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Chapter 1: Part 2 Atomic theory

Chapter 1: Part 2 Atomic theory

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Chapter 1: Part 2 Atomic theory

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  1. Chapter 1:Part 2Atomic theory Introduction: Matter, Measurement and Molecules

  2. Atomic Theory The theory that atoms are the fundamental building blocks of matter came into being during the period 1803 to 1807 in the work of an English schoolteacher, John Dalton.

  3. Dalton’s Postulates • Each element is composed of extremely small particles called atoms. • All atoms of a given element are identical to one another in mass and other properties, but the atoms of a particular element are different from the atoms of all other elements. • Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions. This is the basis of the law of conservation of mass (or law of conservation of matter) which states that the total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process took place. • Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms. This is the basis of the law of constant composition (or law of definite proportions) which states that the relative numbers and kinds of atoms are constant, i.e. the elemental composition of a pure substance never varies.

  4. The Law of Multiple Proportions • Was deduced by Dalton from the preceding four postulates and states that: • If two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in the ratio of small whole numbers. Examples • H2O consists of 2 hydrogens to 1 oxygen • H2O2 consists of 1 hydrogen to 1 oxygen

  5. Cathode Rays and Electrons • Streams of negatively charged particles were found to emanate from cathode tubes. • J. J. Thompson is credited with its discovery in 1897. • He determined the charge/mass ratio of the electron to be 1.78 x 108 Cg-1.

  6. Cathode Ray Tubes (TV’s!) Cathode-Anode+ electons

  7. Millikan Oil Drop Experiment Once the charge/mass ratio of the electron was known, determination of either the charge or the mass of an electron would yield the other. Figure 1.19

  8. Millikan Oil Drop Experiment In 1909, Robert Millikan at the University of Chicago determined the charge on the electron.

  9. Radioactivity The spontaneous emission of radiation by an atom was first observed by Henri Becquerel. It was also studied by Marie and Pierre Curie.

  10. Radioactivity Three types of radiation were discovered by Ernest Rutherford •  particles •  particles •  rays Figure 1.21

  11. Discovery of the Nucleus Ernest Rutherford shot  particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

  12. The Nuclear Atom Some particles were deflected at large angles. This led Rutherford to postulate that the atom had a nucleus. Figure 1.22

  13. The Nuclear Atom • Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom. • Most of the volume of the atom is empty space.

  14. Subatomic Particles • Protons and electrons are the only particles that have a charge. • Protons and neutrons have essentially the same mass. • The mass of an electron is so small we ignore it. Table 1.5

  15. Mass Number (A) 12 6C Symbol of element Atomic Number (Z) Atomic Numbers, Mass Numbers and Isotopes The number of protons in the nucleus of an atom of any particular element is called that element’s atomic number (Z). The mass number (A) of an atom is the total number of protons plus neutrons in that atom.

  16. 11 6 12 6 13 6 14 6 C C C C Isotopes Atoms with identical atomic numbers (Z) but different mass numbers (A), or atoms with the same number of protons which differ only in the number of neutrons are called isotopes. Examples: carbon-14 isotope carbon-12 isotope

  17. Which Isotopes make up carbon in diamonds, graphite, petroleum and coal?Where does 14C originate and why is there so little in nature except for living things?

  18. Atomic Mass Atomic and molecular masses can be measured with great accuracy with a mass spectrometer. Figure 1.23

  19. Average Atomic Mass(commonly called Atomic Mass) • We use average masses in calculations, because we use large amounts of atoms and molecules in the real world. • Average atomic mass is calculated from the fractional abundance of each isotope and mass of that isotope. For example, the average atomic mass of C - made up mostly of 12C (98.93%) and 13C (1.07%) - is 12.01 u.

  20. Mass Spectrum of Chlorine Atoms How many protons does Chlorine have? a) What makes the rest of the masses? b) Why are there 2 different isotopes?

  21. Mass Spectrum of Chlorine Atoms What would the mass spectrum of Chlorine Molecules look like? a) What are the masses? b) Which is least common?

  22. Periodicity When one looks at the chemical properties of elements, one notices a repeating pattern of reactivities. Figure 1.25

  23. Periodic Table • A systematic catalogue of elements. • Elements are arranged in order of atomic number.

  24. Aluminum wrench, Copper pipe, Lead shot, Gold nuggetsBromine L+V vial, Iodine crystals, Carbon black + Diamond+ Graphite pencilSulfur

  25. Periodic Table • The rows are called periods. • Elements in each row are similar sizes. • The columns are called groups. • Elements in columns bond alike and • Have similar chemical properties.

  26. Groups Table 1.7 The above five groups are known by their names.

  27. Periodic Table Nonmetals are on the right side of the periodic table (with the exception of H). Metalloids border the stair-step line (with the exception of Al and Po). Metals are on the left side of the chart.

