Periodic Table 2016 (PDF)

The Periodic Table Honors Chemistry Mrs. Calder  History of the table: J.W. Dobereiner: 1829 Published a system with elements grouped into triads. (3 elements with similar properties) He found that the middle element always had a value between the other two. ...

The Periodic Table Honors Chemistry Mrs. Calder  History of the table: J.W. Dobereiner: 1829 Published a system with elements grouped into triads. (3 elements with similar properties) He found that the middle element always had a value between the other two. Dmitri Mendeleev: 1869 Organized elements by repeating properties based on atomic mass. He left spaces for undiscovered elements & predicted their properties. Henry Moseley: 1913 discovered atomic number and determined the atomic number for the known elements of his time.  The Periodic Law When elements are arranged in order of increasing atomic number, their physical and chemical properties will show a periodic pattern.  The Periodic Table TODAY Arranged by atomic number Periods across (rows) Groups/Families form columns Squares in the Periodic Table symbols and names of the elements information about the structure of their atoms: (differs by table) Atomic number, atomic mass Color codes black = solid; red = gas; blue = liquid Electron configurations sometimes there  Arrangement - columns Groups or Families form vertical columns. Some tables label 1-18 Many have family names: – 1A (1) = Alkali Metals (form bases in water) – 2A (2)= Alkaline Earth Metals (also form bases, don’t dissolve well) – 7A (17)= Halogens (means salt, react to form salts) – 8A (18)= Noble Gases – Others named by top element (oxygen group, carbon group, nitrogen group) – All the “A” columns are called REPRESENTATIVE ELEMENTS  Arrangement - Rows Periods form rows (7 rows) Each period corresponds to an energy level Placed in the order in which shells are filled Row 1 = fills up to 1s2 Row 2 = up to 2p6 Row 3 = up to 3p6 Row 4 = up to 4p6 Row 5 = up to 5p6 Row 6 = up to 6p6 Row 7 = the rest of the elements  Classes of elements: Metals (on left, 80% of table) Good conductors of heat & electricity Solid (except Hg) Luster (reflect light) Ductile (can be drawn into wires) Malleable (can be hammered into sheets) Sample Metals Non-Metals Upper right of the table Properties vary Most are gases, few solids, Bromine=liquid Mostly poor conductors Brittle Metalloids Between metals & nonmetals Show some properties of both Include: B, Si, Ge, As, Sb, Te, At LOOK FOR ZIG ZAG LINE  Electron Configurations Elements in the same group have the same valence number That’s why they behave similarly Table is named by blocks: H 1 1s1 All alkali metals end in s1 = 1 valence electron Li 1s22s1 3 Na 1s22s22p63s1 11 K 1s22s22p63s23p64s1 19 Rb 1s22s22p63s23p64s23d104p65s1 37 Cs 1s22s22p63s23p64s23d104p65s24d10 55 5p66s1 Fr 87 1s22s22p63s23p64s23d104p65s24d105p66s 4f 5d106p67s1 2 14  Elements in the s - blocks s1 He 2 s Alkali metals all end in s1 Alkaline earth metals all end in s2 – really should include He, but it fits better in a different spot, since He has the properties of the noble gases, and has a full outer level of electrons.  Main Groups Based on electron configurations, there are 4 1) Noble gases 2) Representative elements Let’s now take a 3) Transition metals closer look at these. 4) Inner transition metals  Electron Configurations in Groups 1) Noble gases Group 8A (also called Group18) – very stable = typically don’t react – outer s and p sublevels completely full  He All have outer shell full 1s 2 2 Ne 1s22s22p 6 10 Ar 1s22s22p63s23p6 18 1s22s22p63s23p64s23d104p6 Kr 36 1s22s22p63s23p64s23d104p65s24d105p6 Xe 54 1s22s22p63s23p64s23d104p65s24d10 Rn 5p66s24f145d106p6 86 2) Representative Elements are in Groups 1A through 7A – Display wide range of properties, thus a good “representative” – End with s or p, but NOT filled  Elements in the 1A-7A groups 1A 8A are called the representative 2A elements 3A 4A 5A 6A 7A outer s or p filling Electron Configurations in Groups 3) Transition metals are in the “B” columns of the periodic table – “d” sublevel is filling – A “transition” between the metal area and the nonmetal area  Transition Metals - d block Note the change in configuration. s1 s1 d1 d2 d3 d 5 d5 d6 d7 d8 d10 d10  The “d” orbitals fill up in levels 1 less than the period number, so the first d is 3d even though it’s in row 4. 1 2 3 4 3d 5 6 7 4d 5d 4) Inner Transition Metals below the main body of the table, in two horizontal rows (only because would make table too wide) – filling the “f” sublevel – Used to be “rare earth elements” The group B are called the transition elements  These are called the inner transition elements, and they belong here 1 2 3 4 5 6 7 4f 5f f orbitals start filling at 4f, and are 2 less than the period number Periodic Trends ALL Periodic Table Trends Influenced by three factors: 1. Energy Level Higher energy levels are further away from the nucleus. 