A German scientist named Johann Dobereiner put forward his law of triads in 1817.

1. A German scientist named Johann Dobereiner put forward his law of triads in 1817. Each of Dobereiner's triads was a group of three elements. The appearance and reactions of the elements in a triad were similar to each other. At this time, scientists had begun to find out the relative atomic masses of the elements. Dobereiner discovered that the relative atomic mass of the middle element in each triad was close to the average of the relative atomic masses of the other two elements. This suggested that atomic mass might be important to arranging the elements.

2. With this idea in mind nearly 50 years later, in 1864, an Englishman, John Newlands, arranged the known elements into seven rows. Newlands' Octave Arrangementfrom Chemical News [1866] The elements were placed in order of increasing atomic mass and arranged so elements with similar chemical properties are in the same group.

3. Since the properties were repeated every eighth element, Newlands referred to his arrangement as the Law of Octaves. Unfortunately for Newland’s, there were a number of problems with his arrangement and it really only worked up through calcium. About five years later, Dmitri Mendeleev a Russian chemist who was unaware of the work of Newlands, designed his own table. After about 18 months of gathering information and arranging element cards, it was finished. Like Newlands, Mendeleev arranged the 63 known elements in order of increasing atomic mass, having elements with similar chemical properties in the same group.

4. Unlike Newlands his table had groups of varying lengths. He also left gaps in the table for elements he believed had not yet been discovered. He even made predictions about the properties of some of these elements.

5. Although his successful predictions allowed many to accept his periodic idea, there were still anomalies that Mendeleev could not explain. One of these is the order of iodine and tellurium. The problem was finally solved in 1913 by a 25-year-old English physicist named Henry Moseley. Moseley showed that the ordering of X-ray spectral lines was dependent upon the ordering of nuclear charge, that is, in order of the atomic number. When the elements were placed in order of increasing atomic number, the anomalies in Mendeleev’s table were eliminated. Moseley’s work gave rise to the modern periodic law:

6. The properties of the elements are a periodic function of their increasing atomic numbers. Tragically for the development of science, Moseley was killed in battle during World War I, only two years later. Moseley’s revised periodic table looked something like this

7. The Modern Periodic Table • seven horizontal rows called periods • 18 vertical columns called groups or families • groups 1 and 2 and groups 13-18 are called representative elements • groups 3-12 are the transition metals

8. Modern Periodic Table cont. • elements in any group have similar physical and chemical properties • properties of elements in periods change from group to group • symbol placed in a square • atomic number above the symbol • atomic mass below the symbol

9. Metallic Character • Metals are found on the left of the table, nonmetals on the right, and metalloids in between • Most metallic element always to the left of the Period, least metallic to the right, and 1 or 2 metalloids are in the middle • Most metallic element always at the bottom of a column, least metallic on the top, and 1 or 2 metalloids are in the middle of columns 4A, 5A, and 6A

10. Metallic Character • Metals • malleable & ductile • shiny, lustrous • conduct heat and electricity • lose electrons in reactions • Metalloids • Also known as semi-metals • Show some metal and some nonmetal properties • Nonmetals • brittle in solid state • dull • electrical and thermal insulators • gain electrons in reactions

11. Other Important Groups to Know • Group IA  alkali metals • Group IIA  alkaline earth metals • Group VIIIA  noble gases • Group VIIA  halogens – “salt formers” • Group VIA  chalcogens • Group VA  Nitrogen group • Group IVA  IVA group • Group IIIA  IIIA group

12. Other Groups • s & p block filled  representative elements • d block filled  transition metals • f block filled  inner transition metals • 4f  lanthanides • 5f  actinides • f elements that are naturally occurring  rare earth elements

13. What are Periodic Trends • also called “atomic trends” – take place at the atomic level • trends are general patterns or tendencies • they are general not definite – there are exceptions • when looking at trends we look for increases & decreases • across  periodic • down  group

14. Effects on the Trends • Nuclear Charge • the “pull” of the nucleus • proportional to the number of protons in an atom • the greater the number of protons, the stronger the nuclear charge (“pull”) • this generally affects periodic trends

15. Effects on Trends Cont. • Shielding - the electron protection from the nuclear “pull” - shield = an energy level of electrons - we are not concerned with single electrons, only energy levels of electrons - these electrons reduce the nuclear pull - affects group trends

16. Effects on Trends Cont. • Stability - where electron arrangement is compared to stable octet (or other special stabilities) - determines if atom gains or loses electrons - can be used to explain anomalies in trends

17. Trend in Atomic Size • Increases down column • valence shell farther from nucleus because of increased shielding • Decreases across period • left to right because of the nuclear “pull” • adding electrons to same valence shell • valence shell held closer because more protons in nucleus • Illustration

18. Trend in Ionization Energy • Minimum energy needed to remove a valence electron from an atom • 1 mole of electrons in the gaseous state (kJ/mol) • The lower the ionization energy, the easier it is to remove the electron • metals have low ionization energies • Ionization Energy decreases down the group • valence electron farther from nucleus • Ionization Energy increases across the period • left to right • harder to remove an electron from the atom because of the increased nuclear “pull” • Exceptions: Group 3, Group 6 (chalcogens)

19. Ionization Energy Cont. • Li + energy  Li+ + e- • 1st ionization = 520 kJ/mol • Li+ + energy  Li+2 + e- • 2nd ionization = 7297 kJ/mol • Li+2 + energy  Li+3 + e- • 3rd ionization = 11,810 kJ/mol • Notice, each successive ionization energy is greater than the preceding one – there is a greater “pull” between the nucleus and the electron and thus more energy is needed to break the attraction. • Examining ionization energies can help you predict what ions the element will form. • easy to remove an electron from Group IA, but difficult to remove a second electron. So group IA metals form ions with a 1+ charge.

20. Electron Affinity • atoms attraction to an electron • it is the energy change that accompanies the addition of an electron to a gaseous atom • “opposite” of ionization energy • Across a Period • electron affinity increases because of increased “pull” • Down a Group • electron affinity decreases because the electrons are shielded from the pull of the nucleus Exceptions: Nitrogen Group & Noble Gases

21. Ionic Size • cations – lose electrons (positively charged) • anions – gain electrons (negatively charged) • elements gain or lose e- to become stable – being like noble gases (filled outer sublevel) IA - +1 VA - -3 IIA - +2 VIA - -2 IIIA - +3 VIIA - -1 IVA – share VIIIA – 0, stable

22. Ionic Size Cont. • GOOD RULE OF THUMB • anions are always larger than their neutral atom • cations are always smaller than neutral atom • Across a Period • cations decrease (I-III) because of greater pull on electrons • anions decrease (V-VII) because of less pull on electrons and repulsion of the electrons • Down a Group • both cations and anions increase size

23. Electronegativity • the ability of an atom to attract electrons when the atom is in a compound • electron “tug-of-war” • very similar to electron affinity • Across a Period • increases because of increased pull • Down a Group • decreases because of shielding • Fluorine – most electronegative element

24. Reactivity • Reactivity of metals increases to the left on the Period and down in the column • follows ease of losing an electron • Reactivity of nonmetals (excluding the noble gases) increases to the right on the Period and up in the column • Reactivity Video

25. Practice • Which element has a greater ionization energy – Mg or Ba • Which element has a greater atomic radius – N or F • Which element has a great electron affinity – S or Pb • REVIEW