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Chapter 5

Chapter 5. The Periodic Law. Modern Russian Table. Chinese Periodic Table. Orbital Filling Table. Periodic Table with Group Names. 5.1 History of the Periodic Table. Chapter 5. The Periodic Law. Predecessors to the Modern Periodic Table. Dobereiner’s Triads

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Chapter 5

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  1. Chapter 5 The Periodic Law

  2. Modern Russian Table

  3. Chinese Periodic Table

  4. Orbital Filling Table

  5. Periodic Table with Group Names

  6. 5.1 History of the Periodic Table Chapter 5 The Periodic Law

  7. Predecessors to the Modern Periodic Table • Dobereiner’s Triads J.W. Dobereiner classified some elements into groups of three, which he called triads. • similar chemical properties • physical properties varied in an orderly way according to their atomic masses.

  8. Mendeleev’s Periodic Table Dmitri Mendeleev

  9. Predecessors to the Modern Periodic Table • Mendeleev’s Periodic Table Dmitri Mendeleev realized that the chemical and physical properties of the elements repeated in an orderly way when he organized the elements according to increasing atomic mass. • In 1869, Mendeleev published a table of the elements organized by increasing atomic mass. • Mendeleev was a Russian scientist and is often referred to as the “Father” of the Periodic Table. • Periodicity is the tendency to recur at regular intervals.

  10. Mendeleev’s Table"The chemical properties of the elements are a periodic function of their atomic weights"

  11. Vertical columns in atomic weight order Mendeleev's Periodic Table Horizontal rows have similar chemical properties • Mendeleev made some exceptions to place elements in rows with similar properties (tellurium & iodine's places were switched) • Missing Elements: gaps existed in Mendeleev’s table • Mendeleev predicted the properties of the “yet to be discovered” elements (scandium, germanium and gallium)

  12. Problems with Mendeleev’s Table Why didn't some elements fit in order of increasing atomic mass? Moseley helped to clarify some of the problems… Why did elements exhibit periodic behavior?

  13. Henry Moseley 1887-1915 • English physicist who determined the number of positive charges in the nucleus (protons) by measuring the wavelength of the x-rays given off by certain metals in 1913. • He was killed by a sniper in Turkey in August 1915 during WWI. Many people think that Britain lost a future Nobel Prize winner. This is because Nobel Prizes, the most prestigious awards for scientific achievement are awarded only to living people.

  14. Moseley and the Periodic Table • Protons and Atomic Number: X-ray experiments revealed a way to determine the number of protons in the nucleus of an atom The periodic table was found to be in atomic number order, not atomic mass order!!! • This explained tellurium-iodine anomaly • ***Moseley was killed in battle in 1915, during WWI. He was 28 years old The Periodic Law The physical and chemical properties of the elements are periodic functions of their atomic numbers

  15. Discovery of the Noble Gases William Ramsay discovers argon Ramsay discovers krypton and xenon Helium discovered as a component of the sun, based on the emission spectrum of sunlight Ramsay finds helium on Earth Freidrich Dorn discovers radon 1868 1894 1895 1898 1900 Sir William Ramsay

  16. The Lanthanides Early 1900's the elements from cerium (#58) to lutetium (#71) are separated and identified. Also known as the rare earth elements, less than 0.01% naturally occurring. • The Actinides Discovery (or synthesis) of thorium, # 90 to lawrencium #103 • Both groups pulled out of the table for space reasons. • Periodicity: Elements with similar properties are found at regular intervals within the "periodic" table

  17. 5.2 Electron Configuration & The Periodic Table Chapter 5 The Periodic Law

  18. 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 Sublevel Blocks on the Periodic Table

  19. The Properties of a Group: the Alkali Metals • Easily lose 1 valence electron • (Reducing agents) • React violently with water • React with halogens to form salts

  20. The Properties of a Group: the Alkali Earth Metals • Easily loses 2 valence electron • (Reducing agents) • Harder, denser, stronger than Group 1 metals • Higher melting points • Less reactive than Group 1, but too reactive to be found free in nature

  21. 5.3 Electron Configuration & Periodic Properties Chapter 5 The Periodic Law

  22. Determination of Atomic Radius: Half of the distance between nuclei in covalently bonded diatomic molecule "covalent atomic radii" Periodic Trends in Atomic Radius: • Radius decreases across a period Increased effective nuclear charge due to decreased shielding • Radius increases down a group Addition of principal quantum levels

  23. Table of Atomic Radii

  24. How to Achieve an Octet… • Atoms can form ions by gaining or losing electrons to obtain a stable outer configuration • Cation- Positive ion (+) ion • Anion- Negative ion (-) ion • Ions attract (opposites attract)

  25. Predicting Ionization • Metals tend to lose electrons They form cations. Ex: Na, 1s22s22p63s1 becomes Na+1,1s22s22p6 • Nonmetals tend to gain electrons. They form anions. Ex: O, 1s22s22p4 becomes O-2, 1s22s22p6

  26. Electron Transfer: Anions • When an atom gains electrons it increases its negative charge so it becomes negatively charged. There are now more electrons than protons. X + e- = X – Ex: Nitrogen Atom + 7 protons - 7 electrons Nitrogen Ion +7 protons - 10 electrons Neutral -3 charge

  27. Electron Transfer: Cations • When an atom loses electrons, it loses negative charges so it becomes more positively charged. There are now more protons than electrons. X - (e-) = X + Ex: Potassium Atom + 19 protons - 19 electrons Potassium Ion +19 protons -18 electrons Neutral +1 charge

  28. Ionization Energy - the energy required to remove an electron from an atom • Increases for successive electrons taken from • the same atom • Tends to increase across a period Electrons in the same quantum level do not shield as effectively as electrons in inner levels Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove • Tends to decrease down a group Outer electrons are farther from the nucleus

  29. Ionization of Magnesium Mg + 738 kJ  Mg+ + e- Mg+ + 1451 kJ  Mg2+ + e- Mg2+ + 7733 kJ  Mg3+ + e-

  30. Table of 1st Ionization Energies

  31. Another Way to Look at Ionization Energy

  32. Electron Affinity - the energy change associated with the addition of an electron • Affinity tends to increase across a period • Affinity tends to decrease down a group Electrons farther from the nucleus experience less nuclear attraction Some irregularities due to repulsive forces in the relatively small p orbitals

  33. Table of Electron Affinities

  34. Ionic Radii Cations • Positively charged ions • Smaller than the corresponding • atom Anions • Negatively charged ions • Larger than the corresponding • atom

  35. Summation of Periodic Trends

  36. Table of Ion Sizes

  37. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons • Electronegativities tend to increase across • a period • Electronegativities tend to decrease down a • group or remain the same

  38. Periodic Table of Electronegativities

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