Electrons Review and Electrons Review and

Periodic Table TrendsPeriodic Table Trends

Unit 7Unit 7

ElectronsElectrons

Shapes of the orbitals

Electron ConfigurationElectron Configuration

Electrons spin in opposite direction

BackgroundBackground

Electrons can jump between shells (Bohr’s model supported)

The electrons can be pushed so far that they escape the attraction of the nucleus

Losing an electron is called IONIZATION

Remember an ion is an atom that has either a net positive or net negative charge

Q: What would the charge be on an atom that lost an electron? Gained two electrons?

IONIZATION ENERGYIONIZATION ENERGY

o Defn: Ionization energy is the energy required to remove

one outer electron from an atom

o When an electron is taken away, what kind of ion results?

o A positively charged ion (cation)o Review: Oxidation numbers

Metals form (+) or (-) ions?+

Nonmetals form (+) or (-) ions?-

FIRST IONIZATION ENERGY - the energy required to pull off the

first valence electron.

SECOND IONIZATION ENERGY - the energy required to pull off the

second valence electron.

Low IE = not a lot of energy required = cations formed easily

High IE = takes a lot of energy = cations NOT formed; anions

What has a higher ionization energy – a metal or nonmetal?

Nonmetal – Why?

Tend to form anions

Ionization Energy cont.Ionization Energy cont.

a) Moving left to right across the periodic table IONIZATION

ENERGY INCREASES (harder to pull off an electron)

WHY ? More protons are in the nucleus therefore the valence

electrons are strongly attracted to the nucleus which

increases the energy required to remove them.

b) Moving down a group, IONIZATION ENERGY DECREASES

(easier to pull off an electron)WHY ? Shielding effect

Shielding - core e- block the attraction between the nucleus and the valence e-

c) The second ionization energy is greater, third is even greater…

WHY ? Electrons that remain move closer to the nucleus.

IONIZATION ENERGYIONIZATION ENERGY

The ionization energy is the amount of energy needed to strip an electron off of an atom , ion, or molecule. The illustration shows that the metals like lithium, Li, and cesium, Cs, have relatively low ionization energies. This means it takes relatively small amounts of energy to remove electrons from these atoms. Metals tend to lose electrons and form positive ions.The nonmetals like neon, Ne, fluorine, F, and oxygen, O, have relatively high ionization energies. This indicates that the nonmetals have strong attractions for their valence electrons. The nonmetals hold on to their electrons. In fact nonmetals gain electrons to form negative ions.

Ca

K

H

He

Li

B

Be C

NO

F

Ne

Na

MgAl

Si

P S

Cl

Ar

0

500

1000

1500

2000

2500

0 2 4 6 8 10 12 14 16 18 20

Element

Ion

izati

on

en

erg

y (

kJ/m

ol)

Factors Affecting Ionization EnergyFactors Affecting Ionization Energy

ATOMIC RADIUSATOMIC RADIUS

o The approximate distance from the nucleus of an atom to its

valence electrons.

a) Moving left to right across the periodic table

ATOMIC RADIUS DECREASES

WHY ? More protons are being added to the nucleus, valence

electrons are strongly attracted to the nucleus. Electrons

are also being added, but in the same shell at about the

same distance so there is not much of a shielding effect.

b) Moving down a group,

ATOMIC RADIUS INCREASES

WHY ?

Each atom has another energy level, so the atoms get bigger.

Since electrons are being added to distant shells away

from the nucleus the valence electrons are SHIELED

by the inner shell electrons.

Atomic Radius Cont.Atomic Radius Cont.

H

Li

Na

K

Rb

Ca

K

H He

Li

B

Be

C N O F

Ne

Na

MgAl Si P S Cl

Ar

0

50

100

150

200

250

0 2 4 6 8 10 12 14 16 18 20

Element

Ato

mic

Rad

ius (

pm

)

Review Review

The size of an atom is largely determined by its electrons. The electrons are arranged in shells surrounding the nucleus of each atom. The top elements of every group have only one or two electron shells. Atoms of elements further down the table have more shells and are therefore larger in size. Moving across a period from left to right, the outermost electron shell fills up but no new shells are added. At the same time, the number of protons in the nucleus of each atom increases. Protons attract electrons. The greater the number of protons present, the stronger the attraction that holds the electrons closer to the nucleus, and the smaller the size of the shells.

ELECTRONEGATIVITYELECTRONEGATIVITY

o Defn: How strongly the nucleus of an atom attracts the

electrons of other atoms in a bond.

o High electronegativity = wants to gain an electron = easy to

become a negative ion

o Given specific values and found on electronegativity tableDO NOT NEED TO MEMORIZE!

a) Moving from left to right on the periodic table

ELECTRONEGATIVITY INCREASES.

b) Moving down a group

ELECTRONEGATIVITY DECREASES.

Review of ElectronegativityReview of Electronegativity

Atoms of different elements have different attractions for bonding electrons.

Electronegativity is a measure of the attraction an atom involved in a bond has for the electrons of the bond.

An atom with a high electronegativity will tend to attract bonded electrons towards it.

An atom with a low electronegativity will have a very weak attraction for electrons.

Electronegativity values can be useful in predicting which type of bonding is most likely between two elements

Electronegativity cont.Electronegativity cont.

What element has the highest EN?

F

Which would have greater EN, a metal or a nonmetal?

Nonmetal

Which is more electronegative, Cu or S?

S

Br or Ga?

Br

Other Trends…Other Trends…

Melting Points

Metals

USUALLY decreases as you go down a group

Non-metals

USUALLY increases as you go down a group

1

2

3

4

5

6

7

Melting/Boiling Point

Highest in the middle of a period.

F. Melting/Boiling Point

Other Trends…Other Trends…

Reactivity

Defn: how likely or vigorously an atom is to react with other substances

Usually determined by how easily electrons can be removed (ionization energy) and how badly they want to take other atom’s electrons (electronegativity) since it is the transfer/interaction of electrons that is the basis of chemical reactions

Reactivity Cont.Reactivity Cont.

Metals Period: DECREASES from left to right

Group: INCREASES down a group

WHY? The farther left and down the periodic table, the easier it is for elections to be given or taken away, resulting in a higher reactivity

Non-metals Period: INCREASES from left to right

Group: DECREASES down a group

WHY? The farther right and up the periodic table, the higher the electronegativity, resulting in a more vigorous exchange of electrons

Periodic Trend SummaryPeriodic Trend Summary

Based upon an element’s position on the periodic table, it is possible to make predictions regarding its behavior.

All periodic trends can be understood in terms of three basic rules:

1.) Electrons are attracted to the protons in the nucleus of an atom

a) The closer an electron is to the nucleus, the more strongly it is attracted.

b) The more protons in a nucleus, the more strongly an electron is attracted.

Periodic Trend Summary Cont.Periodic Trend Summary Cont.

2.) Electrons are repelled by other electrons in an atom. So if other electrons are between a valance electron and the nucleus, the valence electrons will be less attracted to the nucleus.

This is called the SHIELDING EFFECT.

3.) Completed orbits, and sub-orbits, are very stable. Atoms prefer to add or subtract valence electrons to create complete shells if possible.

Review: IE vs. ENReview: IE vs. EN

A high IE tends to form a:

(positive, negative) ion

A high EN tends to form a:

(positive, negative) ion

Metals have a: (high, low) IE

Nonmetals have a: (high, low) EN