Chemical Bonding

1. Chemical Bonding Ionic, Metallic, and Covalent

2. What is a chemical bond? A chemical bond is a force that holds two or more atoms together.

3. Electronegativity • The ability of an atom to attract electrons in a chemical bond • A table of electronegativities appears on p. 263 of the textbook. • Differences in electronegativity cause molecules to have different types of bonds

4. Polarity • Polarity occurs when the distribution of electron is NOT uniform throughout a molecule. • The electrons spend more time around one nucleus than the other

5. Types of Bonds • Nonpolar covalent bond – electrons are shared equally • Polar covalent bond – electrons are shared unequally • Ionic bond – electrons are completely transferred and ions are held together by electrical forces

6. How to determine bond type • Look up the electronegativities (p. 263) • Find the difference between the electronegativities. • Use the table to classify the bond type. Note: There is some disagreement on the “cut-off” number. Some sources use 2.0 instead of 1.7.

7. Vocabulary - Ionic Bonds • Ion – an atom that has an electrical charge due to the gain or loss of electrons • Monatomic – contains one (1) atom • Polyatomic – contains 2 or more atoms • Oxidation number – the charge on a monatomic atom

8. Positive Ions • Positive ions form when electrons are lost • Groups 1A, 2A, 3A, and the transition metals form positive ions. • Oxidation numbers: • Group 1A is 1+ • Group 2A is 2+ • Group 3A is 3+

9. Negative Ions • Negative ions form when electrons are gained. • Groups 5A, 6A, and 7A form negative ions. • Oxidation numbers: • Group 5A is 3- • Group 6A is 2- • Group 7A is 1-

10. More vocabulary • Ionic Bond – The electrostatic force that holds oppositely charged particles together. • Ionic Compound – a chemical compound formed by an ionic bond. • Salt – a name for an ionic compound.

11. calcite Properties of Salts 1. very hard – each ion is bonded to several oppositely -charged ions 2. high melting points – many bonds must be broken with sufficient force, like atoms are brought next to each other and repel 3. brittle –

12. Naming Binary Ionic Compounds • 1st ion is the positive ion • Use the element name • 2nd ion is the negative ion • Take the element name, but change the ending to –ide. • Examples: • chlorine becomes ____________ • sulfur becomes _____________

13. Naming Examples • NaF • MgCl2 • Al2O3

14. Writing the Formula for a Binary Ionic Compound • Write the symbols for the ions. Include the oxidation numbers (charges). • Criss-cross the charges. The charge of one ion will become the subscript on the opposite ion. • Reduce the subscripts to a lowest-terms ratio. Note: the formula for an ionic compound gives the number of atoms in one formula unit.

15. Formula Examples • aluminum chloride • magnesium oxide

16. Polyatomic Ions • Polyatomic ions are charged particles that have 2 or more atoms covalently bonded. • A covalent bond is a chemical bond formed by the sharing of electrons. • Polyatomic ions should be treated as a single unit. • Example: SO42- is the sulfate ion.

17. Naming Ionic Compounds containing Polyatomic Ions • Cation – positive ion • Anion – negative ion • The cation: • If it is a polyatomic ion, use the name of the ion as the first word. • If it is a monatomic ion, use the element name.

18. The anion: • If polyatomic, use the ion name as the second word. • If monatomic, use the element name modified with the ending –ide. • Examples: • NH4Cl • CaSO4

19. Writing Formulas for Ionic Compounds with Polyatomic Ions • Use the criss-cross technique that is used for the binary compounds. • If the polyatomic ion needs a subscript, then place parentheses around the polyatomic ion. The subscript goes outside of the parentheses.

20. Examples • sodium sulfate • ammonium phosphate • Iron (III) nitrate

21. Metallic Bonds In metals, valence shells of atoms overlap, so v.e– are free to travel between atoms through material. Not so in metals. In insulators (like wood), the v.e– are attached to particular atoms.

22. ductile conduct heat and electricity malleable Properties of Metals All due to free-moving v.e–.

23. Covalent Bonding • A form of chemical bonding where electrons are shared so that each atom has a complete valence shell • One particle of a covalent compound is called a molecule.

24. Types of Covalent Bonds • Single bond – 2 electrons (1 pair) are shared. • Double bond – 4 electrons (2 pairs) are shared. • Triple bond – 6 electrons (3 pairs) are shared.

25. Other Names for Bond Types • A single bond is a sigma bond. • Double and triple bonds contain pi bonds. • Double = 1 sigma, 1 pi • Triple = 1 sigma, 2 pi

26. F F 8 Valence electrons 8 Valence electrons Covalent bonding • Fluorine has seven valence electrons • A second F atom also has seven • By sharing electrons both end up with full orbitals (stable octets)

27. n = 2 - - - - Covalent Bonding - - - - n = 1 - - - - - - - - + - - - - - - - - - - - - O [He]2s22p4 O [He]2s22p4 O2 Sharing of electrons to achieve a stable octet (8 electrons in valence shell).

28. Lewis Structures Lewis structure: a model of a covalent molecule that shows all of the valence electrons 1. Two shared electrons make a single covalent bond, four make a double bond, etc. 2. unshared pairs: pairs of un-bonded valence electrons 3. Each atom needs a full outer shell, i.e., 8 electrons. Exception: H needs 2 electrons

29. Lewis Structures 1) Count up the total number of valence electrons. Gilbert Lewis 2) Connect all atoms with single bonds. - “multiple” atoms usually on outside - “single” atoms usually in center; C always in the center, H always on the outside.

30. 3) Complete octets on exterior atoms (not H, though) - no unpaired electrons 4) Check - valence electrons match with Step 1 - all atoms (except H) have an octet; if not, try multiple bonds - any extra electrons? Put on central atom

31. Example Lewis Diagram: CF4

32. Example Lewis Diagram: CH4

33. Example Lewis Diagram: PF3

34. Naming Binary Covalent Compounds • Identification = Use this system with nonmetal-nonmetal pairs • If the compound contains a metal, it is likely an ionic compound – use the other set of rules for those compounds. Naming Rule • First word = prefix + 1st element name • Second word = prefix + 2nd element name with –ide ending

35. Prefixes • mono- • di- • tri- • tetra- • penta- • hexa- • hepta- • octa- • nona- • deca- Exception: Use mono on the second word only, never the first word.

36. Examples – Writing Names • N2O • SiF4 • N2O5 • NH3

37. Examples – Writing Formulas • sulfur trioxide • dihydrogen monoxide

38. VSEPR Valence Shell Electron Pair Repulsion • This is a model to determine the shape of a molecule. • The model is based on an arrangement (of the atoms) that minimizes the repulsion of shared and unshared pairs.

39. Terms-VSEPR • Shared pair = two electrons that are involved in a covalent bond. • Lone pair = two electrons that are not involved in any bond. These are sometimes called unshared pairs. • Bond angle = the angle formed by the two terminal atoms (on the end) and the central atom.

40. B : N : : O Bonding and Shape of Molecules Number of Bonds Number of Unshared Pairs Covalent Structure Shape Examples -Be- 0 0 0 1 2 2 3 4 3 2 Linear Trigonal planar Tetrahedral Pyramidal Bent BeCl2 BF3 CH4, SiCl4 NH3, PCl3 H2O, H2S, SCl2 C

41. .. .. .. S O O O C O O S N F O O F F F F F F F P S F F F F F F The VSEPR Model The Shapes of Some Simple ABn Molecules SO2 Linear Bent Trigonal planar Trigonal pyramidal AB6 Tetrahedral Trigonal bipyramidal Octahedral Brown, LeMay, Bursten, Chemistry The Central Science, 2000, page 305