Unit 5 Elements and their Properties

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1 Unit 5 Elements and their Properties 1. In 1871, Russian chemist Mendeleev created the forerunner of the modern periodic table. 2. The elements in Mendeleev's table were arranged in order of increasing _atomic mass. 3. What is the Law of Octaves and who proposed it? Newlands when arranged in order of increasing atomic mass, the properties of the elements repeat by eights 4. As a result of Henry Moseley's work, the modern periodic table is arranged according to increasing atomic number. 5. How can the periodic table be used to predict the properties of the elements? Elements in the same family have the same number of valence electrons, thus similar chemical and physical properties 6. Is the following true or false? The subatomic particles that play the most significant role in determining the physical and chemical properties of an element are electrons. true 7. List four things that can be deduced about an element using the periodic table. oxidation number, metal or nonmetal, relative atomic size, relative reactivity 8. The atomic number of an element indicates the number of _protons_ in the nucleus. 9. What determines an element s chemical properties? Electron configuration Are the following statements true or false? If a statement is false correct it to make the statement true. 10. The electron configurations of all elements in Group 1(1A) end in the same orbital notation. true 11. In the electron configuration for the outer energy level of potassium, 4s 1, the coefficient 4 indicates the period number. 12. Atoms with full outer levels are rarely reactive. H. Cannon, C. Clapper and T. Guillot Klein High School

2 13. Elements such as silicon are called metalloids because they have properties of both metals and nonmetals. true 14. Elements with three or fewer electrons in the outer energy level are usually metals. true 15. The electron configurations of hydrogen and helium are not similar, so each element is in a separate column of the periodic table. true 16. The elements in columns 3 through 12 (IIIB through IlB) are called the transition metals. 17. Each time a new principal energy level is started, a new row in the periodic table begins. true 18. The lanthanoid series contains the elements with the 4f sublevel filling. true 19. In the periodic table, a horizontal row of elements is called a period. true Answer the following questions. 20. Describe the general positioning of metals and nonmetals in the periodic table metals on the left nonmetals on the right 21. Describe the properties of nonmetals and give examples. 5 or more valence electrons, high electronegativity, form anions, relatively small atomic radius 22. What is the main reason that atoms react with each other? to achieve pseudo- noble gas configuation 23. State the octet rule There are never more than 8 electrons in the outer level of an atom 24. Why are neon and helium placed in the same column in the periodic table? they both have full outer electron levels 25. Atoms of most nonmetallic elements achieve noble-gas electron configurations by gaining electrons to become_anions_ or negatively charged ions 26. What property of nonmetallic elements makes them more likely to gain electrons than lose electrons? High electronegativity 27. Metallic atoms tend to lose their valence electrons to produce a(n) _cation_ or a positively charged ion. Most nonmetallic atoms achieve a complete octet by gaining or _sharing_electrons. 28. What are valence electrons? Electrons that are gained lost or shared to form chemical bonds 5-2

3 29. The valence electrons largely determine the _chemical and physical properties_ of an element and are usually the only electrons used in _bonding_. 30. Is the following sentence true or false? The group number of an element in the periodic table is related to the number of valence electrons it has. true 31. Complete the table about classifying elements according to their electron configuration. Category Description of Electron Group Configuration Noble gases Full s and p sublevels 8 (18) Representative elements Transition metals Outermost s sublevel and Group B nearby d sublevel contain electrons Inner transition metals Outermost s sublevel and nearby f sublevel contain electrons Group B 32. If an element is in the third row of the periodic table in column 2, what is the final electron configuration? Which element is it? 3s 2 Mg 33. Tin element number 50 - is located in which row of the periodic table? Which column? Give the ending configuration for tin. 5 th row; 5s 2 4d 10 5p Gold is a transition metal. What distinction with regard to electron configuration do all transition metals have in common? What sublevel contains the primary valence electrons for all transition metals? All transition metals have the d sublevel filling, the primary valence electrons are in the s sublevel 35. Where on the periodic table would you find the element with the ending electron configuration 5s 2 4d 10 5p 2? Would you predict it to be a metal, a nonmetal or a metalloid? 5 th row, column 4, metalloid 5-3

4 Periodic Trends 1. Define electronegativity. The tendency of an element to attract electrons from another atom to form a chemical bond 2. How does electronegativity vary as the atomic number of an element increases within the same period of the periodic table? increases 3. What trend do you see in the relative electronegativity values of elements within a group? Within a period? Increase left to right, decrease top to bottom Indicate if statements 4-7 are true or false when considering electronegativity. 4. The electronegativity values of the noble gases are all zero. 5. The element with the highest electronegativity is fluorine. 6. Nonmetals have higher electronegativity values than metals. true 7. Electronegativity values can help predict the types of bonds atoms form. true 8. What trend do you see in the sizes of elements within a group? Within a period? Decrease left to right, increase top to bottom 9. What is the name of the effect that is responsible for differences in atomic radii between elements in the same group? shielding 10. Number the following elements according to relative size from largest (1) to smallest (4) He Ca P Cs 11. Metallic elements easily form _cat_ ions, nonmetallic elements readily form _an_ ions. 5-4

