1. CHEMICAL FORMULAS AND BONDING Ions and Molecules

2. Learning Target ➢ Understand octet rule and how it applies to oxidation state of an ion. ➢ What do all atoms need in order to be stable?

3. OCTET RULE ➢ Elements will gain or lose electrons in order to obtain a noble gas valence (full shell)

4. Remember…. ➢ It takes energy to gain or lose electrons ➢ Nature wants to move towards the path of least resistance….. ……because it requires less energy

5. Gain or Lose Electrons? ➢ Depends on the nearest noble gas. Noble Gases

6. OXYGEN ➢ Oxygen has 6 valence electrons ➢ Is it easier to gain 2e - or lose 6e - ? ➢ Easier to gain 2e -

7. SODIUM ➢ Sodium has 1 valence e - ➢ Is it easier to gain 7e - or lose 1e - ? ➢ Easier to lose 1e -.

8. Try some…… ➢ Sulfur ➢ Fluorine ➢ Potassium ➢ Argon ➢ Carbon ➢ Hydrogen ➢ Calcium Gain Lose Neither Both Lose

9. OXIDATION NUMBER ➢ The possible charge an atom could obtain by gaining or losing electrons or the number of electrons an element will donate/accept in a bond. ● Remember electrons are negative

10. OXYGEN ➢ Oxygen gains 2e - ● So its charge is 2 - O 2- Oxidation Number

11. SODIUM ➢ Sodium loses 1e - ● So its charge is 1 + Na +

12. TRY SOME….. ➢ Magnesium ➢ Boron ➢ Bromine Loses 2e -, Mg 2+ Loses 3e -, B 3+ Gains 1e -, Br -

13. 2 TYPES OF CHEMICAL BONDS WE WILL TALK ABOUT ➢ Ionic compounds ➢ Covalent compounds ➢ 2 terms that go with the ionic and covalent: ➢ Empirical ➢ Molecular

14. Review 1.What is the oxidation number for Rubidium? Selenium? 2. Do metals have (+) or (–) oxidation numbers? 3. Do non metals have (+) or (–) oxidation numbers?

15. Review 1. What rule states that elements will gain or lose electrons to obtain a noble gas valence? o Octet Rule 2. Do non metals have (+) or (–) oxidation numbers? (-) 3. Do metals have (+) or (–) oxidation numbers? (+) 5. What is the oxidation number for barium? Iodine? ➢ Ba 2+ and I -

16. Goal of Today ➢ Know how to write formulas for Ionic Compounds.

18. Ionic Bond ➢ Electrons are everywhere – static is a good example ➢ Positive ion is attracted to a negative ion in an ionic bond ➢ What kind of elements?

19. Ionic Compound ➢ Made up of ions ➢ Electrically neutral ➢ Charges must equal each other ➢ Bonds between metals (+ charge) and nonmetals (- charge)

20. Ionic Bonding ➢ Ionic bonds occur between positive metal ions (cations) and negative nonmetal ions (anions) ➢ Made up of ions ➢ Electrically neutral ➢ Charges must equal each other

21. Properties of Ionic Compounds ➢ Strong bonds ➢ High melting points ➢ Brittle ➢ Soluble in water ➢ Their solutions are good conductors of electricity ➢ Solids at 20°C

22. Formation of Ionic Compounds ➢ Use ● The Octet rule ● Lewis Dot Diagrams ● Crisscross Method

23. Monatomic Cations ➢ Positive 1, 2, or 3 ➢ Transition metals can vary, and even have a charge of 4+ ➢ Use element name ➢ Use a Roman numeral with any metal that varies in charge ➢ Example Copper(I) is Cu +1 and Copper (II) is Cu 2+

24. Monatomic Anions ➢ Can be negative 1,2, or 3 ➢ Change element name to “-ide” ending ➢ Example: chloride

25. Polyatomic Ions ➢ Two or more atoms that are bonded together and carry a single charge ➢ Names are on the handout ➢ Most are negative with one positive ➢ Usually end in “-ate” or “-ite” ➢ Example: NO 3 - is nitrate

26. Formulas for Binary Compounds ➢ Contain a monatomic cation (metal) and a monatomic anion (nonmetal) ➢ Metal is first ➢ Charges must add to “0” ➢ Use subscripts to get the value to “0” ➢ Why is sodium chloride NaCl? ➢ Try some

