Chapter 6: The Periodic Table

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1 Chapter 6: The Periodic Table Name: Per: Test date: In-Class Quiz: Moodle Quiz: preap Learning Objectives Trace the historical development of the periodic table Identify the major groups and key features of the periodic table (metals, nonmetals, metalloids, family, period, transition, etc) Using Lewis structures, explain why elements in the same group have similar properties. Describe and compare properties of each family of elements; specifically Alkali, Alkaline Earth, Transition, Inner Transition, Halogens, and Noble Gases Compare period and group trends of properties such as atom size, ion size, electronegativity, shielding effect, ionization energy Relate period and group trends to electron configuration Identify the four blocks of the periodic table based on electron configuration Given data pertaining to unknown elements, use your knowledge of periodic trends to predict placement in the periodic table. List similarities and differences of elements within a given group Define allotropes and provide examples Explain the importance of the representative elements Provide a basic description of the properties of transition elements, metals, metalloids, nonmetals Predict likely charges for an element based on position on the periodic table Define the metallic bond and give characteristics of metals. Differentiate between a metallic bonding model and other models (covalent, ionic). Make an alloy and define if it is substitutional or interstitial. Concepts to be covered History and Development of Periodic Table Periodic Law modern P. Table arrangement: groups, periods, families, blocks (metals, nonmetals), etc e- configuration and how it fits with P. Table Trends: atomic radius, ionic radius, ionization energy, electronegativity similarities and differences of elements in a given group properties of each family and of each block (alkali, alkaline earth, transition, halogens, noble gases) (s, p, d, f) charge prediction for ions using periodic table basics of metallic bonding Alloys and their types *diagonal relationships* *element uses and where found* *explain why some transition metals form compounds with color and some have magnetic properties* (Starred concepts may or may not be covered) TEKS 3.D: evaluate the impact of research on scientific thought, society, and the environment; 5 Science concepts. The student understands the historical development of the Periodic Table and can apply its predictive power. The student is expected to: 5.A: explain the use of chemical and physical properties in the historical development of the Periodic Table; 5.B: use the Periodic Table to identify and explain the properties of chemical families, including alkali metals, alkaline earth metals, halogens, noble gases, and transition metals; and 5.C: use the Periodic Table to identify and explain periodic trends, including atomic and ionic radii, electronegativity, and ionization energy. 7.D: describe the nature of metallic bonding and apply the theory to explain metallic properties such as thermal and electrical conductivity, malleability, and ductility; 1

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4 Periodic Table History Notes Scientist Contribution Lavoisier Dobereiner Newlands Mendeleev Mosley 4

5 6.1 Development of the Modern Periodic Table 1. What 3 things elements in the same group have in common? 2. How many groups/families are on the periodic table? 3. What happens as you move across a period? 4. How many periods are there on the periodic table? 5. Where are the lanthanide and actinide series found? 6. Who constructed the original periodic table? 7. Mendeleev listed the elements in order of. 8. What was Mendeleev able to predict? 9. The elements in the modern periodic table are arranged in order of. 10. Who made this change? 11. In the modern periodic table, elements with similar properties fall into. 12. What periodic pattern is repeated in each group? 13. What groups are located in the s-block of the periodic table? 14. What groups are located in the p-block of the periodic table? 15. What groups are located in the d-block of the periodic table? 16. What groups are located in the f-block of the periodic table? 17. Which group on the periodic table ends its electron configuration with np 5? 18. Which group on the periodic table ends its electron configuration with ns 1? Match each element in Column A with the best matching description from Column B. Each Column A element may match more than one description from Column B. Column A Column B 19. Strontium a. halogen 20. Chromium b. noble gas 21. iodine c. alkaline earth metal 22. nitrogen d. metalloid 23. argon e. alkali metal 24. rubidium f. representative element 25. silicon g. transition element 6.2 Classification of the Elements Use the periodic table to answer the following questions. 1. How many valence electrons are in an atom of each of the following elements? a. magnesium d. arsenic b. selenium e. iodine c. tin f. iron 2. In which energy level are the valence electrons of the elements listed in question 8? 3. Identify each of the following elements. a. an electron configuration of [Kr]5s 2 4d 10 5p 2 b. five valence electrons in the sixth energy level c. two valence electrons in the first energy level d. three fewer electrons in the fourth energy level than krypton e. an electron configuration ending in 4p 2 4. Without using the periodic table, determine the group, period, and block in which an element with each of the following electron configurations is found. a. [He]2s 2 2p 5 c. [Kr]5s 2 4d 10 5p 3 b. [Ar]4s 2 d. [Ar]4s 2 3d 3 5. Write the electron configuration of the following elements. a. the alkaline earth element in the sixth period b. the halogen in the third period c. the group 14 element in the third period d. the group 5 element in the fourth period e. the group 1 element in the fifth period 5

