Chapter+7+Discussion+Questions

Below are my responses to your discussion questions. It is important that you answer these questions with //**clarity**// and //**precision**//. In some cases, I don't think that what you said is what you intended to say. As your responses are posted, I will post my response.

The atomic radii increases down a group. This is so because more energy levels are being added each time you move down a group. With more energy levels, the outer energy level is farther away from the nucleus causing the atomic radius to be larger. - Colby a.k.a "The Awesomest" I thought this answer was good. With the atomic radii, there are two trends though. The first trend, as mentioned, is that the atomic radius tends to increase from the top to bottom of each column or group. The quantum number of the outer electrons has a lot to do with this trend, and when going down a column, there is a greater probability of the electrons being further away from the nucleus. The other trend observed in atomic radii is that within each row or period, the atomic radius has a tendency to decrease from left to right. When moving across the row, the major fact is the increase in the effective nuclear charge. This effect draws the valence electrons in closer to the nucleus thus causing the atomic radius to decrease. --Samantha
 * 1. Explain the trend for atomic radii down a group.**

Very good answer--it would be good to also mention the shielding effect--inner electrons shield the valence electrons from the attraction to the nucleus.

If ionization energy decreases down a group, it would make it easier to lose electrons. This means that within a group, more electrons means more likely to lose electrons. This is because of 2 reasons. Remember, electrons in the outer energy level are lost. As you go down a group, valence electrons are further from the nucleus, so they have less attraction to the nucleus. Also, the inner electrons shield the valence electrons from the pull of the nucleus.
 * 2. Explain the trend for ionization energy down a group.**
 * The ionization energy decreases down a group. The elements with more electrons get less likely to lose more electrons. - Jay a.k.a "Better than Colby" **

Besides the noble gases, atomic radii decreases as you go across a group from left to right. This happens because, as you go across the group, the attraction of the electron to the protons becomes greater. --Catherine
 * 3. Explain the trend for atomic radii across a period.**

Why is the attraction of the electrons to the protons greater? As you go across a period, electrons are being added to the same energy level, but the positive charge in the nucleus is increasing. This results in greater attraction between the nucleus and valence electrons, making the atom smaller.


 * 4. Explain the trend for atomic radii down a group.**
 * When going down a group, the atomic radii increases. The size begins to get larger as you go down because there are more energy levels as you go down. As you add an energy level, you are adding more outer orbitals (?) which is adding a larger distance from the nucleus. When the electrons are farther from the nucleus, there is less attraction between the nucleus and the outer energy level (The Shielding Effect).**

**-Kiana ..**
You are adding energy levels. For full credit, you should elaborate on the shielding effect--the inner electrons shield the outer electrons from the attraction to the positive nucleus, making the radii even larger.
 * Ahzuree-yeah that is right. I remember it from the stealing chickens story Mrs. Kornegay used.**

The radius of S2- is bigger than S. S2- has 18 electrons and 16 protons. S has 16 protons and 16 electrons. The more protons there are the greater the nuclear charge, which atrracts the outer energy level electrons. -Victoria
 * 5. How does the radius of S** 2- **compare to S? Why?**
 * Yeah yeah. That's right. S2- attracts 16 e- with 2+ left over. **

** -Kiana.. **
You mention more protons and greater nuclear charge; S2- and S have the same number of protons and the same nuclear charge. What is different is that S2- has more electrons, so the nucleus does not attract the outer electrons as strongly, making the ion larger than the atom. The way you specified the number of protons and electrons is good, just be careful in your explanation.

Fe3+ has a smaller radius than Fe2+ because there are more protons in Fe3+ and therefore a greater attraction to hold the elcetrons closer to the nucleus. --Ali I agree this was an easier question so there isn't much to evaluate.-Alex Good evaluation Alex ; ) from Ali How many protons in Fe3+? How many in Fe2+? If they had different numbers of protons, they couldn't be the same element. They both have 26 protons, Fe3+ has 23 electrons and Fe2+ has 24 electrons. Since Fe3+ has fewer electrons, and the same number of protons, the electrons are attracted more strongly to the nucleus and this ion is smaller in radius.
 * 6. How does the radius of Fe** 3+ **compare to Fe** 2+ **. Why?**

