Section 6-1

2. Section 6.1 Development of the Modern Periodic Table • Trace the development of the periodic table. • Identify key features of the periodic table. atomic number: the number of protons in an atom The periodic table evolved over time as scientists discovered more useful ways to compare and organize the elements. Section 6-1

3. Section 6.1 Development of the Modern Periodic Table (cont.) periodic law group period representative elements transition elements metal alkali metals alkaline earth metals transition metal inner transition metal lanthanide series actinide series nonmetals halogen noble gas metalloid Section 6-1

4. Development of the Periodic Table • In the 1700s, Lavoisier compiled a list of all the known elements of the time. Section 6-1

5. Development of the Periodic Table (cont.) • The 1800s brought large amounts of information and scientists needed a way to organize knowledge about elements. • John Newlands proposed an arrangement where elements were ordered by increasing atomic mass. Section 6-1

6. Development of the Periodic Table (cont.) • Newlands noticed when the elements were arranged by increasing atomic mass, their properties repeated every eighth element. Section 6-1

7. Development of the Periodic Table (cont.) • Meyer and Mendeleev both demonstrated a connection between atomic mass and elemental properties. • Moseley rearranged the table by increasing atomic number, and resulted in a clear periodic pattern. • Periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number is called periodic law. Section 6-1

8. Development of the Periodic Table (cont.) Section 6-1

9. The Modern Periodic Table • The modern periodic table contains boxes which contain the element's name, symbol, atomic number, and atomic mass. Section 6-1

10. The Modern Periodic Table (cont.) • Columns of elements are called groups. • Rows of elements are called periods. • Elements in groups 1,2, and 13-18 possess a wide variety of chemical and physical properties and are called the representative elements. • Elements in groups 3-12 are known as the transition metals. Section 6-1

11. The Modern Periodic Table (cont.) • Elements are classified as metals, non-metals, and metalloids. • Metals are elements that are generally shiny when smooth and clean, solid at room temperature, and good conductors of heat and electricity. • Alkali metalsare all the elements in group 1 except hydrogen, and are very reactive. • Alkaline earth metalsare in group 2, and are also highly reactive. Section 6-1

12. The Modern Periodic Table (cont.) • The transition elements are divided into transition metalsand inner transition metals. • The two sets of inner transition metals are called the lanthanide seriesand actinide seriesand are located at the bottom of the periodic table. Section 6-1

13. The Modern Periodic Table (cont.) • Non-metalsare elements that are generally gases or brittle, dull-looking solids, and poor conductors of heat and electricity. • Group 17 is composed of highly reactive elements called halogens. • Group 18 gases are extremely unreactive and commonly called noble gases. Section 6-1

14. The Modern Periodic Table (cont.) • Metalloidshave physical and chemical properties of both metals and non-metals, such as silicon and germanium. Section 6-1

15. The Modern Periodic Table (cont.) Section 6-1

16. A B C D Section 6.1 Assessment What is a row of elements on the periodic table called? A.octave B.period C.group D.transition Section 6-1

17. A B C D Section 6.1 Assessment What is silicon an example of? A.metal B.non-metal C.inner transition metal D.metalloid Section 6-1

18. End of Section 6-1

19. Section 6.2 Classification of the Elements • Explain why elements in the same group have similar properties. valence electron: electron in an atom's outermost orbitals; determines the chemical properties of an atom • Identify the four blocks of the periodic table based on their electron configuration. Elements are organized into different blocks in the periodic table according to their electron configurations. Section 6-2

20. Organizing the Elements by Electron Configuration • Recall electrons in the highest principal energy level are called valence electrons. • All group 1 elements have one valence electron. Section 6-2

21. Organizing the Elements by Electron Configuration (cont.) • The energy level of an element’s valence electrons indicates the period on the periodic table in which it is found. • The number of valence electrons for elements in groups 13-18 is ten less than their group number. Section 6-2

22. Organizing the Elements by Electron Configuration (cont.) Section 6-2

23. The s-, p-, d-, and f-Block Elements • The shape of the periodic table becomes clear if it is divided into blocks representing the atom’s energy sublevel being filled with valence electrons. Section 6-2

24. The s-, p-, d-, and f-Block Elements (cont.) • s-block elements consist of group 1 and 2, and the element helium. • Group 1 elements have a partially filled s orbital with one electron. • Group 2 elements have a completely filled s orbital with two electrons. Section 6-2

