chem

DISCUSSION NO. 3
THE PERIODIC TABLE OF ELEMENTS

Members:
Salvador, Kim Kristal G.
Sanchez, Jhoana Marie V.
Santeco, Francesca Anya C.
Sy, Shane Francheska Lei Dayanne S.
Ureta, Nheimelyn E.
INUR-7 GROUP 10
HISTORY OF THE PERIODIC TABLE
The discoveries of elements stole a long period of time. As a matter of fact, several elements were mentioned in Biblical accounts and Koran. However, it was in the 19th century where they get to discover a lot of element. And as the number of elements grew in number, so does the interest of the scientists in developing its techniques.
Johannes Dobereiner
In the midst of 1817 a German chemist in the name of Johannes Dobereiner discovered that properties of elements like Fe, Co and Ni were similar .And from there he made his very own law of triads which contains groups of three elements based on their atomic masses.
Example;
Alkali Formers Salt Formers Li 7 Cl 35.5 Na 23 Br 80 K 39 I 127

*He arranged from lightest to the heaviest atomic mass.And so he discoveres that the average atomic mass of the lightest element together with the heaviest is very close to the atomic mass of the middle element.
John Alexander Reina Newlands
Then came an English chemist, John Newlands. In 1868, he arranged them based on their atomic masses. He noticed that the chemical properties, melting point, boiling point of the element have a regular pattern. He noticed that the properties would repeat every eight element. Having a musical background, he called this as “The Law of Octave”.His work is not accepted by his colleagues but he fought for it until finally in 1887 he received the davy medal in recognition of his earlier work.

Periodic Property is one which varies regularly when elements are arranged according to a common basis.

Julius Lothar von Meyer (August 19, 1830 – April 11, 1895)
A German chemist, competitor and colleague of Mendeleev in 1869 whom included 56 elements in his tale.at first he classified 28 elements in his periodic table and for the first time he grouped the elements according to their valence but eventually it became inaccurate in terms of the measurements of their atomic weights. He also received a Davy medal in recognition of his work.

Meyer table with a horizontal display of periods in 1864[edit]

Valence IV
Valence III
Valence II
Valence I
Valence I
Valence II
The mass difference
I line

Li
Be
~16
II line
C
N
O
F
Na
Mg
~16
III line
Si
P
S
Cl
K
Ca
~45
IV line

As
Se
Br
Rb
Sr
~45
V line
Sn
Sb
Te
I
Cs
Ba
~90
VI line
Pb
Bi

Tl

~90
Meyer table with vertical display of periods in 1870

Dmitri Ivanovich Mendeleev(1834-1907)
“The father of the Modern Periodic Table”
He has the most broad organization of elements which was in the 1869.The elements were arranged along a horizontal row, from left to right. In the order of their increasing atomic masses. An element that has the same properties would start a new row.Thus, the elements with the same properties will fall under one column. Mendeleeve assumed the fact that there were still undiscovered elements.So he left an extra column in his table for the elements that didn’t seem to fit any particular group. And he was accredited for all of these works.
Thereafter, Mendeleev and Meyer stated a Periodic Law ”properties of elements are periodic functions of their atomic masses.”
But Mendeleev had his shortcomings, it is he was not able to predict the noble gases and arranging them in order of increasing atomic mass proved not to be a good basis.

His Table’s shortcomings:
1. Hydrogen was placed in group 1 although it resembles both characteristics of the alkali metals and the halogens.
2. Some elements were arranged according to their properties and not of their increasing atomic masses ex: cobalt with an atomic mass of 58.9 was placed before nickel with 58.6
3. The position for lanthanides and actinides were not given

“I began to look about and write down the elements with their atomic weights and typical properties, analogous elements and like atomic weights on separate cards, and this soon convinced me that the properties of elements are in periodic dependence upon their atomic weights.”
--Mendeleev, Principles of Chemistry, 1905, Vol. II
Henry Gwyn Jeffreys Moseley (23 November 1887 – 10 August 1915)

