Lecture 1 - introduction to inorganic chemistry PDF

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CHEM1020 Inorganic Chemistry Introduction to Inorganic Chemistry – lecture one

The periodic table Organic Chemistry is the chemistry of carbon and its compounds. Inorganic Chemistry is the chemistry of the rest of the Periodic Table. Inorganic Chemistry therefore involves the study of over 100 elements! To understand the chemistry of so many elements, we need to consider trends and relationships that may exist, as reflected in periodicity. The Periodic Table represents our understanding of, and trends in, chemical (and physical) properties. **The arrangement of elements in the Periodic Table also fits our basic model of atomic structure (i.e. electron configuration) **

We identify two main classes of elements, the metals and non-metals. At the borderline between these two groups is found a ‘grey’ area, where both metallic and non-metallic properties are shared, represented by a limited number of elements classed as the metalloids What’s out there – Common elements on Earth:

Atomic structure and the periodic table

All of the alkali metal are thus expected to have similar chemistries, and are found in the same group of the periodic table, with the same number of valence electrons. Similarly, for other groups e.g. Group 15 – General configuration of Group 15 is:

Question: How many valence electrons does chlorine have? CL – check periodic table – Z = 17

Electron configuration:

S BLOCK - Groups of elements in columns of the periodic table may -

show related chemistry This is particularly true for elements at the left-hand side of the periodic table

- Groups 1 and 2 are called the alkali and alkaline earth -

elements Together, the two groups are referred to as the S-BLOCK

Z = atomic number

Chemistry is not only similar down each group, but next-door neighbours can have something in common. For example, most (but not all) s-block metals react with liquid water to produce hydrogen and a metal hydroxide:

The relationship between the position of an element in the Periodic Table and atomic theory also shows up (as expected) in the metal-water reactions.

P-BLOCK The p-block occupies the right-hand side of the periodic table (groups 13-18, with 3-8 valence electrons)

Trends within periodic table Behaviour and trends of elements within the periodic tale can be related to the number of valence electrons, as indicated earlier. Consider the s- and p- blocks:

The ‘end’ elements of the Periodic Table (Group 18, with 8 valence e- ), the noble gases, have filled shells and (theoretically) are not expected to be reactive – but have some limited chemistry (Xe, for example, forms XeF2, XeF4, XeF6, XeO3 and XeO4, but no hydrides!). Non-metals at the right-hand side of the Periodic Table form (almost) exclusively anions. Note, even non-metals can be thought of formally as a cation (when combined with a very electronegative atom, i.e. F or O). (Called the “ionic limit”!) e.g. in IF7, it is possible to consider the iodine as I7+ (the fluorine is F- ); similarly in Cl2O7, chlorine is formally Cl7+ (O is O2- ).

Atoms as Spheres It is convenient (though not fully correct) to think of atoms as hard spheres. We can calculate (approximately) the radius of each atomic sphere by simply using known covalent bond distances between atoms, although the non-bonded radius (also called the van der Waals radius) is a better choice. Interatomic distances are experimentally measurable – so we feel comfortable in handling this simple model.

ATOM SIZE Leaving out the transition elements (dblock) we see a trend in atomic size across and down the periodic table: 1) Atomic radii increase down a periodic group 2) Atomic radii decrease going across a periodic row

Why do these trends occur? (1) Down a Group The probability of finding an electron further from the nucleus increases with increasing principal quantum number. Thus as the number of electrons increases down a Group, the radii of the atoms increase (2) Across a Row (or Period) The sizes of the s and p orbitals with the same principal quantum number are about the same. Across a row, the electrons in the ns and np orbitals feel the increasing electrostatic attraction from the increasing number of protons in the nucleus, but have poor shielding by other electrons, which leads to a progressively decreasing atomic radius.

Ions as Spheres The concept of (monatomic) ions as spheres (even hard spheres) is an extension of visualising atoms as spheres; it is useful in Chemistry, and we use it heavily. It may even be a good model, as atomic-level pictures of surfaces and surface-adhered compounds do indeed show the ions behaving as spheres.

- Ions (as spheres) show similar trends as atoms, although cations and anions are -

difficult to compare. As with atoms, ions get larger going down a periodic group. Cations with noble gas configurations get smaller going across a periodic row. Anions with noble gas configurations get smaller in the same way.

Note: cations are always smaller than their parent atoms, and anions are always larger than their parent atoms....