25. The s-, p-, d-, and f-Block Elements (cont.) • After the s-orbital is filled, valence electrons occupy the p-orbital. • Groups 13-18 contain elements with completely or partially filled p orbitals. Section 6-2

26. The s-, p-, d-, and f-Block Elements (cont.) • The d-block contains the transition metals and is the largest block. • There are exceptions, but d-block elements usually have filled outermost s orbital, and filled or partially filled d orbital. • The five d orbitals can hold 10 electrons, so the d-block spans ten groups on the periodic table. Section 6-2

27. The s-, p-, d-, and f-Block Elements (cont.) • The f-block contains the inner transition metals. • f-block elements have filled or partially filled outermost s orbitals and filled or partially filled 4f and 5f orbitals. • The 7 f orbitals hold 14 electrons, and the inner transition metals span 14 groups. Section 6-2

28. A B C D Section 6.2 Assessment Which of the following is NOT one of the elemental blocks of the periodic table? A.s-block B.d-block C.g-block D.f-block Section 6-2

29. A B C D Section 6.2 Assessment Which block spans 14 elemental groups? A.s-block B.p-block C.f-block D.g-block Section 6-2

30. End of Section 6-2

31. Section 6.3 Periodic Trends • Compare period and group trends of several properties. principal energy level: the major energy level of an atom • Relate period and group trends in atomic radii to electron configuration. ion ionization energy octet rule electronegativity Trends among elements in the periodic table include their size and their ability to lose or attract electrons Section 6-3

32. Atomic Radius • Atomic size is a periodic trend influenced by electron configuration. • For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element. Section 6-3

33. Atomic Radius (cont.) • For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms. Section 6-3

34. Atomic Radius (cont.) • There is a general decrease in atomic radius from left to right, caused by increasing positive charge in the nucleus. • Valence electrons are not shielded from the increasing nuclear charge because no additional electrons come between the nucleus and the valence electrons. Section 6-3

35. Atomic Radius (cont.) Section 6-3

36. Atomic Radius (cont.) • Atomic radius generally increases as you move down a group. • The outermost orbital size increases down a group, making the atom larger. Section 6-3

37. Ionic Radius • An ionis an atom or bonded group of atoms with a positive or negative charge. • When atoms lose electrons and form positively charged ions, they always become smaller for two reasons: • The loss of a valence electron can leave an empty outer orbital resulting in a small radius. • Electrostatic repulsion decreases allowing the electrons to be pulled closer to the radius. Section 6-3

38. Ionic Radius (cont.) • When atoms gain electrons, they can become larger, because the addition of an electron increases electrostatic repulsion. Section 6-3

39. Ionic Radius (cont.) • The ionic radii of positive ions generally decrease from left to right. • The ionic radii of negative ions generally decrease from left to right, beginning with group 15 or 16. Section 6-3

40. Ionic Radius (cont.) • Both positive and negative ions increase in size moving down a group. Section 6-3

41. Ionization Energy • Ionization energy is defined as the energy required to remove an electron from a gaseous atom. • The energy required to remove the first electron is called the first ionization energy. Section 6-3

42. Ionization Energy (cont.) Section 6-3

43. Ionization Energy (cont.) • Removing the second electron requires more energy, and is called the second ionization energy. • Each successive ionization requires more energy, but it is not a steady increase. Section 6-3

44. Ionization Energy (cont.) Section 6-3

45. Ionization Energy (cont.) • The ionization at which the large increase in energy occurs is related to the number of valence electrons. • First ionization energy increases from left to right across a period. • First ionization energy decreases down a group because atomic size increases and less energy is required to remove an electron farther from the nucleus. Section 6-3

46. Ionization Energy (cont.) Section 6-3

47. Ionization Energy (cont.) • The octet rulestates that atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons. • The octet rule is useful for predicting what types of ions an element is likely to form. Section 6-3

48. Ionization Energy (cont.) • The electronegativityof an element indicates its relative ability to attract electrons in a chemical bond. • Electronegativity decreases down a group and increases left to right across a period. Section 6-3

49. Ionization Energy (cont.) Section 6-3

50. A B C D Section 6.3 Assessment The lowest ionization energy is the ____. A.first B.second C.third D.fourth Section 6-3