5 12. Cations are usually smaller than the neutral atoms from which they form. true 13. Anions are usually larger than the neutral atoms from which they form. 14. Within a period, cations with greater charge have larger ionic radii. true 15. Within a group, cations with greater atomic number have larger ionic radii. 16. Which ion has the larger ionic radius : Ca +2 or Cl -1? 17. _ionization energy is the energy required to overcome the attraction of the nuclear charge and remove an electron from a gaseous atom. 18. Why does ionization energy increase as you move across a period, but decrease as you move down a group? Follows electronegativity trend 19. Describe the trend in atomic radii within a group by completing this sentence: Within a group, as atomic number increases, atomic radius _increases. Explain. Additional energy levels increased shielding 20. Describe the trend in atomic radii within a period by completing this sentence: Within a period, as atomic number increases, atomic radius _decreases. Explain. Increased nuclear charge 5-5

6 Periodic Table Review Worksheet 1. Identify on a periodic table: 2. period d) metalloids g) inert gas 3. transition metal e) alkaline earth h) rare earth 4. c) family or group f) alkali metal i) halogen 5. Which element is more active and WHY? a) oxygen or sulfur b) calcium or barium 6. In what period and what group does selenium (At. No. 34) belong'? 7. In general, in the "main groups", where are the elements which are: a) more metallic b) less metallic c) more non-metallic d) less non-metallic e) most active metals f) most active non-metals 8. Which are the most/least active transition metals? 9. An atom with outermost electron in the fourth energy level is in what period? 10. Which elements tend to gain electrons? Lose electrons? 11. What element does not clearly fit one place on the periodic table? 12. What family would include a neutral atom having a) one electron in its outermost energy level b) two electrons c) 8 electrons d) 7 electrons 13. What atomic structure makes one period different from another? 14. What atomic structure makes one family different from another? 15. How was Mendeleev able to predict the properties of the elements that were missing from his periodic table? 16. Define "ionization energy". Why is it important? 17. Why is it harder to remove one electron from calcium than from potassium? 5-6

7 18. How does ionization energy change as you go from left to right across the periodic table? top to bottom? 19. How does atomic size change from left to right across the periodic table? Top to bottom? 20. How does nuclear charge change from left to right across the periodic table? Top bottom? 21. Which group of elements will have 1 valence electron? 22. Which group of elements will have 7 valence electrons? 23. Which group of elements will have 8 valence electrons? 24. Refer to a table of atomic and ionic radii, if necessary. To answer the following questions. a. Within a period, does the size of atoms generally increase or decrease with increasing atomic number? b. Within a family, does the size of atoms increase or decrease with increasing atomic number? c. When metallic atoms lose electrons, do they form ions that are smaller or larger than the original atoms? d. When nonmetallic atoms gain electrons, do they form ions that are smaller or larger than the original atoms? 25. Circle the larger particle in each of the following pairs. a. Na or Li b. Br or I c. F or F -1 d. Cs or Ba e. K or K +1 f. Ne or Ar 5-7

8 A PERIODIC FUNCT10N OF ATOMIC NUMBER OBJECTIVE: Given a table of atomic radii the student will determine whether or not atomic radius is a periodic function of atomic number by constructing, a graph which plots this property versus atomic number and interpreting the results. PROCEDURE: Using the data on the data sheet, plot atomic radius versus atomic number, atomic number vs. first ionization potential and atomic number vs. boiling point, on separate pages, for the first 50 elements. (atomic number on X-axis) data can be found in the second 6 weeks portion of the teacher website 1. Orient the graph paper provided so that the holes in the margin are in a horizontal position farthest from you. 2. Using a straight-edge draw the axes along the bottom and left edges of the paper. Use pencil lightly first, then go over later in ink. 3. The plot should make the best use of the space provided on the graph paper so you must carefully choose and appropriate scale for each axis in order to make the most efficient use of the paper. The origin on an axis will be determined by the range of numbers to be used. The first number does not have to be zero, but can be whatever is most convenient. However in this case, zero is probably the most appropriate. For example, the atomic numbers to be used run from 1 to 50. The longer side of the graph paper is divided into 100 divisions. How many divisions must you skip between 0 and 1 to use the whole sheet? How should you arrange the Y-axis to use at least half the paper? 4. Once the scale has been determined for each axis, mark it off on the graph paper. Don't crowd marks, or numbers You may skip regular intervals for the sake of neatness. 5. For each atomic number, plot a point corresponding to its atomic radius. Draw a small circle around each point. Use pencil first, then trace over in ink when you re sure it's correct. 6. When all of the points have been plotted, connect all of the points consecutively with straight lines. 7. Directly above or to the outside of each point plotted, write the chemical symbol for that element. 8. Recreate this graph using the graphing tools in Microsoft Excel or another computer spreadsheet program. Both graphs must be turned in for a complete lab 5-8

9 report. Remember the graphs will be the analysis part of your report since you are comparing data. QUESTIONS to be answered in your conclusion 1. Does the pattern which develops repeat at regular intervals, irregular intervals, or at regular intervals of increasing length? 2. Is the atomic radius a periodic function of atomic number? 3. How can you determine how many elements are in each 'period" using your graph'' 4. What group of elements are represented by the peak on your graph? 5. How are the halogens identified by their positions on the graph? 6. How could you determine whether "lonization Potential" were a, periodic function of atomic number? 5-9

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