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Organizing the Elements and Periodic Trends

Learn how the modern periodic table is organized, understand electron configuration patterns, and explore trends in the periodic table.

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Organizing the Elements and Periodic Trends

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  1. Lesson 6 The Periodic Table Anything in black letters = write it in your notes (‘knowts’)

  2. Objectives for Lesson 6 Section 1 – Organizing the Elements Section 2 – Periodic Trends • Describe ways in which the modern periodic table is organized • Understand electron configuration patterns in the periodic table • Describe and explain trends in the periodic table

  3. Section 1 – Organizing the Elements Dmitri Mendeleev (1869) – created 1st modern periodic table.

  4. Mendeleev arranged elements with similar properties. He also left gaps where proposed elements should be. These gaps were later filled in as more elements were discovered. Ga & Ge Discovered later Similar properties

  5. Mendeleev’s table was an accepted success because it predicted the properties of elements that had not yet been discovered. Woo Hoo!

  6. Today’s periodic table is arranged in order of increasing atomic number (not mass). Also, elements with similar chemical properties are placed in the same vertical column.

  7. Columns are called groups or families. Horizontal rows are called periods.

  8. Valence Electrons – Electrons in the highest occupied energy level; maximum of 8. Elements in the same column have similar properties because they have the same number of valence electrons.

  9. Electrons in the s and p orbitals of the outer shell are the valence electrons. 8 is the maximum number of valence electrons

  10. The Octet Rule – Atoms tend to gain or lose electrons to have 8 e- Sodium: 1s22s22p63s1 Magnesium: 1s22s22p63s2 Fluorine: 1s22s22p5 Nitrogen: 1s22s22p3

  11. The noble gases are chemically stable because they have a full outer energy level (valence). Atoms tend to gain or lose electrons to have 8 e- Sodium: 1s22s22p63s1 Magnesium: 1s22s22p63s2 Fluorine: 1s22s22p5 Nitrogen: 1s22s22p3

  12. Electron configurations for Group 1 (valence e- underlined) 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 [Xe]6s1 [Rn]7s1

  13. 1s22s22p5 1s22s22p63s23p5 Get the idea?... Why is it called the Periodic Table of the Elements? The properties of the elements repeat going across each row.

  14. Three broad classes of elements;metals, metalloids, nonmetals

  15. Metals – good conductors of heat and electricity, shiny, most are solid at room temp (except Hg), malleable, ductile Nonmetals – not metals!, most are gases at room temp Metalloids – can show properties of both metals and nonmetals

  16. Practice • Explain why Mendeleev’s table was an accepted success. • Why is the table of elements called the “periodic” table of elements? • State 4 properties of metals. • Explain the reason that elements in the same column have similar chemical properties? • How can you tell if an elements is a metal, nonmetal or metalloid from the periodic table?

  17. 6. Name an element that is part of the • Halogen family • Alkali metal family • Alkaline earth metal family • Transition metals • Inner transition metals • Noble gas family • 7. A horizontal row in the periodic table is called a _____.

  18. 8. Write the electron configuration for • Nitrogen • Chlorine • Rubidium • 9. How many valence electrons are in each element from question 8?

  19. Section 2 – Periodic Trends Atomic size Ionic size (skip!) Ionization Energy Electronegativity

  20. Atomic radius (pm) Atomic number Atomic Size

  21. Atomic size generally decreases from left to right across a period. As Z increases across a row, the +/- electrical attraction increases, making the atom smaller. As Z increases down a group, another energy level is added to the atom which ‘shield’ the outer electrons from this nuclear attraction.

  22. Ion – atom or group of atoms that has a positive or negative charge. Ions are formed when electrons are transferred between atoms.

  23. Cation – ion with a positive charge. Anion – ion with a negative charge.

  24. Metals tend to form cations Nonmetals tend to form anions

  25. Ionization Energy – energy required to remove an electron from an atom. lithium +1 ion lithium atom 1st ionization energy = 520 kJ/mol lithium ion Lithium +2 ion 2nd ionization energy = 7297 kJ/mol

  26. Ionization Energy – energy required to remove an electron from an atom. lithium +1 ion lithium atom 1st ionization energy = 520 kJ/mol lithium ion Lithium +2 ion 2nd ionization energy = 7297 kJ/mol

  27. First ionization energy (kJ/mol) Atomic number

  28. What does 1st ionization energy mean? • Explain why the 2nd ionization energy of Li and Na is so much higher than the 1st ionization energy. • Explain why the 1st ionization energy of Na is smaller than Li. • Why are the ionization energies of the noble gases so large?

  29. Electronegativity – tendency of an atom to attract electrons of another atom. Metals have low e-neg values, Nonmetals have high e-neg values

  30. B<H<C Noble gases do not have e-neg values

  31. Lesson 6 Practice • How does the size of an atom change from left to right a) across a period? b) down a column? • Give the explanation for question 1. • What is an ion and how are they formed? • Metals tend to form _____ ions and nonmetals tend to form _____ ions. • Define Ionization Energy. • Describe the trend in ionization energy in the periodic table.

  32. Define electronegativity. • Which atom is the a) most electronegative, b) least electronegative? • Which atom has the highest ionization energy? • How do electronegativity values differ between metals and nonmetals?

  33. A little more for Chapter 6…

  34. Lesson 6 Quiz Review Terms to know: valence electron, cation, anion, electronegativity, ionization energy (1st & 2nd)

  35. Things to know: Metal, nonmetals, metalloids locations 4 properties of metals metals form cations, nonmetals form anions family names (alkali, alkaline earth, noble, halogens, transition and inner transition) electronegativity and ionization energy trends electron configurations (w/out aufbau diagram)

  36. Possible Short Answer Questions: 1. Why was Mendeleev’s table an accepted success? 2. Why is the periodic table called the “periodic” table? 3. What causes elements in the same column to have similar chemical properties? 4. What is an ion and how are ions formed? 5. Why is the 2nd ionization energy of Na so much larger than the 1st ionization energy?

  37. Energy levels can also be called electron shells

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