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Chapter 12 - Ionic Bonding

Chapter 12 - Ionic Bonding. Bonds. Chemical bond - forces of attraction that hold atoms together. The molecule is more stable than the separate ions/ atoms. Bond energy - the energy required to break (or form) a chemical bond. Energy is released when the bond is formed.

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Chapter 12 - Ionic Bonding

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  1. Chapter 12 - Ionic Bonding

  2. Bonds • Chemical bond - forces of attraction that hold atoms together. • The molecule is more stable than the separate ions/ atoms. • Bond energy - the energy required to break (or form) a chemical bond. • Energy is released when the bond is formed. • Ionic bonds - bond formed by the attraction between oppositely charged particles (anion and cation) • The transfer of one or more electrons from one atom to another

  3. Ionic Compounds • Ionic compounds - a chemical compound that has an equal amount of positive and negative charges so the compound is electrically neutral. Formed by ionic bonds. • Salt - an inorganic compound containing cations other than H+ and anions other than OH- and O2-.

  4. Properties of Salts • All salts are made of ions and are held together by ionic bonds. • Salts form a crystal lattice - an ordered packing 3-dimensional arrangement of atoms or ions in a crystal • Crystalline versus amorphous- crystalline solids have a fixed arrangement of atoms while amorphous solids like talcum powder have haphazard/ random arrangement.

  5. Properties of Salts • Salts do not melt or boil easily because of the strong attraction between ions • Salts are hard and brittle • Salts conduct electricity when melted or dissolved to form an aqueous solution (in water- aq.)

  6. Energy and Ionic Bonding • Removing an electron from an atom requires an input of energy- Ionization energy • Adding an electron to an atom releases energy- Electron Affinity • Lattice energy - the energy released when the crystal lattice of an ionic solid is formed

  7. Nomenclature PO43- phosphate ion HC2H3O2 Acetic Acid C2H3O2- acetate ion

  8. Forms of Chemical Bonds • There are 3 forms of bonding between atoms: • Ionic—complete transfer of 1 or more electrons from one atom to another (one loses, the other gains) • Covalent—some valence electrons shared between atoms • Metallic – holds atoms of a metal together. Electrons form a common pool called sea of electrons. Most bonds are somewhere in between ionic and covalent.

  9. Common Names • A lot of chemicals have common names as well as the proper IUPAC name. • Chemicals that should always be named by common name and never named by the IUPAC method are: • H2O water, not dihydrogen monoxide • NH3 ammonia, not nitrogen trihydride

  10. COMPOUNDS FORMED FROM IONS CATION + ANION ---> COMPOUND Na+ + Cl- --> NaCl A neutral compound requires equal number of + and - charges.

  11. Predicting Charges on Monatomic Ions KNOW THESE !!!! +1 +2 -3 -2 -1 0 Cd+2

  12. Properties of Ionic CompoundsForming NaCl from Na and Cl2 • A metal atom can transfer an electron to a nonmetal. • The resulting cation and anion are attracted to each other by electrostatic forces.

  13. NH4+ Cl- IONIC COMPOUNDS ammonium chloride, NH4Cl

  14. Some Ionic Compounds Ca2+ + 2 F- ---> CaF2 Mg2+ + N-3 ----> Mg3N2 magnesiumnitride Sn4+ + O2- ----> SnO2 Tin (IV) oxide calcium fluoride

  15. Formulas of Ionic Compounds Formulas of ionic compounds are determined from the charges on the ions atoms ions  – Na  +  F :  Na+ : F :  NaF  sodium + fluorine sodium fluoride formula Charge balance: 1+1- = 0

  16. Monatomic Ions

  17. Writing a Formula Write the formula for the ionic compound that will form between Ba2+ and Cl. Solution: 1. Balance charge with + and – ions 2. Write the positive ion of metal first, and the negative ion Ba2+ ClCl 3. Write the number of ions needed as subscriptsBaCl2

