The Periodic Table

1. The Periodic Table

2. History of the Periodic Table • Anton Lavoisier wrote a textbook in 1789, and listed the known elements in a table. He grouped the elements based on similar properties. • Johann Wolfgang Dobereiner published a paper in 1829 describing triads of elements where groups of 3 elements could be arranged together based on similar properties.

3. Leopold Gmelin identified ten triads, three groups of 4 elements, and one group of 5 elements in 1843. • August Kekule in 1858 showed relationships between elements and their bonding strength which he called valency. • StannislaoCannizzaro in 1860 established a universal method for measuring atomic mass. • Julius Lothar Meyer in 1864 arranged the 49 known elements by valency. Elements with similar properties had the same valency.

4. John Newlands in 1864-1865 published a table that arranged elements based on increasing atomic mass. His periodic table repeated properties after 8 elements, and he called this periodicity “The Law of Octaves”, similar to musical note periodicity of 8 notes. • Dmitri Ivanovich Mendeleev published his periodic table in 1869. He arranged the elements in rows by increasing mass then creating a new row when the properties repeated.

5. Meyer also created a similar periodic table in 1870. • Mendeleev however supported his table by ignoring discrepancies in mass and leaving gaps for missing elements. • Meyer did not make any predictions that would validate his reasoning for his table.

6. Mendeleev’s Periodic Table • Arranged by increasing atomic mass • Some elements had “switched positions” where the mass placed them in different columns than their properties indicated • He reasoned that the switched elements masses weren’t measured accurately enough and would later be corrected by better measurements.

7. Some gaps left in the table so elements fit into better positions. • These gaps were explained by Mendeleev as elements that hadn’t yet been discovered. • To support his table, Mendeleev made predictions of properties for these missing elements

8. Mendeleev’s Predictions

9. Actual Element Properties

10. The Modern Periodic Table • Further measurements of atomic mass never corrected the misplaced elements. • Discovery of noble gases supported the periodic table, by being able to be placed in a group that connected halogens and alkali metals. • Mass problems became worse . • Antonius Van den Broek in 1911, first proposes that nuclear charge and electrons responsible for element location on the periodic table

11. Henry Moseley tested Bohr’s hypothesis and Van den Broek’s hypothesis, by experimentally measuring the atomic numbers of Al to Au. • Moseley then arranged elements based on increasing atomic number and found the mass misplacement on the periodic table is resolved • Sadly, Moseley was killed in battle in WWI. • Quantum model of the atom is responsible for the current look of the table. The blocks are arranged based on electron configurations.

12. Electron Configurations & The Periodic Table • Sublevels can easily be identified on the periodic table, and are called blocks. • s-block: Groups 1 & 2 (including He) • p-block: Groups 13-18 (excluding He) • d-block: Groups 3-12 • f-block: Two rows at the bottom, Lanthanide & Actinide series. • The period of an element is determined from its electron configuration.

13. This is determined from the principle quantum number for the s-sublevel & p-sublevel. • Ex. Zr [Kr] 5s2 4d2 The 5 indicates the period! • Ex. Pb [Xe] 6s2 4f14 5d10 6p2 6th period • Group 1: The Alkali Metals ns1 • Exception: Hydrogen, though in Group 1, is NOT an alkali metal! • Very reactive, cannot be found alone in nature. • Very soft, can be cut with a knife. • Low density, some float on water. • Silvery metals with low melting points.

14. Group 2: The Alkaline Metals ns2 • Harder, denser, and higher melting points than alkali metals. • Very reactive, not as much as alkali metals, and are also never found alone in nature. • Hydrogen & Helium 1s1 & 1s2 • Though similar configurations to alkali and alkaline metals, they don’t belong because of their unique properties resulting from the small 1st energy level • Hydrogen is not a metal, and has unique properties that prevent it from being in any group.

15. Helium has a configuration of a filled energy level, which gives it extremely stable and unreactive properties, and so is similar to the noble gases. • Transition Metals, Groups 3-12 ns2 (n−1)d(1−10) • Have electron configurations containing d-sublevels filling and are therefore in the d-block. • The d-sublevel filling has a quantum number 1 less than the period where the element is found. • Some elements have configurations different from the aufbau principle due to shifts in stability. • Transition elements are metals that are less reactive than alkali & alkaline metals, and exhibit a larger range of properties.

