1. Periodic Trends

2. Predicting Periodic Trends A number of physical and chemical properties of elements can be predicted from their position in the periodic table. Among these properties are Ionization Energy, Electronegativity, and Atomic/Ionic Radius. These properties all involve the outer shell (valence) electrons as well as the inner shell (shielding) electrons. Electrons closest to the nucleus experience the full nuclear charge and are held most strongly. As the number of electrons between the nucleus and valence electrons increases, the nuclear charge decreases, due to the “shielding” of these inner shell electrons.

3. Predicting Periodic Trends Cont. The charge felt by the valence electrons is called the effective nuclear charge, Z eff. Going down a group increase the value of n, and increases the number of inner shell electrons. This leads to better shielding and a weaker attraction between the nucleus and the outer shell electrons. Going across a period leads to a larger nuclear charge, as the number of protons increases. There is also increase in the number of valence electrons, but electrons in the same shell are poor at shielding each other. Going across a period generally leads to a stronger interaction between the nucleus and valence electrons.

4. Atomic Radius One-half the distance between the two nuclei of two identical atoms. Period trend: decreases as you move across the period (horizontal). Due to the increase in atomic charge (Z) that draws in the electron cloud. Group trend: increases as you move down the group (vertical). As n increases, the sizes of the orbitals increases.

5. Trends in atomic radius in Periods 2 and 3

6. Arrange the following atoms in order of decreasing atomic radius. Na Al P Cl Mg Which is the largest atom in Group 4? Which is the smallest atom in Group 7? Which is the smallest atom in Period 5? Atomic Radius

7. Ionic Radius Ion: an atom that has a positive or negative charge. Atoms lose or gain electrons to become ions. Ionic radius is the radius of an anion or cation. When neutral atoms are ionized, there is a change in their sizes. If the atom forms an anion the size increases, because the nuclear charge (Z) is unchanged but the electron-electron repulsion increases, due to the added electron, enlarging the electron cloud. If the atom forms a cation, the size decreases. The nuclear charge is unchanged and the decreased electron-electron repulsion shrinks the electron cloud.

8. Ionic Radius Cont. Periodic trend: decrease across period –Groups 1-13: decrease across the period (protons pull in the electron cloud) –Group 14: decrease (protons pull in the electron cloud) or increase (added repulsion in electron cloud increases volume) –Group 15: large increase –Groups 16-18: decrease across the period

9. Ionic Radius Cont. Which of the two species is larger? –N 3- or F - –Mg 2+ or Ca 2+ Which of the two species is smaller? –K + or Li + –P 3- or N 3-

10. Ionization Energy Energy required to remove one electron. Related to how “tightly” the electron is held by the nucleus. The higher the IE, the more difficult it is to remove the electron. Period trend: increases across a period. Electrons in the same set of orbitals do not shield each other very well but the nuclear charge increases, making the electrons more difficult to remove. Group trend: decreases down a group. Electrons become increasingly easy to remove since they are at an increasing distance from the nucleus, with increasing numbers of shielding inner electrons.

11. Ionization energies are measured in kJ mol -1 (kilojoules per mole). They vary in size from 381 (which you would consider very low) up to 2370 (which is very high). Ionization Energy

12. The first ionization energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. This is more easily seen in symbol terms. H (g)  H + (g) + e - All elements have a first ionization energy - even atoms which don't form positive ions in test tubes. The reason that helium (1st I.E. = 2370 kJ mol -1 ) doesn't normally form a positive ion is because of the huge amount of energy that would be needed to remove one of its electrons. Ionization Energy

13. Patterns of first Ionization Energies in the Periodic Table: The first 20 elements First ionization energy shows periodicity. That means that it varies in a repetitive way as you move through the Periodic Table. For example, look at the pattern from Li to Ne, and then compare it with the identical pattern from Na to Ar.

14. Ionization Energy Cont. The amount of energy required to remove a second electron from a 1+ ion is called the second ionization energy. The amount of energy required to remove a third electron from a 2+ ion is called the third ionization energy and so on. The energy required for each successive ionization always increases.

15. 1st IE: Be + 900 kJ  Be +1 + e - 2nd IE: Be +1 + 1760 kJ  Be +2 + e - 3rd IE: Be +2 + 14850 kJ  Be +3 + e - Ionization Energy Cont.

16. Electronegativity Measure of the ability of an atom in a chemical compound to attract a bonding pair of electrons. Depends on: the number of protons in the nucleus, the distance from the nucleus, and the amount of screening by inner electrons. –The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to cesium and francium which are the least electronegative at 0.7

17. Electronegativity Period trend: increases across the period Group trend: decreases down the group

18. Electron Affinity Amount of energy released when an atom gains an electron. – Nonmetal + e -  X - + energy Halogens have a very high electron affinity (because they need only one electron to be stable). The unstability of the halogens makes them very reactive. When they become stable, they release energy (exothermic).

19. Electron Affinity Period Trend: increases across a period. Group Trend: decrease down a group. The greater the distance from the nucleus, the less the attraction.