2. What happens to the number of protons and electrons in atoms across a period on the periodic table?

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1 Name Period Date Honors Chemistry - Periodic Trends Check Your Understanding Answer the following, formulating responses in your own words. (This helps you better understand the concepts) 1. Define shielding effect. In multi-electron atoms, the inner shells of electrons between the positive nucleus and the valence electrons shield the valence electrons from the attractive positive nature of the positive protons. What is the effective nuclear charge (Z eff )? The net nuclear charge (or effective nuclear charge, Z eff ) felt by an outer electron is substantially lower than the actual nuclear charge (Z). We say that the outer electrons are shielded from the full charge of the nucleus by the inner electrons. a. What shielding effect and Z eff do the valence electrons feel as you progress down a group? Shielding effect increases down a group. Z eff decreases b. What shielding effect and Z eff do the valence electrons feel as you progress across a period left to right?\ Shielding effect remains constant across a period. Z eff increases 2. What happens to the number of protons and electrons in atoms across a period on the periodic table? p + and e - are added 1 by 1 to the same energy level (shell) increasing the atomic number a. What happens to the size of nucleus as you move left to right on the periodic table? The nucleus adds protons and may increase in size but this is insignificant compared to the atoms size. b. What effect would this have on the size (volume) of the atom? Radii decreases across a period as p + and e - are added 1 by 1 to the same energy level (shell) c. Make a general statement on the trend in atomic radius across a period on the periodic table. Include your reason for the change in radius. Radii decreases as number of energy levels (shells of electrons) remains constant keeping the electrons nearer the nucleus attraction. The effective nuclear charge increases steadily and e - added to the outermost shell are ineffective in shielding each other and the e - are drawn closer to the nucleus.

2 3. What happens to the number of energy levels in atoms down a group on the periodic table? The number of energy levels increases one at a time like adding nesting dolls. a. Make a general statement on the trend in atomic radius down a group on the periodic table. Include your reason for the change in radius. Radii increase down a group as energy levels (e - shells) are added a. which increases the size of the orbitals (e - shell n = 1, 2, ) taking the electrons further from the nucleus attraction. b. shielding effect decreases effective nuclear charge (+) as additional shells are added and electrons are not held tight. 4. Define ionization energy. (IE) The minimum energy required to remove an electron from the ground state of an isolated gaseous atom or ion. (in kj/mole) a. How would the size (radius) of an atom affect ionization energy? As the radius decreases the effective nuclear charge increases steadily and the e - are drawn closer to the nucleus making them more difficult to remove. b. What would happen to ionization energy across a period from left to right? Why? Ionization energy increases across a period as p + and e - are added 1 by 1 to the same energy level (shell). a. atomic radius deceases b. effective nuclear charge increases steadily and the e - are drawn closer to the nucleus making them more difficult to remove. c. What would happen to ionization energy down a group? Why? Ionization energy decreases down a group as energy levels (e - shells) are added a. which increases the atomic radius taking the electrons further from the positive nucleus attraction. b. shielding effect decreases effective nuclear charge (+) as additional shells are added and electrons are not held tight 5. Below is the table of the 1 st through 5 th ionization energies for seven elements. Use the table and your knowledge on atomic size, valence electrons and ionization energy to answer the following questions. Element Ionization Energy (kilojoules/mole) IE 1 IE 2 IE 3 IE 4 IE 5 Element (A) Element (B) Element (C) Element (D) Element (E) Element (F) Element (G)

3 Determine the number of valence electrons for each of the following elements. (Hint: Identify the easily removed valence electrons first) Element A 3 Element B 2 Element C 3 Element D 1 Element E 2 Element F 4 Element G 1 a. Which of the elements are in the same groups and families? G & D E & B A & C F b. Which of the elements above has the largest atomic radius? G & D c. Which of the elements above has the smallest atomic radius? F & A next smallest d. What elements above would be considered members the alkali metal family? G & D e. What elements above would be considered members of the alkaline earth metals family? E & B 6. The table on page 2 lists successive ionization energies for several elements. The first four ionization energies in kj/ mole for Mg are as follows: IE 1 = 738, IE 2 = 1451, IE 3 = 7733, and IE 4 = s 2 a. Identify the electron (electron config.) that is removed in the first ionization energy of Mg. 3s 1 b. Identify the electron that is removed in its second ionization energy. 2p 6 c. Identify the electron that is removed in its third ionization energy. d. Explain why the second ionization energy is higher than the first, the third is higher than the second, and so on. As e - are removed in successive ionizations, fewer e remain within the atom to shield the attractive force of the nucleus. Each successive e - removed from an ion experiences and increasingly strong, effective nuclear charge (Z eff ) 7. Briefly explain why barium (Ba) has a lower first ionization energy than calcium (Ca). Ba has two more electron shells between the valence e - and the nucleus than does Ca. With more e - shells comes more shielding. Radii has increased from Ca to Ba so it is easier to remove e -. a. Which of the following elements would you expect to have the largest increase in third ionization energy: K, Ca, or Ga? Why? Ca would have the largest 3 rd IE. When 2 e - are lost, Ca has a stable e - configuration like Ar, once the big jump is made the energy needed to remove another electron does not jump as much. The 2 e - lost would also increase the effective nuclear charge and attractiveness for the remaining e - which would need more energy to be liberated.

