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Chapter 6

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Chapter 6

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  1. Chapter 6 The Periodic Table

  2. Periodic Table Activity

  3. The periodic table is arranged by elements with similar properties. • What atomic particle gives them their properties? • Valence e-

  4. Scientists and the Periodic Table • Antoine Lavoisier • 1790’s • First to compile a list of the 23 known elements

  5. John Newlands • 1864 • Law of Octaves • Noticed properties repeated every eight elements • Based on atomic mass • Transition elements messed him up • They all tend to be similar

  6. Meyer and Mendeleev • 1869 • Demonstrated the relationship between atomic mass and elemental properties. • Mendeleev published first and predicted elements. • He knew there were missing elements and predicted their properties. • Mendeleev is known as the Father of the Periodic Table. • What was the problem with Mendeleev’s table?

  7. Moseley • 1913 • Arranged elements by atomic number • Everything worked • This is our current table

  8. The Periodic Table • Metals are to the left of the metalloid line. • Metalloids touch the stair step line. • All but Al • Nonmetals are to the right of the metalloid line • Everything above 92 is a synthetic element • They are short lived and radioactive

  9. Arrangement of Elements • Columns are called groups or families • Elements in each group have similar characteristics • Show the number of valence electrons • Rows are called periods • They transition from reactive metals to nonmetals • Periods show the energy levels for electrons • Main or representative group elements are groups 1, 2, and 13-18.

  10. Groups • Group 1 • Alkali metals • Most reactive group • 1 valence electron • Never found free in nature • Does not include hydrogen • Why is hydrogen there? • 1 valence electron

  11. Group 2 • Alkaline Earth Metals • 2 valence e- • Tend not to be free in nature • Groups 3-12 • Transition elements • All have similar characteristics • Inner Transition Elements are at the bottom of the periodic table.

  12. Groups 13-16 • Named by the element at the top of the group • Arranged by number of valence electrons • Group 17 • Halogens • 7 valence electrons • Most reactive nonmetals • Group 18 • Noble or Inert Gases • 8 valence e- • Stable • Why are they stable? Use electron configuration in your answer.

  13. Example • Pg 162 #7

  14. Periodic Trends • Atomic size • How close an atom can lie to another atom • Atomic radii • Measures from the center of the atom to the edge of the electron cloud • Atomic radii increases as you go down and decreases as you go right.

  15. Ionic Radii • What is an ion? • Increases as you go down. Cations decreases as you move right until you get to anions. At the start of the anions you will have the largest nonmetals then radii will decrease again as you continue to move right.

  16. Ionization Energy • Energy required to remove an e- from an atom in its gaseous state. • Pg. 168 • 1 e- is removed easier but it becomes harder to remove with each following electron. • Decreases as you go down and increases as you move right. • Why does it decrease as you go down? • The electrons are further from the nucleus

  17. Electronegativity • An atoms attraction for an electron in a chemical bond • Pg. 169 • Decreases as you go down and increases as you move right. • Why does Fluorine have the highest electronegativity? • Why does group 18 show no electronegativity?

  18. Homework • Pg. 175 # 40, 47, 50, 63, 65, 67 – 69, 72, 79

  19. What are 2 dominate features for each group? • Number of valence electrons • Size increases as you move down • Why do we have these trends? • The trend is not the explanation it is an observation

  20. Two Concepts to Explain Periodic Trends • 1 • Size increases as you move down. • As you move down you add energy levels. Valence e- are put further and further away from the nucleus. • Energy decreases as you move down. • As the atom gets bigger the electrons are moved further away from the nucleus making it easier to remove electrons. There is also less of an attraction for an e- as the atoms get bigger.

  21. 2 • Zeff • Effective nuclear charge • The positive charge that an e- feels from the nucleus • Tends to increase across a period and decrease down a group.

  22. Why does the Zeff increase as you move across? • You increase a proton and an electron • Think about energy levels. We will use Na, Mg, and Cl as examples. They are all in the same energy level which means their electrons are the same distance from the nucleus. When moving from Na to Mg to Cl you are increasing the positive charge by increasing the number of protons. So, you have more positive charges as you move right but the electrons are still in the same energy level. The more positive charges you have the closer the negatively charged electrons will be pulled. • Inner core e- do not change and act as a shield from the positive pull of the nucleus.

  23. Atomic Radii • Zeff increases and pulls e- toward the nucleus as you move right and up. This is why size increases as you go down and decreases as you move across. • Ionic Radii • Think about how Zeff will increase. Why does ionic radii increase after the metalloid line?

  24. Ionization Energy and Zeff • When Zeff goes up its harder to pull e- away • Electronegativity and Zeff • When the Zeffgoes up there is a greater ability to attract an electron.

  25. Electron affinity • An atoms attraction for an e- in the gaseous state • Generally electron affinity follows the same trend as electronegativity • Expressed in negative values • The more negative the higher the affinity