The History of the Modern Periodic Table Modified from www.thecatalyst.com
During the nineteenth century, chemists began to categorize the elements according to similarities in their physical and chemical properties. The end result of these studies was our modern periodic table!
Johann Dobereiner 1780-1849 In 1829, he classified some elements into groups of three, which he called triads. Model of triads The elements in a triad had similar chemical and physical properties.
John Newlands 1838-1898 In 1863, he suggested that elements be arranged in octaves because he noticed that certain properties repeated every 8th element when placed in order of increasing atomic mass. Law of Octaves
John Newlands Newland s law of octaves failed beyond the element calcium. WHY? Would his law of octaves work today with the first 20 elements?
Elements known at this time
Dmitri Mendeleev 1834-1907 In 1869 he published a table of the elements organized by increasing atomic mass.
Lothar Meyer 1830-1895 At the same time, Lothar Meyer published his own table of the elements organized by increasing atomic mass.
Mendeleev vs. Meyer Both Mendeleev and Meyer arranged the elements in order of increasing atomic mass. Both left vacant spaces where unknown elements should fit. So why is Mendeleev called the father of the modern periodic table and not Meyer?
And the Winner is Mendeleev stated that if the atomic weight of an element caused it to be placed in the wrong group, then the weight must be wrong. Mendeleev was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown. Mendeleev s predictions for Sc, Ga, and Ge turned out to be amazingly close to the actual values!!
After the discovery of these unknown elements, Mendeleev s table was generally accepted.
Mendeleev s 15 minutes of Fame? In spite of Mendeleev s great achievement, problems arose when new elements were discovered and more accurate atomic weights determined. Can you identify the problematic elements? Ar and K Te and I Co and Ni Th and Pa
Henry Moseley 1887-1915 In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements. He then arranged the elements in order of increasing atomic number (nuclear charge).
Henry Moseley There is in the atom a fundamental quantity which increases by regular steps as we pass from each element to the next. This quantity can only be the charge on the central positive nucleus. Moseley s research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.
Glenn T. Seaborg 1912-1999 After co-discovering 10 new elements, he moved 14 elements out of the main body of the periodic table to their current location below the Lanthanide series. These became known as the Actinide series.
Glenn T. Seaborg He is the only person to have an element named after him while still alive. "This is the greatest honor ever bestowed upon me - even better, I think, than winning the Nobel Prize."
Periodic Table Geography
Periods The horizontal rows of the periodic table are called PERIODS.
Groups The vertical columns of the periodic table are called GROUPS, or FAMILIES.
Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties. In other words The elements in any group of the periodic table have similar physical and chemical properties!
Alkali Metals All of the elements in the first Group are considered part of the Alkali Metals Family.
Alkaline Earth Metals All of the elements in the second Group are considered part of the Alkaline Earth Metal Family.
Transition Metals The Transition metals are indicated by the letter B in the Group Number. All of these elements have a partially filled d-orbital.
Inner-Transition Metals These elements are also called the rare earth elements, as they are not commonly found in nature. In fact, many of them were created in particle accelerators.
Halogens All of the elements in Group VII A are considered part of the Halogen Family.
Noble Gases All of the elements in Group VIII A are considered part of the Noble Gas Family, as they are rather inert (i.e. non-reactive) under most conditions.
The periodic table is the most important tool in the chemist s toolbox!
Periodic Trends & Properties We will be looking at the following periodic trends in this chapter: Nuclear Charge Shielding Effect Atomic size (radius) Ionic size (radius) Ionization energy Electronegativity
Nuclear Charge The nuclear charge increases as you move down a group in the periodic table. The Reason: As Moseley discovered, the number of protons in the nucleus is what determines the nuclear charge. The more protons in the nucleus, the larger or stronger the nuclear charge.
Nuclear Charge Nuclear Charge increases from left to right across the periodic table. The Reason: The number of protons in the nucleus increases by one as you move from left to right across the table. Increasing the number of protons in the nucleus results in an increase in nuclear charge.
