The Periodic Table Ch 6

History of the Periodic Table Only 13 elements had been discovered by 1700 As time went on and more elements were discovered (5 in the decade of 1765-1775) a method to organize the elements was needed. Around the world chemists used different mass values for the elements which made it hard to reproduce experiments.

Example: John Newland’s arrangement of the (then known) elements H 1 Li 2 G 3 Bo 4 C 5 N 6 O 7 F 8 Na 9 Mg 10 Al 11 Si 12 P 13 S 14 Cl 15 K 16 Ca 17 Cr 19 Ti 18 Mn 20 Fe 21 Co & Ni 22 Cu 23 Zn 25 Y 24 In 26 As 27 Se 28 Br 29 Rb 30 Sr 31 Ce & La 33 Zr 32 Di & Mo 34 Ro & Ru 35 Pd 36 Ag 37 Cd 38 U 40 Sn 39 Sb 41 Te 43 I 42 Cs 44 Ba & V 45 Ta 46 W 47 Nb 48 Au 49 Pt & Ir 50 Os 51 Hg 52 Tl 53 Pb 54 Bi 55 Th 56

Mendeleev In 1869 Russian scientist Dimitri Mendeleev, demonstrated a connection between atomic mass and elemental properties. Using element cards, it took him 18 months to set it up. Mendeleev developed his table while making a textbook for his students.

(continued) He wrote each of the about 60 elements each on a note card and arranged them in order of increasing atomic mass. Unlike other scientists, Mendeleev left blanks in his table for elements that weren’t discovered yet.

Mendeleev’s periodic table

Discovery After the publication of his Periodic Table missing elements (Ga, Ge, Sc) were found with masses and properties very close to what Mendeleev predicted. Mendeleev noticed that some elements seemed out of place in his table. He thought the masses (from other scientists) were wrong.

Henry Moseley In 1913, Moseley determined the atomic number for each element. Elements are now arranged by atomic number, not mass, in the modern periodic table.

Modern Periodic Table

Periodic Table seven horizontal rows called periods 18 vertical columns called groups or families groups 1 and 2 (1A and 2A) and groups 13-18 (3A – 8A) are called representative elements groups 3-12 are the transition metals

Periodic Table elements in any group have similar physical and chemical properties properties of elements in periods change from group to group symbol placed in a square atomic number above the symbol atomic mass below the symbol

Periodic Table Three main classes of elements: 1. metals 2. nonmetals 3. metalloids

Metals Most elements are metals. Properties: Good conductors of heat and electricity Solid at room temperature (except Mercury) Reflect light (shiny) Lose electrons in reactions

Metallic character Most metallic element always to the left of a period, least metallic to the right, and 1 or 2 metalloids are in the middle Most metallic element always at the bottom of a column, least metallic on the top, and 1 or 2 metalloids are in the middle of columns 4A, 5A, and 6A

Nonmetals Located in the upper right corner of PT Greater variation among these than metals. Most are gases at room temperature. A few are solids (C, S, P) and one is a liquid (Br) Tend to have properties opposite of metals. Gain electrons in reactions.

Metalloids Generally have properties similar to metals and nonmetals. An element in this group may behave like a metal under certain conditions, and then behave like a nonmetal under different conditions. For example, the metalloid Silicon is a poor conductor of electricity, but it becomes a good conductor when it is mixed with another metalloid, Boron.

Main groups Group IA  alkali metals Group IIA  alkaline earth metals Group VIIIA  noble gases Group VIIA  halogens – “salt formers” Group VIA  chalcogens Group VA  Nitrogen group Group IVA  IVA group Group IIIA  IIIA group

Other groups s & p block filled (A-groups)  representative elements d block filled (B-groups)  transition metals f block filled  inner transition metals 4f  lanthanides 5f  actinides f elements that are naturally occurring  rare earth elements

What are periodic trends? also called “atomic trends” – take place at the atomic level trends are general patterns or tendencies they are general not definite – there are exceptions when looking at trends we look for increases & decreases across  periodic down  group

Effect on trends Nuclear Charge the “pull” of the nucleus proportional to the number of protons in an atom the greater the number of protons, the stronger the nuclear charge (“pull”) this generally affects periodic trends

Effect on trends Shielding - the electron protection from the nuclear “pull” - shield = an energy level of electrons - we are not concerned with single electrons, only energy levels of electrons - these electrons reduce the nuclear pull - affects group trends

Effect on trends Stability - where electron arrangement is compared to stable octet (or other special stabilities) - determines if atom gains or loses electrons - can be used to explain anomalies in trends

Trends in Atomic Size Increases down column Decreases across period valence shell farther from nucleus because of increased shielding Decreases across period left to right because of the nuclear “pull” adding electrons to same valence shell valence shell held closer because more protons in nucleus

Ionization Energy Minimum energy needed to remove a valence electron from an atom 1 mole of electrons in the gaseous state (kJ/mol) The lower the ionization energy, the easier it is to remove the electron metals have low ionization energies

Trends in Ionization Energy Ionization Energy decreases down the group valence electron farther from nucleus Ionization Energy increases across the period left to right harder to remove an electron from the atom because of the increased nuclear “pull” Exceptions: Group 3, Group 6 (chalcogens)

Ionization Energy Li + energy  Li+ + e- Li+ + energy  Li+2 + e- 1st ionization = 520 kJ/mol Li+ + energy  Li+2 + e- 2nd ionization = 7297 kJ/mol Li+2 + energy  Li+3 + e- 3rd ionization = 11,810 kJ/mol Notice, each successive ionization energy is greater than the preceding one – there is a greater “pull” between the nucleus and the electron and thus more energy is needed to break the attraction.

Examining ionization energies can help you predict what ions the element will form. easy to remove an electron from Group IA, but difficult to remove a second electron. So group IA metals form ions with a 1+ charge.

Electron Affinity atoms attraction to an electron it is the energy change that accompanies the addition of an electron to a gaseous atom “opposite” of ionization energy (Concept NOT actual trend)

Exceptions: Nitrogen Group & Noble Gases Across a Period electron affinity increases because of increased “pull” Down a Group electron affinity decreases because the electrons are shielded from the pull of the nucleus Exceptions: Nitrogen Group & Noble Gases

Electronegativity the ability of an atom to attract electrons when the atom is in a compound very similar to electron affinity Across a Period increases because of increased pull Down a Group decreases because of shielding F– most electronegative element

Decrease Increase

Ionic size cations – lose electrons (positively charged) anions – gain electrons (negatively charged) elements gain or lose e- to become stable – being like noble gases (filled outer sublevel) IA - +1 VA - -3 IIA - +2 VIA - -2 IIIA - +3 VIIA - -1 IVA – share VIIIA – 0, stable

Ionic size GOOD RULE OF THUMB anions are always larger than their neutral atom cations are always smaller than neutral atom Across a Period: cations decrease (I-III) because of greater pull on electrons anions decrease (V-VII) because of less pull on electrons and repulsion of the electrons Down a Group: both cations and anions increase size

Reactivity Reactivity of metals increases to the left on the Period and down in the column follows ease of losing an electron Reactivity of nonmetals (excluding the noble gases) increases to the right on the Period and up in the column