2. Charge on nucleus (# protons) More charge pulls electrons in closer. (+ and – attract each other) 3. Shielding effect (blocking)  Atomic Radius/Size It is the distance from the center of the nucleus to outermost electron. (or ½ the distance between 2 nuclei when atoms of same element joined) INCREASES as you move down a group because you are adding another shell each time (1s then 2s then 3s and so on.) DECREASES as you move across because adding more protons (more nuclear charge) within a shell pulls harder on the electrons. WHY?  Atomic Size } Radius Measure the Atomic Radius - this is half the distance between the two nuclei of a diatomic molecule.  #1. Atomic Size - Group trends As we increase the H atomic number (or go Li down a group)... each atom has Na another energy level, K so the atoms get bigger. Rb  #1. Atomic Size - Period Trends Going from left to right across a period, the size gets smaller. Electrons are in the same energy level. But, there is more nuclear charge. Outermost electrons are pulled closer. Na Mg Al Si P S Cl Ar  Rb K Period 2 Atomic Radius (pm) Na Li Kr Ar Ne H 3 10 Atomic Number  Ions Some compounds are composed of particles called “ions” – An ion is an atom (or group of atoms) that has a positive or negative charge – Positive and negative ions are formed when electrons are transferred (lost or gained) between atoms  METALS Metals tend to LOSE electrons, from their outer energy level – Sodium loses one: there are now more protons (11) than electrons (10), and thus a positively charged particle is formed = “cation” – The charge is written as a number followed by a plus sign: Na1+ – Now named a “sodium ion”  NON-METALS Nonmetals tend to GAIN one or more electrons – Chlorine will gain one electron – Protons (17) no longer equals the electrons (18), so a charge of -1 – Cl1- is re-named a “chloride ion” – Negative ions are called “anions” Left side metals are attracted to right side non-metals!  Ionization Energy It is the energy required to remove an electron from a gaseous atom It is measured in Joules/atom or K J/mol Removing one electron makes a 1+ ion. The energy required to remove only the first electron is called the first ionization energy. It INCREASES as you move across a period It DECREASES as you move down a group  Multiple Ionization Energies The second ionization energy is the energy required to remove the second electron. – Always greater than first IE. The third IE is the energy required to remove a third electron. – Greater than 1st or 2nd IE. Symbol First Second Third H 1312 Why did these values increase so much? He 2731 5247 Li 520 7297 11810 Be 900 1757 14840 B 800 2430 3569 C 1086 2352 4619 N 1402 2857 4577 O 1314 3391 5301 F 1681 3375 6045 Ne 2080 3963 6276  What factors determine IE The greater the nuclear charge, the greater IE. Greater distance from nucleus decreases IE Shielding effect  Shielding The electron on the outermost energy level has to look through all the other energy levels to see the nucleus. Electrons in same period have same shielding  IE - Group trends As you go down a group, the first IE decreases because... – The electron is further away from the attraction of the nucleus, and – There is more shielding.  IE - Period trends All the atoms in the same period have the same energy level. Same shielding. But, increasing nuclear charge So IE generally increases from left to right. Exceptions at full and 1/2 full orbitals.  Driving Forces Full Energy Levels require lots of energy to remove their electrons. –Noble Gases have full orbitals. Atoms behave in ways to try and achieve a noble gas configuration.  Configuration of Ions Non-metals form ions by gaining electrons to achieve noble gas configuration. They end up with the configuration of the noble gas after them.  Ion Group trends Each step down a group is adding an Li1+ energy level Na1+ Ions therefore get K1+ bigger as you go Rb1+ down (same as atoms) Cs1+  Ion Period Trends Across the period from left to right, the nuclear charge increases - so cations get smaller. Anions also get smaller but they are always bigger than cations in the same level!!!! Notice the energy level changes between anions and cations. N3- B 3+ O2- F1- Li1+ Be2+ C4+  Size of Isoelectronic ions Iso- means “the same” Isoelectronic ions have the same # of electrons Al3+ Mg2+ Na1+ Ne F1- O2- and N3- – all have 10 electrons all have the same configuration: 1s22s22p6 (which is the noble gas: neon)  Size of Isoelectronic ions? Positive ions that have more protons would be smaller (EASY! – more protons pulling on same # electrons) N3- F1- O2- Ne Al3+ Na1+ 13 12 11 10 9 8 7 Mg2+  Electronegativity Discovered by Linus Pauling ability to attract electrons from other atoms in a chemical bond – Higher electronegativity = the atom pulls electrons toward itself more strongly no units for it. It increases left to right (highest on the table is F = 4.0) It decreases as you go down the table (lowest are Cs and Fr)  WHY the Group Trend? further down a group, the farther away from the nucleus, plus more electrons total Thus, more willing to share. Low electronegativity.  WHY the Period Trend? Low # valence e- prefer to lose to become like noble gas Thus, low electronegativity Higher valence e- prefer to gain to become like noble gas Try to take them away from others High electronegativity.