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Atomic Theory

Atomic Theory. Past to Present. Aristotle : Four Element Theory. theory lasted for about 2000 years. not a scientific theory because it couldn’t be tested against observation. Democritus. In 300 BC, Democritus said atoms were indivisible particles.

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Atomic Theory

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  1. Atomic Theory Past to Present

  2. Aristotle: Four Element Theory • theory lasted for about 2000 years. • not a scientific theory because it couldn’t be tested against observation

  3. Democritus In 300 BC, Democritus said atoms were indivisible particles. His theories were the first mention of atoms (atomos). Not a testable theory, but a conceptual model. • no mention of any atomic nucleus or its constituents. • cannot be used to explain chemical reactions

  4. Lavoisier (late 1700s) • law of conservation of mass (the mass of substances produced from a chemical reaction is always equal to the mass of the reactants.) • law of definite proportions (Ex. water is always 11% hydrogen and 89% oxygen)

  5. Proust (1799) • if a compound is broken down into its constituents, the products exist in the same ratio as in the compound. • experimentally proved by Lavoisier laws.

  6. Dalton (early 1800s) • atoms are solid, indestructible spheres (like billiard balls) • provides for different elements (these would be different spheres) • based on the Law of Conservation of Mass • having a molecule (atoms combine in simple whole number ratios) explains the law of constant composition

  7. JJ Thompson (1850s) • raisin bun model • solid, positive spheres, with negative particles embedded in them • 1st atomic theory to include protons (+) and electrons (-)

  8. Rutherford (1905) • showed that atoms have a positive, dense center with electrons outside it. • resulted in planetary model • explains why electrons spin around nucleus • suggests atoms are mostly empty space

  9. Bohr Models

  10. Bohr’s Theory • electrons exist in outer shells of an atom • when electrons absorb energy —> they move to a higher orbital or shell • when electrons release energy as a photon of light —> they fall from a higher shell to a lower one • Bohr based his model of the energy or light emitted by different atoms

  11. each atom has a different spectra of light • each element has a set of lines which represents a photon of light emitted from the atom

  12. prisms separate white light into colours • AS WELL: • matter emits light when it’s heated • light travels as protons • the energy photons carry depends on their wavelength

  13. Separating Mixtures

  14. Physical Separations • do not involve chemical reactions • classified as Mechanical (uses gravity, contact forces or motion) or Non-mechanical (uses heat, magnetism, electricity, dissolving) • Decanting (sedimentation): carefully pour • off liquid and leave sediment • Centrifugation: spinning to • enhance density separation

  15. Filtration: separation based on particle size • Chromatography: a flowing liquid or gas (mobile phase) carries substances at different rates through a stationary phase, e.g. gas chromatography, paper chromatography

  16. Electrophoresis: electric current used to separate particles by surface charge

  17. Distillation: a mixture of liquids is heated until components vaporize. The gas is then condensed back into liquid form and collected; uses difference in BP to separate mixture

  18. Other examples: evaporation, magnetic separation, reverse osmosis, sifting, skimming Recrystallization: rock candy!

  19. A Review of Chemical Nomenclature

  20. Binary Compounds • A binary compound contains the atoms of only two different elements, and a binary ionic compound contains only two types of monatomic ions (charged atoms), e.g. NaCl. • Cation: positively charged ion (Trick: think of ‘t’ in ‘cation’ as +), e.g. Pb4+ • Anion: negatively charged ion, e.g. S2- ***The name of any ionic compound is the name of its metal ion followed by the name of its non-metal ion, e.g. sodium chloride.***

  21. PbS2 • Compounds always have a net charge of 0. They are always neutral. • \ the charges of the ions making up a compound must negate each other (cancel each other out)! 2Al3+ (aq) + 3S2- (aq)  Al2S3 (s) 2(3+) 3(2-) 6+ 6- 0 • Try It: Pb4+ and S2- combine to make ______________ • Note: reduce to lowest terms!

  22. Ion charge of elements depends on group number (column). Elements not included have more than one ion configuration.

  23. Multivalent Ions • Multivalent elements have more than one way of arranging their outer (valence) electrons. They have more than one stable ion form. • e.g. Fe2+ and Fe3+are known as iron(II) and iron(III) ions • The roman numeral in brackets corresponds to the charge on the ion and is pronounced as the number it represents, i.e. iron(II) is pronounced “iron two”.

