Trends the Periodic Table Trends more than 20

Trends & the Periodic Table

Trends • more than 20 properties change in predictable way based location of elements on Periodic Table some properties: – Density – melting point/boiling point – atomic radius – ionization energy – electronegativity

Going down column 1: Period Element Configuration 1 H 1 2 Li 2 -1 3 Na 2 -8 -1 4 K 2 -8 -8 -1 5 Rb 2 -8 -18 -8 -1 6 Cs 2 -8 -18 -18 -8 -1 7 Fr 2 -8 -18 -32 -18 -8 -1 increasing # energy levels as go down

Increasing number of energy levels

Atomic Radius • Atomic radius: defined as ½ distance between neighboring nuclei in molecule or crystal • Affected by 1. # of energy levels 2. Proton Pulling Power

Increasing Atomic Radius Increasing number of energy levels

Cs has more energy levels, so it’s bigger previous | index | next Li: Group 1 Period 2 Cs: Group 1 Period 6

As we go across, elements gain electrons, but they are getting smaller! Family IA or 1 IIA or 2 IIIA or 13 IVA or 14 VA or 15 VIA or 16 VIIA or 17 VIIIA or 18 Element Li Be B C N O F Ne Configuration 2 -1 2 -2 2 -3 2 -4 2 -5 2 -6 2 -7 2 -8

Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius

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Why does this happen. . • Moving from left to right, more protons are added (the atomic number increases) • The nucleus has greater “proton pulling power” – Remember the nucleus is + and the electrons are - so they get pulled towards the nucleus • The more protons your have, the more Proton Pulling Power

Why does this happen. . • Moving from left to right, more electrons are added too, but NOT TO ADDITIONAL ENERGY LEVELS. • This means there is not extra space between nucleus and valence electrons.

Why does this happen. . • The greater proton pulling power causes the electrons to be pulled closer to the nucleus and the radius decreases.

previous | index | next as go across row size tends to decrease a bit because of greater PPP “proton pulling power”

SAMPLE PROBLEM 8. 3 PROBLEM: Using only the periodic table (not Figure 8. 15)m rank each set of main group elements in order of decreasing atomic size: (a) Ca, Mg, Sr PLAN: Ranking Elements by Atomic Size (b) K, Ga, Ca (c) Br, Rb, Kr (d) Sr, Ca, Rb Elements in the same group decrease in size as you go up; elements decrease in size as you go across a period. SOLUTION: (a) Sr > Ca > Mg These elements are in Group 2 A(2). (b) K > Ca > Ga These elements are in Period 4. (c) Rb > Br > Kr Rb has a higher energy level and is far to the left. Br is to the left of Kr. (d) Rb > Sr > Ca Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Space between nucleus and electrons

Francium's one valance electron barely feels the proton pull from the nucleus. It is very far away from the nucleus and has 86 other electrons between it and the nucleus. No wonder it will lose its one electron the easiest. No wonder it’s the most reactive metal

Ionization Energy • = amount energy required to remove a valence electron from an atom in gas phase • 1 st ionization energy = energy required to remove the first e- [farthest from nucleus]

previous | index | next • Cs valence electron lot farther away from nucleus than Li • electrostatic attraction much weaker so easier to steal electron away from Cs • Li has a higher Ionization Energy then Cs

First ionization energies of the main-group elements Trends in the Periodic Table

Increased Ionization Energy (harder to remove an electron) Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreasing Atomic Radius Decreased Ionization Energy (easier to remove an electron)

SAMPLE PROBLEM 8. 4 PROBLEM: Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE 1: (a) Kr, He, Ar PLAN: Ranking Elements by First Ionization Energy (b) Sb, Te, Sn (c) K, Ca, Rb (d) I, Xe, Cs IE increases as you proceed up in a group; IE increases as you go across a period. SOLUTION: (a) He > Ar > Kr Group 8 A(18) - IE decreases down a group. (b) Te > Sb > Sn Period 5 elements - IE increases across a period. (c) Ca > K > Rb Ca is to the right of K; Rb is below K. (d) Xe > I > Cs I is to the left of Xe; Cs is further to the left and down one period.

Electronegativity • ability of atom to attract electrons in bond • noble gases tend not to form bonds [octet rule], so don’t have electronegativity values • Unit = Pauling • Fluorine: most electronegative element = 4. 0 Paulings

Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Increased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius Decreased Electronegativity

Reactivity of Metals • judge reactivity of metals by how easily give up electrons • Why do they lose electrons? Sodium Atom: 1 s 22 p 63 s 1

Reactivity of Metals • judge reactivity of metals by how easily give up electrons Sodium Cation: 1 s 22 p 6

Increasing number of energy levels Increasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Decreased Electronegativity Increased Ionization Energy (harder to remove an electron) Decreasing Atomic Radius More metallic Most reactive metal = Fr (the most metallic)

Reactivity of Non-metals • judge reactivity of non-metals by how easily gain electrons • Why do they gain? Chlorine Atom: 1 s 22 p 63 s 23 p 5

Reactivity of Non-metals • judge reactivity of non-metals by how easily gain electrons • Why do they gain? Chlorine Anion: 1 s 22 p 63 s 23 p 6

Increased Ionization Energy (harder to remove an electron) Increasing number of energy levels Increasing Atomic Radius Decreasing Atomic Radius Increased Electron Shielding Decreased Ionization Energy (easier to remove an electron) Decreased Electronegativity Increased Electronegativity More metallic Most reactive metal = Fr (the most metallic) Most Reactive Nonmetal =F Nonreactive BACK

How do you know if an atom gains or loses electrons? • Atoms form ions to get a valence of 8 (or 2 for H) • Metals tend to have 1, 2, or 3 valence electrons – It’s easier to lose them • Nonmetals tend to have 5, 6, or 7 valence electrons – It’s easier to add some • Noble gases already have 8 so they don’t form ions very easily. • The easier an atom will lose an electron, the MORE METALLIC it is.

Trends in metallic behavior

Positive ions (cations) • Formed by loss of electrons • Cations always smaller than neutral atom • Strong proton pull since more protons than electrons. 2 e 8 e 8 e 2 e Ca Ca 8 e 8 e 2 e Ca+2

Negative ions or (anions) • • • Formed by gain of electrons Anions always larger than neutral atom Electrons repel each other and create space between each other. • Weaker proton pull since less protons than electrons

Ionic vs. atomic radius

Ranking Ions by Size PROBLEM: Rank each set of ions in order of decreasing size, and explain your ranking: (a) Ca 2+, Sr 2+, Mg 2+ PLAN: (b) K+, S 2 -, Cl - (c) Au+, Au 3+ Compare positions in the periodic table, formation of positive and negative ions and changes in size due to gain or loss of electrons. SOLUTION: (a) Sr 2+ > Ca 2+ > Mg 2+ (b) S 2 - > Cl - > K+ (c) Au+ > Au 3+ These are members of the same Group (2 A/2) and therefore decrease in size going up the group. The ions are isoelectronic; S 2 - has the smallest proton pull and therefore is the largest while K+ is a cation with a large proton pull is the smallest. The higher the + charge, the smaller the ion.