The Periodic Table and TrendsTopics 3.1 - 3.3 and 12.1.1 - 12.1.2

SONGPlease have a periodic table out.

IB prefers this one.

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 now atomic #)

• he even left empty spaces to be filled in later

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

• all elements have 1,2,3,4,5,6,7, or 8 valence electrons

Lewis Dot-Diagrams/Structures• valence electrons are represented as dots

around the chemical symbol for the element

Na

Cl

2

2,3 2,5 2,8

2,8,2

3.1.3 Electron arrangement and position on the table 3.1.4 # of electrons in the highest occupied energy level and position on the table

2,1

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Look, they are following

my rule!

• 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 3.2.1• atomic radii– the distance from the nucleus to the outermost

electron• ionic radii

– same measurement as atomic radii but have lost or gained valence electrons

• first ionization energy (kJ mol-1)– the energy needed to remove the outermost/valence

electrons/highest energy electron from a neutral atom in the gaseous phase, leaving behind a gaseous ion• X(g) X+(g) + e-

• Na(g) Na+(g) + e-

don’t forget-- gaseous

• electronegativity– measures the attraction for electrons in a bond

• Linus Pauling (1901 to 1994) came up with a scale where a value of 4.0 is arbitrarily given to the most electronegative element, fluorine, and the other electronegativities are scaled relative to this value.

• melting point• chemical properties

– how elements react with other elements

Trends in the table (3.2.2- 3.2.4)

IB loves the alkali metals, the halogens, and period 3(look at the assessment

statements)

• many trends are easier to understand if you comprehend the following

• the ability of an atom to “hang on to” or attract its valence electrons is the result of two opposing forces – the attraction between the electron and the

nucleus– the repulsions between the electron in question

and all the other electrons in the atom (often referred to the shielding effect)

– the net resulting force of these two is referred to effective nuclear charge

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• 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 in the given level closer to it

» added electrons are not contributing to the shielding effect

H

Li

Na

K

Rb

–IONIC RADII –across period 3

• decreases at first when losing electrons (+ ion)• then suddenly increases when gaining

electrons (- ion)– then goes back to decreasing after that

DARK GREY IS THE SIZE OF THE ION

This is period 2 but same trend as period 3

– alkali metals• cations (+ ions) are smaller than 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 nm

Li+ .078nm

+Li forming a

cation

–halogens (group 7)• anions (- ions) are larger than parent atom

– have gained an electron to achieve noble gas configuration

– effective nuclear charge has decreased since same nuclues now holding on to more electrons

• 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

– IONIZATION ENERGY• recall, the energy to remove an electron

– example: Na(g) Na+(g) + e-

• 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 and therefore easier to remove

• increases across a period– extra electrons are just filling up the same level

(or shell, same thing)– the nucleus is becoming stronger and therefore

the electrostatic force increases making the atom smaller• this makes it harder to remove a valence electron

since it is closer to the nucleus– or another way to look at it… a higher nuclear

charge acting on more contracted orbitals

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More on ionization energy-- Topic 12.1.1 and 12.1.2

• Evidence for levels and sub-levels –using the FIRST IONIZATION energy

for elements• electrons are harder to remove…

–when there are more protons to attract the valence electrons in a given level (higher nuclear charge)

–when the atomic radius decreases and the electrons are closer to the nucleus

–a sub-level (s,p,d,f) is completely filled–a sub-level (s,p,d,f) are half filled

only last sub-level

1s

2s

2p

3s

3p

4s

• Evidence for levels and sub-levels –using SUCCESSIVE (1st, 2nd, 3rd,…)

ionization energy for an element• as more electrons are removed, the effective nuclear

charge increases, pulls the electrons in closer and holds them tighter

– 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

» less electrons are shielding every time you remove a level

Topic 12.1.2

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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

K

–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

Back down to Topic 3

H

Li

Na

K

Rb

• as you go across a period– electronegativity increases

• the atoms become smaller as the nuclear charge increases –attracts electrons as they can be closer to the

nucleus moving from L to R on the table

• next concept requires understanding of concepts covered in later topics– what a metal is– Van der Waals forces– bonding

–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 increases

• 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 (3.3.1 AND 3.3.2)

–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 attracts electrons better since they can be close to the nucleus

1+ charge

1- charge

most reactive

least reactive

•halogens (group 7)

–diatomic molecules such as Cl2, Br2, I2

»can react with halide ions (Cl -, Br -, and I -)» the most reactive ends up as an ion (1- charge)

and is not visible (molecules 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–oxides are

» ionic…..and then covalent bonds–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

Just look for the on the next slide It is basically 3.3.2

http://www.free-smiley-faces.de/free-animated-laughing-Smileys-animiert-froehlich-lachend/www.free-smiley-faces.de_animated-laughing-smiley_07_800x800.gif