An English physicist, who worked on atomic number. Discovered that there is a mathematical relationship between the wavelengths of the X-rays and atomic numbers of the metals that were used as the targets in X-ray tubes and stated that elements could be better related with their atomic numbers. He showed in his work the number of protons increases by 1 from element to element.
And after all these things the periodic law was restated
PERIODIC LAW
“When the elements are arranged in the order of their increasing atomic number, elements with similar properties appear at periodic intervals.”

```````The End May God Bless you!`````

PARTS OF THE PERIODIC TABLE

1. Period - it is the row in the periodic table. Elements in the same period has the same number of energy shells.
2. Group - it is the column in the periodic table. Elements in the same group has the same number of valence electrons.
PARTS OF AN ELEMENT IN A PERIODIC TABLE
1. Atomic number
2. Element symbol
3. Element name
4. Atomic mass number

References: http://www.chem4kids.com/files/elem_pertable.html http://www.youtube.com/watch?v=7mLPC74GHMo

CLASSIFICATION OF ELEMENTS IN THE PERIODIC TABLE
There are three groups that classify the elements in the periodic table. These three groups are: metals, nonmetals, and
Metals
Found in the left side of the periodic table.
They are good conductors of electricity and heat
They are shiny, malleable and ductile
a.) Alkali Metals
They are elements that are located in Group 1 of the periodic Table.
They are very reactive metals
They are soft, malleable, ductile
They are good conductors of heat and electricity
Examples: Lithium, Sodium, Potassium, etc.
b.) Alkaline Earth Metals
They are elements that are located in Group 2 of the periodic Table
They are all found in the Earth’s crust, but not in the elemental form because they are reactive.
They are widely distributed in rock structures
Examples: Beryllium, Magnesium, Calcium, etc.
c.) Transition Metals
They are elements that are located in Groups 3-12 of the periodic table.
They are ductile, malleable and can conduct electricity and heat.
Examples: Titanium, Zinc, God, Mercury, etc.
d.) Inner Transition Metals
They are sometimes called rare earth metals due to their extremely low natural occurrence.
They are as reactive as the alkali metals
Actinides
They are all radioactive
They are very dense metals with distinctive structures.
They combine directly with most nonmetals.

Lanthanides
They have high melting points and boiling points
They are very reactive
They burn easily in air
e.) Other Metals
They are solid and have a relatively high density and they are opaque.
Examples: Tin, Lead, Bismuth, etc.
Non-metals
They are elements that are located in Groups 14, 15 and 16min the Periodic Table
They are poor conductors of electricity and heat
They don’t reflect light
They are very brittle and cannot be rolled into wires or pounded into sheets.
Examples: Hydrogen, Carbon, Oxygen, etc.
Halogens
They are elements that are located in Group 17 of the periodic table.
They are called “salt formers” because compounds containing halogens are called “salts”.
Examples: Fluorine, Chlorine, Iodine, etc.
Noble Gases
They are elements that are located in Group 18 in the periodic table
They are stable because all noble gases have the maximum number of electrons possible in their outer shell.
Example: Helium, Neon, Argon, etc.
Metalloids
They are elements that are located in Groups 13, 14, 15, 16 and 17
They have both properties of both metals and non-metals.
Some of them are semi-conductors and carry an electrical charge in them that are useful in calculators and computers.
Example: Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium.

ELECTRONIC CONFIGURATION USING THE PERIODIC TABLE The electronic configuration shows the arrangement of electrons in an atom’s orbital. There are three important rules in determining the electronic configuration; first and the most basic rule is that electrons occupy first the orbitals with the lowest energy levels before occupying the higher levels, according to the Aufbau principle (from the German Aufbau meaning “building up”).The second rule states that an orbital can hold a maximum of two electrons and that no electrons can fill one orbital with the same spin, according to the Pauli Exclusion principle (formulated by an Australian physicist Wolfgang Pauli in 1925). The third rule states that electrons should fill each orbital singly before any orbital gets a second electron, according to Hund’s rule (formulated by German physicist Friedrich Hund).
