A Chemist’s View of Explosives :

1. A Chemist’s View of Explosives: Ionic Bonding and Nomenclature Notes

2. I. Chemical bond: a mutual electrical attraction between the nuclei and valence electrons of different atoms that binds the atoms together. Another way to describe a chemical bond is to say the attractive forces between atoms or ions in compounds. In ionic compounds it is an attractive force between positive and negative ions. http://www.visionlearning.com/library/module_viewer.php?mid=55

3. In ionic bonding valence electrons are actually TRANSFERRED between a nonmetal and metal. This happens because a non-metallic atom is much more electronegative and it can pull electrons away from the less electronegative metallic atom. In an ionic compound the positive and negative ions combine so that the overall charge is zero.

4. Sometimes the more electronegative atom is not “powerful” enough to completely take away the electrons from another atom so the atoms SHARE electrons. This sharing of electrons is called a covalent bond. http://web.visionlearning.com/custom/chemistry/animations/CHE1.7-an-H2Obond.shtml

5. Ionic Bonding occurs between metals and nonmetals.Covalent Bonding occurs between nonmetals.Bonds (and compounds) form in order to obtain an electron configuration like that of noble gases!

6. Electron Dot Structures: show the placement and transfer of valence electrons. Rules to remember when drawing electron dot structures: II. Formation ofIonic Bonds and Ionic Compounds

7. Only valence electrons are shown. Valence electrons are the electrons in the outermost s and p sublevels. Transition metals could also have d sublevel valence electrons. • Valence electrons are shown as dots and are not drawn randomly! They are arranged around the element's symbol to correspond to the elements electron configuration. (Only 2 dots or electrons per side.) 3. Follow the Octet Rule which sates that atoms form bonds in order to obtain 0 or 8 valence electrons, because of this electron dot structures will show no more than 8 electrons for each atom or ion. Another way to think of the Octet rule: Atoms react by changing the number of their electrons so as to acquire the stable electron configuration of a noble gas.

8. B. Electron Dot Structures for Atoms: Write the element's symbol and place the appropriate number of dots to represent the valence electrons around the symbol. (The electron configuration is given to help you understand the idea of valence electrons.) a.) Ca [Ar]4s2 b.) Li [He]2s1 c.) Be [He]2s2 d.) O [He]2s22p4 e.) Br [Ar]4s23d104p5

9. C. Electron Dot Structures for Ions: Ions form when atoms lose or gain valence electrons.

10. (1.) cations - these form when atoms have LOST valence electrons. a.) Mg ion b.) Li ion c.) Al ion d.) Ba ion

12. (2.) anions-these form when atoms have GAINED valence electrons. a.) S ion b.) Br ion c.) N ion d.) P ion

13. (3.) Transition and Inner Transition Elements-the number of valence electrons for these are harder to predict based on their position on the periodic table because some ofthese elements have valence electrons in the d sublevel. Example: a.) How many valence electrons does an atom of iron have? To answer this question write the electron configuration for iron: Are there any unstable electrons in the d level? When iron ionizes what are the possible ions?

14. b.) How many valence electrons does an atom of titanium have? Electron configuration for titanium: Are there any unstable electrons in the d level? When iron ionizes what are the possible ions?

15. D. Pseudo-noble gas electron configuration-elements that cannot acquire a noble gas electron configuration, but can become somewhat stable with 18 electrons in their outer shell. Examples are: Hg+2, Cd+2, Au+1, Cu+1

16. E. Electron Dot Structures for Ionic Compounds: • Write the electron dot structure for each of the elements involved. • Draw arrows from the electrons of the metallic atom to the non-metallic atom. This shows the transfer of electrons. 3. After the Write the dot diagram for the new ionic compound, including charges.

17. Lewis Dot Structures for Ionic Compounds (compounds held together by ionic bonds – usually a M Example: a.) Sodium and Chlorine Na + Cl  [ Na ]+ + [ • Cl• ]- • •• •• • • • • • •• ••

18. Examples: b.) Magnesium and Oxygen

19. Examples: c.) Aluminum and Oxygen

20. Examples: d.) Calcium and Fluorine

21. Examples: e.) Sodium and Nitrogen

22. F. Characteristics of ionic compounds (compared to molecular compounds) -higher melting points -higher boiling points -generally hard, brittle solids -when melted or dissolved in water they can conduct electricity -shapes are crystalline in nature – square/cube

23. Copper sulfate has a triclinic crystal structure.

24. Lattice Energy: the energy released when one mole of an ionic crystalline compound is formed from gaseous ions. Negative values for lattice energy mean that energy was released when the ionic crystal is formed.