6 6.2-5 Periodic Org 1. What is a group (or family)? What is a period? 2. How can you determine the number of electrons in an element s outer energy level by the group it s in? 3. What is the octet rule? 4. Why do elements that make positive ions occur on the left side of the periodic table while those that make negative ions occur on the right? 5. What is the common name for group 18? 6. Why do the elements of this group usually not form ions? 7. Complete the following table. Group Common Name Charge on Ions of this Group / 3A / 6A / 7A 8. Predict the charges on ions of the following atoms. Ra As Te Cs In At Ga 9. a) In group 1, which element is the most active? b) Metallic activity tends to (increase, decrease) as one goes down Group a) Which element is most active in group 17? b) Nonmetal activity tends to (increase, decrease) as one goes down Group Which element is a metal: Ba (56) or At (85)? 12. Which period is Ca (20) in? 13. What is the number of the group N (7) is in? 14. Which element is an alkali metal: Rb (37) or Al (13)? 15. Which element is a halogen: Na (11) or Cl (17)? 16. Which element is a noble gas: Ne(10) or Br(35) or O(8)? 17. Which element is the most active nonmetal? 18. Which element is the most active metal? 19. Which element would be a positive ion in a compound: Sr (38) or Te (52)? 20. How many electron dots should As (33) have? 21. When Te (52) is an ion in a compound, what charge does it have? 22. How many is an octet of electrons? 23. Which element has 5 valence electrons? B (5) or P (15)? 24. Which element has 18 electrons when it is an ion with a 1 charge? 25. What atomic number would an isotope of U(92) have? 26. How many neutrons does bromine-80 have? 6.3 Periodic Trends Shielding Effect: Why? As you go from left to right across the periodic table, what happens to the shielding effect? Why? As you go from top to bottom down the periodic table, what happens to the shielding effect? Why? How are shielding effect and the size of the atomic radius related? 6

7 Atomic Radius: Why? 1. Identify and explain the trend in atomic size for the following transitions in the periodic table. (a) Moving vertically from Ar to He (b) Moving horizontally from Na to Ar Ionic Radius: Why? 2. Which diagram best represents the relationship between the diameter of a sodium atom and the diameter of a positive sodium ion? a. b. c. 3. In each of the following pairs, pick the larger species. Explain you answer in each case. (a) Cu and Cu 2+ (b) F and F - 4. For each of the following pairs, predict which atom is larger. a. Mg, Sr d. Ge, Br b. Sr, Sn e. Cr, W c. Ge, Sn 5. For each of the following pairs, predict which atom or ion is larger. a. Mg, Mg 2 d. Cl, I b. S, S 2 e. Na, Al 3 c. Ca 2, Ba 2 6. Predict which of the ions, Mg 2 or S 2, is larger. Explain your prediction. Ionization Energy: Why? 7. There is a general increase in the first ionization energy from sodium to argon. Where does this not hold true? 7

8 8. Boron has a lower first ionization energy than beryllium. Why is that? 9. Consider the table of ionization energies for element X shown below. Ionization Energy in kj/mol 1st 2nd 3rd 4th 5th 6th (a) In which group will X be found? Explain. (b) What would the likely charge be on X s ion? 10. Comparing elements from left to right across a period, what general trend would you predict for the energy required to remove a valence electron from an atom? Explain the basis for your prediction. 11. For each of the following pairs, predict which atom has the higher first ionization energy. a. Mg, Na d. Cl, I b. S, O e. Na, Al c. Ca, Ba f. Se, Br 12. For each of the following pairs, predict which atom forms a positive ion more easily. a. Be, Ca d. K, Ca b. F, I e. Sr, Sb c. Na, Si f. N, As Electronegativity: Why? Circle the element with the largest electronegativity and underline the smallest electronegativity: 13) Si N 14) Ca Al 15) Pb Bi Sn 16) O S F Cl 17) Ge Si P As S What elements are excluded from electronegativity? Why? Reactivity: Why? As you go from left to right across the periodic table, what happens to reactivity of the elements? Why? As you go from top to bottom down the periodic table, what happens to reactivity of the elements? Why? How are reactivity and the valence electrons related? 13. For each of the following elements, state whether it is more likely to gain or lose electrons to form a stable octet configuration and how many electrons will be gained or lost. a. K e. Al b. Br f. I c. O g. Ar d. Mg 8

