Periodic Table pt. 2

1. Periodic Table pt. 2

2. Chemical Groups And Periodicity • As scientists first began to discover and classify the elements, patterns and similarities were observed in chemical behaviors of certain groups of elements. • Consider the three metals Li, Na, and K • All 3 metals have similar appearance, melting points, and densities • The most interesting feature is that all 3 metals react with the same elements in a nearly identical manner • As you see in the periodic table, these elements are all listed in the same group, or vertical column.

3. Chemical Groups And Periodicity • Dmitri Mendeleev created the periodic table in in 1869 by arranging the elements from left to right in order of increasing atomic number, and vertically according to their behavior (groups) • In doing so, he observed repetitive patterns in chemical behavior across periods (horizontal rows) • This periodicityis described in the next slide.

4. Chemical Groups And Periodicity 25 Br 18 Ar 19 K 12 Mg 11 Na 26 Kr 20 Be 17 Cl 14 Si 22 Ge Decreasing metallic character Totally unreactive gas Highly reactive, diatomicspecies Less reactive, less conductive metal Slightly conductive semi-metal Highly reactive, highly conductive metal Decreasing metallic character Less reactive, less conductive metal Highly reactive, highly conductive metal Slightly conductive semi-metal Highly reactive, diatomic species Totally unreactive gas

5. WHY? • We must now answer many questions about chemical reactivity. • Why is it that some atoms join together and form molecules, while others can’t? • Why is there such wide variation in the reactivity and physical properties of elements? • Why is there periodic repetition (periodicity) of the chemical/physical properties of elements as we move across the periodic table?

6. Explanation Of Elemental Groups • The existence of periodicity proves a very important point: The number of protons in the nucleus has no effect on chemical behavior. If it were so, chemical behaviors would never repeatgiven that no two elements have the same atomic number. The chemical behavior of an element is dictated by the arrangement of its electrons.

7. Ionization Energy • A direct indication of the arrangement of electrons about a nucleus is given by the ionization energiesof the atom • Ionization energy (IE) is the minimum energy needed to remove an electron (form a cation) completely from a gaseous atom • Ionizations are successive. It becomes increasingly harder to remove additional electrons due to increased attraction to the remaining protons in the nucleus.. 1st Ionization Energy 2nd Ionization Energy IE1 < IE2 < IE3 …….IEn

8. Ionization Energy • By measuring the energy required to remove electrons from an element, you can gain an idea of: • how “willing” an atom is to lose an electron, and relate this to its reactivity • where the electrons are positioned

9. 1st Ionization Energy Shows A Periodic Trend For T very difficult to ionize very easy to ionize

10. Ionization Energy The lower the ionization energy of an element, the more METALLICand REACTIVEit is.

11. What Happens if You Remove ALL of the Electrons? • Let’s take a look at the electron configurations of Lithium (atomic # = 3) and Beryllium (atomic # = 4) Li 3 electrons Be 4 electrons Pair of tightly bound electrons Pair of electrons that are more easily removed Single electron that is easily removed

12. Successive Ionizations Ne 10 electrons Na 11 electrons Same two tightly bound electrons Eight electrons of similar attraction to the nucleus 11th electron enters different “shell”

13. Electrons Reside In “Shells” Of Different Distances From The Nucleus • From these plots, Niels Bohr derived the Bohr model of the atom. In it, electrons reside in shells that orbit at different distances from the nucleus. • Each shell has a finite number of electrons that it can hold • The two electrons closest to the nucleus are the hardest to remove. The furthest electrons are the easiest. • Each shell holds 2n2 electrons, where the n=1 shell is the closest to the nucleus. Na

14. Noble Gas Configurations • Only the outermost electrons are involved in reactions. These are called valence electrons. • The inner-most electrons of an element comprise what is known as a noble gas core. They do not react. • Noble gases are chemically inactive because they have completely filled shells. • Lithium, for example, has a two electron core, which we call a Helium core, and one valence electron. Sodium has a 10-electron, Neon core, and one valence electron; and so on.

15. Electron Configurations • Electrons in a shell exist in different “orbitals”. An orbital is a enclosed volume of space around the nucleus that an electron can move throughout. More detail will be provided in CHEM 105. • There are four types of orbitals: • s orbitals (max capacity: 2 electrons) • p orbitals (max capacity: 6 electrons) • d orbitals (max capacity: 10 electrons) • f orbitals (max capacity: 14 electrons) • S-orbitals exist in all shells. P-orbitals exist in all shells where n>2, d-orbitals exist in all shells n>3, and f-orbitals exist in all shells n>4.

16. Electron Configurations • Electrons in different orbitals have different energies. Energies depend on the value of n. So, a 2s electron has higher energy than a 1s electron. • Orbitals having the same n, but different L (like 3s, 3p, 3d) have different energies. • When we write the electron configuration of an atom, we list the orbitals in order of increasing energy, as described by the diagram on the left.

17. Example • Write the electron configurations of N, Cl, and Ca Energy

18. Noble Gas Configurations • Now that we can assign orbitals to electrons, we can write proper valence electron configurations.

19. ns2 np6 ns1 ns2 np2 ns2 np3 ns2 np4 ns2 np1 ns2 np5 1 2 3 4 5 6 7 ns2 TRANSITION METALS ns2 (n-1)dx