SAMPLE PROBLEMS! 1. From which of the following is it easiest to remove an electron? a. Mg b. Na c. K d. Ca 2. Which of the following influenced your answer to number one the most? a. effective nuclear charge b. valence electrons c. inner shell electrons d. shielding effect 3. As the atomic numbers of the halogens increase, electronegativity: a. decreases b. increases c. remains the same d. breaks 4. Ions formed by the gain of an electron are always more than the atom they came from. a. stable b. unstable c. electronegative d. I can't tell because I don't know what the atom is. 5. Electron affinity within a period or series generally with increasing atomic number. a. increases b. decreases c. gets more positive d. stays about the same 6. What is the most probable oxidation state of Ca? a. -3 b. -2 c. +3 d. none of these 7. You do an experiment which involves adding an electron to an atom of K. You make the following assumption in regard to its stability: a. the ion is more stable b. the ion is less stable c. the stability doesn't change d. You can't tell what happened because this statement doesn't give enough information. 8. Which of the following elements is most nonmetallic? a. O b. F c. S d. Br 9. Which of the following has a larger radius? a. rubidium b. rubidium ion c. they are the same 10. Which of the following is most likely to lose electrons in a chemical reaction? a. magnesium b. chlorine c. fluorine d. oxygen 11. Low ionization energy is characteristic of: a. metals b. non-metals c. metalloids d. liquids 12. Which of the following trends have the same pattern as atomic radius? a. electronegativity b. electron affinity c. ionization energy d. all of these e. none of these f. more than one of these 13. If energy is absorbed during a process, the resulting particle is: a. more stable b. less stable c. about the same d. irrelevant 14. An atom which gained energy during an electron transfer: a. gained an electron b. lost an electron c. lost a proton d. there is insufficient information to answer this question. 15. Which of the following elements would you expect to have a positive electron affinity? a. K b. Cl c. O d. F e. none of these 16. Compared to the stability of the original atom, the stability of its ion formed by the gain of an electron is: a. always less b. sometimes less c. the same d. always greater 17. Which of the following is the smallest? a. Ne b. F - c. Na +1 d. Mg +2 18. Which element would form an ion which is isoelectronic with Ne and would bond with oxygen in a 1:1 ratio: a. F b. Na c. Mg d. Al e. Si 19. Which of the following elements would have the highest ionization energy? a. Ga b. Zn c. As d. K e. Na 20. When nitrogen reacts with a metal, which of the following would be a possible oxidation state for N? a. +3 b. +1 c. -3 d. +4
MORE SAMPLE PROBLEMS! 1. Which is largest? a. K b. Cr c. Zn d. Kr 2. Which has the greatest increase in stability when an electron is added? a. Al b. Cl c. Ni d. Fe e. Fr 3. Which becomes the most unstable when an electron is removed? a. K b. Co c. V d. Cl e. Mn 4. Which is most likely to gain 2 electrons? a. Si b. P c. S d. Cl e. Ar 5. The dominant factor in determining the variation in size of successive atoms in a family is the: a. increase in nuclear charge b. increase in number of valence electrons c. decrease in the radii of electron shells d. increased shielding effect 6. Atom size goes down as you move up a family because of: a. increasing number of protons b. increasing shielding effect c. relatively constant shielding effect d. decreasing shielding effect 7. You do an experiment which involves adding an electron to an atom. You make the following assumption in regard to its stability: a. the ion is more stable b. the ion is less stable c. the stability doesn't change d. You can't tell what happened because this statement doesn't give enough information. 8. Which of the following has a larger radius? a. sulfur b. a likely sulfur ion c. they are the same 9. Which of the following is most likely to gain electrons in a chemical reaction? a. magnesium b. sulfur c. fluorine d. oxygen 10. Which of following pairs would have the greatest difference in electronegativity? a. Cl and O b. K and Cl c. Ti and N d. I and F 11. Which of the following trends have the same pattern as ionization energy? a. electronegativity b. electron affinity c. atomic radius d. a and b e. none of these 12. The energy change when a single electron is added to an atom is called: a. electron affinity b. electronegativity c. ionization d. none of these 13. The decrease in ionization energy between Mg and Al is related to: a. starting a new sublevel b. increasing electronegativity c. the stability of a half full sublevel d. high electron affinity 14. Which of the following is least likely to gain an electron? a. Na b. K c. Fr d. Cs 15. Which of the following would have an extremely high third ionization energy compared to its second ionization energy? a. Mg b. Al c. Si d. Sb 16. An element which loses three electrons in a chemical reaction is most likely a: a. metal b. non-metal c. liquid d. none of these 17. Which of the following elements would you expect to have a negative electron affinity? a. B b. K c. O d. F e. all of these 18. Which of the following elements would you expect to have an ion which is smaller than it's atom? a. Se b. Cl c. S d. O e. none of these 19. The decrease in electron affinity between carbon and nitrogen is primarily due to: a. shielding effect b. effective nuclear charge c. electron pair repulsion d. proton-electron attraction 20. Extremes of the trends we covered are at the of the periodic table: a. middle b. ends c. corners d. bottom
