Chapter 5

1. Chapter 5 Exploring the Periodic Table & Elemental Properties

2. Chapter 5 Objectives • At the end of this section you will be able to… • Explain the shape of the periodic table. • Explain why elements in columns have similar chemical properties. • Determine if an element has metallic, nonmetallic or metalloid properties because of its location on the table. • Label the common families on the table.

3. Chapter 5 Objectives (Continued) • Determine the charge of an elements ion based on its location in the table. • Will understand the filling order of orbitals and energy levels • Write the electron configuration of an element. • Identify periodicity & use the PeriodicLaw to explain these patterns.

4. Intro to the PT • The periodic table is a tool that has no likeness in any other field of science • Despite dramatic changes in technology & the discovery of many new elements, there has been no need to change the basic setup of the table. • The PT allows chemists to function by mastering the properties of only a few elements.

5. Intro II • The discovery of the periodic system used to classify elements is a combined work of a number of scientists, not a sudden realization by a single researcher. • The power of the periodic table lies in the two or three dimensional display of all the known (and unknown) elements in a logical system of precisely ordered rows and columns.

6. Early Attempts to Organize the Elements Triads, Octaves & Tables

7. Dobereiner (1817) • His focus was to organize the elements known at the time • Noticed a trend…the mass of the middle element was approximately the average of the masses of the first and third elements • Just as the mass was equal to the average, so were the chemical/physical properties. • He called these sets of three consecutive elements with similar properties “triads”

8. Newlands (1864) • Was the first to use a sequence of common numbers to organize elements. • He arranged the known elements by mass. • Noticed a trend, elements showed similar properties with those elements seven before and after them.

9. Newlands II • Newlands coined the phrase “Law of Octaves.” • The Law of octaves states that an element will share similar properties with those that are seven places before and after it. • Newlands knew there were elements missing, but he was not able to determine what they were.

10. Mendeleev (1869) • While in the process of writing a chemistry text, he developed an effective way to organize the elements according to their atomic mass and similarities in chemical properties. • By incorporating similar chemical properties he noticed “periodic” similarities between elements…so he placed them in the same columns in his table.

11. Mendeleev II • By incorporating chemical and physical properties into the organization method he was able to accurately predict the masses and properties of the unknown elements. • Although his method of arranging elements by mass did cause some elements to appear out of place (Te & I).

12. Terms from Mendeleev • Periodic: refers to the way that elements show patterns in their chemical properties that tend to repeat at regular intervals. • Periodicity: refers to the repetition of properties in certain, regular intervals.

13. Valence Electrons • Valence electrons are the electrons in the highest energy level (those farthest from the nucleus). • Valence electrons are the only electrons that take part in reactions. • Thus the # valence electrons will determine the properties of the element. • Every Group has a distinct # of valence electrons.

14. Lewis Dot Notation • This is a way to show the number of electrons in the highest energy levels without having to draw the full Bohr Model of the atom

15. To Make a Lewis Dot… • Write the element symbol from the Periodic Table. • Determine the # of Valence Electrons by looking at the group #. • Show the electrons with “X” around the symbol in the 12, 3, 6 & 9 positions. P

16. Valence Electrons for Main Group Elements • 2 3 – 12 13 14 15 16 17 18 H He Li Be TM B C N O F Ne Na Mg Al Si P S Cl Ar Notice that every time you start a new period on the table you begin the trend all over again!

17. Valence Electrons by Groups • Group 1 1 e- • Group 2 2 e- • Group 3-12  2 e- (refer to s/p/d/f blocks) • Group 13  3 e- • Group 14 4 e- • Group 15  5 e- • Group 16  6 e- • Group 17  7 e- • Group 18 8 e-

18. The Modern Periodic Table Noble Gases & Atomic Numbers

19. Ramsay (1894) • Ramsay discovered the Noble Gases, a group of gaseous elements that did not combine with any other elements to form compounds, in 1894. • He Named them Noble Gases since they did not mix with other groups of elements just like nobility does not mix with commoners.

20. Moseley (1913) • Moseley discovered that each element had a distinct number of protons in the nucleus…he called this the atom’s Atomic Number • He arranged the elements in order of increasing atomic number.

