Periodic Trends

1. Periodic Trends

2. Chapter 5 Mantra Electron Configuration determines properties. Electron Configuration is a predictable, repeating pattern, So Properties are a predictable, repeating pattern. All property trends will be examined in the horizontal (period) and vertical (group) according to changes in Effective Nuclear Charge. We will look at 5 periodic trends: Atomic Radius Ionic Radius Ionization Energy Electron Affinity Electronegativity

3. Atomic Radius Atomic Radius can change under different conditions, so to be standard, it is defined as half the distance between the nuclei of two identical atoms bonded together. Horizontal Trend - Atoms tend to be smaller the farther to the right they are found across a period. Reason – Moving right, atoms in a period have the same amount of shielding, but increasing protons increases effective nuclear charge which pulls the electrons in closer. Vertical Trend - Atoms tend to be larger the farther down in a group they are found. Reason – New shells are added moving down, which decreases effective nuclear charge, letting electrons move farther out.

4. Ionic Radius Ions come in two forms, cations and anions Cations – are positively charged because the have lost electrons. Anions – are negatively charged because the have gained electrons. Gaining are losing electrons makes the ion’s radius different from the neutral atom’s radius.

5. Ionic Radius Trends in ionic radii are similar to those of atomic radii. Vertical Trend – Ions tend to be larger the farther down in a group they are found. Reason – New shells are added moving down, which decreases effective nuclear charge, letting electrons move farther out. Horizontal Trend – Ions tend to be smaller the farther to the right they are found across a period. However, there is a jump in size when anions start to form. Reason – Moving right, there is increasing effective nuclear charge which draws the electrons in closer. The jump in size for anions is caused by the gaining of electrons to form the anions. Note: No radii listed for noble gases and other elements which do not form ions.

7. Notice the same trend as Atomic Radius, but with the jump at anions.

8. Note how radii decrease moving to the right and the jump cause by switching to anions.

9. Of the elements Mg, Cl, Na, and P, which has the largest atomic radius? Explain your answer in terms of trends of the periodic table. • Of the elements Be, Mg, Ca, and Sr, which has the largerst atomic radius? Explain your answer in terms of trends of the periodic table. • Of the ions Na+, Mg2+, P3-, S2-, and Cl-, which has the largerst ionic radius? Explain your answer in terms of trends of the periodic table. • Why is there not ionic radius listed for any noble gas? • What causes the jump in ionic radius between elements with atomic number 31 and 34 in period 4?

10. The trends we will look at today all have to do with atoms becoming ions, that is gaining or losing electrons. So lets review about an ion.

11. Ionization Energy Any process that results in the formation of an ion is referred to as ionization. The energy required to remove one electron from a neutral atom of an element is the ionization energy, IE. In other words, Ionization Energy is the energy required to make a cation. Low Ionization Energy means the atoms tend to form cations.

12. Ionization Energy Horizontal Trend – In general, ionization energies of the elements increase going right across each period. Reason – Increasing nuclear charge increases the attraction for electrons making them more difficult to remove. Also, atoms on the right side of the table tend to form anions, while ionization energy is making cations. The higher ionization energy shows that cations of these elements are unstable. Vertical Trend – In general, ionization energies of the elements decrease going down a group. Reason – Decreasing effective nuclear charge makes the electrons easier to remove. Noble Gases have high Ionization Energies since they are stable and would not want to lose an electron.

13. Note the spike in ionization energy for Noble Gases

14. Ionization Energies for more than 1 electron The energy to remove additional electrons can also be measured. A large jump between ionization energies would indicate which cation is most stable. Large jumps in energy occur after the most stable ion. Li+, Be2+, and B3+ are most stable.

15. Electron Affinity Electron Affinity is the energy released after and electron is added to a neutral atom. Since it is the energy released, the sign of Electron Affinity is negative. The more negative, the greater the Electron Affinity and the more easily the atoms makes an anion. It is basically the opposite of Ionization Energy. Horizontal Trend - Electron affinity generally increases right across periods. • Reason - Increasing nuclear charge attracts extra electrons making anions easily • Remember, the elements on the right form anions. Vertical Trend – Electron Affinity general decreases down groups. Reason – Lower effective nuclear charge does not attract extra electrons very much. Exceptions - Noble Gases have low Electron Affinities since they are stable and would not want to gain an electron.

16. Electronegativity Electrongatvity is ability of an atom in a compound to attract electrons from surrounding atoms. It is similar to Electron Affinity and describes if an atom will “share” electrons equally or not in a bond. Elements with a high Electron negativity attract electrons from other atoms. If two atoms with very unequal Electronegativities are bonded together, there will be unequal “sharing” The trend is the same for Electron Affinity with F having the highest electronegativity.

17. Miniquiz • Which will have the highest ionization energy, Li, Be, C or F? • Use the Ionization energies below to determine the most stable ion for Al. IE1 578 IE2 1817 IE3 2745 IE4 11578 IE5 14831 • Do elements on the right of the periodic table have higher Electron Affinities? • Electronegativity increases as you move towards which element? • Will each of these pairs of atoms share equally or not? • C-Si • Na-F • K-Cl