The Periodic Table and Physical Properties

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Topics 3.1 - 3.3 and 12.1.1 - 12.1.2

Get a periodic

table out.

Dmitri Mendeleev 8 February 1834 – 2 February 1907

• Russian chemist and teacher• given the elements he knew

about, he organized a “Periodic Table” based on increasing atomic mass (it’s atomic # now)

• he even left empty spaces to be filled in later (TOK– he was a “scientist” and “risk taker”!)

At the time the elements gallium and germanium were not known. These are the blank spaces in his periodic

table. He predicted their discovery and estimated their properties.

Z

Design of the Table• Groups are the vertical columns.

– elements have similar, but not identical, properties• most important property is that

they have the same # of valence electrons

valence electrons- electrons in the highest occupied energy level

http://images.google.com/imgres?imgurl=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/imgper/econfig.gif&imgrefurl=http://hyperphysics.phy-astr.gsu.edu/hbase/pertab/perlewis.html&h=267&w=512&sz=22&tbnid=__EXctBwlG0J:&tbnh=66&tbnw=128&hl=en&start=1&prev=/images%3Fq%3DElectron%2BDot%2BDiagrams%26svnum%3D10%26hl%3Den%26lr%3D

2

2,3 2,5 2,8

2,8,2

Electron arrangement (SL level – 3.1.3)

2,1

http://www.google.com/url?sa=i&rct=j&q=meaning+spdf+periodic+table&source=images&cd=&cad=rja&docid=-AsBTI9Np9nQ0M&tbnid=bsqX3-sPDM8f1M:&ved=0CAUQjRw&url=http://msjschemclass.blogspot.com/2010_02_01_archive.html&ei=W_UJUeiHB8bKiwKOp4GACg&bvm=bv.41642243,d.cGE&psig=AFQjCNEkXJ2vzpDdrafm3Mmpxi_KeMbChQ&ust=1359693528210963

• B is 1s2 2s2 2p1;– 2 is the outermost energy level – it contains 3 valence electrons, 2 in

the s and 1 in the p• Br is [Ar] 4s2 3d10 4p5

How many valence electrons are present?

• Periods are the horizontal rows– do NOT have similar properties– however, there is a pattern to their properties as

you move across the table that is visible when they react with other elements

Definitions• atomic radii

– the distance from the nucleus to the outermost electron

• ionic radii– same distance, but for ions (atoms that have lost

or gained valence electrons)• first ionization energy (kJ mol-1)

– the energy needed to remove the outermost, or highest energy, electron from a neutral atom in the gaseous phase

• electronegativity– measures the attraction for a shared pair of

electrons• melting point• chemical properties

– how elements react with other elements

Trends in the table

But first, the electron shielding effect

• electrons between the nucleus and the valence electrons repel each other

• ATOMIC RADII – McGraw Hill video – groups (alkali metals and halogens)

• increases downwards as more levels are added– periods across the periodic table (period 3)

• radii decreases – the number of protons in the nucleus increases

» increases the strength of the positive nucleus and pulls electrons closer to it

H

Li

Na

K

Rb

• IONIC RADII – decreases across periods for same reason

as atomic radii (nucleus becomes stronger)– alkali metals

• cations are smaller that the parent atom– have lost an electron (actually, lost an entire level)– therefore have fewer electrons than protons

• radii still increases downwards as more levels are added on

Li 0.152 nm3e and 3p

Li+, 0.078 nm

2e and 3 p

+forming a

cation

–halogens• anions are larger than parent atom

– have gained an electron to achieve noble gas configuration

• radii still increases downwards as more levels are added on

F 0.064 nm9e- and 9p+

F- 0.133 nm10 e- and 9 p+

- forming an

anion

http://www.google.com/url?sa=i&source=images&cd=&cad=rja&docid=6-36i7PgL4iMDM&tbnid=tsNxs705PypEYM:&ved=0CAgQjRwwADgW&url=http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Periodic_Properties_of_the_Elements&ei=fsEMUYukCMazigKB6YH4DA&psig=AFQjCNHc_PVwitHNAAKSKujAesgidWigUQ&ust=1359876862165395

– IONIZATION ENERGY• decreases down a group

– outer electrons are farther from the nucleus and therefore easier to remove

– inner core electrons “shield” the valence electrons from the pull of the positive nucleus

• increases across a period– extra electrons are just filling up the same level– the nucleus is becoming more powerful and

therefore the electrostatic force increases making it harder to remove an electron

