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Periodicity Chapter 8, Sections 5-7. Periodic Repeats periodically Repeats periodically Forms a pattern, like… Forms a pattern, like… Day/Night Day/Night.

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Presentation on theme: "Periodicity Chapter 8, Sections 5-7. Periodic Repeats periodically Repeats periodically Forms a pattern, like… Forms a pattern, like… Day/Night Day/Night."— Presentation transcript:

1 Periodicity Chapter 8, Sections 5-7

2 Periodic Repeats periodically Repeats periodically Forms a pattern, like… Forms a pattern, like… Day/Night Day/Night Cycle of the moon Cycle of the moon TENDS to go up or go down TENDS to go up or go down

3 Oxidation and Reduction Here’s a periodic property you’ve already seen… Here’s a periodic property you’ve already seen… Oxidation – The loss of electrons Oxidation – The loss of electrons Which elements tend to lose electrons? Which elements tend to lose electrons? Reduction – The gain of electrons Reduction – The gain of electrons Which elements tend to gain 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.

4 Atomic Radius Commonly known as covalent radius Commonly known as covalent radius This is 1 / 2 the distance between two nuclei of the same elements that are covalently bonded. This is 1 / 2 the distance between two nuclei of the same elements that are covalently bonded.

5 Atomic Radius Notice, what happens to the atomic size going down a group? Notice, what happens to the atomic size going down a group? Why? (What’s occurring within the atoms’ structures?) 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.

6 Atomic Radius Notice, what happens to the atomic size going across a period? Notice, what happens to the atomic size going across a period? Why? (What’s occurring within the atoms’ structures?) 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.

7 Atomic Radius Example: Example: Refer to a periodic table and arrange the following in order of increasing atomic radius: Br, Se, Te. BrTeSe 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.

8 Ionic Radius v. Atomic Radius What happens to the atom’s size when it turns into an ion? What happens to the atom’s size when it turns into an ion? If a positive ion is formed, what happens to the electrons? 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.

9 Ionic Radius v. Atomic Radius When a negative ion is formed, what happens to the electrons? 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.

10 Ionic Radius v. Atomic Radius

11 In summary… 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 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. 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.

12 Ionization Energy When an atom becomes positively charged, it absorbs energy. When an atom becomes positively charged, it absorbs energy. Breaking attractions = endothermic Breaking attractions = endothermic This energy is called an ionization energy (IE). 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. The first ionization energy is the amount of energy to remove the first electron from an atom. Energy + M  M + + e – IE

13 Ionization Energy What happens to the ionization energy going down a group? Why? What’s going on with the size of the atom?

14 Ionization Energy What happens to the ionization energy going across a period? Why? What’s going on with the size of the atom?

15 Ionization Energy Energy + M  M + + e – The first ionization energy… The first ionization energy… becomes smaller as the atomic radius gets larger, i.e. going down a group. 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. This is because there are fewer attractions between the nucleus and the outermost electrons so less energy is required to remove the electron.

16 Ionization Energy Energy + M  M + + e – The first ionization energy… The first ionization energy… becomes larger as the atomic radius gets smaller, i.e. going across a period. 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. This is because the attractions between the nucleus and the outermost electrons are stronger so more energy is required to remove the electron.

17 Ionization Energy Energy + M  M + + e – In summary, the greater the attraction for the electron, the more endothermic the ionization energy In summary, the greater the attraction for the electron, the more endothermic the ionization energy

18 Ionization Energy There are also successive ionization energies. There are also successive ionization energies. Electrons that are removed after having already taking off electrons are create successive ionization energy. 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 (IE i ). The ionization energy (IE) number will be indicated with a subscript (IE i ). IE 1 + M  M + + e – IE 2 + M +  M 2+ + e – IE 3 + M 2+  M 3+ + e –

19 Ionization Energy ElementIE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 Na498456069109540134001660020100 Mg7361445773010600136001800021700 Al5771815274011600150001831023290 Si787157532204350161001/90023800 P106318902905495062702120025400 S10002260337545656950849027000 Cl12552295385051606560936011000 Ar15192665394557707230878012000 For each element, where are the most distinct jumps in energy?

