PeriodicityChapter 8, Sections 5-7

Periodic

Repeats periodically

Forms a pattern, like… Day/Night Cycle of the moon

TENDS to go up or go down

Oxidation and Reduction

Here’s a periodic property you’ve already seen…

Oxidation – The loss of electrons Which elements tend to lose electrons?

Reduction – The gain of electrons Which elements tend to gain electrons?

metals

nonmetals

So metals oxidize the best and are the best reducing agents.

So nonmetals reduce the best and are the best oxidizing agents.

Atomic Radius

Commonly known as covalent radius

This is 1/2 the distance between two nuclei of the same elements that are covalently bonded.

Atomic Radius

Notice, what happens to the atomic size going down a group?

Why? (What’s occurring within the atoms’ structures?)There are more electrons.

There are more protons.There are more energy levels.

Higher energy levels are further from the nucleus, so the atom gets larger.

Atomic Radius

Notice, what happens to the atomic size going across a period?

Why? (What’s occurring within the atoms’ structures?)There are more electrons.There are more protons.The outer energy level DOES NOT CHANGE!!!

The outer energy level does not force the atom to get larger.The increased attraction between e– and p causes the atom to get smaller.

Atomic Radius

Example:Refer to a

periodic table and arrange the following in order of increasing atomic radius: Br, Se, Te.

Br

Te

Se

Te has to be the largest since it

is in the highest energy level.

Se and Br will be similar in size

because they are in the same

energy level. However, Se will be larger than Br

because it has fewer electrons

and protons attracting each

other.

Br < Se < Te

Ionic Radius v. Atomic Radius

What happens to the atom’s size when it turns into an ion?

If a positive ion is formed, what happens to the electrons?

Often, when losing electrons, the outer energy level is lost as well.

In addition, the number of electrons is now less than the number of protons. So the nuclear attraction is stronger.

Ionic Radius v. Atomic Radius

When a negative ion is formed, what happens to the electrons?

When an atom gains an electron, the number of electrons is now higher than the number of protons in the nucleus.

When there are more electrons than the protons, the nucleus can not attract the electrons as well. Because its attractions are weaker, the atom gets larger.

Ionic Radius v. Atomic Radius

Ionic Radius v. Atomic Radius

In summary… If the number of electrons becomes lower

than the number of protons, the nuclear attraction becomes stronger = the atom gets smaller

If the number of electrons is greater than the number of protons, the nucleus can not attract the electrons as well so the nuclear attractions are weaker = the atom gets larger.

Ionization Energy

When an atom becomes positively charged, it absorbs energy. Breaking attractions = endothermic

This energy is called an ionization energy (IE).The first ionization energy is the

amount of energy to remove the first electron from an atom.

Energy + M M+ + e–

IE

Ionization Energy

What happens to the ionization energy going down a group? Why?

What’s going on with the size of the atom?

Ionization Energy

What happens to the ionization energy going across a period? Why?

What’s going on with the size of the atom?

Ionization Energy

Energy + M M+ + e–

The first ionization energy… becomes smaller as the atomic radius

gets larger, i.e. going down a group. This is because there are fewer attractions

between the nucleus and the outermost electrons so less energy is required to remove the electron.

Ionization Energy

Energy + M M+ + e–

The first ionization energy… becomes larger as the atomic radius gets

smaller, i.e. going across a period. This is because the attractions between the

nucleus and the outermost electrons are stronger so more energy is required to remove the electron.

Ionization Energy

Energy + M M+ + e–

In summary, the greater the attraction for the electron, the more endothermic the ionization energy

Ionization Energy

There are also successive ionization energies.Electrons that are removed after having

already taking off electrons are create successive ionization energy.

The ionization energy (IE) number will be indicated with a subscript (IEi).

IE1+ M M+ + e–

IE2+ M+ M2+ + e–

IE3+ M2+ M3+ + e–

Ionization Energy

Element

IE1 IE2 IE3 IE4 IE5 IE6 IE7

Na 498 4560 6910 9540 13400 16600 20100

Mg 736 1445 7730 10600 13600 18000 21700

Al 577 1815 2740 11600 15000 18310 23290

Si 787 1575 3220 4350 16100 1/900 23800

P 1063

1890 2905 4950 6270 21200 25400

S 1000

2260 3375 4565 6950 8490 27000

Cl 1255

2295 3850 5160 6560 9360 11000

Ar 1519

2665 3945 5770 7230 8780 12000For each element, where are the most distinct jumps in energy?