16. The Halogens, Group 17 ns2 np5 • Nonmetals that are very reactive • Nearly full s- & p-sublevels • The Noble Gases: Group 18 ns2 np6 & He 1s2 • Nonmetals that are not reactive • Full s- & p-sublevels completing an octet, within a period (n). This configuration is known to be extremely stable. • The Main-Group Elements • s-block & p-block of elements. (Groups 1,2 & 13-18) These two show the periodic pattern quickly and with common elements

17. f-block of elements ns2 (n−1)d10 (n−2)f(1−14) • Filling f-sublevel of energy level 2 less than the period number the element is located in • Wedged between groups 3 & 4 in the d-block • A.K.A. “The Rare Earth Metals” • A.K.A. “Innertransition Metals” • Lanthanide Series: Elements 58-71 • Shiny, dense, hard metals with similar reactivity to alkalines • Actinide Series: Elements 90-103 • All are radioactive • First 4, Thorium (Th) to Neptunium (Np) have been found naturally on Earth • The rest are all artificially created in laboratories

18. Periodic Trends • Periodic Law: • The physical & chemical properties of the elements are periodic functions of their atomic numbers. • Element Properties & Their Periodic Trends:* *The trends will be best demonstrated by the main group elements. The trends, though true for both the d-block & f-block, are not as apparent as they are with the main group.

19. Periodic & Group Trends • There are trends for all properties in both Periodic (Horizontal) rows & Group (Vertical) rows • Periodic trends in these notes will refer to the trend as you move left to right across the table. • Group trends in these notes will refer to the trend as you move from top to bottom on the table. • The trends would be the opposite if you move in reverse of these assigned motions.

20. Atomic Radii • I.O.W. The size of the atoms • Atomic Radius is ½ the distance between two identical atoms (same element) that are bonded together. • Periodic Trend: Decreases • Group Trend: Increases • P Trend Reason: Stronger nuclear charge pulls electron energy levels in closer to the nucleus. • G Trend Reason: Electrons are placed into energy levels further away from the nucleus

21. Ionization Energy • Ionization Energy is the energy required to remove one electron from a neutral atom. • A.K.A. The 1st Ionization Energy. The 2nd IE would be energy to remove 2nd electron, etc. • Periodic Trend: Increases • Group Trend: Decreases • P Trend Reason: Increasing Nuclear Charge-holds the electrons more tightly • G Trend Reason: Electrons located further away from the nucleus, and more easily lost

22. Electron Affinity • It is essentially the opposite of Ionization Energy • Electron Affinity is the energy change that occurs when an electron is acquired by a neutral atom. • Periodic Trend: Decreases • Group Trend: Increases • P Trend Reason: Nucleus is closer to the outer edge of the atom and requires less energy to bring new electrons in to the atom • G Trend Reason: Electrons added further away from the nucleus, so requires more energy to bring new electrons in to the atom

23. Ionic Radii • The size of ions • Slight difference between metals & nonmetals • Metals form Cations (+ charge ions) • Nonmetals form Anions (− charge ions) • Period Trend: Decrease* • Group Trend: Increase • P Trend Reason: Stronger attraction from nucleus • G Trend Reason: Electrons further away from nucleus *Change when switching from metals to nonmetals the ions suddenly swell, then decrease again

24. Electronegativity • Property of the ability of atoms to attract electrons to their nuclei. • Period Trend: Increases* • Group Trend: Decreases* • P Trend Reason: Nucleus is closer to outer edge of the electron cloud • G Trend Reason: Electron cloud is further away from the nucleus *Not including the Noble Gases!!! • Most electronegative element is Fluorine, the least electronegative element is Francium

25. Reactivity • General property, typically defined by the amount of energy required to separate the element from compounds. • Metals: Periodic Trend: Decreases Group Trend: Increases • Nonmetals: Periodic Trend: Increases Group Trend: Decreases • P Trend Reason: The distance from nucleus to electrons • G Trend Reason: The distance from nucleus to electrons Remember-Metals want to lose electrons, Nonmetals want to gain electrons!

26. Valence Electrons • The outermost electrons of an atom are the valence electrons. • Main Group Elements:

27. Electron Configurations of Ions • Just like electron configurations of neutral atoms. • Main Group elements gain or lose electrons to have electron configurations that resemble the nearest noble gas elements. Ex. #1: O−2 1s2 2s2 2p6 Ne 1s2 2s2 2p6 Ex. #2: K+ 1s2 2s2 2p6 3s2 3p6Ar 1s2 2s2 2p6 3s2 3p6 • The Transition Metals are a bit more complicated. Many can lose multiple amounts of electrons to reach differing levels of stability depending on other variables the atoms experience.

28. Example of Transition Ions: Fe 1s2 2s2 2p6 3s2 3p6 4s2 3d6 Fe+2 1s2 2s2 2p6 3s2 3p6 3d6 Fe+3 1s2 2s2 2p6 3s2 3p6 3d5 Fe [Ar] Fe+2 [Ar] 2 4s electrons lost Fe+3 [Ar] 2 4s & 1 3d electrons lost