20 Ionization Energy ElementIE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 Na498456069109540134001660020100

21 Ionization Energy ElementIE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 Na498456069109540134001660020100 Why does it take almost nine times the amount of energy as the first ionization energy?

22 Ionization Energy ElementIE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 Mg7361445773010600136001800021700 In order to make a magnesium +2 ion, 2 electrons must be lost… IE 1 + IE 2 = 2181 kJ

23 Ionization Energy ElementIE 1 IE 2 IE 3 IE 4 IE 5 IE 6 IE 7 Mg7361445773010600136001800021700 Why does it take so much energy to take off a 3 rd electron? IE 3 = 7730kJ

24 Ionization Energy Example: Example: Refer to a periodic table and arrange the following in order of increasing ionization energy: As, Br, Sb. BrSbAs 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.

25 Electron Affinity When an atom becomes negatively charged (gains an electron, it releases energy. When an atom becomes negatively charged (gains an electron, it releases energy. Forming attractions = exothermic Forming attractions = exothermic This energy is called electron affinity (EA). This energy is called electron affinity (EA). e – + X  X – + Energy EA

26 Electron Affinity

27 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.

28 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.

29 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.

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

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

32 Electron Affinity e – + X  X – + Energy In summary, the greater the attraction for the electron, the more exothermic the electron affinity. In summary, the greater the attraction for the electron, the more exothermic the electron affinity.

33 Electronegativity Electronegativity is the measure of the tendency for an atom to attract an electron. Electronegativity is the measure of the tendency for an atom to attract an electron. The measure of electronegativity occurs on a scale. The measure of electronegativity occurs on a scale. 0 Not likely to attract an electron 4.00 Very likely to attract an electron

34 Electronegativity

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

36 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.

37 Metallic Character What is the trend for metallic character? What is the trend for metallic character?

38 Explanations Going down the periodic table… Going down the periodic table… As the principle quantum number (energy level) increases, the nuclear attractions to the outermost electrons… As the principle quantum number (energy level) increases, the nuclear attractions to the outermost electrons… decreases

39 Explanations Going down the periodic table… Going down the periodic table… As the principle quantum number (energy level) increases, the nuclear attractions to the outermost electrons… 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… 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

40 Explanations Going across the periodic table… Going across the periodic table… The number of electrons and protons increases while the energy level stays the same… The number of electrons and protons increases while the energy level stays the same… This increases the attractions to the nucleus

41 Explanations Comparing one subshell to another subshell in the same energy level… 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 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. A higher energy subshell is further from the nucleus, so the nuclear attractions are weaker.

42 Explanations Comparing paired v. unpaired electrons of the same subshell… Comparing paired v. unpaired electrons of the same subshell… UNLIKE comparing one subshell to another subshell, the amount of shielding remains the same. UNLIKE comparing one subshell to another subshell, the amount of shielding remains the same. So what happens when two electrons share the same space? So what happens when two electrons share the same space?

43 Explanations Comparing a charged atom to a neutral atom… Comparing a charged atom to a neutral atom… A neutral atom has the same number of e – as p. A neutral atom has the same number of e – as p. A positive ion has e – than p. A positive ion has e – than p. This causes the nuclear attractions to be significantly greater. This causes the nuclear attractions to be significantly greater. fewer

44 Explanations Comparing a charged atom to a neutral atom… Comparing a charged atom to a neutral atom… A neutral atom has the same number of e – as p. A neutral atom has the same number of e – as p. A negative ion has e – than p. A negative ion has e – than p. This causes the nuclear attractions to be significantly weaker. This causes the nuclear attractions to be significantly weaker. more

45 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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