12.1.1

• Evidence for levels and sub-levels – First ionization energy

• electrons are harder to remove…– when there are more protons to attract them– a sub-level (s,p,d,f) is completely filled– a sub-level (s,p,d,f) are half filled

http://www.google.com/url?sa=i&rct=j&q=first+ionization+energy&source=images&cd=&cad=rja&docid=O4XPgoQ4n23uUM&tbnid=8stKgOyumBvN_M:&ved=0CAUQjRw&url=http://staff.norman.k12.ok.us/~cyohn/index_files/ionizationenergynotes.htm&ei=0nYQUYHkCuzgigKV9oD4Dg&bvm=bv.41934586,d.cGE&psig=AFQjCNF06moKjTqNA9CwedXrw-CBodPy_w&ust=1360119872878443

12.1.1

• Evidence for levels and sub-levels – successive (1st, 2nd, 3rd) ionization energy

• as more electrons are removed, the electrostatic pull of the protons holds the remaining electrons closer

• therefore, more energy is required to remove them (even have to use a logarithmic scale to show this)

• large “jumps” are when the electrons are being removed from the next, lower level that are much closer to the nucleus

http://www.google.com/url?sa=i&rct=j&q=ionization+energy&source=images&cd=&cad=rja&docid=RUeqin9x2yNKnM&tbnid=fm5xZiut5BnBvM:&ved=0CAUQjRw&url=http://www.avon-chemistry.com/p_table_lecture.html&ei=bbMMUcewApG0igKt3oCYCA&bvm=bv.41867550,d.cGE&psig=AFQjCNFCaL0lXIfQTXo1TXl1vX0NDNICAg&ust=1359873240414555

4s1 removed

starting to remove the 3p sub-level

starting to remove the 2p

sub-level

starting to remove the 3s

sub-level

starting to remove the 2s sub-level

starting to remove

the 1s sub-level

–ELECTRONEGATIVITY• as you go down a group electronegativity

decreases – the size of the atom increases

» the bonding pair of electrons (-) is increasingly distant from the attraction of the nucleus (+)

» the valence electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electrons

• as you go across a period– electronegativity increases

• the atoms become smaller so the positive nucleus can hold onto the electrons better

http://www.google.com/url?sa=i&rct=j&q=shielding+effect+trend&source=images&cd=&cad=rja&docid=I7WJwI2OeZNfaM&tbnid=cMI-TQz1CuZJnM:&ved=0CAUQjRw&url=http://www.msduncanchem.com/Unit_5/trends_notes_rev.htm&ei=YMIMUcbcE82vigLay4DgBQ&psig=AFQjCNGnN54NXyVjLx7i1hyJJOMcUsyvEA&ust=1359876822976465

–MELTING POINT• group 1 (alkali metals)

– decreases as “sea of negative electrons” are farther away from the positive metal ions

• group 7 (halogens)– increases downwards as the van der Waals’ forces increase

» larger molecules have more electrons which increases the chance that one side of the molecule could be negative

Element Melting Point (K)

Li 453Na 370K 336Rb 312Cs 301Fr 295

increases decreases

• across the table (period 3)– from left to right

• bonding goes from strong metallic to very strong macromolecules (network covalent) to weak van der Waals’ attraction

• CHEMICAL PROPERTIES–groups

• alkali metals– react vigorously with water and air

» 2Na (s) + H2O (l) 2Na (aq) + 2OH- (aq) + H2 (g)» (Li, Na, K… all the same equation)

» reactivity increases downwards» because the outer (valence) electron is in higher energy

levels (farther from the nucleus) and easier to remove – react with the halogens

» halogens’ reactivity increases upwards» smaller size can attract electrons

better» (see next slide)

1+ charge

1- charge

most reactive

least reactive

• halogens–diatomic molecules such as F2, Cl2, Br2, I2

» can react with halide ions (Cl -, Br -, and I -)» the single bond is broken and each atom can gain one

electron to form halide ions (F1-, Cl1-, Br1-, I1-)» the most reactive ends up as an ion (1- charge) and is not

visible (molecules F2, Cl2, Br2, I2 are a visible gas)

» Cl > Br > I  Cl-(aq) Br-(aq) I-(aq)

Cl2 Colorless- no reaction

turns red due to formation of Br2

turns brown due to formation of I2

Br2 no reaction no reaction turns brown due to formation of I2

I2 no reaction no reaction no reaction

–periods• from left to right in period 3

–metals…metaloids…nonmetals–when oxides react with water

»basic…amphoteric (either basic or acidic)…acidic

» Na2O(s) + H2O (l) 2 NaOH (aq) strong base» MgO (s) +H2O (l) Mg(OH)2 (aq) weaker base» P4O10 (s) + 6H2O (l) 4 H3PO4 (aq) weak/strong acid» SO3(g) + H2O (l) H2SO4 (aq) strong acid

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