25. Formula unit: the smallest number of atoms that can make up an ionic compound

26. III. Nomenclature – Ionic Compounds Part 1: Writing Formulas for Ionic Compounds

27. A. Rules for Writing Formulas for Binary Ionic Compounds – these are compounds containing only 1 metal and 1 nonmetal. 1. Write the cation (metal ion) first and the anion (nonmetal ion) second. 2. Determine the smallest whole number ratio of cations to anions that would make the charge 0. To determine the ion formed for main group elements look where it is located on the periodic table. Many elements in groups 3-12 have either a +2 or +3 charge. Noble gases to do not form ions except in rare cases. Hydrogen can either gain, lose, or share an electron depending on the other elements with which it combines.

28. Example: Write the formula for sodium chloride. Na+1 Cl1- NaCl charges equal 0 Example: Write the formula for aluminum oxide. Al3+ O2- Al2O3 charges equal 0

29. Practice: Magnesium phosphide____________________ iron(II) bromide______________________ Calcium oxide __________________________ sodium sulfide _____________________ Copper (II) iodide __________________________ lead (IV) nitride ____________________ Aluminum nitride _______________________ tin (II) chloride _________________ Potassium fluoride __________________________ Copper (I) phosphide ________________ Copper (II) oxide ___________________________ potassium bromide __________________ Iron (III) fluoride _______________________ Tin (II) oxide ______________________

30. Rules for Writing Formulas for Ternary Ionic Compounds – these are compounds containing polyatomic ions. (Look at the chart to figure out the formula and charge. Polyatomic Ions-two or more elements (usually nonmetals) bonded together that have collectively lost or gained electrons and now have a charge. Compounds have a zero charge but a polyatomic ion has a charge. You must memorize several polyatomic ions (look to “Memorize These Ions” sheet).

31. 1. Write the cation first and the anion second. 2. Determine the smallest whole number ratio of cations to anions that would make the charge 0. If a subscript must be added to a polyatomic ion, keep the polyatomic ion in parentheses. Example: Write the formula for sodium phosphate. Na1+ (PO4)3- Na3PO4 Ammonium sulfide (NH4)1+ S2- (NH4)2S

32. Practice: aluminum sulfate _______________ potassium chlorate ______________ Copper (II) acetate ______________________ plumbous nitrate _______________________ Iron (III) oxalate _______________________ magnesium chlorate _____________________ Magnesium dichromate __________________ tin (II) hypochlorite ____________________ Lead (II) perchlorate__________________ tin (II) nitrite _________________________ Ammonium carbonate___________________ iron (II) sulfite _________________________ Sodium cyanide ___________________ Lithium phosphite _____________________

33. Part 2: Writing Names for Ionic Compounds There are two naming systems currently acceptable: • IUPAC – International Union of Practical and Applied Chemistry is the newest system-this system uses Roman numerals to give the charges or oxidation number of positive ions ONLY if the positive ion has variable charges. This is the system we will use.

34. 2. “ous” and “ic” system – oldest system and still very commonly used. May be used ONLY if the positive ion has a variable charge and exhibits only 2 oxidation numbers. Review the 4 ions with which we may use this system. Cu ____ _________ _____ ________ Sn ____ _________ _____ ________ Fe ____ _________ _____ ________ Pb ____ _________ _____ ________

35. A. Rules for Writing Names for Binary Ionic Compounds – these are compounds containing only 1 metal and 1 nonmetal. 1. The correct full name of the cation (metal ion) is written first. (Do not forget about the roman numerals if it is a Cu, Fe, Sn, or Pb ion!) • The last syllable in the anion (nonmetal ion) is dropped and –ide is added. Example: NaCl Sodium Chloride Example: CuS Copper (II) Sulfide or Cupric Sulfide

36. Practice: MgCl2 ___________________________ AlI3 _____________________________ Na3P _____________________________ Ca3N2 _______________________________ FeN _____________________________ PbCl2 ________________________________ CuF ____________________________ CuCl2 _________________________________ ZnS ______________________________ Pb02 _________________________________ Fe203 _____________________________ KI ___________________________________ Cu0 ______________________________ Cs3N ________________________________

