Unit 4 The Periodic Table

1. Unit 4The Periodic Table Know It Love It Amaze Your Friends

2. How do I write names & symbols? • Spelling COUNTS!! • Symbols MUST be written in BLOCK print!! • Symbol’s first letter is always uppercase and the second letter (if one) is always lowercase!! For Example: Co Cobalt (an element) CO Carbon Monoxide (compound of carbon and oxygen combined) NO Scripty / Cursive  A, N, H, etc... l is a lowercase L not an uppercase I

3. Which elements do I have to know? ELEMENTSYMBOL Fluorine F Neon Ne Sodium Na Magnesium Mg Aluminum Al Silicon Si Phosphorus P Sulfur S ELEMENTSYMBOL Hydrogen H Helium He Lithium Li Beryllium Be Boron B Carbon C Nitrogen N Oxygen O

4. Which elements do I have to know? ELEMENTSYMBOL Silver Ag Tin Sn Iodine I Gold Au Mercury Hg Lead Pb Francium Fr Uranium U ELEMENTSYMBOL Chlorine Cl Argon Ar Potassium K Calcium Ca Iron Fe Copper Cu Zinc Zn Bromine Br

5. Dmitri Mendeleev • Russian Chemist • In 1869, he published the first periodic table. • He organized the elements in a way that would help his students learn them more easily. • He made a card game with the information known about each element listed on separate cards. They could then be arranged by the properties the elements had in common.

6. Mendeleev settled on an organization of elements that was based on the masses of the elements. • Mendeleev found that when the elements were arranged in order of atomic mass, many physical and chemical properties of the elements followed repeating patterns. • Worked for most elements, but not all. • Three pairs of elements had to be switched, but Mendeleev thought these masses were measured incorrectly.

7. Ekasilicon • Mendeleev was able to accurately predict the existence of elements not yet discovered. These showed up as gaps in his periodic table. • One such element gap, Mendeleev called ekasilicon. He predicted its mass, density, melting point and color based on its location in the periodic table. • Fifteen years after this prediction, a new element was discovered in Germany and given the name Germanium. Its properties matched the properties of ekasilicon.

8. H.G.J Moseley worked in Rutherford’s lab. • Found that metals produce X-rays when bombarded with energetic electrons and that the frequencies differed for each metal. These frequencies, came from differences in a fundamental property of each element; the amount of positive charge in the nucleus. • The amount of positive charge = the number of protons = the atomic number.

9. Modern Periodic Law • When the elements are arranged in order of increasingatomic number, their physical and chemical properties show a periodic, repeating pattern.

10. Periodic Table Arrangement • Periods – the horizontal rows of elements. • The modern periodic table has 7 periods. • The period tells the number of energy levels used. This number is called the Principal Quantum Number. • Groups – the vertical columns of elements. • The modern periodic table has 18 groups. • The “A” groups tell the number of valence electrons. • Groups are sometimes referred to as families.

11. Li = 1s22s1 Na = 1s22s22p63s1 K = 1s22s22p63s23p64s1 *Li, Na, K all have 1 valence e- *All are found in group 1 or IA *Li in pd. 2 Na in pd. 3 K in pd. 4 B = 1s22s22p1 Al = 1s22s22p63s23p1 Ga = 1s22s22p63s23p64s23d104p1 *B, Al, Ga all have 3 valence e- *All are found in group 13 or IIIA *B in pd. 2 Al in pd. 3 Ga in pd. 4

12. N = 1s22s22p3 P = 1s22s22p63s23p3 As = 1s22s22p63s23p64s23d104p3 *N, P, As all have 5 valence e- *All are found in group 15 or VA *N in pd. 2 P in pd. 3 As in pd. 4 F = 1s22s22p5 Cl = 1s22s22p63s23p5 Br = 1s22s22p63s23p64s23d104p5 *F, Cl, Br all have 7 valence e- *All are found in group 17 or VIIA *F in pd. 2 Cl in pd. 3 Br in pd. 4

13. Groups of the Periodic Table • Elements within a group on the periodic table have similar properties to each other. • This is due to the number of valence electrons. Having the same number of valence electrons makes them bond to similar atoms in the same ratios. • If you know the properties of one element in a group, you know the properties of all the elements in that group!!!