9 14. Which noble-gas configuration is each of the following elements most likely to attain by gaining or losing electrons? a. S e. Fr b. Sr f. N c. Cl g. Ba d. Be 15. For each of the following pairs, predict which atom has the higher electronegativity. a. Mg, Na d. Ca, Ba b. Na, Al e. S, O c. Cl, I f. Se, Br Chapter 6 Concept Review 1. Explain why the word periodic is applied to the table of elements. 2. Why do elements in a group in the periodic table exhibit similar chemical properties? 3. What chemical property is common to the elements in group 18 of the periodic table? Why do these elements have this property? 4. In terms of electron configurations, what does the group number of the representative elements in the periodic table tell you? 5. Describe the group and period trends in the following atomic properties. a. atomic radius b. electronegativity c. first ionization energy d. ionic radius 6. Describe the relationship between the electronegativity value of an element and the tendency of that element to gain or lose electrons when forming a chemical bond. 9

10 Additional Periodic Trends Practice a) Define: Atomic Radius b) Define: Electronegativity c) Define: Ionization Energy d) Define: Ionic Radius For each of the following group of elements, specify which would have the largest and smallest atomic radii: 18) Si N 19) Ca Al 20) Pb Bi Sn 21) O S F Cl 22) Ge Si P As S For each of the following group of elements, specify which would have the largest and smallest electronegativity: 23) Si N 24) Ca Al 25) Pb Bi Sn 26) O S F Cl 27) Ge Si P As S For each of the following group of elements, specify which would have the largest and smallest ionic radius: 28) N 3- N 29) Ba Ba 2+ 30) K 1+ Rb 1+ Ag 1+ 31) N 3- O 2- F 1- Cl 1-32) Ge 4+ Se 2- F 1- As 3- S 2- For each of the following group of elements, specify which would have the largest and smallest ionization energy: 33) C N 34) Ba Sr 35) K Rb Ag 36) N O F Cl 37) Ge Se F As S 1 st Ionization Energy (kj/mol) 2 nd Ionization Energy (kj/mol) 3 rd Ionization Energy (kj/mol) 4 th Ionization Energy (kj/mol) Na Mg ,540 Al ,577 e) Sodium has a significantly higher second ionization energy than Magnesium. Explain 10

11 Practice Problems Section 6.1 Development of the Modern Periodic Table Use each of the terms below just once to complete the passage. octaves elements protons atomic mass properties periodic law atomic number Henry Moseley Dmitri Mendeleev nine eight accepted The table below was developed by John Newlands and is based on a relationship called the law of (1). According to this law, the properties of the elements repeated every (2) elements. Thus, for example, element two and element (3) have similar properties. The law of octaves did not work for all the known elements and was not generally (4) H Li G Bo C N O F Na Mg Al Si P S The first periodic table is mostly credited to (5). In his table, the elements were arranged according to increasing (6). One important result of this table was that the existence and properties of undiscovered (7) could be predicted. The element in the modern periodic table are arranged according to increasing (8), as a result of the work of (9). This arrangement is based on number of (10) in the nucleus of an atom of the element. The modern form of the periodic table results in the (11), which states that when elements are arranged according to increasing atomic number, there is a periodic repetition of their chemical and physical (12). Use the information in the box on the left taken from the periodic table to complete the table on the right. 7 Atomic Mass 13. N Atomic Number 14. Nitrogen Electron Configuration Chemical Name 16. [He]2s 2 2p 3 Chemical Symbol 17. For each item in Column A, write the letter of the matching item in Column B. Column A Column B 18. A column on the periodic table a. metals 19. A row on the periodic table b. group 20. Elements in groups 1, 2, and 13 to 18 c. period 21. Elements that are shiny and conduct d. representative elements electricity e. transition elements 22. Elements in groups 3 to 12 In the space at the left, write true if the statement is true; if the statement is false, change the italicized word or phrase to make it true. 23.There are two main classifications of elements. 24.More than three-fourths of the elements in the periodic table are nonmetals. 25.Group 1 elements (except for hydrogen) are known as the alkali metals. 26.Group 13 elements are the alkaline earth metals. 27.Group 17 elements are highly reactive nonmetals known as halogens. 28.Group 18 elements are very unreactive elements known as transition metals. 29.Metalloids have properties of both metals and inner transition metals. 11

