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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
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