14. Trends of the Periodic Table • Atomic Number • Atomic Mass • Metal / Nonmetal Trend • Atomic Radius • Ionization Energy • Electronegativity

15. Atomic Number • The number of protons in one atom of a given element. • Increases as you move down a group. • Increases as you move across a period.

16. Average Atomic Mass • The average mass of all the isotopes of a given element. • Increases as you move down a group. • Increases as you move across a period.

17. Atomic Number/Avg.Atomic Mass increases I n c r e a s e s

18. Metal / Nonmetal / Metalloid • Elements on the left side of the table are metals. • There are 88 metals on the periodic table. • Elements on the right side are nonmetals. • There are 17 nonmetals on the periodic table. • Elements on the “staircase” between the metals and nonmetals are metalloids. • There are 7 metalloids on the periodic table.

19. Properties of Metals and Nonmetals

20. Properties of Metalloids • Metalloids have properties of both metals and nonmetals. • For example: Silicon Is silver, shiny and a conductor like a metal. Is rough, brittle and has no rxn to acid like a nonmetal.

21. Activity Series • Some metals are more reactive than others and will replace less reactive metals during a reaction. Li, K, Ca, Na, Mg, Al, Zn, Fe, Pb, H, Cu, Hg, Ag DECREASING REACTIVITY       Will Ca replace Zn in a reaction? Yes, Ca is more reactive than Zn Will Zn replace Mg in a reaction? No, Zn is less reactive than Mg

22. Metal vs. Nonmetal Trend Most Active Nonmetal Most Active Metal Metalloids Metals Nonmetals

23. Atomic Radius • The distance from the center of an atom’s nucleus to its outermost electron. • Atomic radius increases as you move down a group on the periodic table. • Because electron clouds are added. • Atomic radius decreases as you move across a period on the periodic table. • Full shells with paired-up electrons have less repulsion, so the take up less space.

24. Fr Atomic Radius

25. Ionization Energy • The amount of energy required to remove a valence electron from an atom. • Decreases as you move down a group on the periodic table. • More “shielding” from full electron shells. • Increases as you move across a period on the periodic table. • More difficult to remove electrons from a full shell.

26. Electronegativity • The attraction for shared electrons in a chemical bond. • Decreases as you move down a group on the periodic table. • Valence electrons get farther from the nucleus because there are more full e- shells. The attraction between the + nucleus and – electrons decreases. • Increases as you move across a period on the periodic table. • More protons, stronger attraction to electrons.

27. Electronegativity/Ionization Energy • Look at the positions of Fluorine and Francium on the periodic table. • Francium doesn’t need another electron to become stable. It actually would be more stable by giving one away. (Has lowest ionization energy due to 6 full shells of shielding.)

28. Electronegativity/Ionization Energy • Fluorine needs one more electron to fill its valence shell. It has the strongest electronegativity. The closer the valence electrons are to the nucleus, the stronger the electronegativity. (The valence e- are closer to the positive nucleus.)

29. Ionization energy & Electronegativity

30. Alkali Metals • Group 1 or 1A. • These elements have 1 valence electron which they will give up when bonding to become more stable. (full valence shell) Therefore, they will become +1 ions. • Alkali means “ashes.” Sodium and potassium are present in the ashes of burned plants.

31. Alkali Metals • Alkali metals are very reactive. Will react with water and air! (Stored in oil.) • They are soft enough to be cut by a knife. • The members of this group are: (Not Hydrogen.) Li [He]2s1 Na [Ne]3s1 K [Ar]4s1 Rb [Kr]5s1 Cs [Xe]6s1 Fr [Rn]7s1

32. Alkaline Earth Metals • Group 2 or 2A. • These elements have 2 valence electrons which they will give up when bonding to become more stable. (full valence shell) Therefore they will become +2 ions. • Obtained from alkaline earths. Earths were substances unchanged by fire.