We can have an analogy pertaining to those rules: A couple with a 6 storey-mansion, with one bedroom in each floor, has ten children. Each kid wants to have their own room, so in order to avoid the fighting of siblings; they first filled the room with one kid (Hund’s rule) starting with the youngest on the first floor of the mansion (Aufbau principle), since the youngest easily gets tired. The second on the second floor and until the fifth floor. Since the couple wants the younger ones to have someone with them, they placed the oldest kid with the youngest, the next oldest in the second and until in the fifth floor, since there are two beds in each room (Pauli Principle).
The periodic table can be used as a guide to determine the electron configuration of an atom. The period number is the value of n (shell). Each shells are divided into subshells and it contains orbitals. Groups 1A and 2A have the s-orbital filled. Groups 3A to 8A have the p-orbital filled. Groups 3B to 2B have the d-orbital filled. The lanthanides and actinides have the f-orbital filled.

In the first period, there are only two elements. Their electrons are in the first principal energy level: n=1. The second period contains a total of eight elements, having two sublevels: s and p; s sublevels contain 2 electrons when filled while p sublevels 6 electrons since p sublevels each contains 3 orbitals. Third is the same as the second period. In the fourth period, the 4s is filled first since it has lower energy level than 3d and so on. In the lanthanides and actinides (the fourteen rare earth elements), f orbitals are filled.

PERIODIC TRENDS

1. Atomic Size – size of an atom
Across the period, the atomic size decreases.
Reason: As you go across a period, electrons are added to the same energy level. Protons are added to the nucleus. Therefore, there is a stronger force of attraction pulling the electrons closer to the nucleus, making it smaller in size.

Down the group, the atomic size increases.
Reason: The number of energy levels increases as you move down, hence, making the atomic the size bigger.

2. Ionic Size – size of an atom after it becomes an ion
Cations are smaller than their neutral atoms.
Reason: Since cations lose electrons, the protons outnumber the electrons. The pull of protons will be stronger.

Anions are bigger than their neutral atoms.
Reason: Since anions gain electrons, the repulsion of electrons make the size bigger.

3. Metallic Character – how readily metals lose their electrons
Across the period, metallic character decreases.
Reason: This occurs as atoms more readily accept electrons to fill a valence shell than lose them to remove the unfilled shell.

Down the group, metallic character increases.
Reason: This is because electrons become easier to lose as the atomic size increases, where there is less attraction.

4. Ionization Energy – energy needed to remove an electron from an atom
Across the period, ionization energy increases.
Reason: The atomic size is smaller. That means, electrons are closer and are more strongly attracted to the nucleus. Therefore, it becomes more difficult to remove the electron.

Down the group, ionization energy decreases.
Reason: Electrons are farther from the nucleus. Thus, making it easier to be removed.

5. Electron Affinity – energy released when an electron is added to a gaseous atom Across the period, electron affinity increases. Reason: Non-metals have the highest electron affinity values, because gaining electrons results in completely filled shells.

Down the group, electron affinity decreases.
Reason: When an electron is added to a larger atom, less energy is released, because the electron can’t move as close to the nucleus as it can in a smaller atom.

Group VIII elements, the noble gases, have electron affinity values near zero, since each atom possesses a stable octet and will not accept an electron readily.

6. Electronegativity – ability of an atom to attract electrons to itself
Across the period, electronegativity increases.
Reason: Since the valence shell is almost full, it is easier to pull an electron to fill the shell.

Down the group, electronegativity decreases.
Reason: The bonding pair of electron is increasingly distant from the attraction of the nucleus.

Noble gases have complete valence shells and do not usually attract electrons.

References: http://www.chem.tamu.edu/class/majors/tutorialnotefiles/trends.htm http://www.chem.tamu.edu/class/majors/tutorialnotefiles/metals.htm http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Periodic_Trends http://en.wikibooks.org/wiki/High_School_Chemistry/Electron_Affinity

References: http://www.chem.tamu.edu/class/majors/tutorialnotefiles/trends.htm http://www.chem.tamu.edu/class/majors/tutorialnotefiles/metals.htm http://chemwiki.ucdavis.edu/Inorganic_Chemistry/Descriptive_Chemistry/Periodic_Table_of_the_Elements/Periodic_Trends http://en.wikibooks.org/wiki/High_School_Chemistry/Electron_Affinity