12 Section 6.2 Classification of the Elements Match each element in Column A with the element in Column B that has the most similar chemical properties. Column A Column B 1.arsenic (As) a. boron (B) 2.bromine (Br) b. cesium (Cs) 3.cadmium (Cd) c. chromium (Cr) 4.gallium (Ga) d. cobalt (Co) 5.germanium (Ge) e. hafnium (Hf) 6.iridium (Ir) f. iodine (I) 7.magnesium (Mg) g. iron (Fe) 8.neon (Ne) h. nitrogen (N) 9.nickel (Ni) i. platinum (Pt) 10.osmium (Os) j. scandium (Sc) 11.sodium (Na) k. silicon (Si) 12.tellurium (Te) l. strontium (Sr) 13.tungsten (W) m. sulfur (S) 14.yttrium (Y) n. zinc (Z) 15.zirconium (Zr) o. xenon (Xe) Answer the following questions. 16. Why do sodium and potassium, which belong to the same group in the periodic table, have similar chemical properties? 17. How is the energy level of an element s valence electrons related to its period on the periodic table? Give an example. 18. Into how many blocks is the periodic table divided? 19. What groups of elements does the s-block contain? 20. Why does the s-block portion of the periodic table span two groups? 21. What groups of elements does the p-block contain? 22. Why are members of group 18 virtually unreactive? 23. How many d-block elements are there? 24. What groups of elements does the d-block contain? 25. Why does the f-block portion of the periodic table span 14 groups? 26. What is the electron configuration of the element in period 3, group 16? 27. Which group is the alkali metals? 28. Which group forms salts with metals? 29. Which group is inherently nonreactive? 30. Which group forms a 2+ charge? 12

13 Section 6.3 Periodic Trends Circle the letter of the choice that best completes the statement or answers the question. 1. Atomic radii cannot be measured directly because the electron cloud surrounding the nucleus does not have a clearly defined a. charge. b. mass. c. outer edge. d. probability. 2. Which diagram best represents the group and period trends in atomic radii in the periodic table? a. c. b. d. 3. The general trend in the radius of an atom moving down a group is partially accounted for by the a. decrease in the mass of the nucleus. c. increase in the charge of the nucleus. b. fewer number of filled orbitals. d. shielding of the outer electrons by inner electrons. 4. A(n) is an atom, or bonded group of atoms, that has a positive or negative charge. a. halogen b. ion c. isotope d. molecule 5. An atom becomes negatively charged by a. gaining an electron. b. gaining a proton. c. losing an electron. d. losing a neutron. 6. Which diagram best represents the relationship between the diameter of a sodium atom and the diameter of a positive sodium ion? a. b. c. Answer the following questions. 7. What is ionization energy? 8. Explain why an atom with a high ionization-energy value is not likely to form a positive ion. 9. What is the period trend in the first ionization energies? Why? 10. What is the group trend in the first ionization energies? Why? 11. State the octet rule. 12. What does the electronegativity of an element indicate? 13. What are the period and group trends in electronegativities? 13