21. Moseley II • By arranging the elements according to increasing atomic number he was able to correct the “?” in Mendeleev’s Table. • The discovery of the atomic number showed that there was a fundamental quantity that increased by regular steps when we pass from one element to another.

22. Moseley III • The Atomic # and the fact that it increased by one for each element allowed us to determine the number of missing elements. • After his discovery Chemists accepted the atomic # as the principal method for arranging the elements on the periodic table.

23. The Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic number.

24. The Periodic Table Is an arrangement of elements in order of atomic number so that elements with similar properties fall into the same column (also called a group or Family).

25. Capacities of Energy Levels • Here is a quick formula to remember the # of electrons that an energy level can hold… # e- in an EL = 2(EL#)2

26. The Shape of the Periodic Table • The PT is shaped the way it is because there are four major regions that are placed relative to one another. • We call these regions “Blocks” but they actually help us determine the type of orbital the element’s electrons are in… especially their valence electrons.

27. The “s” Block • Is the first two columns in the PT. • All of the elements in these two columns have their first and second outer shell electrons in the “s” orbital. • The “s” orbital is shown to the left. • It is a 3-D Sphere with the nucleus @ the center.

28. The “p” Block • The “p” block is the group of six columns (# 13 – 18) • The elements in this block have their third – eighth valence electrons in the “p” orbital that has the shape seen off to the right.

29. The “d” Block • The “d” Block contains all the elements in columns 3 – 12. • The energy level of the “d” block is always one less than the period number. • After the “s” block fills the electrons in the 4th – 7th energy levels drop down to a the previous energy level.

30. The “d” Block II • Think of it this way… • The 4s fills up with the 19th & 20th electrons…the 21st – 30th electrons do not have sufficient to get themselves into the 4p yet & there is still room in the 3rd EL (remember the 3rd EL holds 18 electrons). • The electrons that lack the energy to get to the 4p go back down to the 3rd energy level to generate the 3d orbitals.

31. The “f” Block • The “f” block has such a confusing shape to its orbitals that I am not even going to place it here. • The thing that you need to remember is that the lanthanides and actinides (the two rows of elements at the bottom of the PT) are what makes up the “f” block.

32. The “f” Block II • The # for the “f” block is always equal to the period # - 2 • You will see that they fit in between column 3 & 4 in period 6 & 7.

33. Why only 8 valence e-? • If you look at the 4th EL you will see that there are 18 elements in the period. BUT there are only 8 valence electrons. • WHY? Because the elements in groups 3 – 12 are in the lower, 3rd EL. Therefore there are only 8 electrons on the highest energy level…EL 4.

34. The Diagonal Rule

35. Filling Order of Sublevels • Use the Diagonal Rule to help you determine the order in which electrons are placed in to each sublevel.

36. Ionization They all Want to be Noble

37. The Octet Rule • Elements are most stable when they attain eight electrons in their outer energy level.

38. Stability via Ionization • In order to get the octet all elements except the Noble Gases must lose or gain electrons. • If the element has between 1 & 3 electrons in the valence shell then it will lose electrons. • If the element has between 5 & 7 electrons then the element will gain electrons.

39. Stability via Ionization II • If the element has 4 valence electrons it is caught in the middle so whether it gains or loses electrons depends on the elements it comes into contact with during the period of time when it ionizes.

40. The Major Classes Metals, Nonmetals & Metalloids

41. The 3 Classes of Elements • Metals - Elements that lose electrons when forming compounds. • Nonmetals- Elements that gain electrons when forming compounds. • Metalloids - elements that have properties of both metals & nonmetals.

42. Location

43. The Metalloids • The following elements are the Metalloids… • Boron • Silicon • Germanium • Arsenic • Tellurium • Some texts include Polonium & Astatine but, they really are not true metalloids

44. Metal Properties • Metals are… • Lustrous - they are shiny • Ductile - they can be drawn into thin strands • Malleable - they are able to be shaped • Conductive - transfer heat and electricity very efficiently • Lose electrons when forming compounds

45. Nonmetal Properties • Nonmetals are… • Gases or brittle solids • Have dull surfaces • Are insulators - they do not transfer heat and electricity well.

46. Metalloid Properties • Metalloids have a mixture of M & NM properties… • They may be good conductors but are also brittle. • May be lustrous but not malleable