Ionization of Transition Metals

1. Ionization of Transition Metals • K • Ca • Sr • Ti • V • Cr • Mn • Fe

2. Oxidation Numbers: Examples • CaBr2 • CO • CO2 • Mg3N2 • P4O10 • (NH4)2S • BeF2 • SO2 • H2O • CH4 • NH4Cl • NaH • CaH2 • KCl • RbNO3 • SrSO4

3. CHAPTER 7 Chemical Bonding

4. Chemical Bonds • Attractive forces that hold atoms together in compounds are called chemical bonds. • There are two main types of chemical bonds • Ionic bonds – resulting from electrostatic attraction between cations and anions • Covalent bonds – resulting from sharing of one or more electron pairs between two atoms

5. Lewis Dot Formulas • Valence electrons • Electrons which are involved in chemical bonding • These are usually the outermost electrons • These electrons are most important chemically • Schematic representations of valence electrons in atoms

6. Lewis Dot Formulas: Single Atoms • We show only electrons in the outermost occupied shell • An electron pair is represented as a pair of dots • An unpaired electron is represented as a single dot Li F Na Cl

7. Formation of Ionic Compounds • Consider reaction between metallic sodium and gaseous chlorine • Electron configurations of the elements Na Cl • Sodium atom has low ionization energy and easily looses the only 3s electron forming a cation of Na+ • Chlorine atom has highly negative electron affinity and readily gains an electron becoming an anion of Cl–

8. Formation of Ionic Compounds Na – e– Na+ Cl + e– Cl– Na + Cl  Na+ + Cl– • Na+ cations and Cl– anions are electrostatically attracted to each other resulting in an extended ionic lattice • The high energy of the lattice overcomes all other factors involved in the formation of NaCl from elemental sodium and chlorine

9. Formation of Ionic Compounds Na + Cl  Na+ Cl– • We can write this equation using Lewis dot formulas • The complete equation is 2 Na + Cl2 2 Na+ Cl– • It can also be written as

10. Alkali Metals + Halogens • Reaction with halogens leading to the formation of ionic halides M+X– is a general chemical property of alkali metals • It is also a general chemical property of halogens 2M(s) + X2 2MX(s)

11. Ionic Bonding • Electrostatic interaction • Non-directional • The central ion attempts to maximize the number of interactions with the ions of opposite charge • Formation of an ionic compound involves loss of electrons by metal (oxidation) and gain of electrons by nonmetal (reduction)

12. Alkali Earth Metals + Halogens Ca + F2 • The remainder of the IIA metals and VIIA nonmetals react similarly: M(s) + X2 M2+(X–)2 (s)

13. Formation of Ionic Compounds Li + O2 • The remainder of the IA metals and VIA nonmetals react similarly: 4M(s) + O2(g)  2(M+2O2–) (s) 2M(s) + X(s)  M+2X2– (s) X = S, Se, Te, Po

14. Formation of Ionic Compounds Mg + O2 • The remainder of the IIA metals and VIA nonmetals react similarly: 2M(s) + O2(g)  2(M2+O2–) (s) M(s) + X(s)  M2+X2– (s) X = S, Se, Te, Po

15. Example Write the reaction between calcium and nitrogen. Show what happens to valence electrons using Lewis dot formulas. Learn Table 7-2

16. Covalent Bonding • If the difference in electronegativity of two elements is not large enough, an electron cannot be transferred completely from one atom to the other • It becomes shared between both atoms and a covalent bond is formed

17. Formation of H2 Molecule • When two H atoms are indefinitely far from each other, they do not interact • If the separation decreases to a certain distance, the 1s electron of each H atom is attracted by the nucleus of the other H atom, as well as by its own nucleus • If electrons from different atoms can occupy the same orbital, they will form a covalent bond

18. H2 Molecule • We can use Lewis dot formulas to show covalent bond formation • The covalently bonded atoms are held at a distance corresponding to the lowest total energy

19. Covalent Bond • We say that the covalent bond is formed by the overlap of atomic orbitals • The covalently bonded atoms are held together by a pair of shared electons • The distance between their nuclei corresponds to the lowest total energy • Below this equilibrium distance the nucleus-nucleus and electron-electron repulsions become too large, pushing the nuclei back to the equilibrium distance

20. HF Molecule • F is more electronegative than H • In this molecule the electron pair will be shifted towards the F atom

21. F2 Molecule

22. H2O Molecule

23. NH3 Molecule

24. NH4+ Ion • Lewis formulas can also be drawn for polyatomic ions

25. Bonding & Nonbonding Electrons • Representative elements usually attain stable noble gas electron configurations in most of their compounds • Electrons which are shared among two atoms are called bonding electrons • Unshared electrons are called lone pairs or nonbonding electrons • Lewis dot formulas are based on the octet rule

26. The Octet Rule S = N - A • S = total number of electrons shared in bonds • N = total number of electrons needed to achieve a noble gas configuration • 8 for representative elements • 2 for H atoms • A = total number of electrons available in valence shells of the atoms • A is equal to the periodic group number for each element • A-S = number of electrons in lone pairs

27. Examples • F2 • H2O • CH4 • CO2

28. Covalent Bonding • Covalent bonds are formed when atoms share electrons • If the atoms share 2 electrons a single covalent bond is formed • If the atoms share 4 electrons a double covalent bond is formed • If the atoms share 6 electrons a triple covalent bond is formed

29. The Octet Rule • For ions we must adjust the number of electrons available, A: • Add one e- to A for each negative charge • Subtract one e- from A for each positive charge • Example: NH4+

30. Assignments & Reminders • Go through the lecture notes • Read Sections 7-1 through 7-5 • Read Sections 4-5 & 4-6 of Chapter 4 • Homework #3 due by Oct. 10 • Monday (10/10) and Tuesday (10/11) – lecture quiz #3 based on Chapters 5&6