Since the hint said to look at electron configuration, let's look at the electron configuration of iron: [Ar]4s23d6. When iron forms a 2+ ion, it loses its two valence electrons (4s). When iron forms a 3+ ion, it loses the two 4s electrons and one 3d electron. This makes all of the sublevels full or half full, making it more stable. I agree with your electron configurations and how you explained why iron can form a +2 or +3 ion. I had to think about what you said, but i understand! :) ~Rebekah They are both transition metals and are fairly ractive so they lose electons because they are metals. Also they both are in an electron configuration were to go down an energy level or clear more than one oribital takes less energy than taking just one. For example lead fills up the electron energy level 6p with the last two oribitals having only one electron, so it would be simpiler to remove both electrons or all four electrons than to remove just one.- Alex
 * 7. Explain why iron can form a +2 ion or a +3 ion. Hint: look at electron configurations**
 * 8. Explain why both tin and lead can form +2 or +4 ions. Hint: look at electron configurations **

Since the hint said to look at electron configuration, let's look at the electron configuration of tin: [Kr]5s 2 4d 10 5p 2. I hope the answer is now obvious. Tin can lose the two 5p electrons to form a 2+ ion. This results in all sublevels being full. OR it can lose the two 5p electrons AND the two 5s electrons to form the 4+ ion. This makes its outer s and p sublevels full.

Group 18 elements are now commonly known as the ‘noble gases’ though ‘inert gases’ fit them just as well. Inert, in chemistry terms, means having little to no ability to react. Before the 1960’s, the group 18 elements were believed to be incapable of forming chemical compounds. There are reasons for this. The elements of group 18 are stable. All their //s// and //p// subshells are completely filled up, so they don’t need to gain or lose electrons. The inert gases, or ‘noble gases’, also all have large first ionization energies, and they don’t naturally bond together to make chemical compounds. There can’t be molecules of XeF4 unless someone forced them together. Out of the inert gases only Xe, Kr, and Ar have been forced into stable molecular compounds. Radon isn’t messed with because of its radioactive nature, and all the others have to large first ionization energy. It was Neil Barlett, in 1962, who reasoned that Xe might be able to react with a substance with an extremely high ability to remove electrions, and this reasoning has lead to finding molecular compounds such as XeF2, XeF6, KrF2, and HArF. Back before the 1960’s, no one had thought that the noble gases were capable of forming molecular compounds, and this lead to the connection to the word ‘inert’. The gases showed no reaction ability, and back then, it only made sense to say that the group 18 elements were ‘inert’. That was why before Barlett the noble gases were commonly known as the inert gases. --Samantha Your answer is correct. I think that you explained it well. They were called inert gases because they are exceptionally unreactive. Neil Bartlett reasoned that the ionization energy of Xe might be low enough to allow it to form compounds. Bartlett synthesized the first noble-gas compound. Again, your answer was correct and stated well. -Victoria Excellent answer and excellent response! Great job!
 * 9. Until the 1960s, group 18 elements were called the inert gases. Why was this name appropriate?**

The transition metals are adding electrons to a d sublevel. The s in a higher energy level is already filled. For instance in the fourth period transition metals, electrons are being added to 3d, but 4s is already filled. The increase in nuclear charge still causes the atomic radii to decrease, but it does not have as big of an impact as when electrons are being added to the outer energy level. Excellent answer--you really know your stuff!
 * 10. As we move across a period in the periodic table, why do the sizes of the transition metals change more gradually than those of the representative elements?**

Metallic character refers to the physical and chemical characterisics of a metal. One of those characteristics is that metals tend to lose electrons easier than other elements. The whole point of first ionization energy is to take one electron with as little energy as possible, so with an easier to lose electron, less energy will be used. Ahzuree I agree that metals have a low ionization energy, meaning it is easier for them to lose electrons. Also, metallic character and ionization energy increase down a group. In contrast, ionization energy increases left to right across a period while metallic character decreases across a period. --Ali Both responses had some unclear or incorrect statements. You are right, metals tend to lose electrons. The easier it is to remove an electron (lower ionization energy) the more metallic character of an element. Metallic characater increases down a group and decreases across a period while ionization energy decreases down a group and increases across a period.
 * 11. How are metallic character and first ionization energy related?**

Sodium and potassium are the alkali metals that play a vital role in physiology. They are major parts of the blood plasma and intracellular fluid, they are the vital carriers of the charges in the normal cellular functions and they are the two main ions that help regulate the heart. ~Rebekah Exactly--I hope every reads this entire section!
 * 12. Which alkali metals play a vital role in human physiology? What is that role? (p. 284)**