Chapter 6 - The Periodic Table and Periodic Law

1. Objectives: • Identify different key features of the periodic table. • Explain why elements in a group have similar properties. • Relate the group and period trends seen in the periodic table to the electron configuration of atoms • Why this is important: • The periodic table is one of the most useful reference tools available in chemistry. Understanding its organization and interpreting its data will aid in understanding chemistry concepts. Chapter 6 - The Periodic Table and Periodic Law www.privatehand.com/flash/elements.html

2. Development of the Periodic Table • In 2003, there were 118 elements known. • The majority of the elements were discovered between 1735 and 1843. • How do we organize all the different elements in a meaningful way that will allow us to make predictions about undiscovered elements?

3. A variety of scientists tried to arrange the known elements to reflect the trends in chemical and physical properties but their systems did not allow for newly discovered elements to fit in their charts. • 1869 Dmitri Mendeleev and Lothar Meyer separately arranged the elements in order of increasing atomic mass and into columns with similar properties • That seemed to work to organize most of the elements. http://www.chemistryexplained.com/images/chfa_03_img0535.jpg http://www.chemistrydaily.com/chemistry/upload/a/a1/Dmendeleev.jpg

4. Mendeleev is given more credit than Meyer because he published his findings first and he left spaces for elements that were not yet discovered. • Some of the elements that he predicted were scandium, gallium, and germanium. • In 1871, Mendeleev noted that arsenic (As) properly belonged underneath phosphorus (P) and not silicon (Si), which left a missing element underneath Si. He predicted a number of properties for this element. In 1886 Germanium (Ge) was discovered. The properties of Ge matched Mendeleev’s predictions.

5. Mendeleev’s table was not completely correct. • Arranging elements by atomic mass caused some elements to be put in the wrong groups so that the properties did not exactly match up • 1913 English chemist Henry Moseley arranged elements in order of increasing atomic number • Problems with order of elements were solved and there was a clear repeating pattern of properties of the elements in their groups. • The PERIOIDC LAW states there is a “periodic” repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number. • Periodic means: happening or reoccurring at regular intervals (definition from Webster’s Dictionary) http://www.rsc.org/education/teachers/learnnet/periodictable/scientists/moseley.jpg

6. Boxes are arranged in order of increasing atomic # • Elements are grouped into columns by similar properties • Scientists keep adding elements that were discovered • The final adjustment was when physicist Glenn Seaborg had the inner-transition elements pulled below the rest of the periodic table and into 2 separate rows (This occurred in the late 1940s) The Modern Table http://www.lbl.gov/Science-Articles/Research-Review/Magazine/1994/seaborgium-mag.html

7. Horizontal rows are called periods • There are 7 periods

8. Vertical columns are called groups or families. • Elements are placed in columns by similar properties. • b/c of the similar numbers of valence e- they contain

9. 1 1A 2 2A 18 8A 13 3A 4A 14 15 5A 16 6A 17 7A 3 3B 4 4B 5B 5 6B 6 7B 7 8B 8 9 8B 8B 10 1B 11 2B 12 The Different Groups of Elements 1 – 18 system used by all chemists A & B system is an older American System A elements are representative elements B elements are transition elements

10. They have a wide range of physical & chemical properties. • They have the whole range of possible valence electrons (1 to 8) • Also called s and p block elements • Here are some important groups: Representative or Main Group Elements

11. Group 1 (1A) contains the alkali metals (remember to NOT include hydrogen) • Group 2 (2A) contains the alkaline earth metals

12. Called this because these metals react w/ water to form alkaline (basic) solutions. • Highly reactive metals that lose their 1 valence electron to form 1+ ions • Soft enough to be cut with a knife. • They are stored in oil to prevent reactions with oxygen and water in the air. Alkali Metals

13. Called this because most of these metals react with oxygen to form compounds called oxides (the alchemists called them “earths” because of this) and the oxides react w/ water to form alkaline (basic) solutions • Not as reactive (but do react easily) & harder than group 1 metals • They lose their 2 valence electrons to form 2+ ions Alkaline Earth Metals

14. Named for the first element in each group. • They do have mixed groupings of elements because each column contains nonmetals, metalloids, and metals. • Many of the elements in these groups form various charges Groups13 - 16