Shielding Effect Within any group of elements, the shielding effect increases as you move down the column. The reason: The elements located in higher periods have more electrons, and therefore require more [higher] energy levels. More energy levels translates into more inner electrons and thus more shielding.
Shielding Effect Within any period, the shielding effect decreases from left to right across the table. The Reason: Across the period, the nuclear charge increases but the number of energy levels remains the same. The increasing nuclear force draws the outermost electrons ever closer to the nucleus
Atomic Radius Within any group of elements, the atomic radius will increase from top to bottom. The reason: The elements located in higher periods have more electrons, and therefore require more [higher] energy levels. More energy levels translates into a larger radius.
With more nuclear charge, and no increase in shielding, the outermost electrons are drawn closer to the nucleus, and hence the atomic radius decreases. Atomic Radius Within a period, the atomic radius decreases from left to right across the periodic table. The reason: Elements within the same period have the same number of energy levels and therefore the same shielding effect. As you move toward the right, the nuclear charge increases.
Ionic Radius of Cations Cations are ALWAYS SMALLER in size than the neutral atom from which they originate. The Reason: The nuclear charge does not change, but the atom loses 1 or more electrons. Each valence electron experiences more electrostatic attraction toward the nucleus. As a result, the ion is smaller than the original atom. Lithium Lithium ion, Li +1
Ionic Radius of Anions Anions are ALWAYS LARGER in size than the neutral atom from which they originate. The Reason: The atom gains extra electrons, but the number of protons attracting those electrons remains the same. Each valence electron experiences less electrostatic attraction toward the nucleus. As a result, the ion is larger than the original atom. Oxygen Oxygen ion, O -2
Ionization Energies Within Periods Group 1 Metals Have 1 valence electron. It requires a small amount of energy to remove that single electron. (Low 1 st ionization energy) + 520 kj/mol In order to remove a second electron, it requires removing an electron from a stable (noble) configuration. (High 2 nd ionization energy) + 7,297 kj/mol
Ionization Energies Within Periods Group 2 Metals Have 2 valence electrons. It requires a small amount of energy to remove the first and second electrons. + 900 kj/mol + 1,757 kj/mol In order to remove a third electron, it requires removing an electron from a stable (noble) configuration. + 14,840 kj/mol
Trends in Ionization Energy Ionization Energy increases as you move left to right across the periodic table. The Reason: As you move from left to right within a period: Atomic radius decreases Shielding Effect does Nuclear charge increases not change As a result, valence electrons are drawn closer to the nucleus and held more tightly. This means it requires more energy to remove an electron.
Trends in Ionization Energy Ionization Energy decreases as you move top to bottom on the periodic table. The Reason: As you move from top to bottom in a group: Atomic radius increases Nuclear charge increases Shielding Effect increases As a result, valence electrons are farther from the nucleus and more loosely held. This means it requires less energy to remove an electron.
Electronegativity of Metals Metal Atoms Compared to non-metals, metals have: Larger atomic radii Weaker nuclear charge As a result, metal atoms tend to lose their valence electrons when reaching a stable configuration. Thus, metals tend to have lower electronegativities. H 2.1 Li 1.0 Na 0.9 K 0.8 Be 1.5 Mg 1.2 Ca 1.0 B 2.0 Al 1.5 Ga 1.6 C 2.5 Si 1.8 Ge 1.8 N 3.0 P 2.1 As 2.0 O 3.5 S 2.5 Se 2.4 F 4.0 Cl 3.0 Br 2.8
Electronegativity of Non-Metals Non-Metal Atoms Compared with metal atoms, non-metals have Smaller atomic radii Stronger nuclear charge As a result non-metals tend to gain valence electrons when reaching a stable configuration. Thus, non-metals tend to have higher electronegativities. H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 K Ca Ga Ge As Se Br 0.8 1.0 1.6 1.8 2.0 2.4 2.8