37. B. Rules for Writing Names for Ternary Ionic Compounds - these are the compound containing polyatomicions. (It is imperative that you know the correct names of these ions!!!!) 1. The correct full name of the cation (metal ion or polyatomic ion) is written first. (Do not forget about the roman numerals if it is a Cu, Fe, Sn, or Pb ion!) 2. The correct full name of the anion (polyatomic ion or nonmetal ion) is written second. If the anion is a polyatomic ion do not change the ending. If the anion is a nonmetal ion then the ending is dropped and –ide is added. Example: KNO3 potassium nitrate Example: Cu2CrO4copper (I) chromate or cuprous chromate

38. Practice: Na3PO4_______________________ Al2(SO4)3 _________________________________ CuNO3 _______________________ PbCO3 __________________________________ Li2SO3 ________________________ CaCr207 _______________________________ NH4Cl_________________________ CsClO4 __________________________________ K2CN ________________________ Fe(HSO3)3 _______________________________

39. Mixed Review: Write the name for the following: CuO BaO CaCl2 NaBr K2O Mg3N2 AgOH Pb(ClO)4 CaSO3 Sr(NO3)2 Write the formula for the following: potassium iodide iron (II) chloride sodium sulfide aluminum sulfide copper (II) nitride potassium oxide lead (IV) oxalate magnesium phosphite sodium bicarbonate

40. Part 3: Writing Names and Formulas for Acids A. Rules for Writing Names for Acids • Acids: compounds that produce hydrogen ions when dissolved in water; an anion bonded to a hydrogen cation is an acid. Any compound starting with hydrogen is an acid. • Even though acids usually contain only nonmetals, they are treated as ionic compounds because the hydrogen present is the H+1 ion.

41. Rules for writing the name of an acid: Does the compound contain an oxygen atom? If No – always start with hydro, then the root of the second element, and end with -ic. Example: HCl hydrochloric acid If Yes – oxygen is present – start with the root of the anion (usually a polyatomic ion). If it ends in –ate, change it to –ic. If the anion ends in –ite, change it to –ous. Example: HNO3 Nitric acid

42. Common Roots for the Polyatomic Ions: Acetate acet- Bromate brom- Chlorate chlor- Nitrate nitr- Nitrite nitr- Chromate chrom Oxalate oxal- Sulfate sulfur- Sulfite sulfur- Phosphate phosphor- Perchlorate perchlor- Hyporchlorite hypochlor- Carboante carbon-

43. Practice writing the name for the following acids: Practice: HNO2 HCl H2SO4 HBr H3N HF HI H2S

44. B. Rules for Writing the Formula for an Acid: • Hydrogen will always be the cation (H+1). Write the cation first and the anion second. • Determine the smallest whole number ratio of cations to anions that would make the charge 0. Example: phosphoric acid H1+ (PO4)3- H3PO4 Example: phosphoric acid H3PO4

45. nitric acid_________________ acetic acid______________________ carbonic acid _______________ chloric acid ___________________ chlorous acid__________________ phosphorous acid _____________

46. Let’s Review: Binary Ionic Compounds (BIC): 2 elements, one is a metal and one a nonmetal, will end with –ide, use periodic table to look up ions formed to determine the formula Ternary Ionic Compounds (TIC): 3 or more elements, at least one is a metal and at least one is a nonmetal, these will contain a polyatomic ion (memorize the polyatomic ions), use the ions charge to determine the formula Acids (A): will begin with H, hydrogen’s charge when forming an acid is H+1, If the acid does NOT contain oxygen then start the name with hydro-, followed by the root of the second element, and end with –ic. If the acid does contain oxygen start with the root of the anion (usually a polyatomic ion) and then change the ending to –ic if it was –ate and –ous if it was –ite.

47. CaCO3Mg3P2 Cu(NO2)2 CuCl HCl Magnesium permanganate Barium fluoride Iron (III) nitrate Sulfuric acid Lead (IV) fluoride

48. Part 4: Percent Composition, Empirical, and Molecular Formula Problems A. Molar Mass Practice: • Find the molar mass of ammonium sulfate (also called the formula mass): • Find the molar mass of copper (II) chloride:

49. B. Percent Composition by Mass – the percent, by mass, of each element in a compound. • If you have a box containing 100 golf balls and 100 ping pong balls, which type of ball contributes the most to the mass of the box? • The same principle applies to finding the % composition of a compound. Different elements have different masses and this must be taken into consideration.

50. How to find the percent composition of a compound: • Write a correct formula for the compound • Find the molar mass of the compound 3. Divide the total atomic mass of EACH ELEMENT by the molar mass 4. Multiply by 100 to convert your results to a percent 5. Since you have no significant figures to go by, express your answer to TWO decimal places with the % sign.