Chapter 6 The Periodic Table

1. UNIT 3 – Periodic Table & Bonding Chapter 6 – The Periodic Table Chapter 7 – Ionic Bonding Chapter 8 – Covalent Bonding Chapter 6 The Periodic Table Anything in black letters = write it in your notes (‘knowts’)

2. Objectives for Chapter 6 • 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. 6.1 – Organizing the Elements Dmitri Mendeleev (1869) – created 1st modern periodic table; arranged elements based on atomic mass and chemical properties.

4. Mendeleev arranged elements with similar properties in the same row. 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. Valence Electrons – Electrons in the highest occupied energy level; maximum of 8. Chapter 7 Elements in the same column have similar properties because they have the same number of valence electrons.

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

9. 1s22s22p5 1s22s22p53s23p5 Get the idea?... Why is it called the Periodic Table of the Elements? The properties of the elements repeat going across each row.

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

11. 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

12. 6.2 – Classifying the Elements Columns are called groups or families. Horizontal rows are called periods.

15. Chapter 6 Practice • Explain why Mendeleev’s table was an accepted success. • Why is the table of elements called the “periodic” table of elements? • How is the modern periodic table arranged? • State 4 properties of metals. • What is the explanation for the reason 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?

16. 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 • A horizontal row in the periodic table is called a _____. • Write the electron configuration for • a) Nitrogen

17. 9. Write the electron configuration for • Nitrogen • Chlorine • Rubidium 10. How many valence electrons are in each element from question 1?

18. Chapter 6 ASSIGNMENT (p. 166-173) #1-17

19. 6.3 – Periodic Trends Atomic size Ionic size 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, more energy levels are in 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 + ions (cations) Nonmetals tend to form - ions (anions)

25. Ionic Size Cations are smaller than the atoms they formed from Anions are larger than the atoms they formed from.

26. Ionic Size

27. Ionization Energy – energy required to remove an electron from an atom.

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

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

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

33. Chapter 6 ASSIGNMENT (p. 182) #18-25

34. Chapter 6 ASSIGNMENT (p. 188) #26, 27, 30, 34, 39-42, 45-47, 49-51, 53, 62, 64, 75

35. A little more for Chapter 6…

36. Energy levels can also be called electron shells

37. Each shell corresponds to a period on the table. 2 8 8 18 18 32 32

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

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

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

41. 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)

42. 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?