A. Periodic Trends – Understanding and being able to model the structure of atoms helps us understand the properties of the atoms. Fortunately, there are some trends across periods and down families/groups.

Atomic Radius – usually measured by dividing a covalent or metallic bond by 2. Down a group: The atomic radius INCREASES down a group.

(a) As you go down a group, the outer electrons are entering increasingly higher ENERGY LEVELS that are increasingly further from the nucleus. NOTE: because energy levels CONVERGE the difference in atomic radius between Rb and Cs is less that the difference between Li and Na. (b) When electrons are far from the nucleus, they do not experience the full attraction of the positive nucleus. (as distance , attraction ). Additionally, the inner CORE electrons SHIELD the outer VALENCE electrons from the positive nucleus.

VALENCE

Sb: 1s22s22p63s23p64s23d104p65s24d105p3

CORE

Picture from: 1

ESTIMATE OF EFFECTIVE NUCLEAR CHARGE: Zeff = Z – # core electrons

This is not exact, but it does provide a ballpark figure for the effect of the repulsion by inner, core electrons on the outer, valence electrons.

MODEL: Let’s take a look at Cs and K. First – calculate Zeff for each. K: Zeff = +19-18 = +1 Cs: Zeff = +55-54 = +1 When we combine shielding with effective nuclear charge we see that Cesium’s 6s1 electron is experiencing the same +1 nuclear charge as potassium’s 4s1 electron but from a much further distance!

1 http://wps.prenhall.com/wps/media/objects/602/616516/Chapter_05.html Across a period: Even though more electrons are being added across the group, they are being added to the same energy level, which in theory is roughly the same distance away. However, energy level n=2 for boron is not the same distance from the nucleus as n=2 for fluorine. Fluorine’s n=2 is CLOSER to the nucleus due to an increase in nuclear charge. All of the 2p electrons in fluorine are experiencing Zeff = 9-2 = +7. Boron’s 2p electron is experiencing Zeff = 5-2 = +3. They both have the same number of shielding electrons (2), but Z = 9 for fluorine and Z = 5 for boron. The greater nuclear charge pulls in the electrons, making fluorine’s atomic radius SMALLER than boron.

Ionic Radius Cations: Atomic radius ion < atomic radius of the atom. Electrons are lost in this oxidation process. For representative elements, typically all of the outer energy level electrons are lost. The outer electrons are now in a lower energy level.

Atom Most common ion Proton:electron ratio Na: 1s22s22p63s1 Na+: 1s22s22p6 11:10 Al: 1s22s22p63s23p1 Al3+: 1s22s22p6 13:10

Way cool! The ion’s Hey – I can predict that Na radius is smaller would become + 1 because because its outer then it would have a noble energy level is n=2 gas configuration! instead of n=3

Check out aluminum ion– it has the same W hatev...check out the p:e electron configuration ratio. Al3+ > Na+ which means as the sodium ion. the same number of electrons are They must be pulled in by a higher Z AND isoelectronic! Al3+ ionic radius is less than Na+!

Anions – Atomic radius ion > atomic radius of the atom. More electrons, p:e ratio decreases, radius increases due to the electron electron repulsion of the added electron.

U do it 1. (a) Place the following atoms in order of increasing atomic radius: Br, As, Rb, Ca. Explain your rationale.

(b) You are packing for a trip to Venus. Venusians use ions to barter, so you need to bring as many ions with you as possible, but you are limited by your baggage space. Place the following ions in increasing Venusian value for your trip: S2─, Ca2+, K+, P3─, Cl─.

Ionization energy – Judging from the name, I hope you can see that this is the energy required to form an ion. In this case, we are referring to forming CATIONS through the successive removal of electrons. The reactant can be either an atom or an ion, but it must be in the GASEOUS state. The first ionization energy involves the first electron; the second ionization involves the second and so on.

Picture2 M(g)  M+(g) + e─ 1st ionization energy M+(g)  M2+(g) + e─ 2nd ionization energy M2+(g)  M3+(g) + e─ 3rd ionization energy

Electrons are removed from the highest energy level first – not the last filled. The ionization process is ENDOTHERMIC so the ionization energies are positive and typically measured in KJ/mol. The 1st I.E < 2nd I.E < 3rd I.E etc. The difference between successive IONIZATION OF electrons is maybe double or triple. The difference in energy between the last valence electron and the first core electron is typically five times. ionization energy increases LEFT TO RIGHT ACROSS a period and UP a group. Remember – fluorine loves electrons more than anything – so point in the direction of F and you will be heading on the right path! As you move across a period, the number of protons increases, thus there is an increase in proton-electron attraction which in turn results in an increase in the amount of energy needed to pluck that little bugger off!

Check out the following figure to compare some of the terms we have discussed so far3:

U do it 2. What are the differences in the electron configuration for each of the following states of aluminum?

Ground state atom: 1s22s22p63s23p1 Excited State atom: 1s22s22p63s13p2 Ion: 1s22s22p6

Electron Affinity – This is the energy involved in adding an electron to an element or ion in its gaseous state. For some elements this is an ENDOTHERMIC UNFAVORABLE process and for others (like fluorine!) it is a EXOTHERMIC FAVORABLE process. There is some confusion with this one! The more NEGATIVE the electron affinity value, the greater the actual AFFINITY for an electron.

2 http://www.800mainstreet.com/4/0004-000-IE.GIF 3 http://earth.ast.smith.edu/james/a111/lectures/figs/04-09.jpg The general trend follows ionization energy. Electron affinity becomes more NEGATIVE (more favorable) UP a group because electrons are added to energy levels closer to the nucleus with less shielding. Electron affinity becomes more NEGATIVE (more favorable) LEFT TO RIGHT ACROSS a period because there are more protons, electrons are added to the same energy level. The added attraction makes it favorable to add an electron (and unfavorable to remove one – hence a higher ionization energy) Unlike ionization energy that continues to increase to the noble gases, electron affinity becomes ZERO (or positive) when you get to the noble gases. They simply do not want electrons!

Electronegativity – Linus Pauling and others came up with the concept of electronegativity to help with bonding models. Electron affinity is the measure of the energy involved in adding an electron to an atom or ion in the gaseous state. Electronegativity is a relative value that measures the ability of an atom to attract electrons in a chemical bond. Fluorine is the MOST electronegative element and francium/cesium is the LEAST Picture4

4 http://tptc.iit.edu/Center/research/PhaseDiagram/Content/periodic table/electronegativity.jpg