  18. Writing BINARY FORMULAS • 1. CANCEL CHARGES TO GET A TOTAL CHARGE OF ZERO. • 2. CROSS- MULTIPLY TILL YOU GET THE LOWEST COMMON MULTIPLE. • 3. WHAT YOU MULTIPLIED WITH IS YOUR SUBSCRIPT. • 4. DO NOT WRITE CHARGES ON YOUR FINAL FORMULA. • 5. METAL IS WRITTEN FIRST AND THEN NON- METAL. • SIMPLEST FORMULA IS WRITTEN

  19. Learning Check Write the correct formula for the compounds containing the following ions: 1. Na+, S2- a) NaS b) Na2S c) NaS2 2. Al3+, Cl- a) AlCl3 b) AlCl c) Al3Cl 3. Mg2+, N3- a) MgN b) Mg2N3 c) Mg3N2

  20. Solution 1. Na+, S2- b) Na2S 2. Al3+, Cl- a) AlCl3 3. Mg2+, N3- c) Mg3N2

  21. Naming Compounds Binary Ionic Compounds: • 1. Cation first, then anion • 2. Monatomic cation = name of the element • Ca2+ = calciumion • 3. Monatomic anion = root + -ide • Cl- = chloride • CaCl2 = calcium chloride

  22. Naming Binary Ionic Compounds • Examples: NaCl ZnI2 Al2O3 sodium chloride zinc iodide aluminum oxide

  23. Learning Check Complete the names of the following binary compounds: Na3N sodium ________________ KBr potassium ________________ Al2O3 aluminum ________________ MgS _________________________

  24. Name the ionic compounds BaCl2 NaF Ag2O CuBr CuBr2 FeO Fe2O3 MgS Al2O3 CaI2 11. K2S 12. CrCl2 13. CrCl3 14. CaO 15. Ba3P2 16. Hg2I2 17. Na2O 18. BeS 19. MnO 20. Mn2O3

  25. Transition Metals Elements that can have more than one possible charge MUST have a Roman Numeral to indicate the charge on the individual ion. 1+ or 2+ 2+ or 3+ Cu+,Cu2+ Fe2+, Fe3+ copper(I) ion iron(II) ion copper (II) ion iron(III) ion

  26. Names of Variable Ions These elements REQUIRE Roman Numerals because they can have more than one possible charge: anything except Group 1A, 2A, Ag, Zn, Cd, and Al (You should already know the charges on these!) Or another way to say it is: Transition metals and themetalsin groups 4A and 5A (except Ag, Zn, Cd, and Al) require a Roman Numeral. FeCl3(Fe3+) iron (III) chloride CuCl (Cu+ ) copper (I) chloride SnF4 (Sn4+) tin (IV) fluoride PbCl2 (Pb2+) lead (II) chloride Fe2S3 (Fe3+) iron (III) sulfide

  27. Examples of Older Names of Cations formed from Transition Metals(you do not have to memorize these)

  28. Learning Check Complete the names of the following binary compounds with variable metal ions: FeBr2 iron (_____) bromide CuCl copper (_____) chloride SnO2 ___(_____ ) ______________ Fe2O3 ________________________ Hg2S ________________________

  29. Polyatomic Ions NO3- nitrate ion NO2- nitrite ion

  30. Polyatomic ions Contain more than one atom. Can be positive (only one is Ammonium NH4+), or negative (lots of them). Stay together during a chemical reaction. Form ternary (3 elements present) or quaternary compounds (4 elements).

  31. Polyatomic Ions You can make additional polyatomic ions by adding a H+ to the ion! CO3-2 is carbonate HCO3– is hydrogen carbonate H2PO4– is dihydrogen phosphate HSO4– is hydrogen sulfate

  32. Ternary Ionic Nomenclature Writing Formulas • Write each ion, cation first. Don’t show charges in the final formula. • Overall charge must equal zero. • If charges cancel, just write symbols. • If not, use subscripts to balance charges. • Use parentheses to show more than one of a particular polyatomic ion. • Use Roman numerals indicate the ion’s charge when needed (stock system)

