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Chapter 7

Chapter 7. REALLY Important!!!. 7.1 – Ionic Compounds: Ions for s and p block elements:. 7.1 – Ionic Compounds: Ions – d Block. *Intro Classes can put Roman Numerals on ALL d-block elements *Honors Class must know which elements NEED Roman Numerals. cobalt (II) = Co +2

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Chapter 7

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  1. Chapter 7 REALLY Important!!!

  2. 7.1 – Ionic Compounds: Ions for s and p block elements:

  3. 7.1 – Ionic Compounds: Ions – d Block *Intro Classes can put Roman Numerals on ALL d-block elements *Honors Class must know which elements NEED Roman Numerals cobalt (II) = Co+2 cobalt (III) = Co+3 copper (I) = Cu+1 copper (II) = Cu+2 chromium (II) = Cr+2 chromium (III) = Cr+3 chromium (VI) = Cr+6 iron (II) = Fe+2 iron (III) = Fe+3 Silver (I) = Ag+1 platinum (II) = Pt+2 platinum (IV) = Pt+4 mercury (II) = Hg+2 Zinc (II) = Zn+2 cadmium (II) = Cd+2 manganese (II) = Mn+2 manganese (IV) = Mn+4 nickel (II) = Ni+2 nickel (III) = Ni+3 gold (III) = Au+3

  4. 7.1 – Polyatomic IonsThe Tough Stuff! THE 8 -ATES: • Carbonate CO3-2 • Nitrate NO3-1 • Sulfate SO4-2 • Chlorate ClO3-1 • Chromate CrO4-2 • Bromate BrO3-1 • Phosphate PO4-3 • Iodate IO3-1 Rules with –ates: 1 more oxygen than –ate = per … ate 1 less oxygen than –ate = …ite 2 less oxygens than –ate = hypo … ite

  5. 7.1 – Other Polyatomic Ions Ammonium NH4+1 Hydronium H3O+1 Peroxide O2-2 Hydroxide OH-1 Cyanide CN-1

  6. 7.1 – Ionic Compounds – Names and Formulas Naming Binary Ionic Compounds: A metal + a nonmetal = IONIC Name = cation name anion name (w/-ide ending) -Use Roman Numerals if needed Examples:

  7. 7.1 – Ionic Compounds – Names and Formulas Writing Binary Ionic Formulas: Remember Chapter 6? New way: Use the swap technique – Number on charge tells you how many of the OTHER element Examples:

  8. Examples

  9. 7.1 – Ionic Compounds – Names and Formulas Naming Ionic Compounds w/more than 2 elements: Name = cation name anion name (at least one will be a polyatomic ion) Examples:

  10. 7.1 – Ionic Compounds – Names and Formulas Writing Ionic Formulas for Ionic Compounds with More Than 2 Elements: Use the swap technique Examples:

  11. 7.1 – Molecular Compounds 2 Nonmetals Bonded USE PREFIXES! Mono= Di= Tri= Tetra= Penta= Hexa= Hepta= Octa= Nona= Deca=

  12. 7.1 - Molecular Compounds Name = prefix first element prefix second element-ide Prefix is the quantity of that element Mono is not needed in front of the FIRST element only Examples:

  13. 7.1 – Acid Naming • Acids are an specific molecular substance • We will discuss more in Ch. 15 • Two types: • Binary Acids • Oxyacids

  14. 7.1 – Acid Naming • Binary Acids • Consist of ONLY two elements • Usually H + one of Halogens (F, Cl, Br,I) • Hydro…Root of element…ic Acid • Examples: • HF = • HCl = • HBr = • HI =

  15. 7.1 – Acid Naming • Oxyacids • Acids that contain three elements (H, O, and usually a third nonmetal element) • Related to the 8-ate polyatomic ions:

  16. 7.1 – Acid Naming • Oxyacids • Examples: • HClO3 • HClO2 • HClO • HClO4

  17. 7.1 – Review • Practice Problems • Ions • Naming Ionic Compounds • Writing Ionic Compound Formulas • Naming Molecular Compounds • Writing Molecular Compound Formulas • Acid Naming • Binary • Oxyacids

  18. 7.2 – Oxidation Numbers -Show distribution of electrons -Negative means “stronger element / ‘grabbing’ electrons” and positive means “weaker element / losing electrons” Rules: 1. Uncombined element has oxidation number = 0 2. Monatomic ion has oxidation number = charge 3. If in a compound: A. Start with the element on the right. It has oxidation number = charge if it was a negative ion. B. If more than 2 atoms, go next to the element on the left, it has oxidation number = charge if it was a positive ion. C. Hydrogen can be –1 if it is on the right or +1 if it is the one on the left. 4. Sum of all oxidation numbers in a neutral compound = 0 5. Sum of all oxidation numbers in a polyatomic ion = charge on ion

  19. 7.2 – Oxidation Numbers Practice Problems:

  20. 7.3 – Molar Mass Molar Mass = Sum of Average Atomic Masses for all elements in a compound from periodic table. Unit = g/mol Examples:

  21. Using Molar Mass Can be used as a conversion factor O2 = 32 grams / 1 mole OR 1 mole / 32 grams Examples:

  22. Percent Composition Percent by mass of an element in a compound (Mass element / molar mass) x 100 = % composition Examples:

  23. Empirical Formulas Empirical Formula = The simplest (most reduced) formula Example: The empirical formula of glucose (C6H12O6) is…. Steps: 1. Find moles of each element (divide given by molar mass of element) 2. Divide all moles by the smallest number of moles 3. Write formula with subscripts 4. If .5s multiply all by 2

  24. Empirical Formulas Calculations Examples:

  25. Molecular Formulas This is the NON-Reduced formula (ex – glucose = C6H12O6) Need to have the Empirical Formula and the molecule’s Molar Mass (AKA Molecular Mass) Step One:Molecular Mass = x Empirical Form Mass Step Two: x multiplied by the Empirical Form

  26. Molecular Formulas Examples:

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