1. Periodicity Chemistry 11

2. Periodic Trends in Atomic Size The radius of an atom can not be measured directly. The radius of an atom can not be measured directly. We use the atomic radius which is half of the distance between the nuclei of 2 atoms in a homonuclear (same nucleus) diatomic molecule. We use the atomic radius which is half of the distance between the nuclei of 2 atoms in a homonuclear (same nucleus) diatomic molecule.

3. Bromine

4. Remember… The radius decreases across a period from left to right (in the periodic table) The radius decreases across a period from left to right (in the periodic table) The radius increases as you go down a group. The radius increases as you go down a group.

5. Example of Increasing Atomic Radius: Consider the alkail metals (Group 1). Consider the alkail metals (Group 1). All of these elements have a single s electron outside a filled p sublevel. All of these elements have a single s electron outside a filled p sublevel. Electrons in these inner p levels are much closer to the nucleus than the outer s electrons and effectively “ shield ” the outer s electrons from the positive charge of the nucleus. Electrons in these inner p levels are much closer to the nucleus than the outer s electrons and effectively “ shield ” the outer s electrons from the positive charge of the nucleus.

6. Example of Decreasing Atomic Radius Consider the 3 rd period where electrons are being added to the third principal energy level. Consider the 3 rd period where electrons are being added to the third principal energy level. The added electrons should be relatively “ poor ” shields for each other since they are all at about the same distance from the nucleus. The added electrons should be relatively “ poor ” shields for each other since they are all at about the same distance from the nucleus.

7. Example of Decreasing Atomic Radius : Only the 10 electrons in the inner 1s 2, 2s 2, and 2p 6 sublevels are expected to shield the outer electrons from the nucleus. Only the 10 electrons in the inner 1s 2, 2s 2, and 2p 6 sublevels are expected to shield the outer electrons from the nucleus. This means that the effective nuclear charge should increase moving across the period. This means that the effective nuclear charge should increase moving across the period. As the charge of the nucleus increases, outer electrons are pulled in more tightly, therefore decreasing the atomic radius. As the charge of the nucleus increases, outer electrons are pulled in more tightly, therefore decreasing the atomic radius.

8. Ionic Radius The radii of cations and anions derived from atoms in the main group elements also follow trends. The radii of cations and anions derived from atoms in the main group elements also follow trends. The radii of both cations and anions decrease from left to right across a periodic table. The radii of both cations and anions decrease from left to right across a periodic table.

9. Ionic Radius Positive ions ( cations ) are smaller than the metal atoms from which they are formed. Positive ions ( cations ) are smaller than the metal atoms from which they are formed. Example : Example : Na+ has a radius of 0.95 nm while the Na atom has a radius of 0.186 nm.

10. Ionic Radius Negative ions ( anions ) are larger than the nonmetal atoms from which they are formed. Negative ions ( anions ) are larger than the nonmetal atoms from which they are formed. Example: Example: The radius of Cl ‑ has a radius of 0.181 nm while the chlorine atom has a radius of 0.099 nm.

11. Ionic Radius The cation is smaller than the corresponding metal ion because the excess of protons in the ion draws the outer electrons in closer to the nucleus. The cation is smaller than the corresponding metal ion because the excess of protons in the ion draws the outer electrons in closer to the nucleus. In contrast, an extra electron in an anion adds to the repulsion between outer electrons, making a negative ion larger than the corresponding nonmetal atom. In contrast, an extra electron in an anion adds to the repulsion between outer electrons, making a negative ion larger than the corresponding nonmetal atom.

12. Ionization Energy Is a measure of how difficult it is to remove an electron from a gaseous atom. Is a measure of how difficult it is to remove an electron from a gaseous atom. Energy must always be absorbed to bring about ionization, so ionization energies are always positive quantities. Energy must always be absorbed to bring about ionization, so ionization energies are always positive quantities.

