PERIODIC TABLE

History of the periodic table

• . Johann Dobereiner (1780-1849).•first known chemist to recognize periodic patterns.

•suggested a “triad” system for classification, but found it limiting.

•John Newlands.•* In 1864, he suggested that elements should be arranged in “octaves,” based on the idea that certain properties seemed to repeat every eighth element; still not very successful.

Atomic Mass• Avogadro had suggested the idea of a mole. • Which was a specific number of atoms, he had no idea of actual the amount, 6.022x1023.

• He said equal volumes of gases at the same pressure and temperature would have the same number of particles, thus 1 mole of any gas at STP = 22.4 L

• By reacting moles of elements with each other you can determine what one mole of every element is.

• If you take the mass of one mole of an element you have the atomic mass.

Mendeleev and Meyer

•Julius Meyer and Dmitri Mendeleev both independently conceived the present periodic table.

•Mendeleev is given credit because of his ability to predict accurately the properties of undiscovered elements.

Dmitri Mendeleev

• Element #101 was named in his honor • first to come up with a periodic table • ~1870 there were 63 elements known to man, he organized the elements in order of their atomic mass, he saw a pattern form with the # of elements that can be bonded to that element.

More Mendeleev

• He arranged his table so that columns of elements with repeating patterns formed, he left spaces in his table where it appeared an element wasn’t discovered yet

• He then predicted the properties of these elements by looking at the other elements in the family and period of that element

Medeleev’s Table

It may not sound like much, but…

• He predicted what undiscovered elements would be like giving specifics!!!

• And he was Right!• This is the basis for acceptance of a scientific theory.

• If the model is correct then the scientist can assume what would happen if…

1871 DiscoveryProperty Ekaaluminum

(Mendeleev’s prediction)

Gallium

atomic mass 68 69.72

density (g/cm³) 6.0 5.904

melting point (°C)

Low 29.78

oxide's formula Ea2O3 (density - 5.5 g cm-3)

(soluble in both alkalis and acids)

Ga2O3 (density - 5.88 g cm-3)

(soluble in both alkalis and acids)

chloride's formula

Ea2Cl6 (volatile) Ga2Cl6 (volatile)

1882 DiscoveryProperty Ekasilicon Germanium

atomic mass 72 72.61density (g/cm³) 5.5 5.35melting point (°C) high 947color grey greyoxide type refractory dioxide refractory dioxideoxide density (g/cm³) 4.7 4.7oxide activity feebly basic feebly basicchloride boiling point under 100°C 86°C (GeCl4)chloride density (g/cm³) 1.9 1.9

Dmitri Mendeleev

Problem with Mendeleev’s table

• Mendeleev arranged his table by atomic mass

• a few elements appeared to be slightly out of place, Mendeleev put them in the right place and guessed that their atomic masses were incorrectly measured

• Actually, he was arranging them by the wrong number.

Atomic Number and Henry Moseley

• Moseley discovered that the charge of a nucleus is always divisible by the same number (which we now call of the charge of a proton).

• This gave him an integral number for each atom he called the atomic number.

• He corrected the periodic law when he stated that periodicity is a result of increasing atomic number, not increasing atomic mass.

Henry Moseley

Aufbau Principle• As protons are added one by one to the nucleus to build up the elements, electrons are similarly added to hydrogen like orbitals

• The lowest energy level is the 1 s orbital• From this we can begin to diagram the orbitals in an atom

•1 s __•2 s __ p __ __ __•3 s __ p __ __ __ d __ __ __ __ __

+ nucleuselectrons

Aufbau exceptions• You see the first exceptions to the Aufbau principle in the 3d orbital.

• Chromium is [Ar]3d54s1

• As opposed to the expected [Ar]3d44s2

• It then goes back to “normal”, until copper.• Copper is [Ar]3d104s1

• There are more examples is the d orbitals below, and in the f orbitals

Hund’s Rule

• The lowest energy configuration (ground state) for an atom is the one having the maximum number of unpaired electrons allowed by the Pauli Exclusion Principle (two electrons per orbital).