14 Families of the table 1) Give the location of the alkali metals, alkaline earth metals, halogens, noble gases, transition metals, and inner transition metals. 2) List the properties of the alkali metals 3) List the properties of the alkaline earth metals 4) List the properties of the halogens 5) List the properties of the noble gases 6) Why are transition metals and inner transition metals so diverse? Even More Practice Section 6.1: 1) Describe the development of the modern periodic table. Include contributions made by Newlands, Mendeleev, and Mosely. 2) Indicate the locations of groups, periods, metals, nonmetals, and metalloids on the blank table. 3) Describe the general characteristics of metals, nonmetals, and metalloids 4) Identify each of the following as a representative element or a transition element: a. Lithium b. Platinum c. Promethium (Pm) d. Carbon 5) For each of the given elments, list two other elements with similar chemical properties: a. Iodine b. Barium c. Iron 6) An unknown element has a chemical behavior similar to silicon and lead. The unknown element has a mass greater than that of sulfur but less than that of cadmium. Use the periodic table to determine the identity of this unknown element. Section 6.2: 1) Explain why elements in the same group on the periodic table have similar chemical properties 2) Given each of the following valence electron configurations, determine which block of the periodic table the element is in a. s 2 p 4 b. s 1 c. s 2 d 1 d. s 2 p 1 12

15 3) Describe how each of the following are related: a. Group number and valence electrons for the representative elements b. Energy level of valence electrons and period number 4) Without using the periodic table; determine the group, period, and block of an atom with a noble gas configuration of [Ne]3s 2 3p 4 5) A gaseous element is a poor conductor of heat and electricity, and is extremely nonreactive. Is the element likely to be a metal, nonmetal, or metalloid? Where would the element be located on the periodic table? Explain 6) Label the s, p, d, and f blocks on the blank periodic table. Section 6.3: 1) Use arrows and labels to compare group and period trends for atomic and ionic radii, ionization energies, and electronegativities on the periodic table on the left 2) What charge does group 1, group 2, and group 13 form when an ion? 3) What charge does group 15, group 16, and group 17 form when an ion? 4) Why does group 18 not form charges at all? 5) Explain how the period and group trends in atomic radii are related to electron configuration 6) Which has the largest atomic radius: a. Nitrogen b. Antimony c. Arsenic 7) Which of the above has the smallest atomic radius: 8) For each of the following properties indicate whether fluorine or bromine has a larger value: a. Electronegativity b. Ionic radius c. Atomic radius d. Ionization energy 9) Explain why it takes more energy to remove the second electron from a lithium atom than it does to remove the first 13

16 PEAT: Performance Enhancing Assessment Task Objectives Name: Period: 1. Classify elements in the appropriate family on the periodic table 2. Explain different properties of elements based on their placement in the periodic table 3. Identify trends of the periodic table Questions: 1. An unknown element has four valence electrons and an atomic mass between 26 and 33 amu, predict the element that this most likely is. 2. Explain your reasoning 3. Identify the atomic radius of your unknown element, given the information in the diagram below: 4. Explain your reasoning 5. Based on the information below, predict the noble gas configuration of your unknown element: B: [He]2s 2 2p 1 C: [He]2s 2 2p 2 N: [He]2s 2 2p 3 O: [He]2s 2 2p 4 Ga: [Ar]4s 2 3d 10 4p 1 Ge: [Ar]4s 2 3d 10 4p 2 As: [Ar]4s 2 3d 10 4p 3 Se: [Ar]4s 2 3d 10 4p 4 6. Explain your reasoning 7. Your unknown element reacts with oxygen in a specific manner. Given the information below, predict the formula for your compound with oxygen. B: B2O3 C: CO2 N: N2O5 O: O2 Ga: Ga2O3 Ge: GeO2 As: As2O5 Se: SeO2 8. Explain your reasoning 14

