Ch. 6 The Periodic Table

1. Ch. 6 The Periodic Table *Inquiry Activity pg. 154 (R47)* What do you do with dead chemists? Barium

3. History • J.W. Dobereiner – 1829 – used properties of elements to sort them into groups • Dmitri Mendeleev – 1869 – arranged the elements in order of increasing atomic mass • Henry Moseley – 1913 – in the modern periodic table, elements are arranged in order of increasing atomic number • Periods – horizontal rows – properties change from L to R but repeat at each period (periodic law) – 1-7 • Groups – vertical columns – similar properties – same # of e’s in outer energy level – 1-18

4. Atomic Mass Atomic Mass Atomic Mass Name Name Name Döbereiner’s Triads Johann Döbereiner ~1817 Calcium 40 Barium 137 Average 88.5 Strontium 87.6 Chlorine 35.5 Iodine 127 Average 81.3 Bromine 79.9 Sulfur 32 Tellurium 127.5 Average 79.8 Selenium 79.2 Döbereiner discovered groups of three related elements which he termed a triad.

5. Dmitri Mendeleev

6. Henry Moseley • Periodic Law

7. Metals • Good conductors of heat and electricity • High luster or sheen • All are solid at room temp. except mercury • Ductile – can be drawn into wires • Malleable – can be hammered into thin sheets w/o breaking • Examples = pg. 159

8. Nonmetals • Poor conductors except carbon • Upper right-hand corner of periodic table • Gases or brittle solids at room temp., bromine is a liquid

9. Metalloids • Metallic and nonmetallic properties • Most border the stair-step line • Pg. 160 6.1 assessment

10. Groups • Alkali metals – group 1 softer than most metals – most reactive esp. w/ oxygen and water – 1 e- in outer energy level • Alkaline earth metals – group 2 – 2 e- in outer • Halogens – group 17 or 7A – forms salt w/ a metal – 7 e- in outer • Noble gases – group 18 or 8A – stable – almost nonreactive (inert) (Representative elements 1-2, 13-17) • Transition metal – presence of e-’s in d orbital – groups 3-12 – occur in nature uncombined, often form colored compounds – iron triad (steel), coinage metals (copper, silver, gold) • Inner transition metal – presence of e-’s in f orbital – periods 6 and 7 b/w groups 3 and 4 – lanthanides 58-71, actinides 90-103=radioactive and unstable

11. CP 6.1 pg. 167 • PP 8-9 pg. 167 • 6.2 assessment pg. 167 • *Element Research*

12. Metallic Characteristic metallic character increases nonmetallic character increases metallic character increases nonmetallic character increases

13. Trends in atomic size • Atomic radius = one half of the distance b/w the nuclei of 2 atoms of the same element when the atoms are joined • Measured in picometers, 1 trillion picometers in 1 meter or 1012 • Atomic size increases from top to bottom w/in a group, decreases from L to R across a period

14. Ions • Ions = a charged atom, one that has gained or lost an electron • when sodium loses an electron to form with chlorine, in order to become stable) it becomes a positive one because it lost a negatively-charged electron • Cation = ion w/ a positive charge • In the same bonding, chlorine gains the electron sodium lost; thus, it becomes a negative one because it gained a negatively-charged electron • Anion = ion w/ a negative charge

15. Trends in Ionization Energy • Ionization energy = the energy required to remove an electron from an atom • First ionization energy – energy required to remove the 1st electron • First ionization energy decreases from T to B w/in a group, increases from L to R across a period • Li, Na, K, relatively easy to remove the first, difficult to remove a second thus they are 1+ ions • The lower the ionization energy, the easier it is to remove an e- (*Predicting reactivity demo pg 174*)

16. Shielding • Shielding = The shielding effect describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. • When more electrons are involved, each electron (in the n-shell) feels not only the electromagnetic attraction from the positive nucleus, but also repulsion forces from other electrons in shells from 1 to n. This causes the net force on electrons in outer shells to be significantly smaller in magnitude; therefore, these electrons are not as strongly bonded to the nucleus as electrons closer to the nucleus. • Shielding Demo – 6 volunteers

17. Trends in Ionic Size • When an atom loses e- it becomes smaller • Cations are smaller than the atoms from which they form, anions are larger””””

18. Trends in Electronegativity • Electronegativity = the ability of an atom of an element to attract electrons when the atom is in a compound • Values decrease from T to B w/in a group. For representative (1,2, 13-18), values increase from L to R across a period • Fluorine (4.0) highest, cesium (0.7) lowest

19. Periodic Table with Electronegativities

20. Electron affinity • The Electron affinityof an atom or molecule is defined as the amount of energy released when an electron is added to a neutral atom or molecule to form a negative ion.

21. Summary of Periodic Trends Shielding is constant Atomic radius decreases Ionization energy increases Electronegativity increases Nuclear charge increases 1A 0 Nuclear charge increases Shielding increases Atomic radius increases Ionic size increases Ionization energy decreases Electronegativity decreases 2A 3A 4A 6A 7A 5A Ionic size (cations) Ionic size (anions) decreases decreases

22. Look at Fig. 6.22 pg. 178 • Listing Elements (Ions) Activity – pg. 172 • 6.3 section assessment pg. 178 • Quick Lab – Periodic trends in ionic radii pg. 175 • Work on Packet