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Periodic Table

Periodic Table. In 1700, only 13 elements had been discovered. As chemists began using the scientific method to search for elements, the rate of discovery increased. How would chemists known when they found them all? They needed a logical way to organize the elements.

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Periodic Table

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  1. Periodic Table

  2. In 1700, only 13 elements had been discovered. • As chemists began using the scientific method to search for elements, the rate of discovery increased. • How would chemists known when they found them all? • They needed a logical way to organize the elements. •  Chemists used the properties of elements to sort them into groups

  3. J.W. Dobereiner (1780-1849) • published a classification system in which elements were grouped in triads. • triad: a set of three elements with similar properties. • all the known elements couldn’t fit in triads

  4. Dmitri Mendeleev (1869) • arranged elements in a periodic table in order of increasing atomic mass. • Left spaces for he predicted that elements would be discovered to fill in those spaces • Predicted their properties based on their location in the table • Had problems organizing by atomic mass

  5. Other Scientist • Lothar Meyer – later in 1869, published a nearly identical table as Mendeleev •  H.G.J. Moseley (1913) – determined the atomic number for each known element • Modern periodic table is organized by atomic number 

  6. Periodic Law • When elements are arranged in order of increasing atomic number, there is a periodic repetition of their physical and chemical properties

  7. Modern Periodic Table

  8. Modern Periodic Table • 1. What are groups? • Vertical columns • 2. What is another name for a group? • families • 3. What are the periods? • Horizontal rows • 4. What group is oxygen in? • 16 • 5. What period contains potassium? • 4

  9. PT • There are three main categories in which the elements are divided: metals, nonmetals, and metalloids. The metals are to the left (6) of the jagged line. The nonmetals are to the right (7) of the jagged line. The metalloids are along (8) the jagged line.

  10. 9. What category do most of the elements fall into? • metals • 10. What are some characteristics of metals? • Shiny, good conductors of heat and electricity, most are solid at room temperature, malleable, and ductile • 11. What are some characteristics of nonmetals? • Not shinny, poor conductors of heat and electricity, most are gases at room temperature, not malleable, not ductile

  11. Metalloids • 12. Metalloids have characteristics of both metals and nonmetals. There are only eight metalloids. Which elements are metalloids? • Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polnium, Astatine

  12. Family Names • 13. Which group is called the alkali metals? • 1 • 14. Which group is called the alkaline earth metals? • 2 • 15. Which groups are called the transition metals? • 3-12 • 16. Where are the inner transition metals found? • At the bottom of the table • 17. Which group is called the halogens? • 17 • 18. Which group is called the noble gases? • 18

  13. Block Names • 19. Where’s the s-block? • 1-2 • 20. Where’s the p-block? • 13-18 • 21. Where’s the d-block? • 3-12 • 22. Where’ s the f-block? • At the bottom

  14. Notes on Valence Electrons

  15. Notes on Valence Electrons • The chemical properties of elements are determined by the number of their valence electrons. • 23. What are valence electrons? • Electrons in the highest principle energy level, the outermost e-

  16. Valence Electrons • To find the number of valence electrons, find the highest quantum number in the electron configuration. Add up the electrons that have that quantum number. • 24. How many valence electrons does sulfur have? • 6 (#16) • 25. How many valence electrons does cesium have? • 1 (#55) • 26. How many valence electrons does francium have? • 1 (#87)

  17. Valence Electrons • 27. Write out the shortcut electron configuration for H, Li, Na, and K and determine the number of valence electrons. • H: 1s1 1 ve- • Na: [Ne]3s11 ve- • Li: [He]2s1 1 ve- • K: [Ar]4s1 1 ve-

  18. VE- Continued • 28. All four of those elements are in what group? • 1 • 29. What do all four have in common? • All have one valence electron • 30. This is not a coincidence. All elements in the same group have the same number of valence electrons. Therefore, what else do they have in common? • Same chemical property

  19. VE- from Periodic Table • 31. How many valence electrons do elements in group 2 have? • 2 • 32. How many valence electrons do elements in group 13 have? • 3 • 33. How many valence electrons do elements in group 14 have? • 4 • 34. How many valence electrons do elements in group 15 have? • 5 • 35. How many valence electrons do elements in group 16 have? • 6 • 36. How many valence electrons do elements in group 17 have? • 7 • 37. How many valence electrons do elements in group 18 have? • 8

  20. D-Block • 38. All elements in the d-bock have the same number of valence electrons. How many? 2

  21. Periodic Trends

  22. Atomic Radius • 1. How is atomic radius defined? • Half the distance between the nuclei of the two atoms of the same element when the atoms are joined • 2. What is the atomic radius trend within a period? Explain this trend. • Increases to the left • LR: With increased number of protons, there is a stronger pull of the outer shell to the nucleus • 3. What is the atomic radius trend within a group? Explain this trend. • Increases to the bottom • Top  Bottom: As you add energy levels you increase the shielding effect from protons, inward pull • 4. Which has the largest atomic radius: C, F, Be, or Li? • Li > Be> C > F

  23. Ionic Radius • 5. What is an ion? • Atom with a + or – charge, due to the loss or gain of electrons, respectively • 6. When atoms lose electrons, they become positive and smaller. Why? • Greater nuclear pull • 7. When atoms gain electrons, they become negative and larger. Why? • Greater electron repulsion

  24. Ionic Radius • 8. What is the ionic radius trend within a period? • Increases left for cations and anions • 9. What is the ionization radius trend within a group? Increases as you go down

  25. Ionization Energy • 10. What is ionization energy? • Energy required to remove an e- from an atom • 11. Write the chemical equation which demonstrates what is going on when a positive ion is formed? • Li  Li+ + e-

  26. Ionization Energy • 12. A high ionization energy indicates that the atom has a strong hold over its electrons. • 13. A low ionization energy indicates that the atom has a weak hold over its electrons. • 14. Once you remove one electron, is it harder or easier to remove another one? Why? • Harder; protons have a stronger pull on fewer electrons.

  27. Ionization Energy • 15. For each element there is a huge jump in the ionization energy required at some point in taking away electrons. How does this relate to valence electrons? • Once valence electrons are gone, almost impossible to take another because that is when it is most stable • 16. What is the ionization energy trend within a period? Explain. • Increase to the right; • nuclear charge increases to the right and shielding effect remain constant • 17. What is the ionization energy trend within a group? Explain. • Increase going up; • as size increases, nuclear charge has a smaller effect on valence electrons, therefore less energy to remove

  28. Electronegativity • 18. What is electronegativity? • The ability of an atom of an element to attract electrons when the atom is in a compound • 19. Why do noble gases not have electronegativity values? • Don’t form many compounds • 20. What is the electronegativity trend within a period? • Increases toward the right • 21. What is the electronegativity trend within a group? • Increases toward the top • F has the greatest electronegativity

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