2 Periodic Repeats periodically Forms a pattern, like… Day/NightCycle of the moonTENDS to go up or go down
3 Oxidation and Reduction Here’s a periodic property you’ve already seen…Oxidation – The loss of electronsWhich elements tend to lose electrons?Reduction – The gain of electronsWhich elements tend to gain electrons?metalsSo metals oxidize the best and are the best reducing agents.nonmetalsSo nonmetals reduce the best and are the bestoxidizing agents.
4 Atomic Radius Commonly known as covalent radius This is 1/2 the distance between two nuclei of the same elements that are covalently bonded.
5 Atomic RadiusNotice, 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.
6 Atomic RadiusNotice, 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.
7 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.SeBrTeTe has to be the largest since itis in the highest energy level.Br < Se < Te
8 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.
9 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.
11 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 smallerIf 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 EnergyWhen an atom becomes positively charged, it absorbs energy.Breaking attractions = endothermicThis 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
13 Ionization EnergyWhat happens to the ionization energy going down a group?Why?What’s going on with the size of the atom?
14 Ionization EnergyWhat 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… 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.
16 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.
17 Ionization EnergyEnergy + M M+ + e–In summary, the greater the attraction for the electron, the more endothermic the ionization energy
18 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+ M e–IE3+ M2+ M e–
19 Ionization Energy For each element, where are the most IE1IE2IE3IE4IE5IE6IE7Na498456069109540134001660020100Mg7361445773010600136001800021700Al5771815274011600150001831023290Si787157532204350161001/90023800P106318902905495062702120025400S10002260337545656950849027000Cl12552295385051606560936011000Ar15192665394557707230878012000For each element, where are the mostdistinct jumps in energy?
20 Ionization Energy Element IE1 IE2 IE3 IE4 IE5 IE6 IE7 Na 498 4560 6910 9540134001660020100
21 Ionization Energy Element IE1 IE2 IE3 IE4 IE5 IE6 IE7 Na 498 4560 6910 9540134001660020100Why does it take almost nine times the amount of energy as the first ionization energy?
22 Ionization Energy IE1 + IE2 = 2181 kJ Element IE1 IE2 IE3 IE4 IE5 IE6 Mg7361445773010600136001800021700In order to make a magnesium +2 ion, 2 electrons must be lost…IE1 + IE2 = 2181 kJ
23 Ionization Energy IE3 = 7730kJ Element IE1 IE2 IE3 IE4 IE5 IE6 IE7 Mg 7361445773010600136001800021700Why does it take so much energy to take off a 3rd electron?IE3 = 7730kJ
24 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.AsBrSbSb has to have the smallest ionization energy since its outer energy level is the furthest away.Sb < As < Br
25 Electron AffinityWhen an atom becomes negatively charged (gains an electron, it releases energy.Forming attractions = exothermicThis energy is called electron affinity (EA).e– + X X– + EnergyEA
27 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.
28 Electron AffinityNotice 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 AffinityThe 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 AffinityIn general, what is the trend for electron affinities headed across the periods?
31 Electron AffinityGoing down a group, why does the electron affinity magnitude become smaller?
32 Electron Affinitye– + X X– + EnergyIn summary, the greater the attraction for the electron, the more exothermic the electron affinity.
33 ElectronegativityElectronegativity 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 electron4.00Very likely to attract an electron
35 Metallic CharacterMetallic character includes all of the properties of metals.Conductivity of electricityConductivity of heatLusterDuctilityMalleabilityReactivity with waterReactivity with acids
36 Metallic CharacterThe 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 CharacterWhat is the trend for metallic character?
38 Explanations Going down the periodic table… As the principle quantum number (energy level) increases, the nuclear attractions to the outermost electrons…decreases
39 Explanations Going down the periodic table… As the principle quantum number (energy level) increases, the nuclear attractions to the outermost electrons…decreasesAs 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… The number of electrons and protons increases while the energy level stays the same…This increases the attractions to the nucleus
41 ExplanationsComparing 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.
42 ExplanationsComparing 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?
43 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.fewerThis causes the nuclear attractions to be significantly greater.
44 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.moreThis causes the nuclear attractions to be significantly weaker.
45 Explanations Going down a group Going across a period Charged atoms Principle quantum numberOrShielding effectCreate weaker attractionsGoing across a periodSame energy level but greater attractionsbetween p and e–.Two different subshellsShielding effect from inner subshell creating weaker attractionsSame subshellUnpaired e–’s v. Paired e–’sPaired e–’s repelCharged atomsNegative ionMore e–’s than creating weaker attractionsPositive ionFewer e–’s than p creating stronger attractions