The Periodic Table of Elements

Organization of the Periodic Table

• Arranged to make predicting the properties of the elements easier.

• The order is based on the # of protons an atom of that element has in its nucleus. – This order is known as the periodic law

(arranged this way, similarities in their properties will occur in a regular pattern).

Arrangement of the P.T.o.E

• The periodic table is arranged in:– Horizontal rows (period)

• # of protons/electrons increase as you move from left to right. Because of this… you can determine the electronic configuration of the atom. Look at the location on the periodic table. What period is the element in? Where in the row is the element located?

Arrangement of the P.T.o.E

Aufbau’s electronic shell filling by period & energy.• Period 1= 1s• Period 2 = 2s, 2p• Period 3 = 3s, 3p• Period 4 = 4s, 3d, 4p• Period 5 = 5s, 4d, 5p• Period 6 = 6s, 4f, 5d, 6p• Period 7 = 7s, 5f, 6d, 7p

– The 1st 2 columns- electrons always occupy the s-orbital. The last 6 columns- electrons always occupy the p-orbital.

Arrangement of the P.T.o.E

• Vertical columns (family or group)– all have the same #of valence electrons. These elements

have similar properties. • The elements in the 1st column are called the alkali metals and

each have 1 valence electron. • The elements in the 2nd column are called the alkaline earth

metals and each has 2 valence electrons. • The elements in the 7th column are called the halogens and

each has 7 valence electrons. • Finally, the elements in the 8th column are known as the inert or

noble gases. They have a full octet or 8 valence electrons and are not likely to react.

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Atoms, Ions and Isotopes

• Some Atoms form Ions:– When atoms have partially full outermost

energy levels, they may undergo ionization (the gaining or losing of valence electrons).

– As atoms gain or lose electrons they no longer have the same # of electrons as protons. The charges no longer cancel out and you are left with a charged atom or ion.

Ex: Li loses an electron to form Li+

Mg loses 2 electrons to form Mg2+

F gains an electron to form F-

Ion Classification

• Atoms who lose electrons, become positively charged (like Li and Mg), and are called cations.

• Atoms who gain electrons, become negatively charged (like F), and are called anions.

• ** HINT** Cations: Alkali metals +

Alkaline-earth metals 2+

Anions: Group VI (O to Po) 2-

Halogens (F to At) -

• The Noble/ Inert gases do not make ions.

Atomic Number

• Atomic # (Z) - how many protons are in a atom. – Remember: atoms are electronically neutral

so they have an equal # of protons as electrons. Thus, the atomic # also tells us how many electrons the atom has.

Ex: Pu Z=94

Rf Z=104

Atomic Mass• Atomic mass # (A) = the # of protons +

neutrons in an atom. – Because the bulk of the atoms mass is provided by

the protons and neutrons, we only consider their masses when calculating the atomic mass #.

• Ex: F has 9 protons and 10 neutrons, so A= 19.

ISOTOPES

• Atoms of the same element will ALWAYS have the same # of protons, however, they may differ in the # of neutrons. – This means that different atoms

of the same element may/will have different masses. These are isotopes (atoms having the same # of protons but different # of neutrons).

Isotopes cont.• Some isotopes are more common than

others. – H is present on the earth and the sun. The most

common isotope is protium, then deuterium and finally, tritium.

Ex: 3 isotopes of H (why A isn’t just 1 but 1.00794).Protium (only has one proton in the nucleus) A=1

Deuterium (1 proton + 1 neutron) A=2

Tritium (1 proton + 2 neutrons) A=3Calculating # of neutrons: A – Z= # of neutrons

Atomic Mass Units

• The mass of a single element is extremely small (in the one trillionth of a billionth range) and is very difficult to work with. So instead we express the mass of atoms in atomic mass units (amu).

• One amu is equal to 1/12th of the mass of a carbon-12 atom. – This isotope has exactly 6 protons and 6

neutrons so the mass of each has to be about 1.0 amu.

Atomic Mass Unit cont.• The atomic mass of an element is often

listed as the average atomic mass as found in nature.

• This is a weighted average of the isotopes for that particular element. The more commonly found isotopes have a greater effect on the averages mass than the more rare isotopes.

Ex: Cl 24% Cl-37 and 76% Cl-35, thus the average atomic mass (35.45 amu) is much closer to 35 than 37.