  28. Molecules and Chemical Formulae The subscript to the right of the symbol of an element tells the number of atoms of that element in one molecule of the compound. Notice how the composition of each compound is given by its chemical formula. Figure 1.29

  29. Diatomic Molecules Figure 1.28 These seven elements occur naturally as molecules containing two atoms. Which are gases?

  30. Molecular Compounds Molecular compounds are composed of molecules and almost always contain only nonmetals.

  31. Types of Formulae Empirical formulae give the lowest whole-number ratio of atoms of each element in a compound, e.g. HO. Molecular formulae give the exact number of atoms of each element in a compound, e.g. H2O2. Structural formulae show which atoms are attached to which within the molecule, e.g. H-O-O-H.

  32. Picturing Molecules Different representations of the methane (CH4) molecule. Figure 1.30

  33. Common Ionic Charges • When atoms lose or gain electrons, they become ions. • Cations are positive and are formed by elements on the left side of the periodic chart. • Anions are negative and are formed by elements on the right side of the periodic chart. Figure 1.31

  34. Ionic Compounds Ionic compounds (such as NaCl) are generally formed between metals and nonmetals. Figure 1.32

  35. Writing Ionic Formulae • Because compounds are electrically neutral, one can determine the formula of a compound by: • writing the value of the charge on the cation as the subscript on the anion. • writing the value of the charge on the anion as the subscript on the cation. Note: if the subscripts are not in the lowest whole number ratio, simplify it, e.g. Ca2O2 would become CaO.

  36. Chemical NomenclaturePositive Ions (Cations) a) Cations formed from metal atoms have the same name as the metal, e.g. Na+ is the sodium ion. b) If a metal can form different cations, the positive charge is indicated by a Roman numeral in parentheses following the name of the metal, e.g. Au+ is the gold(I) ion and Au3+ is the gold(III) ion. c) Cations formed from nonmetal atoms have names that end in -ium, e.g. NH4+ is the ammonium ion.

  37. Chemical NomenclatureCommon Cations Table 1.8

  38. Chemical NomenclatureNegative Ions (Anions) a) The names of the monatomic anions are formed by replacing the ending of the name of the element with -ide, e.g. O2-is the oxide ion, OH-is the hydroxide ion, N3- is the nitride ion. b) Polyatomic anions containing oxygen (called oxyanions) have names ending in -ate for the most oxidized form e.g. SO42-is the sulfate ion or –ite for the more reduced form, e.g. SO32- is the sulfite ion. c) Anions derived by adding H+ to an oxyanion are named by adding the prefix hydrogen e.g. HCO3-is the hydrogen carbonate ion (bicarbonate ion) or dihydrogen as in dihydrogen phosphate H2 PO4-2 .

  39. Chemical NomenclatureCommon Anions Table 1.9

  40. Common Oxy-anionNames, Formulae & Charges

  41. K-Iron-CNsalts:Left: KFe2+(CN)3 (Potassium ferrocyanide)Right:KFe3+(CN)4 (Potassium ferricyanide)

  42. Chemical NomenclatureMore on naming oxyanions Figure 1.34 Examples: ClO4-perchlorate ion (one more O atom than chlorate) ClO3-chlorate (most common oxidized form) ClO2-chlorite ion (one less O atom than chlorate) ClO-hypochlorite ion (one O atom less than chlorite) Which is the most oxidized ion? Which is most reduced?

  43. Figure 1.36 Chemical NomenclatureName and Formulae of Acids 1. Acids containing anions whose names end in -ide are named by changing the -ide ending to -ic, adding the prefix hydro- to this anion name, and then following with the word acid. • Acids containing anions whose names end in -ate or -ite are named by changing the -ate ending to -ic and the -ite ending to -ous and then adding the word acid.

  44. Chemical NomenclatureIonic Compounds Names of ionic compounds consist of the cation followed by the anion name, e.g. Cu(ClO4)2 is copper(II) perchlorate, and CaCO3 is calcium carbonate.

  45. Chemical NomenclatureBinary Molecular Compounds Table 1.10 It’s all Greek prefixes to me! 1. The name of the (more reduced) element farther to the left in the periodic table is written first (usually a metal), eg. MnO2 , NaCl, FeFe2 S4 , K(MnO4) . 2. If both elements are in the same group in the periodic table, the one having the higher atomic number (more reduced) is written first eg. SO2, SiC . 3. The name of the second element is given an -ide ( or –ite or –ate) ending. 4. Greek prefixes are used to indicate the number of atoms of each element. Examples N2O4 is dinitrogen tetroxide P4S10 is tetraphosphorus decasulfide.