 1,445 kJ

Core e- 3rd I.E. 7,730 kJ

Large jump in I.E. occurs when a CORE e- is removed.

Ionization Energy

Al 1st I.E. 577 kJ

2nd I.E. 1,815 kJ

3rd I.E. 2,740 kJ

Core e- 4th I.E. 11,600 kJ

Successive Ionization Energies

Large jump in I.E. occurs when a CORE e- is removed.

Ionization Energy

#3. Trends in Electronegativity

Electronegativity is the tendency for an atom to attract electrons to itself when it is chemically combined with another element.

They share the electron, but how equally do they share it?

An element with a big electronegativity means it pulls the electron towards itself strongly!

Electronegativity Group Trends

Electronegativity Decreases Down a Group Why? Atomic size increases and valence electrons

are farther from the nucleus. More energy levels increases shielding. So the

pull from the positive nuclear charge is less. In General:

Non-Metals have high ElectronegativitiesMetals have low Electronegativities

Electronegativity Period Trend

Metals are at the left of the table.They let their electrons go easilyThus, low electronegativityAt the right end are the nonmetals.They want more electrons.Try to take them away from othersHigh electronegativity.

Electronegativity

Trend is also opposite of atomic size

The smaller the atom, the more electronegative it is because of a greater nuclear charge.

Exception: Noble gases are not included because they generally do not want to gain electrons. They are already stable.

Ionic Radius

Cations (+ ions) the ionic radius is

smaller than the original atom.

Why? There is an increased attraction

for the fewer electrons that remain.

Ionic Radius

Na Na+

Ionic Radius

For Anions (– ions) the ionic radius is

larger than the original atom.

Why? The nuclear attraction is less for

an increased number of electrons.

Extra electrons repel each other and

spread out – larger!)

© 2002 Prentice-Hall, Inc.

Cl Cl-1

Ion Group trends

Each step down a group is adding an energy level

Ions therefore get bigger as you go down, because of the additional energy level.

Li1+

Na1+

K1+

Rb1+

Cs1+

Ion Period Trends

Across the period from left to right, the nuclear charge increases - so they get smaller.

Notice the energy level changes between anions and cations.

Li1+

Be2+

B3+

C4+

N3-O2- F1-

#5. Electron Affinity

The energy change when an electron is added to an atom forming an anion.

Different than Electronegativity because EN is attraction for electrons in a bond!

X(g) + e- X-(g) + Energy

A negative energy indicates energy is released when an electron is added

Electron Affinity Trends

Trends are the same as for electronegativity.

Exceptions occur at ½ full and full sublevels.

Units are in kJ/mole

Li Be B C N O F Ne

-60 >0 -27 -122 0 -141 -328 >>0

Practice

Which atom is larger H or He?

Which atom has a greater ionization energy, Ca or Sr?

Which atom is more electronegative, F or Cl?

Which atom has the larger radius?

Be or Ba

Ca or Br

Ba

Ca

Examples

Which atom has the higher 1st I.E.?

N or Bi

Ba or Ne

N

Ne

Examples

Which particle has the larger radius?

S or S2-

Al or Al3+

S2-

Al

Examples