1. CHAPTER 5 – THE PERIODIC TABLE Read introduction page 158 Early 1800’s German chemist J.W. Dobereiner discovered a triad relationship between elements

2. 1864 – Newlands discovered the Law of Octaves – repeating pattern of chemical reactivity (every 8 elements) 2 bonus points, find the mistake in the description of Law of Octaves on the top of page 161

3. 1869 – Russian chemist/teacher Dmitri Mendeleev created the first periodic table He arranged the table in order of increasing mass and chemical properties He left blank spaces and predicted the existence of undiscovered elements

4. Moseley discovered atomic number by studying the frequency of x-rays produced when metals were bombarded with high energy electrons. He hypothesized that this was due to a different positive charge in each nucleus.

5. The Periodic Law – When elements are arranged in order of increasing atomic #, their physical and chemical properties show a periodic pattern. Do not memorize lots of facts about elements – instead learn to predict an element’s properties by its position on the periodic chart.

6. Vertical columns – families have the same outer electron configuration –Have similar physical properties and chemistry H Li Na K Rb Cs 1S 1 [He] 2S 1 [Ne] 3S 1 [Ar] 4S 1 [Kr] 5S 1 [Xe] 6S 1 Identify the s, p, d, and f block

7. Horizontal rows are called periods – row number is the energy level for the s and p block elements. d energy level is n – 1 f energy level is n – 2 Can read electron configuration directly from the chart Group1A – alkali metalslose one electron 2A – alkaline earth metals lose two electrons 7A – halogens gain one electron 8A – noble gasesdo not react

8. Chemical reaction – competition for electrons Electronegativity – a measure of an atom’s ability to compete for electrons

9. F has highest electronegativity Cs and Fr have lowest As we approach F on the chart, electronegativity increases

10. Three groups of Elements MetalsNon-metalsMetalloids lose electrons Shinny (luster) Malleable Conduct electricity (gain electrons) No metallic luster Are not malleable Do not conduct electricity B, Si, Ge, As, Sb, Te, At Have properties intermediate of metals and non-metals Are semiconductors

11. Periodic Trends Atoms get larger as go down a family group

12. Why? More energy levels and electrons, higher energy levels are further from the nucleus Atoms get smaller (diameter) as go across a period (row). Why? Greater positive pull on electrons and same amount of shielding by inner electrons.

13. Ion Size The greater the net + charge, the smaller the ion The greater the net – charge, the larger the ion

14. Ionization Energy – the energy needed to remove an electron Li Li +1 + e - Ionization energy = 8.64 x 10 -19 J/atom

15. Electron Affinity – the energy change that occurs when an atom gains an extra electron. The greater the negative number the greater the electron affinity.

16. Complete Chapter questions Pages 188-189 1-23, 25, 26, 29, 30, 33