Ions

• Ions have a different number of electrons than the neutral atom

• The number of electrons in an ion is calculated from the normal number of electrons for the particular atom minus the charge on the ion

• Fe2+: 26 – (+2) = 24 electrons• F-: 9 – (-1) = 10 electrons

Electron Configurations of Ions

• Any ion can be made, and an electron configuration can then be experimentally determined

• Under normal conditions, the only ions made are ones that make the system “stable” and consume a minimum of energy in the process.

• “Stable” means the ion only undergoes actions that do not change its electronic state

• Main-group elements: noble gas configurations• Transition elements: more complicated

Electron Configurations of Ions

• Atoms become ions (under normal circumstances) to lower their electronic energy. No matter what else happens the electronic configuration that is achieved serves to lower the electronic energy as much as possible.

• That is, the actual electron configuration of an ion lowers the electronic energy of the system better than any other electron configuration.

Do we have to calculate electronic energies to figure out electron

configurations of ions?

• No! Since we have a (blissfully) naïve view of electronic energy, we’ll reduce the problem to a set of rules we can follow and (usually) not look stupid.

Ionic Electron Configurations: Anions

• Main-group nonmetals accept electrons into their lowest-energy unoccupied atomic orbital.

– N: 1s2 2s2 2p3 = [He] 2s2 2p3

– N3–: 1s2 2s2 2p6 = [He] 2s2 2p6

– O: 1s2 2s2 2p4 = [He] 2s2 2p4

– O2–: 1s2 2s2 2p6 = [He] 2s2 2p6

– F: 1s2 2s2 2p5 = [He] 2s2 2p5

– F–: 1s2 2s2 2p6 = [He] 2s2 2p6

Ionic Electron Configurations: Main-Group Metals

• Main-group metals donate electrons from the atom’s highest-energy occupied atomic orbital.

– Na: 1s2 2s2 2p6 3s2 = [Ne] 3s1 – Na+: 1s2 2s2 2p6 = [Ne]

– Mg: 1s2 2s2 2p6 3s2 = [Ne] 3s2 – Mg2+: 1s2 2s2 2p6 = [Ne]

– Al: 1s2 2s2 2p6 3s2 3p1 = [Ne] 3s2 3p1

– Al3+ 1s2 2s2 2p6 = [Ne]

Ionic Electron Configurations: Transition Metals

• Transition metals lose their valence-shell s-electrons before

losing their d-electrons. Electrons with the highest n-quantum

number are lost first. Wanna lose $1000 fast? Make a $1000 bet

that this rule is never broken.

– Fe: 1s2 2s2 3s2 3p6 4s2 3d6 = [Ar] 4s2 3d6

– Fe2+: 1s2 2s2 3s2 3p6 3d6 = [Ar] 3d6

– Fe3+: 1s2 2s2 3s2 3p6 3d5 = [Ar] 3d5

Wait a minute!

Scandium, for example, fills the 4s subshell before adding the 3d electron.

Doesn’t this imply that the 3d electron is more easily ionizable than the 4s electrons?Shouldn’t the 3d electron be removed first when making a cation?

No! Here are two reasons why:

1 – The “last electron in” for an atom is in its particular spot because a proton and an electron were added. Ionizing an atom just removes an electron. Thus you cannot expect one process to be the reverse of the other.

2 – The 3d electron of scandium actually is at a higher energy than the 4s. (Remember scandium’s PES.) So the 4s electrons really are the easiest electrons to remove.

Predicting IonsAny ion can be made. The ones that occur “naturally” are the ones that lower the electronic energy of the system the most.

Having only full shells or subshells of electrons (that is, a noble gas electron configuration) lowers the energy the most, so the ion with a full-shell electron configuration (that is achieved most easily) is the ion that generally occurs. This is most true for s-block and p-block elements.

d-block and f-block elements usually find many other stable electron configurations, so we will generally justify electron configurations we find experimentally rather than predict electron configurations for any particular element.

Ionic Radii• Cations are smaller than their parent atom.

• Anions are larger than their parent atom.

1 IE

2 IE, 3 IE, 4 IE, etc.

2 IE, 3 IE, 4 IE, etc.

2 IE, 3 IE, 4 IE, etc.

2 IE, 3 IE, 4 IE, etc.

Electron Affinity• Energy change that occurs when an electron is

added to an isolated atom in the gaseous state.

• Abbreviation is Eea.

• Value of Eea results from interplay of nucleus

electron attraction, and electron–electron repulsion.

• Values are generally negative, but with significant exceptions.

Electron Affinity Data

Ionic Bond Formation:Macroscopic View

Ionic Bond Formation:Microscopic ViewSodium Chloride (NaCl):Sodium Chloride (NaCl):

• Thermodynamics allows us to predict whether a reaction will or will not occur. (Chapter Eight)

• For this kind of reaction, the total energy change of the system must be significantly negative.

• For NaCl:

• Umm. . .wha?

Ionic Bond Formation:Energetics View

mol

kJ

mol

kJ

mol

kJClENaIE ea 2.1476.3488.495

The additional factor

• The additional factor that allows the compound to be formed is the added stability from the attraction of the cation to the anion. That is, the ionic bond.

• In effect, it is the lowering of energy that occurs when isolated cations and anions condense into a crystal together.

• It is called the lattice energy (U).• Lattice energy is largely a Coulombic effect.• For NaCl: U = +787 kJ/mol.

Lattice Energy

• Coulombic force:

• Lattice energy is positive because it is energy needed to break lattice, but we use the negative of this energy

• Gets bigger as charges increase• Gets bigger as ions get smaller

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zzkdFU 21

Predicting Ionic Compound Formation

• The total energy of reaction can be calculated with a Born-Haber cycle.

• We devise a path from reactants to products and determine the energy change of each step. The sum of these steps is the total energy change of the reaction.

• The reaction does NOT need to occur with these actual steps.

Born-Haber Cycle Example: NaCl

Born-Haber Cycle:Why not Na2+ and Cl2-?

Born-Haber Step +1/-1 ions +2/-2 ions

1. Sublimation of sodium +100 +100

2. Dissociation of chlorine

+125 +125

3. Ionization of sodium +500 +4500

4. Electron Affinity of Chlorine

-350 +100

5. Lattice Energy -800 -3200

Sum -425 +1625

Coulombic Lattice EnergiesThe Coulombic nature of the lattice energy means that, in order to have a strong ionic bond, you need small, highly-charged ions.

Octet Rule

One result of this Born-Haber treatment is to show that main-group elements engage in chemical reactions that provide it with eight valence electrons in the easiest fashion possible.

We will see this rule preserved even when we expand our bonding world to include three types of bonding. What will change is our concept of “valency.”

Thus group one elements tend to make +1 ions, group seven elements make -1 ions, etc.

“Ates” chart• Symbol O’s Charge• Chlorine (Cl) 3 1• Iodine (I) 3 1• Sulfur (S) 4 2• Nitrogen (N) 3 1• Carbon (C) 3 2• Silicon (Si) 4 4• Phosphorus (P) 4 3• Chromium (Cr) 4 2• Manganese (Mn) 3 1• Arsenic (As) 4 3• Bromine (Br) 4 1• Selenium (Se) 4 2• Tellurium (Te) 4 2• Boron (B) 3 3

So what does it all mean?

• Most of the similarities can be summed up in the octet rule

• Main-group elements undergo reactions that leave it with a noble gas electron conifguration.

• Main-group elements tend to take the easiest path to this configuration

• The tendency is stronger the lower the IE is and the higher the Eea is.

Why?