EVEN MORE SAMPLE PROBLEMS! 1. When a non-metal gains an electron and becomes a negative ion, it becomes: a. smaller b. larger c. remains the same size d. depends on the metal e. none of these 2. The largest contributing factor to the change in electronegativity within a period is : a. effective nuclear charge b. valence electrons c. inner shell electrons d. shielding effect 3. As the atomic numbers of halogens increase, their electronegativity: a. decreases b. increases c. remains the same d. breaks 4. The most significant factor in explaining fluctuations in ionization energy through the transition metals is: a. nuclear charge b. shielding effect c. number of electrons d. electron affinity 5. Elements within a family generally attract electrons with increasing atomic number. a. more b. less c. about the same 6. The most significant factor in determining changes in ionization energy within a family is: a. number of protons b. shielding effect c. valence electrons d. the strong force 7. What is the most probable oxidation number of Oxygen? a. -3 b. -2 c. +3 d. +4 e. +6 8. You do an experiment which involves removing an electron from an atom. You make the following assumption in regard to its stability: a. the ion is more stable b. the ion is less stable c. the stability doesn't change d. You can't tell what happened because this statement doesn't give enough information. 9. Which of the following elements has the lowest ionization energy? a. Na b. Mg c. Al d. K e. Ca 10. Which of the following elements is most metallic? a. O b. S c. As d. Se 11. Which of the following has the largest atomic radius? a. Ni b.sc c. Fe d. Zn e. Bi 12. Low ionization energy is characteristic of: a. metals b. non-metals c. gases d. liquids 13. What factor causes the increase in atomic radius the most between two atoms in a family? a. effective nuclear charge b. start of a new sublevel c. shielding effect d. octet rule 14. Which of following pairs would have the least difference in electronegativity? a. Cl and O b. K and Cl c. Ti and F d. I and F 15. When a metal atom loses an electron and becomes a positive ion, it becomes: a. smaller b. larger c. remains the same size d. depends on the metal e. none of these 16. The energy change when a single electron is removed from an atom is called: a. electron affinity b. electronegativity c. ionization energy d. none of these 17. The greatest increase in ionization energy for the element Na comes between which two electrons? a. 2nd and 3rd b. 3rd and 4th c. 1st and 2nd d. none of these 18. Which element has the highest electronegativity? a. Cl b. S c. Br d. P e. I 19. What factor influence your answer in #18 the most? a. Atomic radius b. shielding effect c. valence state d. electron affinity e. nuclear charge 20. Which of the following elements would you expect to have a positive electron affinity? (meaning, it absorbs energy when it takes an electron) a. B b. Kr c. O d. F e. all of these f. none of these
A Few Questions for Review! 1. Which is larger - K or Ca? P or As? K and As 2. From which of the following is it harder to remove an electron? Mg or Si? Fe or Ru? Si and Fe 3. Which of the following pulls on an electron in a bond the tightest? S or Se? P or S? S and S 4. Explain why P is smaller than Al. (How can it be smaller if it has more stuff in it?) Has more effective nuclear charge the protons pull the electrons in 5. Why is it so difficult to remove an electron from Cl and so easy to remove one from K? (even though K has more protons) Chlorine has a much greater effective nuclear charge, and less shielding effect, than K the nucleus is stronger in its attractive effect - 6. What is the relationship between size and effective nuclear charge within a series or period? As effective nuclear charge increases within a period, then the size decreases 7. What is the relationship between electronegativity and effective nuclear charge within a series or period? As effective nuclear charge increases within a period, the electronegativity increases 8. Discuss trends within both periods and families for atomic size, ionic size, ionization energy, electron affinity, electronegativity, and reactivity. Explain why they occur. All trends occur because of effective nuclear charge and shielding effect as we go across the periodic table within a period, shielding effect stays the same, effective nuclear charge increases, and size decreases, while ionization energy, electron affinity, and electronegativity increase. As we go down a group, shielding effect increases, while effective nuclear charge stays the same, so size increases, and ionization energy, electron affinity, and electronegativity decrease. Atoms always become larger when they take electrons, and become smaller when they lose electrons. Metals with fewer valence electrons react better, and non-metals with more valence electrons react better. 9. Where do you find exceptions to the ionization energy trend within periods? Why do they occur? When you start a new sublevel or pair electrons, exceptions occur new sublevels are unstable and so is pairing electrons 10. Discuss the sign of an atom s electron affinity and the stability of the ion it forms if it takes an electron! If an atom takes an electron, and has a positive electron affinity, that means it takes in energy and is becoming less stable. If an atom takes an electron, and has a negative electron affinity, that means it lost energy and is becoming more stable. 11. Which is more stable, F or F -1? Cl or Cl -1? Ne or Ne -1? Discuss the proper energy sign for what is occurring in going from each neutral element to its corresponding ion - meaning, is energy lost or gained? F -1, Cl -1, and Ne Energy is lost for the first two (- energy), and gained for the third (+ energy) 12. Where do you expect exceptions to the filling order of electrons to occur? In the transition metals violations of the Aufbau principle
