The Periodic Table and TrendsTopics 2 and 3

SONGPlease have a periodic table out.

IB prefers this one.

Dmitri Mendeleev 1834 – 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.

Henry Moseley 1887 – 1915

• arranged the elements in increasing atomic numbers (Z)– properties now recurred

periodically

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

3 5 8

2

1

What two blocks will always be the highest occupied level?

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

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 2s and 1 in the 2p• 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

Trends in the table IB loves the alkali metals and the

halogens

• 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

This is a simple, yet very good picture. Do you understand it?

– the distance from the nucleus to the outermost electron

– cannot measure the same way as a simple circle due to electrons are not in a fixed location

– therefore measure distance between two nuclei and divide by two

• ATOMIC RADII

–groups • increases downwards as more levels are added• more shielding

–periods across the periodic table• 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 because they are still in the same level

H

Li

Na

K

Rb

McGraw Hill video

Looking at ions compared to their parent atoms

• atoms tend to gain or loose electrons in order to have the electron configuration of a noble gas– do atoms become smaller or larger when

they do this?

IONIC RADII

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

Li 0.152 nm

Li+ .078nm

+Li forming a

cation

–anions (- ions) are larger than parent atom• have gained an electron to achieve noble gas

configuration• effective nuclear charge has decreased since

same nucleus now holding on to more electrons• plus, the added electron repels the existing

electrons farther apart (kind of “puffs it out”)

F 0.064 nm9e- and 9p+

F- 0.133 nm10 e- and 9 p+

-

– trends• across a period

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

(- ion)–then goes back to decreasing after just like

neutral atoms because of more protons pulling in the outer level

• down a group (same as neutral atoms)– increases as new levels are added–more levels shielding

DARK GREY IS THE SIZE OF THE ION

– IONIZATION ENERGY• the minimum energy (kJ mol-1) needed to

remove an electron from a neutral gaseous atom in its ground state, leaving behind a gaseous ionX(g) X+(g) + e-

• first ionization energy- energy to remove first electron

• second ionization energy- energy to remove second electron

• third ionization energy- and so on…

don’t forget-- gaseous

• 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– the nucleus is becoming stronger (effective

nuclear charge) and therefore the valence electrons are pulled closer• atomic radii is decreasing • this makes it harder to remove a valence electron

since it is closer to the nucleus

– or another way to look at it… a stronger nuclear charge acting on more contracted orbitals

http://www.google.com/url?sa=i&rct=j&q=ionization+energy&source=images&cd=&cad=rja&docid=f4emPodIHJdDZM&tbnid=a-sl1LeUwYM5KM:&ved=0CAUQjRw&url=http://www.kentchemistry.com/links/bonding/IE.htm&ei=nSQPUZPGDsqAiwK-mIH4Cw&bvm=bv.41867550,d.cGE&psig=AFQjCNGHK8rc758xjeEVDYV0pBZiiKj0Ng&ust=1360033302390412

• the increase in ionization energy is not continuous across the table

• electrons are also harder to remove…–a sub-level (s,p,d,f) is completely filled–a sub-level (s,p,d,f) are half filled

only last sub-level shown

notice discontinuity as move across period 3

• ELECTRON AFFINITY (Ea)– the change in energy (kJ/mole) of a

neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion (anion)

– in other words, the neutral atom's likelihood of gaining an electron

– example• F(g) + e- F-(g) will release 328 kJ/mole of

energy– the more negative Ea, the greater the attraction

for the electron

• trends–across a period

• in general, Ea become more negative from L to R

–down a group• in general, becomes less negative

This one same as “IB textbook”

• ELECTRONEGATIVITY–measures the attraction for a shared pair

of 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.

• trends (same as ionization energy and for the same reasons)

• 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 bonding pair of electrons (-) are shielded because of core electrons (-) interfering with the nucleus’ (+) hold on valence electronsH

Li

Na

K

Rb

• as you go across a period– electronegativity increases

• the atoms become smaller as the effective nuclear charge increases –easier to attract a shared pair of electrons

as they will be in a level closer to the nucleus moving from L to R on the table

• next concepts require understanding of concepts covered in later topics (this year and even senior year)

• only need to know the trends, not the reason why until later

–MELTING POINT• down group 1 (alkali metals)

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

• down group 17 (halogens)– increases 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 453

Na 370

K 336

Rb 312

Cs 301

Fr 295

increases increases

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

• increases until group 14 (think diamonds) then decreases starting at group 15

– 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 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)

Cl2Colorless- no reaction

turns red due to formation of Br2

turns brown due to formation of I2

Br2 no reaction no reactionturns 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

Look at the blue arrows! Senior year…

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