Questions...

From email...
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On Exam 4, Fall 2007, you have 2 answers highlighted for number 7. I don't understand how Li and P can both be the smallest; can we circle more than one answer on the exam?
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Depending upon your explanation, I would have accepted either of those answers. (That's one of the reasons more recent exams have these comparisons as short answer questions.) If you look at the electron configuration, P has more than a full shell of additional electrons, so it might seem like Li is the smaller atom, but because atomic size decreases left-to-right across the Periodic Table and P is much farther right than Li, P might be smaller. Looking at the actual data (Figure 7.22 in your textbook, page 307), P has a radius of 110pm and Li has a radius of 157pm, so it looks like in this case the left-right trend makes more of a difference than the up-down trend.

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On Exam 4, Fall 2006, you circled A for question number 7 for being the largest ions. I thought the largest ion was the lowest negative charge, and the smallest ion was a positive charge. I am not sure how to figure out what ion is the largest?
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It's not only a matter of charge, we also have to look at the size of the parent atom. For a given element, the higher the charge the smaller the ionic radius and the lower the charge the larger the atomic radius, so, for example, Ge⁴⁺ is smaller than Ge²⁺ which is smaller than Ge which is smaller than Ge^2-. Within a row, this trend is pretty reliable, but as we move far up or down the P.T. things can change. Fr⁺¹ is smaller than Fr, and F^-1 is larger than F, but Fr⁺¹ is much larger than F^-1 because the parent Fr atom is SO huge compared to the parent F atom that the change in size when they form ions doesn't make up for the original difference in size. Looking at the ions in this question, F^-1, Li⁺¹ and Al⁺³ are all definitely small, so it comes down to comparing Pb²⁺ and Br^-1. Pb²⁺ has over a full shell of additional electrons, but it's a cation. Looking at the atomic radii, Pb = 180pm and Br = 115pm, so a Pb atom is bigger than a Br atom, but a Pb²⁺ ion should be smaller than 180pm and a Br^-1 ion should be larger than 115pm. How much smaller and larger will determine the answer to this question... I accepted either answer for this question, but looking at real data, Pb²⁺ has a radius of around 140pm and Br^-1 has a radius of around 180pm, so the correct answer should be Br^-1.

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Have we talked about number 11 and 16 on the Fall 2006 exam or is it just something we should study and know?
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We have addressed these, but maybe not in exactly these terms. #11 is based upon forming stable electron configurations, so being able to write a correct electron configuration and then adding or removing electrons to give full shells, full subshells, or half-full subshells will demonstrate which ions are (relatively) stable. #16 is an application of VSEPR, lone pairs are more "sterically demanding" than bonding pairs so the repulsion in each of these molecules will affect the bond angle. We talked about this comparing methane, ammonia and water bond angles in class.

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Then I also had a question on the most polar bonds. Is the most polar the furthest apart on the periodic table?
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In general, yes, but... The polarity of a bond is a function of the difference in electronegativity of the elements involved, so it's better to look at it from that perspective, with fluorine being the most electronegative. For example, if we're comparing a P-Cl bond to a Si-Cl bond, the Si-Cl bond is a little more polar. If instead we compare a C-S bond to an O-S bond, the O-S bond is quite polar while the C-S bond is barely polar at all, even though C and S are farther apart on the P.T. than O and S.

BTW, atomic and ionic radii numbers came from your textbook and from WebElements.com, it's an interesting website if you haven't checked it out. It's much more information-based than explanation-based, but it's a handy one to keep in mind.

Lewis structures...

We are getting in to Lewis structures and looking at how the electrons are distributed in ionic and molecular/covalent substances. Lewis structures are all about practice. The rules I typically use for Lewis Structures are a little bit different from those listed in the book, so:

Lewis Structures – electron counting method

1. Add up total valence electrons in the molecule or ion

2. Draw a skeleton structure using all single bonds (usually the least electronegative atom is central, hydrogen is NEVER the central atom, some structures have multiple “central” atoms)

3. Fill the octet of all peripheral atoms (hydrogen exception…)

4. Place any extra electrons on the central atom, pair up if possible

5. Check formal charge (find missing or extra electrons…)

6. Minimize formal charge distribution (if possible) by forming multiple bonds (resonance?)

7. Check formal charge

My favorite thing about Lewis Structures is that there are a couple places in the "rules" where you can check yourself and find mistakes early without going through the entire process. Formal charge is an extremely useful tool (perhaps even more useful for those of you who will have to take organic chemistry at some point...) so make sure you're comfortable calculating formal charge.

There's a new OWL assignment posted, due next Wednesday.