The History of the Modern Periodic Table

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 Model of triads 1780 - 1849 In 1829, he classified some elements into groups of three, which he called triads. The elements in a triad had similar chemical properties and orderly physical properties. (ex. Cl, Br, I and Ca, Sr, Ba) Model of triads 1780 - 1849

John Newlands Law of Octaves 1838 - 1898 In 1863, he suggested that elements be arranged in “octaves” because he noticed (after arranging the elements in order of increasing atomic mass) that certain properties repeated every 8th element. Law of Octaves 1838 - 1898

John Newlands 1838 - 1898 Law of Octaves Newlands' claim to see a repeating pattern was met with savage ridicule on its announcement. His classification of the elements, he was told, was as arbitrary as putting them in alphabetical order and his paper was rejected for publication by the Chemical Society. 1838 - 1898 Law of Octaves

John Newlands 1838 - 1898 Law of Octaves His law of octaves failed beyond the element calcium. WHY? Would his law of octaves work today with the first 20 elements? 1838 - 1898 Law of Octaves

Dmitri Mendeleev In 1869 he published a table of the elements organized by increasing atomic mass. 1834 - 1907

Lothar Meyer At the same time, he published his own table of the elements organized by increasing atomic mass. 1830 - 1895

Elements known at this time

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, or both?

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. (He corrected the atomic masses of Be, In, and U) was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown.

After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleev’s predictions for Sc, Ga, and Ge were amazingly close to the actual values, his table was generally accepted.

However, in spite of Mendeleev’s great achievement, problems arose when new elements were discovered and more accurate atomic weights determined. By looking at our modern periodic table, can you identify what problems might have caused chemists a headache? Ar and K Co and Ni Te and I Th and Pa

Henry Moseley In 1913, through his work with X-rays, he determined the actual nuclear charge (atomic number) of the elements*. He rearranged the elements in order of increasing atomic number. *“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.” 1887 - 1915

Henry Moseley His research was halted when the British government sent him to serve as a foot soldier in WWI. He was killed in the fighting in Gallipoli by a sniper’s bullet, at the age of 28. Because of this loss, the British government later restricted its scientists to noncombatant duties during WWII.

Glenn T. Seaborg After co-discovering 10 new elements, in 1944 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. 1912 - 1999

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." 1912 - 1999

Periodic Table Geography

The horizontal rows of the periodic table are called PERIODS.

The elements in any group of the periodic table have similar physical and chemical properties! 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.

Alkali Metals

Alkaline Earth Metals

Transition Metals

InnerTransition Metals These elements are also called the rare-earth elements. InnerTransition Metals

Halogens

Noble Gases

The s and p block elements are called REPRESENTATIVE ELEMENTS.

The periodic table is the most important tool in the chemist’s toolbox!

Periodic Trends

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 Metallic vs Non-metallic

Nuclear Charge Defined as the number of protons in the nucleus 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 The shielding effect describes the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. Shielding Effect

Shielding Effect The inner electrons shield the outer electrons from the electrostatic attraction of the protons in the nucleus.

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 remains constant 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.

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.

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. 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

3 Rules for Ionic Radii Cations are ALWAYS SMALLER in size than the neutral atom. Anions are ALWAYS LARGER in size than the neutral atom. Negative anions are ALWAYS LARGER than positive cation

Ionic Radius of Cations Cations are ALWAYS SMALLER in size than the neutral atom. 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

Anions are ALWAYS LARGER in size than the neutral atom. Ionic Radius of Anions Anions are ALWAYS LARGER in size than the neutral atom. 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 ion, O-2 Oxygen

Ionization Energy Ionization energy – the amount of energy needed to remove an electron from an atom. (In order to make an ion) First ionization energy – energy needed to remove one valence electron. Second ionization energy- energy needed to remove a second valence electron. Third ionization energy- energy needed to remove the third valence electron. Etc…..

Successive Ionization Energy Group 1 Metals Have 1 valence electron. It requires a small amount of energy to remove that single electron. (Low 1st ionization energy) + 520 kJ/mol In order to remove a second electron, it requires removing an electron from a stable (noble) configuration. (High 2nd ionization energy) + 7,297 kJ/mol

Successive Ionization Energy 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

Ionization Energy in a Group Ionization Energy decreases as you move top to bottom on the periodic table.

As you move from top to bottom in a group: Trends in Ionization Energy 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.

Ionization Energy in a Period Ionization Energy increases as you move left to right across the periodic table.

Reactivity of Metals Metals tend to have lower electronegativities. 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. 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 Si 1.8 P S Cl K 0.8 Ca Ga 1.6 Ge As Se 2.4 Br 2.8

Reactivity Reactivity is different for metals and non metals Down a family metals increase in reactivity

Reactivity Reactivity is different for metals and non metals Across a period metals decrease in reactivity

Reactivity Reactivity is different for metals and non metals Across a period nonmetals increase in reactivity

Reactivity Reactivity is different for metals and non metals Down a group nonmetals increase in reactivity

Reactivity