Periodicity Chapter 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 So metals oxidize the best and are the best reducing agents. nonmetals 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 Br < Se < Te Example: Refer to a periodic table and arrange the following in order of increasing atomic radius: Br, Se, Te. 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. Se Br Te Te has to be the largest since it is in the highest energy level. 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? In addition, the number of electrons is now less than the number of protons. So the nuclear attraction is stronger. Often, when losing electrons, the outer energy level is lost as well.

Ionic Radius v. Atomic Radius When a negative ion is formed, what happens to the electrons? 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. When an atom gains an electron, the number of electrons is now higher than the number of protons in the nucleus.

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 For each element, where are the most 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 12000 For 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 IE1 + IE2 = 2181 kJ Element IE1 IE2 IE3 IE4 IE5 IE6 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 IE3 = 7730kJ 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 Sb < As < Br Example: Refer to a periodic table and arrange the following in order of increasing ionization energy: As, Br, Sb. 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. As Br Sb Sb has to have the smallest ionization energy since its outer energy level is the furthest away. 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– + Energy EA

Electron Affinity

Electron Affinity Notice 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. Not likely to attract an electron 4.00 Very 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. fewer This causes the nuclear attractions to be significantly greater.

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. more This causes the nuclear attractions to be significantly weaker.

Explanations Going down a group Going across a period Charged atoms 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