  33. Ternary Ionic Nomenclature Sodium Sulfate Na+ and SO4-2 Na2SO4 Iron (III) hydroxide Fe+3 and OH- Fe(OH)3 Ammonium carbonate NH4+ and CO3–2 (NH4)2CO3

  34. Learning Check 1. aluminum nitrate a) AlNO3 b) Al(NO)3 c) Al(NO3)3 2. copper(II) nitrate a) CuNO3 b) Cu(NO3)2 c) Cu2(NO3) 3. Iron (III) hydroxide a) FeOH b) Fe3OH c) Fe(OH)3 4. Tin(IV) hydroxide a) Sn(OH)4 b) Sn(OH)2 c) Sn4(OH)

  35. Naming Ternary Compounds • Contains at least 3 elements • There MUST be at least one polyatomic ion (it helps to circle the ions) • Examples: NaNO3 Sodium nitrate K2SO4Potassium sulfate Al(HCO3)3Aluminumbicarbonate or Aluminumhydrogen carbonate

  36. Learning Check Match each set with the correct name: 1.Na2CO3a) magnesium sulfite MgSO3b) magnesium sulfate MgSO4c) sodium carbonate 2 . Ca(HCO3)2a) calcium carbonate CaCO3b) calcium phosphate Ca3(PO4)2 c) calcium bicarbonate

  37. Mixed Practice! Name the following: • Na2O • CaCO3 • PbS2 • Sn3N2 • Cu3PO4 • HgF2

  38. Mixed Up… The Other Way Write the formula: • Copper (II) chlorate • Calcium nitride • Aluminum carbonate • Potassium bromide • Barium fluoride • Cesium hydroxide

  39. PROPERTIES OF IONIC COMPOUNDS DEMOS-  ELECTRICAL CONDUCTIVITY- • In a well plate, put a table salt (NaCl) solution in one of the wells. Dip the two ends of an electrical conductometer into the well- see if it lights up. • Repeat the process, substituting dry salt, sugar, sugar solution and then distilled water for the aqueous NaCl solution. Rinse and dry the conductometer every time.

  40. Properties (contd.) MELTING POINTS- Heat about 10 gm of table salt (ionic) in a test tube for a maximum of 2-3 minutes. What happens? Repeat with heating sugar (covalent). What happens? What does this tell you about the melting points of ionic compounds versus covalent compounds? Which bond is stronger?

  41. Naming Molecular (Covalent) Compounds All are formed from two or more nonmetals. CO2 Carbon dioxide Ionic compounds generally involve a metal and nonmetal (NaCl) BCl3boron trichloride CH4 methane

  42. Molecular (Covalent) Nomenclaturefor two nonmetals • Prefix System (binary compounds) 1. Less electronegative atom comes first. 2. Add prefixes to indicate # of atoms. Omit mono- prefix on the FIRST element. Mono- is OPTIONAL on the SECOND element . 3. Change the ending of the second element to -ide.

  43. PREFIX mono- di- tri- tetra- penta- hexa- hepta- octa- nona- deca- NUMBER 1 2 3 4 5 6 7 8 9 10 Molecular Nomenclature Prefixes

  44. Molecular Nomenclature: Examples • CCl4 • N2O • SF6 • carbon tetrachloride • dinitrogen monoxide • sulfur hexafluoride

  45. More Molecular Examples • arsenic trichloride • dinitrogen pentoxide • tetraphosphorus decoxide • AsCl3 • N2O5 • P4O10

  46. Learning Check Fill in the blanks to complete the following names of covalent compounds. CO carbon ______oxide CO2 carbon _______________ PCl3 phosphorus _______chloride CCl4 carbon ________chloride N2O _____nitrogen _____oxide

  47. Learning Check 1. P2O5 a) phosphorus oxide b) phosphorus pentoxide c) diphosphorus pentoxide 2. Cl2O7a) dichlorine heptoxide b) dichlorine oxide c) chlorine heptoxide 3. Cl2a) chlorine b) dichlorine c) dichloride

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