4 b. The first four ionization energies in kj/ mole for Mg are as follows: IE 1 = 738, IE 2 = 1451, IE 3 = 7733, and IE 4 = How many valence electrons does magnesium have and explain how you came to your conclusion using these energy values? Mg has 2 valence e - because the 3 rd IE is so much larger than the IE 1 and IE 2. Since more energy is needed to remove the 3 rd e -, Mg must have achieved a stable e - state at the loss of 2 e Define electron affinity (EA) The energy change that occurs when an e - is added to a gaseous atom or ion. For most atoms, the energy released when an e - is added. (The values are negative or also known as - H). a. What would happens to the amount of energy released due to electron affinity across a period? Why? Electron affinity (Energy Release increases) increases across a period (becomes more negative) as p + and e - are added 1 by 1 to the same energy level (shell). a. atomic radius deceases b. effective nuclear charge increases steadily and the e - are drawn closer to the nucleus making it easier to add e - to unfilled sublevels. b. What would happen to the amount of energy released electron affinity down a group? Why? Electron affinity (Energy Release decreases) decreases down a group (becomes less negative) as energy levels (e - shells) are added a. which increases the atomic radius taking the electrons further from the nucleus attraction. b. shielding effect decreases effective nuclear charge (+) as additional shells are added and electrons are not held tightly. 9. Define electronegativity (EN) Tendency for an atom to attract e - closer to itself when forming a chemical bond with another atom. In a chemical bond, the atom with the greater electronegativity more strongly attracts the bond s electrons. a. What would happen to electronegativity across a period from left to right? Why? Electronegativity increases across a period as p + and e - are added 1 by 1 to the same energy level (shell). b. atomic radius deceases b. effective nuclear charge increases steadily and the e - are drawn closer to the nucleus making it easier to add e - to unfilled sublevels. b. What would happen to electronegativity down a group? Why? Electronegativity decreases down a group as energy levels (e - shells) are added b. which increases the atomic radius taking the electrons further from the nucleus attraction. b. shielding effect decreases effective nuclear charge (+) as additional shells are added and electrons are not held tightly. 10. Define ionic size The size atoms become when losing or gaining electrons. One-half the distance between two ionically bound atoms in their mist stable arrangement.

5 a. How do metal ions compare to their original atomic size and Why? Metal Atoms that lose e - and form positive ions become smaller. The lost e - is a valence e - and the atom may lose a shell. The repulsion between the remaining e - is lessened and allows the effective nuclear charge to pull the remaining e - closer. b. How do nonmetal ions compare to their original atomic size and Why? Nonmetal Atoms that gain e - and form negative ions become larger. The repulsion between the added e - and existing e - is increased and the effective nuclear charge cannot hold onto the e - tightly. c. What would happen to the ionic size of atoms across a period from left to right? Why? Ionic radii increase down a group as energy levels (e - shells) are added a. which increases the size of the orbitals (e - shell n = 1, 2, ) taking the electrons further from the nucleus attraction. b. shielding effect decreases effective nuclear charge (+) as additional shells are added and electrons are not held tightly. d. What would happen to the ionic size of atoms down a group? Why? Ionic radius tends to gradually decrease across a period for the positive ions, then beginning in group VA or VIA the much larger negative ions also gradually decreases. Ionic radius decreases as number of energy levels (shells of electrons) remains constant keeping the electrons nearer the nucleus attraction. The effective nuclear charge increases steadily and e - added to the outermost shell are ineffective in shielding each other and the e - are drawn closer to the nucleus. 11. The electron configurations for the following atoms are given for use in answering questions a through i below: A. 1s 2 2s 2 2p 6 3s 2 B. 1s 2 2s 2 2p 6 3s 1 C. 1s 2 2s 2 2p 6 D. 1s 2 2s 2 2p 5 E. 1s 2 2s 2 2p 3 Mg Na Ne F N (Hint: Identify the elements first) a. Which of the above would you expect to have the lowest ionization energy? B. Na b. Which of the above would you expect to be the smallest atom? C. Ne c. List the five configurations in predicted order of increasing ionization energy (lowest to highest). B. Na A. Mg E. N D. F C. Ne d. Predict the configuration that should have the highest second ionization energy. B. Na ( C. Ne 2 nd ) e. Predict the configuration that should have the lowest second ionization energy. A. Mg f. Which of the above would you expect to release the least electron affinity energy? (least negative) A. Mg or C. Ne (or B. Na )

