Chapter 3 – Periodicity

1. Chapter 3 – Periodicity Lindsey Reeder 2011-2012

2. Periodic Trends/Physical Properties Sections 3.1-3.2

3. Main Trends Used To Explain Periodicity • Effective Nuclear Charge • Atomic Radii • Ionic Radii • Ionization Energy • Electronegativity • Melting Point • Reactivity

4. Effective Nuclear Charge • Nuclear Charge – determined by atomic number, increases as one proton and one electron are added atoms of successive elements • Numerically, nuclear charge = (number of protons)-(number of inner electrons) • Outer electrons do not experience full attraction of protons in the nucleus, due to shielding and repulsion from inner electrons. • Effective charge of outer electrons < Full nuclear charge of atom • As you move from left to right (on the Periodic Table), effective nuclear charge increases, because one proton and one electron are added, but the number of inner electrons remains the same. • As you move from top to bottom, effective nuclear charge remains approximately the same, even though there is an additional energy level between the outer electrons and the nucleus.

5. Atomic Radii • Atomic radius – half the distance between neighboring nuclei, or the distance from the nucleus to the outermost electrons of the Bohr atom • As you move from left to right, the attraction between the nucleus and the outer electrons increases, so there is a decrease in atomic radii. • As you move from top to bottom, atomic radii increases as the number of occupied electron shells increases.

6. Ionic Radii • Cations are smaller than their parent atoms, because they lose their outermost electron shells. (Ex: Na+ vs. Na) • Anions are larger than their parent atoms, because they gain electrons in their outermost electron shells, increasing electron-electron repulsion and thus the radius of the outer shell. • As you move from left to right, ionic radii decrease, then increase, then decrease. The jump in values is due to the transition from cations to anions. Within the cations (groups 1-4) and the anions (groups 4-7), ionic radii decreases the same way that atomic radii decreases. • It’s best to ignore the jump, so as to avoid confusion. Just note that anions are larger than cations, but both decrease in ionic radii from left to right. • As you move from top to bottom, ionic radii increases the same way that atomic radii increases.

7. Ionization Energy • Ionization energy – the energy required to remove one valence electron • As you move from left to right, ionization energy increases, as atoms are less prone to “giving up” electrons. • Atoms on the right side of the Periodic Table are (almost) always negative ions. Therefore, unless an immense amount of energy is supplied, it is extremely rare for such atoms to lose electrons. • As you move from top to bottom, ionization energy decreases, because additional energy levels decrease the attraction between valence electrons and the nucleus, thus making it easier for valence electrons to be removed.

8. Successive Ionization Energies • The trends in the previous slide explain first ionization energy. • However, it should be noted that, once one valence electron is removed, it takes more energy to remove the second valence electron, and even more to remove the third, and so on. • After all electrons in the outermost energy level have been remove, there is a large “jump” in ionization energy, because electrons in the new outermost energy level are closer to the nucleus, and thus experience a greater pull from the nucleus.

9. Anomalies in First Ionization Energies • B < Be and Al < Mg – B and Al each have 1 electron in their p orbitals, whereas Be and Mg have full s orbitals. • The shielding of the s orbitals make it easier to remove a p orbital electron, and atoms favor full orbitals as opposed to partially filled orbitals. • The full s orbitals of Be and Mg act like inner energy levels, in that it is more difficult to remove an s orbital electron if that s orbital is full. • O < N, S < P, Se < As, and Te < Sb – Atoms tend to favor half-filled orbitals, because, according to Hund’s Rule, all electrons in a half-filled orbital spin in the same direction. • Thus, if an atom has 4 electrons in its p orbital (O, S, Se, and Te), it would want to get rid of one of these electrons to obtain this half-filled p orbital. • Once it has a half-filled p orbital, it would require more energy to remove another electron. • This principle can also be applied to the d orbital.

10. Electronegativity • In essence, electronegativity is the same concept as ionization energy, just phrased differently. • Electronegativity – the tendency of an atom (or molecule) to attract an electron towards itself • As you move from left to right, electronegativity increases, as atoms are more likely to want electrons (less likely to give them up). • As you move from top to bottom, electronegativity decreases, for the same reasons as ionization energy. • Additional energy levels reduce the attraction between valence electrons and the nucleus, so adding another valence electron is not terribly important to these atoms. • The more negative the numerical value, the “higher” the electronegativity.

11. Melting Point • As you move down Group 1, melting point decreases, because elements possess metallic structures held together by attractive forces between delocalized atoms. These forces decrease with distance (increasing atomic radii). • As you move down Group 7, melting point increases, because the molecular structures of these elements are held together by van der Waals’ intermolecular forces, which become stronger as the number of electrons in the molecule increases. • Generally, melting point increases across a period, reaching a maximum at Group 4, then decreases to reach a minimum at Group 0/8 (noble gases).

12. Reactivity • Reactivity – the ability for a substance to react, that is, to gain or lose a valence electron • The lower the absolute value of the oxidation number of the element, the more reactive it is. • Therefore, both K and Br are reactive, even though they are at opposite ends of the Periodic Table • As you move from top to bottom, reactivity increases for the same reasons as stated to explain the decrease in ionization energies and the increase in electronegativities

13. Chemical Properties Section 3.3

14. Alkali Metals • Alkali metals (Group 1) are very reactive. • They react with water to form a metal oxide and hydrogen gas. • 2Li(s) + 2HOH(l)  2LiOH(aq) + H2(g) • Alkali metals form ionic compounds with non-metals. • When such a reaction occurs between an alkali metal and a halogen (Group 7), the resulting compound is known as a halide. • 2Li(s) + Cl2(g)  2LiCl(s)

15. Halogens • Halogens are very reactive non-metals that form ionic compounds with metals and covalent compounds with other non-metals. • As previously mentioned, halides form when halogens react with alkali metals.

