Ions, Ionization E, Electron Affinity

1. Ions and their electron configuration
1. Why do ions form?  Atoms (except the noble gases) can lower their energy, thus becoming more stable, by gaining or losing electrons in order to have full s and p subshells.  The goal for all atoms is to have a full s and p subshell! NOT a full shell - the d and f don't need to be full to be stable.  Noble gases don't have full d or f subshells.  It is all about getting to an s²p⁶ configuration!!!  This configuration will lower energy and provide stability.  We often say the motivation or goal for the atoms is to "be like a noble gas."
2. Let's examine Na
1. Na - 1s²2s²2p⁶3s¹   not stable
2. Ne - 1s²2s²2p⁶    stable 
3. So sodium's goal in life is to lose that one outer electron to be like Ne thus Na atom becomes Na +1 ion since it now has 11 positive protons but only 10 negative electrons - the resulting charge is +1.
4. Na⁺  - 1s²2s²2p⁶    stable 
3. Let's examine Cl
1. Cl - 1s²2s²2p⁶3s²3p⁵
2. Ar - 1s²2s²2p⁶3s²3p⁶
3. So chlorine's goal is to gain one electron so it can be like Ar thus Cl atom becomes Cl -1 ion since it now has 17 protons but 18 electrons
4. Cl ^- -1s²2s²2p⁶3s²3p⁶
4. Following similar reasoning (atoms trying to be like noble gases with full s and p subshells) we can predict likely charges (called oxidation states) for all the elements except the transition metals.  THIS IS A MUST. Check out here and click on oxidation states for more information. You must learn the oxidation states for all the main group elements.
5. In general:
1. Metals tend to lose electrons and become cations. (positive ions)
2. Nonmetals tend to gain electrons and become anions (negative ions)
6. Transition metal ions electron configurations
1. When these metals lose electrons they come out of the s subshell BEFORE the d subshell.
2. Zn atom:  [Ar] 4s²3d¹⁰    but Zn²⁺ is [Ar] 3d¹⁰ because the 2 lost electrons were removed from the 4s subshell
3. Mn atom: [Ar] 4s²3d⁵   but Mn²⁺ is [Ar] 3d⁵ and Mn⁷⁺ is [Ar]  because the removed electrons came from the s subshell first
7. You should be able to predict the charge and then write out the electron configuration for any main group ion. You should be able to write the electron configuration for any transition metal ion given the charge. Practice here, click on activities then ionic electron configuration activity.
2. Ionic Radii (ionic size)
1. Remember Z_eff is the effective nuclear charge - it is what the electrons actually "feel" from the protons in the nucleus, how much attraction there is between electrons and the positive protons in the nucleus.  Since negative electrons repel each other, Z_eff decreases the farther you get from the nucleus.
2. If electrons are lost then there are more protons than electrons, the Z_eff is large, the electrons "feel" the attraction from the + protons in the nucleus and they are pulled closer - so the ion is smaller than the atom from which it came
3. If electrons are gained then there are less protons than electrons, the Z_eff is smaller, the electrons don't "feel" the attraction from the + protons in the nucleus as much as the electrons repel each other (negatives repel) and they are pushed apart - so the ion is larger than the atom from which it came.
4. Cations are thus smaller than anions of the same period. The more positive an ion, the larger Z_eff, the smaller the ion. The more negative the ion, the smaller the Z_eff, the larger the ion.  As the proton to electron ratio increases, Z_eff increases and size decreases
5. Put in order of increasing size:  Kr, Mo⁶⁺, Br ^-, Se^2-, and Sr²⁺ and figure out how many protons and electrons each has.
6. Answer:  Mo⁶⁺is smallest, next is Sr^{2+ Kr}  Br ^-  Se^2-  .  Number of protons is 42, 38, 36, 35, and 34 respectively.  Number of electrons is 36 for all of them!!!
7. Isoelectronic = same number of electrons.  You should be able to put an isoelectronic series in order of size.
8. Go watch two movies by clicking here, then click on activities, and finally click on Effective nuclear charge movie and Gain and loss of e movie.