  24. Example 1: What is the formula for tin(IV) sulphide? • Example 2: What is the name of Fe2S3? Iron(III) oxide!

  25. Do p. 98 Practice Problems Solutions 1a) Li2S c) AlCl3 e) SnI2 b) Cr2O3 d) PbS f) ZnBr2 2a) zinc oxide d) sodium iodide b) lead(IV) chloride e) potassium sulfide c) copper(II) chloride f) chromium(II) oxide

  26. Polyatomic Ions • Polyatomic ion: a charged group of covalently bonded atoms (similar to a molecule but with a charge). • Relatively stable, often remain intact in chemical reactions • Many are oxyanions: an atom of one element bonded to some number of oxygen atoms. • E.g. nitrite NO2- nitrate NO3- sulfite SO32- sulfate SO42- • lesser number of O atoms takes the –itesuffix • greater numer of O atoms takes the –ate suffix

  27. More than two oxyanions in a series, e.g. • ClO- hypochlorite (hypo- meaning “less than”) • ClO2- chlorite • ClO3- chlorate • ClO4-perchlorate (per- meaning “more than”) • The prefix bi- before the name of a polyatomic ion indicates an H+ was added to it, e.g. • CO32- carbonate ion becomes HCO3- bicarbonate ion (hydrogen carbonate)

  28. Exceptions: • OH- is hydroxide (only polyatomic ion with the –idesuffix) • Cr2O72- is called dichromate ion • Most polyatomic ions we will encounter are anions (negatively charged) with the exception of ammonium, NH4+ • A table of polyatomic ions will be provided to you. • Polyatomic ions are bracketed in ionic compounds

  29. Example 1: What is the formula of potassium sulphite? • Example 2: What is the name of Cr(HSO4)2?

  30. Do Practice Problems, p. 100, 1-2 1a) BaSO4 d) Sn(C2O4)2 b) AgNO3 e) Al2(Cr2O7)3 c) HgBr2 f) KF 2a) zinc hydroxide d) sodium acetate b) tin(II) oxide e) magnesium iodide c) copper(II) hypochlorite f) iron(II) dichromate

  31. Binary Molecular Compounds • Molecular compounds are made up of two non-metal atoms which share electrons to become stable compounds. • The names of all binary compounds have an –ide suffix. • Prefixes are used to indicate the number of atoms. • The term mono- is understood but not used for the first element in the pair, e.g. CO2 is carbon dioxide, not monocarbon dioxide.

  32. Example: What is the formula for xenon tetrafluoride? Example: What is the name of P4S10? Prefixes for Molecular Compounds XeF4 4 phosphorous and 10 sulphur, so… tetraphosphorousdecasulphide.

  33. Do Practice Problems on p. 101 1a) NO c) N2O4 b) NO2 d) N2O3 2a) phosphorous pentachloride b) sulphur dioxide c) carbon monoxide d) diphosphorouspentaoxide

  34. Hydrates • When many salts crystallize out of aqueous solution they incorporate water molecules in a fixed ratio and pattern into their ionic crystal lattice. These are called hydrates. • Since most salts are destined to be dissolved in water, this does not present a problem. • Water is an integral part of the hydrate, so it must be accounted for in the name. We use the prefixes for molecular compounds combined with the term –hydrate to indicate the number of associated water molecules.

  35. Gently warming a hydrated salt will remove the associated water from the crystal. • Anhydrous refers to the salt without the water. • Dessicants are anhydrous salts that can absorb water from the air (hygroscopic) to form hydrates. They are used to… • Keep moisture out of dried foods, e.g. beef jerky • Dry flowers for crafts, e.g. silica salt beads (silica gel) • Keep running shoes, guitars, etc. dry in their boxes before they are purchased

  36. Crystals of CuSO4·5H2O Anhydrous CuSO4 powder

  37. Example 1: What is the formula for copper(II) sulphate heptahydrate? • Example 2: What is the name of NaCH3COO, 3H2O?

  38. Do practice problems on page 102. 1a) BaCl2 , 2H2O b) Na2CO3 , H2O c) Fe(NO3)3 , 9H2O d) Ba(OH)2, 8H2O 2a) cobalt chloride hexahydrate b) iron(III) chloride tetrahydrate c) sodium dichromate dihydrate d) magnesium sulphate heptahydrate

  39. Acids • Acids are a special type of molecular compound that can be induced to form ions. Their names are based on the name of the anion formed and whether it contains oxygen. • Acids can be thought of as one or more H+ ions bonded to an anion.

  40. No oxygen: prefix hydro- precedes the name of the ion and the suffix –ic replaces the –ide in the anion’s name • hydro_________ic acid • Contains oxygen: suffix –ic replaces –ate in the ion’s name or the suffix –ous replaces –ite • e.g. hydrogen sulphate is sulfuric acid and hydrogen sulphite is sulphurous acid

  41. Note: The term ‘acid’ can refer to the compound or to its solution. • H2SO4(l)and H2SO4(aq) are both called sulphuric acid

  42. Example 1: What is the formula for hydrobromic acid? • Example 2: What is the name of HNO2?

  43. Test on Oct. 23, 2012 1a) HFl b) HClO c) H3PO4 d) H2S 2a) acetic acid c) carbonic acid b) sulphurous acid d) hydroiodic acid • Do Practice Problems on page 103. • Complete the Review questions on p. 105-106 • Study questions from p. 70-72, p. 83-85, p. 94-95

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