13. Which family would have an extremely high 3rd ionization energy relative to the second ionization energy? Why? Group II Once they lose two valence electrons, to remove any additional electrons would require those electrons to be removed from a full octet below. 14. Why is the 2nd ionization energy of Ca significantly higher than the first ionization energy when both electrons come from the outer energy level, and Ca wants to give both electrons away? Once an electron is removed from any atom, the atom becomes positive. Since opposites attract, the positive ion will more greatly attract the second electron that one is attempting to remove. 15. Why is the first ionization energy for O less than that for N if the general trend is to increase from left to right? Oxygen has one more electron than nitrogen nitrogen has exactly a half full p sublevel, with all electrons unpaired. For oxygen, the extra electron must pair itself with another in the p sublevel. Since electrons repel, this additional electron is a bit easier to remove. 16. Are positive ions larger or smaller than their atoms? Why? Are negative ions larger or smaller than their atoms? Why? When positive ions are created, they are done so by removing entire valence shells from atoms making them smaller. When negative ions are created, they are done so by adding additional electrons to valence shells but without the addition of additional protons to pull on them, the atom swells in size, or increases because there is no increasing of effective nuclear charge. 17. Which is a smaller atom - Na or Al? Al is smaller more effective nuclear charge 18. Which is larger - K +1 or Cl -1? Why? Same number of electrons, but Cl -1 is larger because it has less protons to pull on the electrons 19. What are some factors that determine the ionization energy of an atom? Shielding effect and effective nuclear charge as well as unpairing of electrons and starting new sublevels 20. Which is larger--k or Na? P or S? K and P 21. From which of the following is it harder to remove an electron? Mg or Ba? Fe or Zn? Mg and Zn 22. Which of the following pulls on an electron in a bond the tightest? S or O? Li or Na? O and Li 23. Explain why Cl is smaller than S. (How can it be smaller if it has more stuff in it?) Because of effective nuclear charge They both have the same amount of shielding effect, but since Cl has more protons, it pulls on the outer shells tighter -
24. Why is it so difficult to remove an electron from F and so easy to remove one from K? (even though potassium has more protons) F has more effective nuclear charge, and less shielding effect 25. What is the relationship between size and ionization energy? As size increases, ionization energy decreases it is an inverse relationship 26. What is the relationship between size and electronegativity? As size increases, electronegativity decreases it is an inverse relationship 27. As we progress down a group, the atoms have more protons - but the size of the atom increases. What effect or phenomenon best explains this? Shielding effect even though the atoms have more protons, the full energy levels shield its effectiveness 28. As we progress across a period, the atoms are adding more electrons to energy levels, but the atoms get smaller - What effect or phenomenon best explains this? Effective nuclear charge 29. More protons mean more pull on electrons - so why does Ba have a lower ionization energy than Ca? Because it has a greater shielding effect there are more full energy levels shielding the nucleus from holding onto valence electrons 30. Why is aluminum slightly larger than magnesium, if it has more protons? It places its one additional electron in the p sublevel, causing the atom s radius to extend slightly 31. Why is the first ionization energy for O less than that for N if the general trend is to increase from left to right? Oxygen has one more electron than nitrogen nitrogen has exactly a half full p sublevel, with all electrons unpaired. For oxygen, the extra electron must pair itself with another in the p sublevel. Since electrons repel, this additional electron is a bit easier to remove. 32. Why about magnesium s electron configuration makes it less likely to form a anion (negative ion) than sodium? With a full sublevel (the s sublevel), taking an electron would require it to place that electron into the p sublevel, which is more unstable than the s sublevel, which is where sodium would place it. 33. Why does oxygen gain 2 electrons if the first electron affinity is large and negative (releases energy) while the second electron affinity is actually positive, meaning it has to absorb energy to take it? Because the stability that it gains from meeting the octet rule is greater than the energy expended in taking the second electron. Also, when it bonds to a metal, forming an ionic compound, energy is released, causing overall stability. 34. What gas are all incandescent light bulbs filled with and why? Think non-reactive - they don t want the metal filament to react too fast! Noble gases or nitrogen, since nitrogen is fairly non-reactive