17 Chapter 6/7 Test Review vs 1 NAME: Period ) Be able to match the 5 scientists who contributed to the arrangement of the periodic table. Lavoisier Dobereiner Newlands Mendeleev Moseley 2) Be able to identify elements based on their electron configuration (incl. noble gas) and write the electron configuration (incl. noble gas) for any element. a) What element has the electron configuration of: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 b) What element has the electron configuration of: Uranium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 c) What element has the electron configuration of: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 Barium d) What element has the electron configuration of: [Rn]7s 2 5f 12 e) Give the full and noble gas configuration for: Palladium Aluminum 3)Be able to determine the number of valence electrons for any element. Be able to draw Lewis dots for any element. a) Given the following elements, give the # of valence electrons (VE). Then draw the Lewis dots for each. O K Si He b) Given the following electron configurations, give the # of VE. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 4) Be able to identify which element has a higher shielding effect. Differentiate between group and period shielding effects. a) How does shielding changes as you go down a group? How does it change going L-> R across a period? b) Identify which element of each pair has the higher shielding effect. Li or K Se or I Ra or Be 5) Be able to give properties of metals, nonmetals, and metalloids. Be able to identify an element as one of those three. a) List four properties for each: Metal c) Which category does each of the following elements fall into: Hg Sr P Po Nonmetal b) What sets metalloids apart? 6) Be able to identify the 4 major trends. Be able to compare elements based on atomic radii, electronegativity, ionization energy, and ionic radii. a) How does each of the following trends change as you go L-> R across a period: Atomic Radii Ion Size Electronegativity Ionization Energy b) How does each of the following trends change as you go down a group: Atomic Radii Ion Size Electronegativity Ionization Energy c) Of the following sets, which particle is the largest? S P Cl Na Al Cl O O 2- Ba Ba 2+ Si 4+ N 3- d) Of the following sets, which has the highest electronegativity? Li, Cs, Be Sn, Sb, Sr P, O, Ne e) Of the following sets, which has the highest ionization energy? Cs, Cl, Fe C, Pb, Sn I, Zr, Sr 8) Be able to identify an element as representative element, transition, inner transition, and the S P D F blocks. a) Identify each of the following as Representative, Transition, or Inner Transition, and which block it comes from: Au Be N Th 15

18 9) Be able to identify an element as one of the 4 main families. Be able to identify an element as one of the four families based on properties. a) Of the four main families (Alkali, Alkaline Earth, d) Identify which family it falls into based on a description Halogens, Noble Gas) which do each of these elements fall of its properties: into? It can be cut with a butter knife, must be stored in F oil/kerosene Mg Na Never found pure in nature, very high melting point Ar Unreactive with all elements, found naturally in our b) What is the most reactive metal family? Element? atmosphere c) what is the most reactive nonmetal family? Element? When exposed to metal it will form salts, highly corrosive 10) Be able to name and write the formula of any monoatomic or polyatomic ion. a) Name each of the following ions: b) Write the formula for each of the following (including F 1- charges): Al 3+ Sodium Ion SO4 2- Bromide Ion ClO 1- Bromate Ion CO4 2- Hyposulfite Ion PO3 3- Perchlorate Ion Nitrite Ion 11) Remember the following vocabulary: Periodic Law Shielding Effect Electronegativity Ionization Energy Ion Cation Anion Polyatomic Ion Monoatomic Ion Group Period Metallic Bond Sea of Electrons Alloy Interstitial Alloy Substitutional Alloy Be able to label a periodic table with: atomic radius trend, electronegativity trend, ionic radius trend, group, period, spdf blocks, alkali metals, alkaline earth metals, halogens, noble gases, transition metals, inner transition metals, representative elements. 1 st Ionization Energy 2 nd Ionization Energy 3 rd Ionization Energy 4 th Ionization Energy Na Mg ,540 Al ,577 Rb a) Magnesium has significantly higher ionization energy than Aluminum. Explain b) What does the ionization energy tell you about the number of valence electrons? c) Why does a sodium atom decrease in size when it goes from a neutral atom to an ion? d) Sodium and Rubidium are both in column 1. Explain the differences in the ionization energy. 16

19 1) B Al Ga In a) Most metallic? b) Least metallic? c) Lowest electronegativity? d) Highest ionization energy? e) Largest atomic radius? f) Which are non-metals? g) Which are conductors? 2) Ca Sr Ba Mg a) Most metallic? b) Least metallic? c) Lowest electronegativity? d) Highest ionization energy? e) Largest atomic radius? f) Which are non-metals? g) Which are conductors? h) List them in order from small to large. 3) Si S Al P a) Highest metallic character? b) Highest electron affinity? c) Lowest I.E.? d) Smallest? e) Which are metalloids? f) List in order of increasing size g) Which are non-metals? h) Lowest electronegativity? PERIODIC TRENDS PRACTICE QUESTIONS 4) Define atomic radius: In each group, arrange the following elements in order of decreasing atomic size: a) sulfur, chlorine, aluminum, and sodium. b) carbon, germanium, lead, silicon c) cesium, lead, bismuth, barium 5) Define ionization energy: In each group, arrange the following elements in order of increasing ionization energy: a) Be, Mg, Sr b) Bi, Cs, Ba c) Na, Al, S 6) In each of the following pairs, circle which element is the most electronegative. a) chlorine, fluorine c) magnesium, neon b) carbon, nitrogen d) arsenic, calcium h) Highest shielding effect? i) List them in order from small to large. j) Would they gain or lose electrons to form ions? What charge would the ion have? k) Would these be considered cations or anions? l) Would the ion be larger or smaller than the atom? i) Would they gain or lose electrons to form ions? What charge would the ion have? j) Would the ion be larger or smaller than the atom? k) Lowest shielding effect? l) Would the ions be cations or anions? m) What is the special name given to the group in which these elements belong? n) What are some characteristics of this group? i) Which would rather gain electrons? j) Which would rather lose electrons? k) Which ones would form cations? Would these be larger or smaller than the neutral atom? l) Which ones would form anions? Would these be larger or smaller than the neutral atom? m) lowest shielding effect? 7) Give the name and symbol for the element found at each of the following locations on the periodic table: a) Group 1, period 4. c) Group 16, period 3 b) Group 13, period 3. d) Group 2, period 6 8) Use the chart to answer the following questions: a) Which group does element Y most likely belong in? How many valence electrons does it have? b) Which group does element Z most likely belong in? What is its most likely charge? 54