• The convention is to represent the spins of electrons as up or down arrows

Periodic Table

• The modern design of the periodic table is built off of these orbital diagrams

You can determine the orbital by the placement

Li

HHe

For these elements the last electrons are filling s orbitals

for these elements the last electrons are filling

p orbitalsFor these elements the lastelectrons are filling d orbitals

For the elements the last electronsare filling f orbitals

Electrons configurations are written to shorten orbital diagrams• The electron configuration for potassium is 1s22s2p63s2p64s1

• This can be abbreviated to [Ar]4s1

• This is the same as saying everything that is in Ar, 1s22s2p63s2p6 +4s1

Total order

• 1 s __• 2 s __ p __ __ __• 3 s __ p __ __ __ d __ __ __ __ __

• 4 s __ p __ __ __ d __ __ __ __ __ f __ __ __ __ __ __ __

• 5 s __ p __ __ __ d __ __ __ __ __ f __ __ __ __ __ __ __

• 6 s __ p __ __ __ d __ __ __ __ __

• 7 s __ p __ __ __

1

2 3

4 5

6

7

8

9

10

11

12

13

14

15

16

17

18

19

Aufbau Principle is the order of filling the orbitals

• 4s is filled before 3d• The (n+1)s orbitals always fill before the n d orbitals, explained by the penetration effect.

• The penetration effect is looking at the forces on the interior electrons in an atom. Not only is there an attraction to the nucleus, but the outer electrons repulsion is pushing these orbitals closer to the nucleus, making them smaller.

Lone electron in the 5d or 6 d• Another exception• Lanthanum’s last electron fills the 5d orbital. [Xe]6s25d1

• Ce has its last electron fill in the 4f, and the one from 5d (La) jumps to 4f. [Xe]6s24f2

• The next elements all fill up the 4f until Gadolinium which fills up the 5d again.

• A similar “jumping” of one electron occurs with the actinide series

Different Periodic Tables

• The electron “jumping” is why there are two forms of the periodic table in common usage.

• One has lanthanum and actinium in with the d filling elements.

• The other has lanthanum and actinium with the f filling elements

To Determine the Condensed electron configuration

• Start at the previous noble gas (lower atomic number)

• Write out every element in order of atomic number to your given

• Determine which shell all elements are filling. The number of elements in that shell is your superscript number.

For Iron•Previous Noble Gas----Ar•Write all elements•K Ca Sc Ti V Cr Mn Fe•These fill 4 s•These fill 3 d•So it is• [Ar]3d64s2

Practice

• Write out the condensed electron configuration and orbital diagram for…

• Nitrogen• Tin• Molybdenum• Rutherfordium

Effective Nuclear Charge (Zeff)•Zeff is the net positive charge experienced by a specific electron in a multi electron atom.

• As electrons get further away, they will be “shielded” from the nucleus by other electrons.

• What is the period trend, with explanation?

Ionization Energy• The energy required to remove an electron.•What type of ion is produced?

•Is the process endothermic or exothermic?

•What is the group trend, with explanation?

•What is the period trend, with explanation?

Exceptions• Normally moving left to right across a period will increase in ionization energy because it is being held in by a stronger nucleus.

• However there is an observed dip in ionization energy going from Be to B.

• Why?• It is harder to remove an electron from 2p than 3s, these aren’t hydrogen orbitals.

• This trend continues for Mg-Al, then you see the trend for Zn-Ga, and Cd-In, and Hg-Tl

Another exception

• There is another between group 15 and 16 (nitrogen group to oxygen group).

• This exception is because in group 15 you have one electron in each of the 3 p orbitals, group 16 has a 4th electron that is easier to remove.

Successive Ionization Energies

•Ionization energy doesn’t have to stop at one electron, you can remove multiple electrons, to get a 2nd ionization energy I2, 3rd ionization energy I3 …

•Why are second ionization energies more than first ionization energies?

•Electrons are more easily removed until what occurs?

Trends in Ionization Energy

•The first ionization energy for phosphorus is 1060 kJ/mol, and that for sulfur is 1005 kJ/mol. Why?

•Consider atoms with the following electron configurations:

•  1s22s22p6

• 1s22s22p63s1

• 1s22s22p63s2

• Which atom has the largest first ionization energy, and which one has the smallest second ionization energy? Explain your choices.

Atomic Radius• ~relative size of an atom. • As you move down a period you are increasing the number of energy levels so the atom is getting larger.

• As you move left to right across a period, a more positive nucleus pulls everything in closer, making the atom smaller.

• It increases down and to the left.

Electron Affinity

•Electron affinity is the change in energy when an atom captures a bonding electron.

•Energy is normally released, so the number is normally reported as a negative.

•What is the group trend, with explanation?

•What is the period trend, with explanation?