Ionization Energy

Element

IE1 IE2 IE3 IE4 IE5 IE6 IE7

Na 498 4560 6910 9540 13400 16600 20100

Ionization Energy

Element

IE1 IE2 IE3 IE4 IE5 IE6 IE7

Na 498 4560 6910 9540 13400 16600 20100

Why does it take almost nine times the amount of energy as the first ionization energy?

Ionization Energy

Element

IE1 IE2 IE3 IE4 IE5 IE6 IE7

Mg 736 1445 7730 10600 13600 18000 21700

In order to make a magnesium +2 ion, 2 electrons must be lost…

IE1 + IE2 = 2181 kJ

Ionization Energy

Element

IE1 IE2 IE3 IE4 IE5 IE6 IE7

Mg 736 1445 7730 10600 13600 18000 21700

Why does it take so much energy to take off a 3rd electron?

IE3 = 7730kJ

Ionization Energy

Example:Refer to a

periodic table and arrange the following in order of increasing ionization energy:

As, Br, Sb.

Br

Sb

As

Sb has to have the smallest ionization energy since its outer energy level is the

furthest away.

As and Br will be similar in size

(and IE) because they are in the same energy

level. However, As will be larger in size than Br so

it will have a lower IE than Br.

Sb < As < Br

Electron Affinity

When an atom becomes negatively charged (gains an electron, it releases energy. Forming attractions = exothermic

This energy is called electron affinity (EA).

e– + X X– + EnergyEA

Electron Affinity

Electron AffinityNotice that the alkaline earth metals would need to add a subshell to hold another electron. Creating a higher energy subshell would be an endothermic process so gaining an electron won’t occur.

Electron Affinity

Notice that the noble gases would also need to add a subshell to hold another electron. Creating a higher energy subshell would be an endothermic process so gaining an electron won’t occur.

Electron Affinity

The halogens, on the other hand, can use the added electron to complete the subshell. This is a highly exothermic process so gaining an electron is very likely.

Electron Affinity

In general, what is the trend for electron affinities headed across the periods?

Electron Affinity

Going down a group, why does the electron affinity magnitude become smaller?

Electron Affinity

e– + X X– + Energy

In summary, the greater the attraction for the electron, the more exothermic the electron affinity.

Electronegativity

Electronegativity is the measure of the tendency for an atom to attract an electron.

The measure of electronegativity occurs on a scale.

0Not

likely to attract

an electron

4.00Very

likely to attract

an electron

Electronegativity

Metallic Character

Metallic character includes all of the properties of metals. Conductivity of electricity Conductivity of heat Luster Ductility Malleability Reactivity with water Reactivity with acids

Metallic Character

The properties of metals are

created by their bonds…

metallic bonds which

are produced when the electron

clouds of the atoms fuse together to

make an electron sea.

Metallic Character

What is the trend for metallic character?

Explanations

Going down the periodic table… As the principle quantum number

(energy level) increases, the nuclear attractions to the outermost electrons…

decreases

Explanations

Going down the periodic table… As the principle quantum number

(energy level) increases, the nuclear attractions to the outermost electrons…

decreases

As more energy levels fall in between the nucleus and the outermost electrons they shield (hinder) the nuclear attractions to those electrons…

the shielding effect

Explanations

Going across the periodic table… The number of electrons and protons

increases while the energy level stays the same…

This increases the attractions to the nucleus

Explanations

Comparing one subshell to another subshell in the same energy level… A full subshell will

shield another subshell from nuclear attractions, making the nuclear attractions weaker.

A higher energy subshell is further from the nucleus, so the nuclear attractions are weaker.

Explanations

Comparing paired v. unpaired electrons of the same subshell… UNLIKE comparing

one subshell to another subshell, the amount of shielding remains the same.

So what happens when two electrons share the same space?

Explanations

Comparing a charged atom to a neutral atom… A neutral atom has

the same number of e– as p.

A positive ion has                e– than p.

This causes the nuclear attractions to be significantly greater.

fewer

Explanations

Comparing a charged atom to a neutral atom… A neutral atom has

the same number of e– as p.

A negative ion has                e– than p.

This causes the nuclear attractions to be significantly weaker.

more

Explanations

Going down a group

Principle quantum number

Or Shielding effect

Create weaker attractions

Going across a period

Same energy level but greater attractions

between p and e–.

Two different subshells

Shielding effect from inner

subshell creating weaker attractions

Same subshell

Unpaired e–’s v. Paired e–’s

Paired e–’s repel

Charged atoms

Negative ion

More e–’s than creating weaker

attractions

Positive ion

Fewer e–’s than p creating stronger

attractions