16. Characteristics of Period 3 Oxides • Bonding • Na2O(s), MgO(s), Al2O3(s): giant ionic • SiO2(s) (metalloid): giant covalent – mix between ionic and covalent bonding structure • P4O10(s)/P4O6(s), SO3(l)/SO2(g), Cl2O7(l)/Cl2O(g): molecular covalent • Electrical conductivity (in molten state) • Na2O(s), MgO(s), Al2O3(s): high • SiO2(s) (metalloid): very low • P4O10(s)/P4O6(s), SO3(l)/SO2(g), Cl2O7(l)/Cl2O(g): none • Acid-base character • Na2O(s), MgO(s): basic • Al2O3(s), SiO2(s): amphoteric – showing both acidic and basic properties • P4O10(s)/P4O6(s), SO3(l)/SO2(g), Cl2O7(l)/Cl2O(g): acidic

17. Additional trends across Period 3 (HL) Section 13.1

18. Characteristics of Period 3 Chlorides • Bonding • NaCl(s), MgCl2(s): giant ionic • AlCl3(s)/Al2Cl6(g), SiCl4(l), PCl5(s)/PCl3(l), S2Cl2(l), Cl2(g): molecular covalent • Electrical conductivity (in molten state) • NaCl(s), MgCl2(s): high • AlCl3(s)/Al2Cl6(g): poor • SiCl4(l), PCl5(s)/PCl3(l), S2Cl2(l), Cl2(g): none • Chlorine reacts slowly with water in a reversible reaction to produce a mixture of hydrochloric (HCl(aq)) and chloric(I) (HOCl(aq)) acid • Cl2(aq) + H2O(l)  HCl(aq) + HOCl(aq) • This reaction is know as a disproportionation reaction because chlorine is simultaneously oxidized and reduced.

19. Characteristics of Period 3 Chlorides (cont.) • Acid-base character • NaCl(s): neutral • MgCl2(s): weakly acidic • AlCl3(s)/Al2Cl6(g), SiCl4(l), PCl5(s)/PCl3(l), S2Cl2(l), Cl2(g): acidic • When ionic chlorides are added to water, the water molecules break up their lattice structure when they dissolve. • Positive ions are attracted to the partially charged negative oxygen atom, while negative ions are attracted to the partially charged positive hydrogen atoms in the water molecule. • Ions surrounded by water molecules in this way are said to be hydrated. • Covalent chlorides are broken up, or hydrolysed, when they are added to water, which is similar to hydration.

20. Complexes and Ligands • Aluminum dissociates into ions when added to water (Al3+ and 3Cl-). • The aluminum ion has a high charge density because of its high charge and small ionic radii, and thus attracts the negative oxygen ends of water molecules. • The water molecules form dative covalent (or coordinate) bonds with the aluminum ion, which are bonds using lone pairs of electrons to form covalent bonds. • This results in an octahedral complex ion, [Al(H20)6]3+. • The number of dative covalent bonds from the ligands to the central ion is called the coordination number. • The water molecules surrounding the aluminum ion are known as ligands. • The hydrated aluminum ion is acidic, because it attracts the O–H of the water molecules and releases an H+ ion to form an acidic solution. • Alkali solutions are formed when an OH- is released.

21. Transition Metals (HL) Section 13.2

22. Electron Configuration • Chromium (Cr) and Copper (Cu) have unusual electron configurations, due to the stability of half-filled and filled d sub-levels. • Their 3d orbitals and 4s orbitals are so close together, that their electrons often fluctuate between orbitals and sublevels. • Cr: 3d54s1 instead of 3d44s2 • Cu: 3d104s1 instead of 3d94s2

23. Physical and Chemical Properties • Physical Properties • High electrical and thermal conductivity • High melting point • High malleability (easily beaten into shape) • High tensile strength (can bear heavy loads) • High ductility (can be drawn into wires) • Chemical Properties • Have multiple oxidation numbers • Form a variety of complex ions • Give off color, and form colored compounds • Can act as catalysts

24. Scandium and Zinc • Although they are part of the d-block of elements, scandium and zinc are not transition metals. • Transition elements form one of more ions with a partially filled d sub-level. • Sc becomes Sc3+, losing its two 4s and one 3d electrons. Zn becomes Zn2+, losing its two 4s electrons while maintaining a full 3d sublevel.

25. Other Characteristics of Transition Metals • All transition metals show both the +2 and +3 oxidation states. • The maximum oxidation states of these elements increases in steps of +1 and reaches a maximum at manganese. Thereafter, the maximum oxidation state decreases in steps of -1. • Oxidation states above +3 show covalent character. • Compounds with high oxidation states tend to be oxidizing agents.

26. Color of Transition Metals • Transition metals are colored due to their partially-filled d sub-levels. • When light passes through a solution containing a transition element complex, one 3d electron can be excited to the 4s sub-level. A photon of light is absorbed and light of the complementary color is transmitted. • Complementary color pairs • green-purple • blue-yellow • red-pale blue

27. Color of Transition Metals (cont.) • The color of the complex depends on the following: • The nuclear charge/identity of the central ion. • The charge density of the ligand. • Higher charge density  higher energy of light received and transmitted • The oxidation number of the central ion. • Higher oxidation number  lower energy of light received and transmitted • The shape/geometry of the complex ion.

28. Transition Metals As Catalysts • Heterogeneous catalysis – catalyst in a different state from the reactants • Transition metals provide a surface for the reactant molecules to come together with the correct orientation. • Homogeneous catalysis – catalyst in the same state as the reactants • Transition metal are particularly effective in redox reactions. • Many reactions in the human body need homogeneous catalysts to occur, including Fe2+ in the heme group of hemoglobin and Co3+ in vitamin B12.