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Isotopes, Atomic Numbers, and Mass Numbers

Isotopes, Atomic Numbers, and Mass Numbers. Atomic number (Z) = number of protons in the nucleus. Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). By convention, for element X, we write Z A X. Isotopes have the same Z but different A.

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Isotopes, Atomic Numbers, and Mass Numbers

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  1. Isotopes, Atomic Numbers, and Mass Numbers • Atomic number (Z) = number of protons in the nucleus. • Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). • By convention, for element X, we write ZAX. • Isotopes have the same Z but different A.

  2. The Atomic Mass Scale • 1H weighs 1.6735 x 10-24 g and 16O 2.6560 x 10-23 g. • We define: mass of 12C = exactly 12 amu. • Using atomic mass units: • 1 amu = 1.66054 x 10-24 g • 1 g = 6.02214 x 1023 amu

  3. Isotope Calculation Review • The atomic masses listed on the periodic table are average atomic masses • They are determined by calculating the weighted mean. • Average atomic mass = Σ (isotope mass)( relative abundance)

  4. Isotope Calculations Example 1 • Using the isotope information for Silicon. Find the average atomic mass.

  5. Isotope Calculations Example 2 Silver consists of two isotopes 107Ag and 109Ag. Its average atomic mass is 107.87. Calculate the percentage of each isotope in naturally occurring silver. (Assume that the masses are 107.00 and 109.00 respectively.)

  6. Naming & Formula Writing

  7. Background: Periodic Table • Some of the groups in the periodic table are given special names. • These names indicate the similarities between group members: • Group 1: Alkali metals. • Group 2: Alkaline earth metals. • Group 16: Chalcogens. • Group 17: Halogens. • Group 18: Noble gases.

  8. Background: Molecules • Definition: a group of two or more atoms held together by a covalent chemical bond. • Typically a covalent bond is between two non-metals (This is a general rule of thumb.) • Examples: Water (H2O), Bromine (Br2), ammonia (NH3), Vinegar (HC2H3O)

  9. Background: Ions • Definition: An atom or group of atoms that have an overall positive or negative charge • Monatomic ions: atoms that have lost or gained electrons. • Charge related to position on Periodic Table for monoatomic ions • Cation: Positive ion (Typically metal) • Anion: Negative ion (Typically non-metal)

  10. Background: Formulas • Empirical formula: shows the lowest whole number ratio of the atoms in the compound. • Molecular formula: shows the exact number of each kind of atom in the compound. • Structural formula: shows how the atoms in the molecule are bonded together

  11. Background: Formula of Molecular Compounds

  12. H H C C C H C H C C H H Background Practice: Benzene ? ? Empirical Formula Structural Formula Molecular Formula

  13. Background: Predicting Charges

  14. I. Ionic Compounds • Are formed because of the strong electrostatic attraction between cations and anions. • Binary ionic compounds are always between a metal and a non-metal. • Other ionic compounds must contain a polyatomic ion • Examples: Table salt (NaCl), baking soda (NaHCO3), Epsom salts (MgSO4)

  15. Common Cation Charges

  16. Naming Ionic Compounds • Simply write the name of the cation first • Group 1,2 elements, Al3+, Zn2+, Ag1+, Ga3+, In3+ are simply named • Polyatomic cations are also simply named • Other metals can have more than one charge, so the name must indicate the charge with a roman numeral. • Cu1+ is copper(I) • Cu2+ is copper(II) • Then write the name of the anion • Polyatomic anions are simply named • Remember the name of a monatomic anion ends in –ide. • oxygen forms the anion oxide (O2-) • nitrogen forms the anion nitride (N3-)

  17. Lots of examples • KCl Potassium chloride • Mg 3N2 Magnesium nitride • Na2SO4 Sodium sulfate • (NH4)2CO3 Ammonium carbonate • CuO Copper(II)oxide • Cu2O Copper(I)oxide • FePO4 Iron(III)phosphate

  18. Formula Writing Ionic Compounds • Identify the compound as ionic • Find the formula and charge of the cation and the anion. • Use subscripts to indicate the number of each ion needed to have an overall neutral charge. “Drop and Swap” • Reduce the subscripts to the lowest whole number ratio.

  19. Ionic Naming Examples • Sodium fluoride Na1+F1- NaF • Calcium nitride Ca2+N3- Ca3N2 • Barium nitrite Ba2+NO21- Ba(NO2)2 • Lead(II)hydroxide Pb2+OH1- Pb(OH)2 • Manganese (IV) Sulfide Mn4+S2- Mn2S4 MnS2

  20. Now You Try

  21. Hydrates • Hydrates are compounds that contain discrete water molecules as part of the crystal lattice structure. • CuSO4•5H2O is called copper(II)sulfate pentahydrate. • You will use the Greek prefixes to indicate the number of water molecules in the compound.

  22. Prefixes for Hydrates

  23. Naming Polyatomic Ions With Oxygen

  24. Example • Selenate is SeO42- • What is selenite? • Answer: • Bromate is BrO3- • What is hypobromite? • Answer:

  25. Oxygen and Hydrogen Containing Polyatomic Compounds • Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi- (one H+), dihydrogen (two H+), etc., to the name as follows: • CO32- is the carbonate anion • HCO3- is the hydrogen carbonate (or bicarbonate) anion. • H2PO4- is the dihydrogen phosphate anion.

  26. II. Naming Acids

  27. II. Naming Acids • A helpful mnemonic for naming oxyacids I don’t feel well because I “ate” something “ic”ky! For example carbonate (CO32-) makes carbonic acid

  28. Practice With Acids • HCN • HNO3 • HNO2 • HClO4 • HClO3 • H2SO3 • HCl • HBr • HI

  29. III. Binary MolecularDiatomic Elements

  30. More about the Elements • Allotropes: Different forms of the same element • Some well know allotropes are: Carbon: Graphite, Diamond, “Bucky Balls” Oxygen: Oxygen gas, Ozone Tin: White (metallic tin), Gray Tin • Diatomic Elements: Elements that exist as molecules with two atoms.

  31. Formula • Empirical formula: shows the lowest whole number ratio of the atoms in the compound. • Molecular formula: shows the exact number of each kind of atom in the compound. • Structural formula: shows how the atoms in the molecule are bonded together

  32. Formula Writing and NamingBinary Molecular Compounds • Identify the molecular compound because there are two non-metals. • The most metallic element is usually written first (i.e., the one to the farthest left on the periodic table). Exception: NH3. • If both elements are in the same group, the lower one is written first. • Use prefixes to indicate the number of a particular atom in the compound. mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca. • Truncate the name of the last element and then add –ide • Example: NCl3 is nitrogen trichloride

  33. Now you Try

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