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Unit 3 Language of Chemistry Part 1

Unit 3 Language of Chemistry Part 1

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Unit 3 Language of Chemistry Part 1

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  1. Unit 3Language of Chemistry Part 1 Zumdahl: Chapter 4 Holt: Chapter 3

  2. ATOMS: The Building Blocks of Matter Objectives Law of conservation of mass Law of definite proportions Law of multiple proportions Dalton’s Atomic Theory How Dalton’s Atomic Theory relates to 1, 2, & 3

  3. Atomic Theory Foundations Law of Conservation of Mass – mass is neither created or destroyed during a chemical or physical change Law of Definite Proportions – a compound contains the same proportions by mass regardless of the size of the sample Example: NaCl – always 39.34% Na & 60.66% Cl

  4. Atomic Theory Foundations Law of Multiple Proportions – if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers Example CO2 and CO : ratio of oxygen is always 2:1

  5. JJ Thomson Showed atoms could emit negative particles “plum pudding” model Electrons were embedded in a positively charged spherical cloud

  6. Rutherford • Shoots alpha particles (Helium atoms) at gold foil • Expected to pass right through • Particles are deflected • Leads to idea of a dense positively charged center with e- orbiting around it

  7. Ernest Rutherford • Gold foil experiment

  8. Dalton’s Atomic theory • All matter is composed of atoms • Atoms of an element have the same size, mass and properties; atoms of a different element have different sizes, masses and properties • Atoms cannot be divided, created or destroyed • Atoms of different elements combine in simple whole number ratios • Chemical reactions combine, separate, or rearrange atoms

  9. Modern Atomic Theory • Atoms can be divided • Atoms of the same element can have different masses • All else remains the same

  10. Structure of the Atom Objectives Discovery of the Electron Rutherford’s Experiments Protons, Neutrons, Electrons

  11. Atomic Structure • Electron = no mass; negative charge • Proton mass = hydrogen atom; positive • Neutron mass = hydrogen atom; no charge • Dalton’s Model • JJ Thompson’s Plum Pudding Model

  12. The Electron • Mass of 9.109 x 10-31 kg • Negative charge

  13. The Proton • Mass = 1.673 x 10-27 kg • Positive charge

  14. The Neutron • Mass = 1.675 x 10-27 kg • No charge

  15. Comparing Theories

  16. Isotopes II.A.2(c) – compare the characteristics of isotopes of the same element IV.B.2(d) – calculate the weighted average atomic mass of an element from isotopic abundance, given the mass of each contributor

  17. Isotopes Def: atoms of the same element that have different masses Example: hydrogen protium – 1 proton in nucleus deuterium – 1 proton; 1 neutron tritium – 1 proton; 2 neutrons *Nuclide – general term for any isotope

  18. Writing Isotopes Example – Uranium 235

  19. Average Mass Number Def: the weighted average of the atomic masses of the naturally occurring isotopes of an element Like calculating a “weighted” grade (decimal % of each isotope x mass of that isotope)

  20. Sample Calculation

  21. The Periodic Table IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements

  22. Atomic Number • Atomic Number • # of protons in nucleus • Element Symbol • Element Name • Atomic Weight • Electron Configuration 3 Li Lithium 6.941 [He]2s1

  23. Calculating Protons, Electrons & Neutrons IV.B.2(c) – use the periodic table to determine the atomic number, atomic mass, mass number, and number of protons, electrons and neutrons in isotopes of elements

  24. Mass Number Def: the number of protons and neutrons in the nucleus of an isotope mass # - atomic # = # of neutrons Example – oxygen Mass # (16) – atomic # (8) = # of neutrons (8)

  25. Formulas II.A.2(a) – use the IUPAC symbols of the most commonly referenced elements III.A.1(a) – distinguish between chemical symbols, empirical formulas, molecular formulas and structural formulas III.A.1(b) – interpret the information conveyed by chemical formulas for numbers of atoms of each element represented III.A.(d) – provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids

  26. Chemical Symbols

  27. Molecular Formulas

  28. Empirical Formulas

  29. Structural Formulas

  30. Percent Composition III.A.1(c) – calculate the percent composition of a substance given its formula or masses of each component element in a sample