THE BIG BAD REVIEW QUESTIONS! YEAH! 1. Where do you find exceptions to the atomic radius trend within periods? Why do they occur? Exceptions to atomic radius occur whenever we start a new sublevel, and within the transition metals. This occurs because new sublevels, such as the p sublevel, extend beyond the s. Also, within the transition metals, you are increasing the shielding effect of the nucleus, which can reduce its ability to pull in electrons. 2. Discuss the relationship between the last electrons placed in an atom and the position of each of the sublevel regions on the periodic table. The block that the element is in tells us which sublevel the last electron is located there are 4 blocks the s, p, d, and f blocks 3. Which family would you expect to have an extremely high 3rd ionization energy relative to its first and second ionization energies? Why? Group II has an extremely high 3 rd ionization energy because once you remove the 2 valence electrons, the third electron comes from the full energy level beneath, which requires large amounts of energy. 4. Why is there a discrepancy between the ionization energy of phosphorous and sulfur? As we pair electrons in the p sublevel, they repel each other and are slightly easier to remove. 5. Why is there a discrepancy between the atomic size of zinc and gallium? When we add an electron to zinc, in order to form gallium, we must start a new sublevel the p sublevel. The p extends beyond the s slightly, making the atom slightly bigger. 6. Where do you find exceptions to the electron affinity trends within periods? Why do they occur? Trends occur between sublevels, and when electrons have to be paired. This is because starting a new sublevel is unstable, and higher energy, and because electrons do not want to be paired. 7. Why is there a discrepancy between the electronegativity of chromium and manganese? In chromium, one of the electrons in the 4s moves down into the empty d orbital below, in order for all electrons to be unpaired. 8. Discuss how shielding effect and effective nuclear charge affect each of the four trends we graphed. This is important! Shielding effect reduces the ability of the nucleus to hold onto electrons, causing atomic size to increase, ionization energy, electron affinity, and electronegativity to decrease. As we increase the number of protons, effective nuclear charge increases, causing atomic size to decrease, ionization energy, electron affinity, and electronegativity to increase. 9. Why does manganese not have an electron affinity? Discuss its electron configuration! With a half-full d sublevel, with all electrons unpaired, and a full s sublevel, manganese does not want to take an electron and have to pair it up. 10. List the following in order of increasing atomic radius: S -2, Cl -1, K +1, Na +1 Na +1, K +1, Cl -1, S -2 11. List the following in order of increasing ionization energy: S, Cl, Ar, K, Ca K, Ca, S, Cl, Ar
12. Half full sublevels are stable - give an example of this using two elements and a trend! If you look at N and O, you will see that there is an exception with the ionization energy O has a slightly lower ionization energy, because it has one more electron than N, which had to be paired up with another electron, causing them to repel each other so a half full sublevel is stable because all electrons are unpaired - 13. Explain the electron affinity discrepancy between fluorine and chlorine! Meaning, why does fluorine have a less negative electron affinity than chlorine, when it should want electrons more? Because Fluorine has such an incredibly small atomic radius, electrons have a difficult time being taken in by fluorine or at least more difficult than chlorine Fluorine still takes electrons EXTREMELY well, but not as well as chlorine 14. Ionization energy should increase as we progress through the periodic table, but it actually fluctuates throughout the transition metals - what phenomenon best explains this? Shielding effect is changing we are actually adding electrons to the d sublevel, which is below the valence level, which actually changes the amount of shielding around the nucleus - 15. List the following in order of increasing electron affinity - N, Na, I, I -1! I -1, N, Na, I 16. Sulfur has a negative first electron affinity, but a positive second electron affinity, even though it wants two electrons total. Explain this! As it takes in electrons, it becomes negative negative atoms repel incoming negative electrons, so it is hard to get it to take the second electron! 17. Why does the ionization energy decrease between Be and B, if B has more protons? Explain! Because we are starting a new sublevel the p sublevel which is higher energy and more unstable than the s 18. An atom releases energy during a process. Is it now more or less stable? Any time energy is released in any process, the resulting atom or molecule is ALWAYS more stable 19. Why is the second ionization energy for magnesium higher than the first, if it wants to give away two electrons total? Because once you remove the first electron, the atom becomes positive, attracting the second electron and making it more difficult to take away 20. Why is ionization energy ALWAYS positive? It always takes energy to remove a negative electron that is being attracted to a nucleus 21. What element is isoelectronic with S -2? With Ca +2? With Ti +4? Argon 22. How many valence electrons does V have? How about Co? How about Ag? What do these elements have in common with each other? 2 electrons they all have 2 electrons in their outer shell, as they violate the Aufbau principle 23. What is so unique about the transition elements and their valence states? They all technically have 2 electrons in their outer shell, because they all filled the s sublevel of the higher energy level before filling the d of the one below it - 24. If Br -1, Kr, and Rb +1 all have the same number of electrons, then why is Rb +1 the smallest of the three?
