Presentation on theme: "The History of the Modern Periodic Table. During the nineteenth century, chemists began to categorize the elements according to similarities in their."— Presentation transcript:
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 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)
John Newlands Law of Octaves 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.
John Newlands 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.
John Newlands 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?
Dmitri Mendeleev In 1869 he published a table of the elements organized by increasing atomic mass.
Lothar Meyer At the same time, he published his own table of the elements organized by increasing atomic mass.
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?
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. Mendeleev...
After the discovery of these unknown elements between 1874 and 1885, and the fact that Mendeleevs predictions for Sc, Ga, and Ge were amazingly close to the actual values, his table was generally accepted.
However, in spite of Mendeleevs 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.
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 snipers 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
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
The horizontal rows of the periodic table are called PERIODS.
The vertical columns of the periodic table are called GROUPS, or FAMILIES. The elements in any group of the periodic table have similar physical and chemical properties!
Periodic Law When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.
Alkaline Earth Metals
InnerTransition Metals These elements are also called the rare-earth elements.
The s and p block elements are called REPRESENTATIVE ELEMENTS.
The periodic table is the most important tool in the chemists toolbox!
Nuclear Charge Shielding Effect Atomic size (radius) Ionic size (radius) Ionization energy Electronegativity Metallic vs Non-metallic We will be looking at the following periodic trends in this chapter: Periodic Trends & Properties
Nuclear Charge 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. 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:
Nuclear Charge The number of protons in the nucleus increases by one as you move from left to right across the table. Nuclear Charge increases from left to right across the periodic table. The Reason: Increasing the number of protons in the nucleus results in an increase in nuclear charge.
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
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: More energy levels translates into more inner electrons and thus more shielding. The elements located in higher periods have more electrons, and therefore require more [higher] energy levels.
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.
Within a period, the atomic radius decreases from left to right across the periodic table. The reason: Atomic Radius 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.
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 The nuclear charge does not change, but the atom loses 1 or more electrons. LithiumLithium ion, Li +1 Cations are ALWAYS SMALLER in size than the neutral atom. The Reason: Each valence electron experiences more electrostatic attraction toward the nucleus. As a result, the ion is smaller than the original atom.
Anions are ALWAYS LARGER in size than the neutral atom. Oxygen Oxygen ion, O -2 The Reason: Ionic Radius of Anions Each valence electron experiences less electrostatic attraction toward the nucleus. The atom gains extra electrons, but the number of protons attracting those electrons remains the same. As a result, the ion is larger than the original atom.
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…..
Group 1 Metals Have 1 valence electron. It requires a small amount of energy to remove that single electron. (Low 1 st ionization energy) Successive Ionization Energy kJ/mol + 7,297 kJ/mol In order to remove a second electron, it requires removing an electron from a stable (noble) configuration. (High 2 nd ionization energy)
Group 2 Metals Have 2 valence electrons. It requires a small amount of energy to remove the first and second electrons. Successive Ionization Energy In order to remove a third electron, it requires removing an electron from a stable (noble) configuration kJ/mol + 1,757 kJ/mol + 14,840 kJ/mol
Ionization Energy in a Group Ionization Energy decreases as you move top to bottom on the periodic table.
Trends in Ionization Energy Atomic radius 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. The Reason: As you move from top to bottom in a group: Nuclear charge increases Shielding Effect increases
Ionization Energy in a Period Ionization Energy increases as you move left to right across the periodic table.
Reactivity of Metals Metal Atoms Compared to non-metals, metals have: 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 0.8 Ca 1.0 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8 As a result, metal atoms tend to lose their valence electrons when reaching a stable configuration. Metals tend to have lower electronegativities. Larger atomic radii Weaker nuclear charge
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