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Atomic Theory

Atomic Theory. Defining the Atom. The Greek philosopher Democritus (460 B.C. – 370 B.C .) was among the first to suggest the existence of atoms (from the Greek word “atomos”) He believed that atoms were indivisible and indestructible

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Atomic Theory

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  1. Atomic Theory

  2. Defining the Atom • The Greek philosopher Democritus (460 B.C. – 370 B.C.) was among the first to suggest the existence of atoms (from the Greek word “atomos”) • He believed that atoms were indivisibleand indestructible • His ideas did agree with later scientific theory, but did not explain chemical behavior, and was not based on the scientific method – but just philosophy

  3. Dalton’s Atomic Theory (experiment based!) All elements are composed of tiny indivisible particles called atoms Atoms of the same element are identical. Atoms of any one element are different from those of any other element. John Dalton (1766 – 1844) Atoms of different elements combine in simple whole-number ratios to form chemical compounds In chemical reactions, atoms are combined, separated, or rearranged – but never changed into atoms of another element.

  4. Dalton’s Atomic Theory (1808) • Elements are composed of extremely small particles called atoms. All atoms of a given element are identical, having the same size, mass and chemical properties. The atoms of one element are different from the atoms of all other elements. • Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same. • Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions. 2.1

  5. 2 Law of Multiple Proportions 2.1

  6. 16 X + 8 Y 8 X2Y Law of Conservation of Mass 2.1

  7. Structure of the Nuclear Atom Dalton Wasn’t Exactly Correct….. • One change to Dalton’s atomic theory is that atoms are divisible into subatomic particles: • Electrons, protons, and neutrons are examples of these fundamental particles • There are many other types of particles, but we will study these three

  8. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

  9. J.J. Thomson, measured mass/charge of e- (1906 Nobel Prize in Physics) 2.2

  10. Cathode Ray Tube 2.2

  11. Thomson’s Atomic Model J. J. Thomson Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

  12. Modern Cathode Ray Tubes Television Computer Monitor • Cathode ray tubes pass electricity (electrons) through a gas that is contained at a very low pressure.

  13. Mass of the Electron Mass of the electron is 9.11 x 10-28 g The oil drop apparatus 1916 – Robert Millikan determines the mass of the electron: 1/1840 the mass of a hydrogen atom; has one unit of negative charge

  14. Measured mass of e- (1923 Nobel Prize in Physics) e-charge = -1.60 x 10-19 C Thomson’s charge/mass of e- = -1.76 x 108 C/g e- mass = 9.10 x 10-28 g 2.2

  15. (Uranium compound) 2.2

  16. 2.2

  17. (1908 Nobel Prize in Chemistry) • particle velocity ~ 1.4 x 107 m/s (~5% speed of light) • atoms positive charge is concentrated in the nucleus • proton (p) has opposite (+) charge of electron (-) • mass of p is 1840 x mass of e- (1.67 x 10-24 g) 2.2

  18. Rutherford’s Findings • Most of the particles passed right through • A few particles were deflected • VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions: The nucleus is small The nucleus is dense The nucleus is positively charged

  19. Rutherford’s Model of the Atom atomic radius ~ 100 pm = 1 x 10-10 m nuclear radius ~ 5 x 10-3 pm = 5 x 10-15 m “If the atom is the Houston Astrodome, then the nucleus is a marble on the 50-yard line.” 2.2

  20. The Rutherford Atomic Model • Based on his experimental evidence: • The atom is mostly empty space • All the positive charge, and almost all the mass is concentrated in a small area in the center. He called this a “nucleus” • The nucleus is composed of protons and neutrons (they make the nucleus!) • The electrons distributed around the nucleus, and occupy most of the volume • His model was called a “nuclear model”

  21. Chadwick’s Experiment (1932) H atoms - 1 p; He atoms - 2 p mass He/mass H should = 2 measured mass He/mass H = 4 · In 1932 James proved the existence of neutral particles in an atom · James said that the neutrons were just about the same weight as protons · He discovered this by using alpha rays, which are charged, and therefore repelled by considerable electrical forces present in the nuclei of heavier atoms · Chadwick led the way to the starting of penetrating and splitting the nuclei of atoms. · Also led the way to the fission of uranium 235, which eventually created the atomic bomb

  22. HOMEWORK 1. Describe JJ Thompson’s CRT (Cathode Ray Tube) experiment & how it showed that atoms contain particles he called “electrons.” 2. Describe JJ’s model of the atom. 3. Explain Rutherford’s scattering experiment and what it helped to prove. Also, how did it disprove Thompson’s model? • Describe Rutherford’s atomic model. • What led to Chadwick’s discovery.

