Unit 3 – The Periodic Table

Mendeleev (1869) arranged elements in order of atomic mass He then noticed certain properties ....... Ex.) unreactive gas Ex.) reactive metal that combines 1:1 with chlorine .....re-appeared periodically Each element was placed below the preceding one with similar properties noble gases

Elements are now arranged by atomic number Vertical columns are called groups --contain elements with similar chemical properties because they have similar electron arrangements same group = same number of valence electrons Horizontal rows are called periods --properties of the elements change predictably same period = same number of principal energy levels (period no. = no. of PELs or shells)

Periodic Properties -- change in a predictable way as you move across or down the table 1. Atomic Radius: half the distance between 2 nuclei in the solid state -- measured in picometers (pm) = 10-12 meter Related to: a. number of e- energy levels: radii: Li less small Fr more large b. strength of attraction between nucleus and outer e- number of protons (nuclear charge): Radii: Li Less large Ne More small

Ionic Radius – radius after electron gain or loss metals lose electrons + charge attracts electrons to nucleus, radius gets smaller non-metals gain electrons – charge repels electrons, radius gets larger

2. Ionization Energy Periodic Properties, continued --the amount of energy required to remove the most loosely held electron from an atom -- measured in kilojoules per mole (kJ/mol) [see ref table S] -- depends on radius a. large radius = weak attraction for outer e- = low ionization energy b. small radius = strong attraction for outer e- = high ionization energy 3. Electronegativity -- the tendency to gain electrons in a bond -- arbitrary scale (no units) [see ref table S] -- depends on radius large radius = low electronegativity small radius = high electronegativity

4. Metallic Character Periodic Properties, continued -- ease of electron loss -- depends on radius large radius = high metallic character small radius = low metallic character

Trends in Periodic Properties Electronegativity Electronegativity

Types of Elements 1. Metals -- luster -- malleable -- good conductors Metals have a weak hold on their electrons -- low ionization energy and electronegativity -- react with non-metals by losing electrons, forming + ions -- metallic solids (except Hg) 2. Non-Metals no luster, brittle, poor conductors Non-metals have a strong hold on their electrons -- high ionization energy and electronegativity -- react with metals by gaining electrons, forming - ions -- react with other non-metals by sharing electrons to form covalent bonds

Non-metals can be: a. molecular solids -- C P S Se I2 At b. liquids -- Br2 c. gases -- N2 O2 F2 Cl2 H2 3. Metalloids -- properties midway between the above -- used as semiconductors -- all solids: B Si Ge As Sb Te 4. Noble Gases -- do not easily gain, lose, or share electrons -- all gases: He Ne Ar Kr Xe Rn

The Elements by Group 1 and 2 -- very reactive metals (except H) -- not found pure in nature -- 1 more reactive than 2 -- within each group, reactivity increases as you move down the group -- atoms lose e- easily: easiest for larger elements ( group 1, lower elements) -- usually form ionic compounds 3 -- 12 --Transition elements -- less reactive metals -- many found pure -- have incomplete inner e- shell: : similar properties : ions can have several different charges

13 – less reactive metals (except B, metalloid) Al – self protective metal 14 – increasing metallic character as you move down the group C Si Ge Sn Pb non-metal metalloids metals C – basis of life; organic chemistry -- 3 allotropes (different forms of the same element)

Allotropes have different properties due to their different structures: Si + Ge – used in computer circuits Pb – easily worked metal subtly toxic; formerly used in pipes, paint, gasoline, solder 15 – Also increasing metallic character as you move down the group N2 – very unreactive; found pure N N strong triple bond must react (“fix”) nitrogen before it can be used found in protein; used in fertilizers and explosives

O2 -- reactive gas; found pure due to photosynthesis 16 mostly nonmetals O2 -- reactive gas; found pure due to photosynthesis -- allotropes O2 and O3 (ozone) ozone in the atmosphere absorbs harmful UV energy light O2  O• + O• O2 + O•  O3 chlorofluorocarbons like Freon (CF2Cl2) reach the ozone layer and destroy it light CF2Cl2  CF2Cl + Cl• Cl• + O3  ClO• + O2 map showing ozone depletion over Earth’s southern hemishpere

S – less reactive than O2 -- an impurity in fossil fuels and metal ores S + O2  SO2 SO2 is an irritating pollutant that combines with moisture in the air to produce acid rain: SO2 + H2O  H2SO3

17 – The Halogens -- all are diatomic, reactive, not found pure -- reactivity decreases as you the group (easier for small atoms to gain e-) -- Fluorine is the most electronegative -- melting and boiling points increase as you the group Fluorine gas Chlorine gas Bromine liquid Iodine solid

18 – The Noble Gases --very stable; full outer electron levels He Ne Ar -- no compounds Kr Xe Rn -- can react with F2 or O2 -- used when reactions are undesirable Ar – to fill light bulbs He – non-flammable; for balloons and blimps helium hydrogen