PreLab Experiment 06

PreLab Experiment 06 - because of full orbitals because of...

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Andrew T. Hedrick Chem 32 Section S Experiment 6 – The Periodic Behavior of Metals Prelaboratory Questions 1. Describe how each of the periodic properties of metals given in Table 1 varies  according to the metal’s position in the Periodic Table. For example, covalent  radius increases from top to bottom in a given Group. 1 st  Ionization potential increases at the end of a suborbital. For example, Zinc has  a relatively high 1 st  ionization energy in this table because it lies at the end of a  full d orbital and would not like to form an ion. Conversely, Aluminum has the  lowest 1 st  ionization energy because it comes at the beginning of a p orbital and  would like to lose that electron to instead have a full s orbital. Basically, ionization  energy (energy needed to remove the 1 st  electron) increases as you get closer 
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Unformatted text preview: because of full orbitals because of full s orbitals in the former case and all ½ full p orbitals in the latter. Generally a group closer to fluorine will have, on the whole, a higher electronegativity than the group to the left of it. Also (but less a factor in magnitude) an element in one of the upper periods but in the same group will have a higher electronegativity. Generally covalent atomic radius decreases as you go up in period and as you get closer to the right, for these reasons: elements have fewer energy levels as you go up and as you go left there are more protons to pull in the electrons. Generally ionic radius depends on charge. The more positive a charge is the smaller the ionic radius, because there are more protons than electrons and the remaining electrons are pulled closer (Z eff is larger) to the nucleus....
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