15. Group 17 (7A) contains the halogens • Group 18 (8A) contains the noble gases

16. Called this because halogen means “salt formers” b/c they react with metals to form salts (ionic compounds) • Physically F & Cl are gases at room temp., Br is a liquid but it evaporates easily, and Iodine is a solid that sublimes easily • Astatine is the odd-ball of the group b/c it’s radioactive w/ no known uses • Chemically they are the most reactive nonmetals • They have 7 valence e- so they will share or gain 1 e- and they tend to form 1- ions. Halogens

17. Last naturally occurring elements to be discovered b/c they are colorless & unreactive • Very stable with full valence electrons = 8 • (except He w/ 2) • With lots of energy you can get Xe, Kr and Ar compounds (There are no known He or Ne compounds) • In 1962 the first compound of the noble gases was prepared: XeF2, XeF4, and XeF6. • To date the only other noble gas compounds known are KrF2 and HArF. Noble Gases

18. Transition elements (metals) d-block f-block

19. These are called the inner transition elements and they belong here

21. Make up the majority of elements on the periodic table • Have a wide variety of uses & effect the economy • The variation of their physical properties is b/c of their electron configurations & b/c unpaired d-electrons can move into valence shells. • The more unpaired d-electrons, the greater the hardness & higher the melting & boiling points • Most lose electrons to become positively charged ions • Cu, Ag, Au, Pt, and Pd are the only ones unreactive enough to be found alone in nature Transition Metals

22. Lanthanide series – follow element lanthanium • All silvery metals w/ high melting points • Actinide series – follow element Actinium • All are radioactive & only 3 exist in nature Inner Transition Metals

23. In a class by itself, it is a unique element • Most often occurs as a colorless diatomic gas, H2 • It is placed in Group 1 b/c it has one valence e- and will easily lose its 1 electron when reacting w/ other nonmetals to become a 1+ ion (H+ is a proton) • But it shares many properties w/ the halogens and will sometimes gain e- when bonding w/ a metal to become a 1- ion (the hydride ion, H-) • It is the most abundant element in the universe (90% by mass) HYDROGEN

24. Metals, Nonmetals, and Metalloids

25. What are some common properties of metals? Have luster (shine) when smooth & clean. Good conductors of heat & electricity Most are solid @ room temp. Most are: Ductile = drawn into wires. Malleable = hammered into sheets. Most lose electrons to become cations Most of the elements on the periodic table are classified as metals Metals

26. What are some common properties of nonmetals? At room temp: Some are brittle & dull solids Gases Poor conductors = good insulators Most tend to gain e- to become anions They have a wide variety of melting & boiling points Nonmetals

27. Chemical & Physical Properties of both metals & non-metals • Example: Si has a metallic luster but it is brittle. • They are semiconductors b/c they do conduct electricity but not as well as metals • Silicon (Si) & Germainum (Ge) are 2 most important for computer chips & solar panels Metalloids or Semimetals

28. These are defined as electrons in the atom’s highest energy levels. • Examples: • Sodium • Has 11 electrons but only one valence electron • Cesium (Cs) • Has 55 electrons but only one valence electron • Bromine • Has _____ electrons but only _____ valence electron • Valence electrons help determine the chemical properties of an element and how it will bond to form compounds. Valence Electrons REVIEW

29. Examples – Look at the electron configuration of each element. What do you notice? Na [Ne]3s1 K [Ar]4s1 Cs [Xe]6s1 F 1s22s22p5 Cl [Ne]3s23p5 Br [Ar]4s23d104p5 Ne 1s22s22p6 Ar [Ne] 3s23p6 Kr [Ar]4s23d104p6 Mostelements in the same group have the same ending electron configuration = same number of valence electrons

30. Valence electrons and groups: • Group 1 elements have 1 valence electron • Group 2 elements have 2 valence electrons • Group 13 elements have 3 valence electrons • Group 14 elements have 4 valence electrons • Group 15 elements have 5 valence electrons • Group 16 elements have 6 valence electrons • Group 17 elements have 7 valence electrons • Group 18 elements have 8 valence electrons (except He, it only has 2) • Most of the transition metals have 2 valence electrons, there are a lot of exceptions!!