Most number of protons most effective nuclear charge most pull on electrons 25. Why is I -1 bigger than Xe if they both have the same number of electrons? I -1 has less protons, and therefore less pull on those electrons so it is bigger! 26. Why are Bi, Po, Pb, and Sb metals if they are in the p block of elements? They have so much shielding effect, and so many energy levels, that the valence electrons are so loosely held and far away that they are easily lost hence, a metal! 27. Sn is in group IV, but it is a metal - why is this? It has so much shielding effect that the electrons are extremely far away from the nucleus, and easy to lose! 28. Does K or As have a greater effective nuclear charge? What effects in trends do we witness because of this? As has a greater effective nuclear charge they both have 4 full energy levels shielding the nucleus, but As has more protons pulling through 29. Sr has more electrons and protons than Ca, and electrons attract protons, but it is a larger atom - explain! Sr has more shielding of the nucleus, causing it to not be able to as effectively pull in electrons it also has 5 energy levels, where Ca has 4. 30. Carbon, Nitrogen, and Oxygen - which has the highest effective nuclear charge? Take exceptions into account! Oxygen 31. Cr and Cu have different electron configurations listed on the periodic table than what we would expect - why? Electrons move to empty orbitals in order to be unpaired 32. Why do transition metals have multiple oxidation states? Because electrons in both the s and d sublevels are unstable, and the atom wants to lose electrons from both sublevels 33. Why does Titanium form a +3 or +4 ion? Why does Iron form a +2, +3, or +6 ion? Because it can lose the two electrons in the 4s, and one in the 3d, or both in the 4s and 3d Because it can lose the two electrons in the 4s, or the 2 in the 4s and one in the 3d, leaving all d electrons unpaired 34. Nitrogen has 5 valence electrons and needs 3 more. It has a positive electron affinity. Why is this? Phosphorous, right below it, also has 5 valence electrons, and also needs 3 more. It has a negative electron affinity, though. Why is this? Nitrogen has a half-full sublevel, where all electrons are unpaired, so it resists taking another electron and having to pair electrons up this is why it has a positive electron affinity. However, phosphorous, with a half-full sublevel as well, has its valence electrons in the third energy level. The third energy level has an empty d sublevel that P could put electrons into and avoid pairing up electrons. This is why it has a negative electron affinity, and wants electrons more than N. 35. Why does electron affinity actually decrease between groups I and II? Don t more protons mean more attractions to electrons? Group II has paired electrons in the s sublevel, and a full s sublevel to take electrons would require them to be put into the p sublevel, which is more unstable than the s
36. If Al has more protons than F, then why does it have a lower electron affinity? Shouldn t the protons pull on electrons more? Because it has to do with shielding effect it has more shielding effect, which reduces the ability of the nucleus to hold onto electrons - 37. What does the oxidation state of an element tell us? It tells us the most probable charge that an atom will form when it reacts with another element 38. An atom loses an electron. Does this process always require energy, always give off energy, sometimes require energy, sometimes give off energy, or do we not know unless we know what atom we are talking about? It always takes energy to remove a negative electron away from a positive nucleus - 39. Why are there no electronegativities listed for most of the noble gases? Why does Cr have a higher electronegativity than Mn, if it has less protons pulling on electrons? The noble gases don t bond, so there is no tendency to attract electrons in a bond Cr has all unpaired electrons, which means it does not want to readily accept an electron and pair them up 40. Why is ionization energy ALWAYS POSITIVE if some elements become stable when they lose an electron? Because it always takes energy to remove electrons from atoms the stability comes when we ADD the electron to the non-metal in a bond, and lots of energy is released 41. There are plenty of exceptions throughout the transition metals this is PRIMARILY due to what? Shielding effect we increase the electrons in the d sublevel of the energy level beneath, which increases shielding effect 42. Which of the following would have the smallest ionic size or radius, if it were to form a likely ion? Cl, K, Ca, P, or S? Ca it has the most protons pulling on the 18 electrons that all of these atoms would have if they formed an ion