8303117-038-Transition-Metals-1 - Chem Factsheet September...

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Transition Metals 1: Definitions and Properties C hem F actsheet September 2002 Number 38 1 To succeed in this topic you need to: Be able to produce full electronic configurations for atoms and ions; Understand the principle of oxidation numbers; Understand types of bonding and molecular shapes; Have an understanding of atomic structure. After working through this Factsheet you will: Be able to produce full electronic configurations for d-block atoms and common ions; Know the definition of ‘transition metal’; Know the common properties of the transition metals; Be able to write the formulae and represent the shapes of complex ions; Have an understanding of why many transition metal compounds are coloured; Know some common uses of transition metals in industry. Definition and electronic configurations The transition metals are d-block elements, which can form one or more stable ions with partially filled d-orbitals . The d-block elements are situated between the s-block and p-block elements on the periodic table. p-block d-block s-block You need to be able to produce the electron configuration of d-block atoms. You will always be given the atomic number (on the periodic table!) and should use the standard procedure for working out electron configuration (see Factsheet 01 - Atomic Structure). The configurations are given below. Note that the configurations for chromium and copper appear slightly anomolous: for chromium the 3d orbital has half shell stability, so this electron configuration is preferable. for copper, the 3d orbital enjoys full shell stability, hence the given electron configuration is more favourable. The change in chemistry cross the d-block series is very slight, as all have two electrons in their outermost shell (except for chromium and copper). The atomic radius and ionisation energies also vary little across the series as increased nuclear charge is offset by increased shielding of the outer electrons. Ions of d-block elements In ion formation, 4s electrons are lost before 3d electrons even though 4s fills before 3d when assigning electrons. This is because when the 3d orbital does start to fill, the 4s electrons are repelled further away from the nucleus.
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This note was uploaded on 03/08/2011 for the course CHEM 101 taught by Professor Hard during the Spring '11 term at UT Arlington.

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8303117-038-Transition-Metals-1 - Chem Factsheet September...

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