31. Electron dot structures – used to visually represent valence electrons in a shorthand method. • We use an element’s symbol to show what element we are talking about and the dots represent the atom’s valence electrons. • In writing these structures the dots are placed one at a time on the four sides of the symbol and then paired up until all are used. Depicting Valence Electrons Review

32. Examples of electron dot structures: • Magnesium • Sulfur • Rubidium (Rb) • Bromine • Oxygen

33. Section 6.3 - Periodic Trends Objectives: Compare period & group trends for shielding, atomic radius, ionic radius, ionization energy, & electronegativity

34. What does this mean? • The valence e- are blocked from the full positive charge of the nucleus (effective nuclear charge) by the inner (core) e- • As the average number of core e- increases, the effective nuclear charge decreases • This idea of shielding will play a large role in a lot of the trends Shielding (or screening)

35. Mg

36. Trend within the period (left to right): Generally decreases • Why: • B/c the number of energy levels & core e- stays the same but the nucleus is increasing • This increase the attraction between the nucleus and valence e- • Trend down a group: Generally increases • Why: • B/c the number of energy levels & core e- increases • This makes the valence e- farther from the nucleus and more blocked by the inner e- Shielding

37. ½ the distance between adjacent nuclei of identical atoms either in crystal form (metals) or in molecular form (nonmetals) • Trend within the period (left to right): Generally decreases • Why: • B/c the number of energy levels & core e- stays the same but the nucleus is increasing • This increase the attraction between the nucleus and valence e- • This attraction pulls the e- closer to the nucleus and makes the atom smaller • Ex: Na vs. S Atomic Radius

38. Trend down a group: Increases • Why: • B/c the number of energy levels increases & core e- increases • Each energy level is larger than the next • This makes the valence e- farther from the nucleus and more blocked by the inner e- - EX: Na vs. K

40. Examples – Place each group of elements in order of increasing atomic radius: • S, Al, Cl, Mg, Ar, Na • K, Li, Cs, Na, H • Ca, As, F, Rb, O, K, S, Ga

41. Examples – Place each group of elements in order of increasing atomic radius: • S, Al, Cl, Mg, Ar, Na • Ar < Cl < S < Al < Mg < Na • K, Li, Cs, Na, H • H < Li < Na < K < Cs • Ca, F, As, Rb, O, K, S, Ga • F < O < S < As < Ga < Ca < K < Rb

42. Ionic Radius – The distance between the nucleus and the outermost electron in ions (can’t be determined directly) • Trend between atom & ion and Why: • Cations are smaller than original atom • b/c losing e- the atom has unequal positive charge that attracts the valence e- closer to the nucleus • Anions are larger than original atom and cations • b/c adding negative e- adds to the repulsion between other valence e-, pushing them apart

43. Trend within the period (left to right): Representative Elements • Cations the size decreases • Anions the size drastically increases compared to the positive ions and then decreases across the period • Trend down a group: Increases for both cations & anions • Why: • Same reason as atomic radii trend Ionic Radius Continued

45. For ions of the same charge, ion size increases down a group. • All the members of an isoelectronic series have the same number of electrons. • As nuclear charge increases in an isoelectronic series the ions become smaller: • O2- > F- > Na+ > Mg2+ > Al3+

46. Examples – Choose the larger species in each case: • Na or Na+ • Br or Br- • N or N3- • O- or O2- • Mg2+ or Sr2+ • Mg2+ or O2- • Fe2+ or Fe3+

47. Ionization Energy: The energy required to remove an electron from a gaseous atom (also called First Ionization Energy, I1) Na (g) + 496 kJ Na+ (g) + e- • The second ionization energy, I2, is the energy required to remove an electron from a (1+) gaseous ion: Na+ (g) + 4562 kJ Na2+ (g) + e- • NOTICE: • Ionization Energy increases for each electron removed from the same element • The larger ionization energy, the more difficult it is to remove the electron.

48. Variations in Successive Ionization Energies • There is a sharp increase in ionization energy when a core electron is removed. • Notice the large increase after the last valence electron is removed. This chart can be used to determine the number of valence electrons in an atom of an element.

49. Trends within the periods: Increases • Why: Electrons are harder to remove from smaller atoms because they are closer to the nucleus and there is an increased nuclear charge • Trends down a group: Decreases • Why: Electrons are easier to remove from large atoms because they are farther away from the nucleus so there is less energy needed to remove them. • Notice the trends in ionization energy is inversely related to trends in atomic radii.