20 Periodic Table Review vs 2 Directions: Put all letters that apply in the blanks next to each question on the right Transition metals 2. Metalloids J 3. Period 3, Group 13 A C 4. Halogens B G 5. Reacts violently with water L H 6. 6 valence electrons D K E 7. Is a stable electron element 8. Lanthanides 9. Alkaline earth metals 10. Alkali metals I 11. Halogens 12. Ending electron configuration of p 5 Directions: Identify which family each of the following belongs to. 13. Electron configuration of 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d Typically forms an anion with a 1- charge 15. Electron configuration of [Kr] 5s 2 4d 10 5p Typically forms a cation by losing two electrons 17. Contains Be, Mg and Ba 18. Contains U, Es and Am Directions: Identify what is meant by each of the following periodic trends. 19. Shielding effect: 20. Atomic radius: 21. Ionic radius: 22. Electronegativity: 23. Explain the difference between first and second ionization energy. 24. Draw arrows to represent increasing trends for each of the following periodic trends. Atomic Radius Electronegativity First ionization energy 25. Circle the larger particle in each of the following pairs. a. Na Li d. Br I b. F F -1 e. Cs Ba c. K K +1 f. Ne Ar 26. Predict the ion charge when given the following electron configurations of neutral atoms. 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 55

21 d. 1s 2 2s 2 2p 5 e. 1s 2 2s 2 2p 1 f. 1s 2 2s 2 2p Circle the element that has the highest first ionization energy. a. Li Na b. Cs Ba c. F Ne d. Kr Rb e. Cl Br f. S Cl 28. List the following atoms in order of increasing atomic radius: N, Au, Al: 29. List the following atoms in order of decreasing atomic radius: Cl, K, Cu: 30. Explain how cations and anions are formed. Explain whether metals and nonmetals form cations or anions. 31. How does the formation of a cation or anion affect the size of the atom? 32. For each of the following pairs, circle the atom or ion that would have the larger radius. a. S O c. S 2- O 2- b. Na 1+ K 1+ d. Ca Ca For each of the following groups of elements, circle the smaller ion. e. Na K f. F F 1- a. K 1+ Ca 2+ b. C 4+ C 4- c. O 2- F 1- Cl 1- O 2- d. Fe 2+ Fe 3+ Ag 1+ Pb List the following atoms in order of increasing electronegativity: O, Al, Ca 35. List the following atoms in order of decreasing electronegativity: Cl, K, Cu 36. Explain the trend in reactivity as you move down a group. 37. The table below shows the first three ionization energies for atoms of four elements from the third period of the periodic table. The elements are numbered randomly. Use the information below to answer the following questions. (a) Identify element 3 and write its electron configuration. Explain your reasoning. (b) What is the chemical symbol for element 2? (c) A neutral atom of which of the four elements has the smallest radius? 38. (a) In general, there is an increase in the first ionization energy from Li to Ne. (b) The first ionization energy of B is lower than that of Be. Explain why. (c) The first ionization energy of O is lower than that of N. Explain why. (d) Predict how the 1 st ionization energy of Na compares to those of Li and of Ne. Explain. 56

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