1. The Periodic Table..unlocking it’s mysteries and making it useful
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2. Why is the periodic law a law and not a theory?
The Periodic Law: If elements are arranged in order of increasing atomic mass, a pattern can be seen in which similar properties repeat(recur) on a regular (periodic) basis.
Read pages List the people involved in the development of the periodic table. Answer these questions.
Why is the periodic law a law and not a theory?
Why was Mendeleev’s table given more credit than Meyer’s?

3. The Modern Periodic Table
The 2 most important things to know about the periodic table:
Elements in the same chemical family have similar chemical properties because they have similar outer electron configurations. (This is among the most important of all chemical concepts)
As we move down a chemical family or across a period, there are regular changes in these properties called periodic trends.

7. The characteristics of an element are mostly determined by it’s valence electrons. These are the outermost electrons of the element and are the ones that are involved in bonding.
All elements are most stable when the achieve a full 8 electrons in the s and p orbitals of an energy level. Elements, such as the noble gases, have this “stable octet” arrangement and are very unreactive (“happy” as they are). This is also the arrangement “desired” by many ions.

8. electron configuration Group 2 - Alkaline metals
Relate the behaviour of the families (vertical groups) with their electron configurations…
Family
electron configuration
chemical properties
Why?
Group 1 - Alkali metals
Group 2 - Alkaline metals
Groups 17 - Halogens
Group 18 - Noble gases
Transition elements

9. Trends in the Periodic Table
As you move across a period or down a family, there are regular changes in the properties of the elements. These are called periodic trends.
Atomic or Ionic size
Ionization Energy
Electronegativity
All three of these properties, and much of chemistry in general, are due to the electrostatic attraction/repulsion between charged particles.

10. Atomic size - this is the radius of the atom as estimated by the distance of it’s nucleus from the nucleus of another identical atom.
i.e. for non-metals O2 it would be the half of the distance between the 2 nuclei.
for metal such as Na it would be half of the distance between 2 nuclei in the metal crystal lattice

12. As we move down a family, the sizes of atoms increase.
In general..
As we move down a family, the sizes of atoms increase.
As we move across a period, the sizes of atoms decrease.

13. Why these trends in atomic radius?
The number of energy levels present- “effective shielding”
The amount of nuclear charge “felt” by the outer electrons- Zeff (effective nuclear charge)

14. Ionization Energy - this is the minimum energy required to remove an electron from a gaseous (in gas state) atom or ion. This tells us how strongly an atom holds onto it’s outermost electrons. This is an important property because it tells us how easy it is to form a positive ion (cation) and is related to the reactivity of the atom in certain situations.
atom (g) + IE ion e-
The greater the ionization energy (IE), the more tightly the atom holds onto it’s electron.

15. In general, as you move down a family, the IE decreases, and it becomes easier to remove an electron.
As you move across a period, the IE increases, and it becomes harder to remove an electron.

16. Now explain why…….

17. Electronegativity - is defined as the relative ability of a bonded atom to attract shared electrons to itself. Atoms with relatively high electronegativities (EN) tend to pull bonded electrons closer to their nuclei.

18. As you move across a period, electronegativity increases.
As you move down a family, electronegativity decreases.

19. Now explain why…….