27. Write the formulas for the following ➢ potassium iodide ➢ barium chloride ➢ lithium bromide ➢ calcium hypochorite ➢ chromium (III) sulfide ➢ gold (III) bisulfate

28. Review 1.What are negative ions called? 2. What are positive ions called? 3. Write the formula for niobium(V) phosphate

29. Goal of Today ➢ Understanding naming of Ionic Compounds and begin our “Ionic Bonding Puzzle Lab”

30. Naming Ionic Compounds ➢ Use NaCl as a good example ➢ Metal first, then nonmetal ➢ Ends in “-ide” if binary (2 elements only) ➢ Use polyatomic name ➢ Use Roman numeral if necessary ➢ Suspect every transition metal to possibly have a Roman numeral

31. Review Questions 1. Write the formulas for: beryllium carbonate silver nitrate 2. Name the following ionic compound: Cu(HSO 4 ) 2

32. GOAL OF TODAY Understand how covalent bonds form and know how to draw Lewis structures to represent covalent compounds.

33. COVALENT BONDING ➢ A covalent bond is formed by a shared pair of electrons between two atoms. ➢ A group of atoms that are united by covalent bonds is called a MOLECULE ➢ Most of what you see is covalently bonded

35. Describing a molecular bond ➢ Molecular formula: tells you how many atoms and which kind are in each molecule. Glucose C 6 H 12 O 6 ➢ Empirical formula: gives you the ratio of the atoms in a molecule. Glucose: C 1 H 2 O 1

36. Covalent Bonding ➢ Single bonds share 2 electrons. Example: Ammonia (H-N-H) H ➢ Double bonds share 4 electrons: Formaldehyde H-C=O H ➢ Triple bonds share 6 electrons Ethyne H-C=C-H

37. Exit Questions 1.What is the difference between ionic bonds and covalent bonds? 2.Draw the Lewis structure for N 2 H 2

38. GOAL OF TODAY ➢ Know that there are exceptions to the octet rule. ➢ Know how to determine and notate polarity of chemical bonds.

39. Exceptions to the Octet Rule ➢ Some elements are satisfied with fewer than 8 electrons (6 or 4) ➢ Some structures can only be drawn with 7 electrons ➢ Compounds with Beryllium (Be) and Boron (B) may have less than an octet BCl 3 ➢ 3 rd Row elements or lower may exceed octet rule. SF 4. ➢ Odd number of electrons cannot follow the octet rule. I.e., NO or CO ➢ These substances can be short-lived and reactive – called free radicals

40. Chloroform CHCl 3 Draw the Lewis structure for chloroform

41. ➢ Draw the Lewis structure for PF 3

42. Properties of Covalent Bonds ➢ Electrons are not necessarily shared equally. This depends on the electronegativity of an atom (atom’s attraction for electrons)

43. Polar and Nonpolar Covalent Bonds ➢ Polar bonds: 1 atom is significantly more electronegative than the other one. One side of the bond is slightly positive the other negative. Example: Water ➢ Nonpolar bonds: Both atoms have similar electronegativities. Example:

44. Bond Type by Electronegativity ➢ First, Find the difference of electronegativities of the atoms.Then apply the information of the table below Electronegativity Difference Bond type < 0.4Non-polar covalent Between 0.4 and 2.0Polar covalent > 2.0Ionic bond

45. Goal for Today ➢ Review bond polarity ➢ Know the difference in naming ionic and covalent compounds.

46. Look at the table on page 241. ➢ What is the electronegativity difference between hydrogen and oxygen (H 2 0)? ● 1.4 (polar covalent) ➢ What is the electronegativity difference between sodium and chlorine (NaCl)? ● 2.1 approximately (ionic) ➢ What is the electronegativity difference between 2 nitrogen atoms (N 2 )? ● 0 (non polar covalent)

47. Properties of Molecular Substances Weak bonds Can be solids, liquids, or gases at 20°C Lower melting points Poor conductors of electricity Soft, not hard and brittle Some are soluble in water.

48. Naming Molecular Compounds Similar to naming ions Numerical prefixes are used. Example: CO 2 =Carbon Dioxide Do not use a prefix for one - ide is added to the more electronegative element Some elements have common names: Diatomics like O 2 = oxygen; not dioxide NH 3 = ammonia; not Nitrogen tetrahydride H 2 O is water not dihydrogen monoxide

49. HYDRATES Ionic compounds that absorb water into their solid structures. Example: CuSO 4. 5 H 2 O (s) Copper (II) Sulfate pentahydrate