  31. % Composition Sample Problems

  32. Zumdahl – Ch 5 Holt – Ch 7 III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids NOMENCLATURE

  33. Oxidation Numbers • Pure element have an oxidation number of zero • Fluorine has an oxidation of -1 in all compounds • Oxygen has an oxidation number of -2 except when in a peroxide when its oxidation number is -1 • Hydrogen has an oxidation number of +1 except when bonded to a metal • The algebraic sum of the oxidation numbers of all atoms in a neutral compound is zero • The algebraic sum of the oxidation numbers of all atoms in a polyatomic ion is equal to the charge of the ion • These rules apply to covalently bonded atoms

  34. Oxidation Numbers Cont’d • Group 1 = +1 • Group 2 = +2 • Transition Metals = varies • Group 15 = varies (-3 usually) • Group 16 = varies (-2 usually) • Group 17 = varies (-1 usually)

  35. General Rules for Naming Chemical Compounds • Cation (pronounced cat-ion) • is named first • has a positive charge/oxidation number • name does NOT change • Anion (pronounced an-ion) • is named second • has a negative charge/oxidation number • ending is “-ide”

  36. Binary Ionic CompoundsType I • Metal is from Group 1 or 2 Examples: NaCl –sodiumchloride KF –potassiumfluoride CaBr2–calciumbromide Li2O –lithiumoxide

  37. Binary Ionic Compounds Type II • Metal is a transition metal • Many metals form more than one type of positive ion • Iron: Fe2+ or Fe3+ • Copper: Cu1+ or Cu2+ • Lead: Pb2+ or Pb4+ • Tin: Sn2+ or Sn4+ • Roman numeral indicates charge on the metal ion (NOT the number of atoms) • Look at 2nd element to determine charge on transition metal

  38. Examples of Type II EX: FeCl2 –iron(II) chloride FeCl3–iron(III) chloride CuO –copper(II) oxide Cu2O –copper(I) oxide SnF4–tin(IV) fluoride SnO2–tin(IV) oxide SnO –tin(II) oxide **NOTE: sum of oxidation numbers = 0

  39. Exceptions to the Rule Aluminum – always Al3+ Cadmium – always Cd2+ Zinc – always Zn2+ Silver – always Ag+ No Roman numerals needed

  40. Binary Covalent Compounds orBinary Molecular Compounds • Formed between two non-metals • first element name doesn’t change • Second element ends with “-ide” • Don’t use “mono-” on first name • Use prefixes to indicate number of atoms of element

  41. Mono– 1 Di– 2 Tri- 3 Tetra- 4 Penta- 5 Hexa- 6 Hepta- 7 Octa- 8 Nona- 9 Deca- 10 Undeca – 11 Dodeca - 12 Prefixes

  42. Examples of Binary Covalent Compounds CO –carbonmonoxide CO2–carbondioxide CCl4–carbontetrachloride SiO2–silicondioxide SeBr2–seleniumdibromide BF3–borontrifluoride P2O4–diphosphoroustetroxide P4O10–tetraphosphorousdecoxide

  43. Polyatomic Ions • Group of atoms with a shared charge III.A.1(c) – use the names, formulas, and charges of commonly referenced polyatomic ionsl

  44. Polyatomic Rule #1 • Common form ends with –ate Example:nitrate, chlorate, sulfate

  45. Polyatomic Rule #2 • If oxygen decreases by one, O-1, changes ending to “-ite” Example: nitrite, chlorite, sulfite

  46. Polyatomic Rule #3 • If oxygen decreases by two, O-2, add prefix “hypo-” • Keep ending “-ite” Example: hypochlorite

  47. Polyatomic Rule #4 • If oxygen increases by one, O+1, add prefix “per-” • Keep ending “-ate” Examples: perchlorate, permanganate

  48. Chlorate Bromate Iodate Nitrate Permanganate Carbonate Silicate Selenate Phosphate Arsenate ClO3- BrO3- IO3- NO3- MnO4- CO32- SiO32- SeO32- PO43- AsO43- Polyatomic Ions

  49. Acetate Hydroxide Bicarbonate Bisulfate Ammonium Thiocyanate Thiosulfate Oxalate Chromate Dichromate C2H3O2- OH- HCO3- HSO4- NH4+ SCN- S2O32- C2O42- CrO42- Cr2O72- More Polyatomic Ions

  50. Binary Acids III.A.1(d) - provide the interconversion of molecular formulas, structural formulas, and names, including common binary and ternary acids