6 g. Which of the above would you expect to release the most electron affinity energy? (most negative) D. F h. Which of the following would you expect to have the highest electronegativity? D. F i. Which of the following would you expect to have the largest ionic size? E. N Accepts 3 e - to achieve Ne e - configuration. 7 p + plus 10 e - are not as attractive to all of its e - and the 10 e - repel each other. Use the periodic table and your knowledge of periodic trends to answer the following questions. 12. Which atom in each pair has the larger atomic radius? (Z eff ) K a. Li or K Br e. Cl or Br Ca b. Ca or Ni Ba f. Be or Ba Ga c. Ga or B Si g. Si or S C d. O or C Au h. Fe or Au 13. Which ion in each pair has the smaller ionic radius? (Z eff ) Ca 2+ _ a. Ca 2+ or O 2- K + d. K + or Cs + Ba 2+ b. Ba 2+ or I Fe 3+ e. Fe 2+ or Fe 3+ NOTE Al 3+ c. Al 3+ or P 3- F - f. F - or S Which atom or ion in each pair has the larger ionization energy? O a. Na or O Ne e. I or Ne Be b. Be or Ba V f. K or V F c. Ar or F Ca g. Ca or Fr Cu d. Cu or Ra Se h. W or Se 15. For the following pairs of ions, tell which one of each pair has the largest ionic radius. A1 3+ a. A1 3+, B 3+ Na + d. Na +, Al 3+ S 2- b. S 2-, O 2- O 2- e. O 2-, F Br c. Br, Cl S 2- f. S 2-, N Indicate which element in the following pairs has the greater ionization energy carbon a. boron, carbon oxygen c. sulfur, oxygen magnesium b. magnesium, calcium silicon d. sodium, silicon 17. Indicate which element of the following pairs has the greater electronegativity. Gallium a. calcium, gallium chlorine c. chlorine, sulfur Oxygen b. lithium, oxygen bromine d. bromine, arsenic

7 18. For each choice of words, circle the one that best completes the sentence. a. The shielding effect increases with increasing atomic number within a (period, group). b. In any (period, group), the number of electrons between the nucleus and the outer energy level is the same. c. Within a group, the nucleus has a stronger ability to pull on the outermost electrons in elements of (high, low) atomic number. d. The more (positive, negative) the value of electron affinity is, the more easily the atom can take on an extra electron. e. From left to right across a period, electron affinity values become more (positive, negative). f. The halogens have the (most, least) negative electron affinity values. g. The (electronegativity, electron affinity) of an atom is the tendency of an atom to attract electrons to itself when it is chemically combined with another element. h. The (noble gases, halogens) are the most electronegative group. 19. Use the positions of the elements in Periodic Table to make predictions about the following. a. Which atom has the more negative value for electron affinity, sodium or chlorine? b. Which atom is more electronegative: fluorine or lithium? c. Which atom has a greater ionization energy: nitrogen or bismuth? d. Which atom exerts a stronger pull on its outermost electrons: calcium or arsenic? e. Which atom has a larger atomic radius: fluorine or barium? f. Which atom has a greater nuclear charge: boron or selenium? 20. Use Electron Affinities table to answer the following questions. a. Which halogen is most likely to gain an electron? Cl b. Which elements need to absorb energy when they gain electrons? Noble Gases, Be, Mg, N c. With which element would calcium be more likely make a compound: oxygen or fluorine? d. With which element would nitrogen be more likely make a compound: magnesium or chlorine?

8 21. At a. Name the halogen with the least-negative electron affinity. Li b. Name the alkali metal with the highest ionization energy. Ar (Cl, too) c. Name the element in Period 3 with the smallest atomic radius. C d. Name the Group 14 element with the largest electronegativity. 22. F a. Name the most chemically active halogen. 1s 2 2s 2 2p 5 b. Write its electron configuration. 1s 2 2s 2 2p 6 for 1- ion c. Write the configuration of the most-stable ion this element makes. 23. Referring only to the periodic table, answer the following questions on periodic trends. In a. Which has the larger radius, Al or In? Ca b. Which has the larger radius, Se or Ca? Ca c. Which has a larger radius, Ca or Ca 2+? nonmetals_ d. Which has greater ionization energies as a class, metals or nonmetals? Cl e. Which has the greater ionization energy, As or Cl? _negative ion_ f.. An element with a large negative electron affinity is most likely to form a positive ion, a negative ion, or a neutral atom? small g. In general, which has a stronger electron attraction, large atoms or small atoms? O h. Which has greater electronegativity, O or Se? O i. In the chemical bond between Se and O, to which atom is the electron pair more closely drawn? 6 j. How many valence electrons are there in a neutral atom of Se? 24. Identify all of the ions at right that do not have noble gas stability. K + S 2- Ca + I - Al 3+ N Copper s electron configuration might be predicted to be 3d 9 4s 2. But in fact, its configuration is 3d 10 4s 1. The two elements below copper in Group 11 behave similarly 3d 10 4s 1 a. Yes b. True c. Which configuration is apparently more stable? Is the d subshell completed in the atoms of these three elements? Every element in Period 4 has four levels of electrons established: True or False.

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