3. Ionization Energy (IE or E_i)
1. Defined as the energy needed to remove an electron.
2. Look at some examples:
1. Na atom - e = Na⁺ This is awesome for Na as it wants to be +1 anyway! So the IE needed is very small.
2. F atom - e = F⁺   This sucks. F wants to be -1 not +1.  So the IE is very high.
3. Ne atom - e = Ne⁺  This is the worst. Ne is already stable.  So the IE is the highest.
3. In general IE increases g and h in the periodic table. Don't worry about the irregularities. 
4. Now why does IE increase up a column?  Well the electrons for small elements are really close to the nucleus so removing them requires more energy than removing electrons that are really far away from the nuclear.  Electrons far away don't feel the attraction to the nucleus as much due to shielding (low Z_eff) while electrons close to the nucleus feel the positive protons well (high Z_eff)
5. Alkali metals have the lowest IE, Fr has the lowest of all elements.  Noble gases have the highest IE, helium has the highest of all elements. This basically means that if you want an electron, Fr will give it up easiest while He will fight to the death. Watch the Periodic Trends IE movie, click on activities.
6. IE is measured in Joules like other energies. 
4. Electron Affinity (EA or E_ea)
1. Defined as the DE when an electron is added. 
1. If adding an e is good, the energy goes down so DE is negative.  This is good and the result is stable, lower E.
2. If adding an e is bad, the energy goes up so DE is positive.  This is bad and the result is unstable, higher E.
3. You can think of EA as "love for an electron"  If adding the e is good, there is lots of love.
2. Examples:
1. Na atom + e = Na ^-   This sucks.  The change is energy is tiny.  Little love here.
2. F atom + e = F ^-    This is great.  F wants to be -1.  The change is energy is huge.  It goes down a lot.  DE is a large negative number.  This is good. Lots of love.
3. Ne atom + e = Ne ^-  This is horrible.  Ne was already happy.  No love at all - the worst.  DE is a positive number. The energy went up.  Ne -1 ion is unstable.
3. In general EA increases g and h in the periodic table EXCEPT the noble gases. Don't worry about the irregularities. 
4. Again the smaller atoms are more affected by changes of adding an electron - the energy changes are larger than for large atoms.  This is why EA increases up a column. Watch the Periodic Trends EA movie, click on activities.
5. Smallest EA is the noble gases since they have none.  Largest is F. 
5. Ionic Bonds and Solids
1. Metals give electrons to nonmetals creating metal cations and nonmetal anions which are attracted to each other (electrostatic attraction) because opposites attract and this is an ionic bond.
2. Na (3s¹) +   Cl (3s²3p⁵)    can make    Na⁺ (2s²2p⁶)  and Cl ^– (3s²3p⁶)   This is great because now both Na and Cl have a full s and p like a noble gas.  Happy Happy Joy Joy.    THIS IS CHEMISTRY!  Forming bonds to lower energy all because the motivation of the atoms is to be s²p⁶ !!!
3. These ions make a nice organized 3D pattern called a crystal lattice structure.
4. Predicting ionic formulas.  This is IMPORTANT.  Ions make compounds so that the total charge adds up to zero.  For example Na is +1 and Cl is -1 so we need one of each to add up to zero.  The formula is NaCl.  Now consider Ba and Cl.  Ba is +2 while Cl is -1 so we need two Cl for every Ba and the formula is BaCl₂.   Check out here and practice by going to the oxidation states and ionic formulas section of the help page.
6. Section 6.8 Lattice energy = energy needed to break an ionic bond.  Usually a large positive value in Joules.  Ionic bonds are strong! Lattice energy increases when the ions are small (more charge to volume so attraction greater) and when the charges are higher (+1 is less attractive than +3 for example - again the charge to volume ratio is greater)  So an ionic bond with small ions would have greater lattice energy than one with large ions if the charges are the same.  And an ionic bond between +3 and -3 ions would have greater lattice energy than a similarly sized +1 and -1 charged ionic bond.
7. Section 6.6  Octet Rule
1. Elements want to have s²p⁶ to be stable and low in energy.  This is 8 electrons! 
2. H and He want to have 2 electrons since they are small and only have a s subshell on the first shell.
3. Row 2 always obeys the octet rule.  Rows 3 and more can break it. 

Good luck.