13. Ionization Energy The 1 st ionization energy is the energy change for the removal of the outermost electron from a gaseous atom to form a + 1 ion: The 1 st ionization energy is the energy change for the removal of the outermost electron from a gaseous atom to form a + 1 ion: M (g) → M + (g) + e -  E 1 = first ionization energy The more difficult it is to remove electrons, the larger the ionization energy. The more difficult it is to remove electrons, the larger the ionization energy.

14. Ionization Energies For Main Group Elements The ionization energy increases across the periodic table from left to right. The ionization energy increases across the periodic table from left to right. The ionization energy decreases moving down the periodic table. The ionization energy decreases moving down the periodic table. The smaller the atom (atomic radius), the more tightly the electrons are held to the positively charged nucleus and the more difficult they are to remove. The smaller the atom (atomic radius), the more tightly the electrons are held to the positively charged nucleus and the more difficult they are to remove.

15. Ionization Energies For Main Group Elements Conversely, in a large atom, such as that of a Group 1 metal, the electron is relatively far from the nucleus, so less energy has to be supplied to remove it from the atom. Conversely, in a large atom, such as that of a Group 1 metal, the electron is relatively far from the nucleus, so less energy has to be supplied to remove it from the atom.

16. Ionization Energies For Main Group Elements…Exceptions For example, the ionization energy of B (801 kJ/mol) is less than that of Be (900kJ/mol). For example, the ionization energy of B (801 kJ/mol) is less than that of Be (900kJ/mol). This happens because the electron removed from the boron atom comes from the 2p as opposed to the 2s sublevel for beryllium. This happens because the electron removed from the boron atom comes from the 2p as opposed to the 2s sublevel for beryllium. Since 2p is higher in energy than 2s, it is not too surprising that less energy is required to remove an electron from that sublevel Since 2p is higher in energy than 2s, it is not too surprising that less energy is required to remove an electron from that sublevel

17. Electron Affinity (EA) The energy change resulting from the addition of an electron to a gaseous atom. The energy change resulting from the addition of an electron to a gaseous atom. For example, energy is released when a fluorine atom gains an electron to become an anion. For example, energy is released when a fluorine atom gains an electron to become an anion. EA quantities are therefore, negative values. EA quantities are therefore, negative values.

18. Electron Affinity (EA) F(g) + e- → F- (g) F(g) + e- → F- (g) EA generally increases moving left to right of the periodic table. EA generally increases moving left to right of the periodic table. This is because they become smaller and the nuclear charge increases…WHY? This is because they become smaller and the nuclear charge increases…WHY? Moving down a group, EA generally decreases with increasing atomic size. Moving down a group, EA generally decreases with increasing atomic size.

19. Electronegativity Measures the ability of an atom to attract to itself, the electron pair forming a bond or, to attract electrons to itself when it is chemically combined with another element. Measures the ability of an atom to attract to itself, the electron pair forming a bond or, to attract electrons to itself when it is chemically combined with another element.

20. Electronegativity The least electronegative is cesium, Cs, with a value of 0.7. The least electronegative is cesium, Cs, with a value of 0.7. It is unable to attract electrons and therefore loses electrons to a more electronegative atom, thereby becoming a cation. It is unable to attract electrons and therefore loses electrons to a more electronegative atom, thereby becoming a cation.

21. Electronegativity The most electronegative element is fluorine, F, with a value of 4.0. The most electronegative element is fluorine, F, with a value of 4.0. These units are based on the Pauli Electronegativity Scale. These units are based on the Pauli Electronegativity Scale. Fluorine is very electronegative and attracts (or gains) the bonding electrons of the element it is chemically combining with. Fluorine is very electronegative and attracts (or gains) the bonding electrons of the element it is chemically combining with.

22. Electronegativity…Remember Moving left to right across the periodic table, the electronegativity of the main group elements increases. Moving left to right across the periodic table, the electronegativity of the main group elements increases. The electronegativity decreases as we travel down a group. The electronegativity decreases as we travel down a group.