CHAPTER 5 – THE PERIODIC TABLE

• Read introduction page 158

• Early 1800’s German chemist J.W. Dobereiner discovered a triad relationship between elements

• 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

• 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

• 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.

• 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.

Vertical columns – families have the same outer electron configuration – Have similar physical properties and chemistry

H Li Na K Rb Cs1S1 [He] 2S1 [Ne] 3S1 [Ar] 4S1 [Kr] 5S1 [Xe] 6S1

Identify the s, p, d, and f block

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

Group 1A – alkali metals lose one electron

2A – alkaline earth metals lose two electrons

7A – halogens gain one electron

8A – noble gases do not react

• Chemical reaction – competition for electrons

• Electronegativity – a measure of an atom’s ability to compete for electrons

• F has highest electronegativity

• Cs and Fr have lowest

• As we approach F on the chart, electronegativity increases

Three groups of Elements

Metals Non-metals Metalloids

•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

Periodic TrendsAtoms get larger as go down a family group

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.

Ion Size

• The greater the net + charge, the smaller the ion

• The greater the net – charge, the larger the ion

Ionization Energy – the energy needed to remove an electron

Li Li+1 + e-

Ionization energy = 8.64 x 10 -19 J/atom

Electron Affinity – the energy change that occurs when an atom gains an extra electron.

The greater the negative number the greater the electron affinity.

Complete Chapter questions• Pages 188-189 1-23, 25, 26, 29,

30, 33