33. Alkaline Earth Metals • Alkaline Earth Metals are also very reactive. (Not as reactive as group I or 1A) • They have higher densities and melting points than the alkali metals. • Found in mineral deposits. • The members of this group are: Be [He]2s2 Mg [Ne]3s2 Ca [Ar]4s2 Sr [Kr]5s2 Ba [Xe]6s2 Ra [Rn]7s2

34. Transition Metals • Groups 3 – 12 • Play an important role in living organisms, are extremely valuable as strong structurally useful materials. • Vary greatly in properties and abundance. • Most have high densities and high melting points. • Number of valence electrons varies, therefore ionic charge varies.

35. Inner Transition Metals • Lanthanide and Actinide Series elements (“f” block elements) • Properties similar to Transition Metals • Many are man-made • Many are naturally radioactive (Large, unstable nuclei. Ratio of neutrons to protons is high) • Have two valence electrons. (after 6s2, 7s2)

36. The Boron Group • Group 13 or 3A. • These elements have 3 valence electrons which they will give when bonding to become more stable. (full valence shell) Therefore they will become +3 ions. • The members of this group are: B [He]2s22p1 Al [Ne]3s23p1 Ga [Ar]4s24p1 In [Kr]5s25p1 Tl [Xe]6s26p1

37. The Carbon Group • Group 14 or 4A. • These elements have 4 valence electrons which they can give up, share or accept to become more stable. (full valence shell) Therefore they can become +4 or –4 ions. • The members of this group are: C [He]2s22p2 Si [Ne]3s23p2 Ge [Ar]4s24p2 Sn [Kr]5s25p2 Pb [Xe]6s26p2

38. The Nitrogen Group • Group 15 or 5A. • These elements have 5 valence electrons. They will accept 3 more to become stable. (full valence shell) Therefore they become –3 ions. • The members of this group are: N [He]2s22p3 P [Ne]3s23p3 As [Ar]4s24p3 Sb [Kr]5s25p3 Bi [Xe]6s26p3

39. The Oxygen Group • Group 16 or 6A. • These elements have 6 valence electrons. They will accept 2 more to become more stable. (full valence shell) Therefore they become –2 ions. • The members of this group are: O [He]2s22p4 S [Ne]3s23p4 Se [Ar]4s24p4 Te [Kr]5s25p4 Po [Xe]6s26p4

40. The Halogen Group • Group 17 or 7A. • These elements have 7 valence electrons. They will accept 1 more to become more stable. (full valence shell) Therefore they become –1 ions. • Comes from the Greek word which means “salt former.”

41. The Halogen Group • The Halogens are very reactive nonmetals and exist in elemental form as diatomic molecules. (F2, Cl2, Br2, I2, At2) • The members of this group are: F [He]2s22p5 Cl [Ne]3s23p5 Br [Ar]4s24p5 I [Kr]5s25p5 At [Xe]6s26p5

42. The Noble Gases • Group 18 or 8A. • These elements have 8 valence electrons.(Except helium, which has 2) They are already stable, so they do not accept or receive electrons. They do not typically form ions. • These elements are very stable.

43. The Noble Gases • Members of this group include: Helium [He] 1s2 Neon [Ne] 1s2 2s22p6 Argon [Ar] 1s2 2s22p63s23p6 Krypton [Kr] 1s2 2s22p6 3s23p64s23d104p6 Xenon [Xe] 1s2 2s22p6 3s23p64s23d104p65s24d105p6 Radon [Rn] 1s2 2s22p6 3s23p64s23d104p6 5s24d105p66s24f145d106p6

44. Hydrogen (The one and only!) • Has one valence electron, but is not an Alkali metal. • Hydrogen is a nonmetal that exists as a gas under normal conditions. • Colorless, odorless, and composed of H2 molecules. (Diatomic, like the halogens) • Most of the Earth’s hydrogen is combined with oxygen as water and carbon in organic compounds called hydrocarbons.

45. 1 IA Fr Oxygen Group Halogens 18 VIIIA  Alkali Metals Nitrogen Group 2 IIA 17 VIIA  Alkaline Earth Metals 13 IIIA 14 IVA 15 VA 16 VIA 1 2 3 4 5 6 7 Carbon Group 1 2 3 4 5 6 7 Boron Group Transition Metals 3 8 9 11 12 4 5 6 7 10 Noble Gases  Lanthanides  Actinides 