  23. mass p+ = mass no = 1840 x mass e- Mass (amu) 0 1 1 2.2

  24. A X Mass Number Element Symbol Z Atomic Number 2 3 1 H (D) H (T) H 1 1 1 235 238 U U 92 92 Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes are atoms of the same element (X) with different numbers of neutrons in their nuclei 2.3

  25. 2.3

  26. 14 11 C C 6 6 How many protons, neutrons, and electrons are in How many protons, neutrons, and electrons are in ? ? Do You Understand Isotopes? 6 protons, 8 (14 - 6) neutrons, 6 electrons 6 protons, 5 (11 - 6) neutrons, 6 electrons 2.3

  27. Ions: Atoms with different number of protons (p+) and electrons (e-) 11 protons 11 electrons 11 protons 10 electrons Na+ Na 17 protons 18 electrons 17 protons 17 electrons Cl- Cl cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion.

  28. How many protons and electrons are in ? 27 3+ Al 13 Do You Understand Ions? 13 protons, 10 (13 – 3) electrons 78 2- How many protons and electrons are in ? Se 34 34 protons, 36 (34 + 2) electrons 2.5

  29. Practice: • Write the elemental symbol if 10 p+, 11no, 10e- 21Ne 2) Write the elemental symbol if 20p+, 20no, 18e- 40Ca2+ 2.5

  30. Noble Gas Halogen Alkaline Earth Metal Period Alkali Metal Group 2.4

  31. Chemistry In Action Natural abundance of elements in Earth’s crust Natural abundance of elements in human body 2.4

  32. H2 H2O NH3 CH4 A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds A diatomic molecule contains only two atoms H2, N2, O2, Br2, HCl, CO A polyatomic molecule contains more than two atoms O3, H2O, NH3, CH4 2.5

  33. 11 protons 11 electrons 11 protons 10 electrons Na+ Na 17 protons 18 electrons 17 protons 17 electrons Cl- Cl An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. 2.5

  34. A monatomic ion contains only one atom Na+, Cl-, Ca2+, O2-, Al3+, N3- A polyatomic ion contains more than one atom OH-, CN-, NH4+, NO3- 2.5

  35. 2.5

  36. 2.6

  37. molecular empirical H2O A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance H2O CH2O C6H12O6 O3 O N2H4 NH2 2.6

  38. ionic compounds consist of a combination of cations and an anions • the formula is always the same as the empirical formula • the sum of the charges on the cation(s) and anion(s) in each formula unit must equal zero The ionic compound NaCl 2.6

  39. 1 x +2 = +2 1 x +2 = +2 2 x +3 = +6 1 x -2 = -2 3 x -2 = -6 2 x -1 = -2 Formula of Ionic Compounds Al2O3 Al3+ O2- CaBr2 Ca2+ Br- Na2CO3 Na+ CO32- 2.6

  40. 2.6

  41. 2.7

  42. Polyatomic Ions

  43. WRITING FORMULAE (DIATOMIC & POLYATOMIC) Polyatomic 1) Potassium Nitrate 2) Sodium Sulfate 3) Potassium Dichromate 4) Ammonium Phosphate 5) Copper I Carbonate 6) Iron (III) Cyanide 7) Silver Sulfite 8) Tin (II) Nitrite 9) Calcium Hydroxide 10) Boron Acetate Mixed 11. Sodium Sulfide 12. Sodium Bicarbonate 13. Potassium Oxide 14. Potassium Permanganate 15. Magnesium Chloride 16. Calcium Chlorate 17. Iron (II) Phosphate 18. Copper Nitride 19. Ammonium Phosphide 20. Aluminum Acetate

  44. Chemical Nomenclature • Ionic Compounds • often a metal + nonmetal • anion (nonmetal), add “ide” to element name barium chloride BaCl2 potassium oxide K2O magnesium hydroxide Mg(OH)2 potassium nitrate KNO3 2.7

  45. Transition metal ionic compounds • indicate charge on metal with Roman numerals iron(II) chloride FeCl2 2 Cl- -2 so Fe is +2 FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide 2.7

  46. Hydrates: • ionic compounds with Water Molecules contained in their crystalline structure. CuCl2 · 2 H2O Copper Chloride Dihydrate Fe(SO3) · 5 H2O Iron (II) Sulfate Pentahydrate Cu(NO3)2 · 4 H2O Copper Nitrate Tetrahydrate 2.7

  47. Hydrates: • USE PREFIXES TO DENOTE THE # OF WATER MOLECULES 1 = mono 6 = hexa 2 = di 7 = hepta 3 = tri 8 = octa 4 = tetra 9 = nona 5 = penta 10 = deca

  48. Molecular compounds • nonmetals or nonmetals + metalloids • common names • H2O, NH3, CH4, C60 • element further left in periodic table is 1st • element closest to bottom of group is 1st • if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom • last element ends in ide 2.7

  49. TOXIC! Laughing Gas Molecular Compounds HI hydrogen iodide NF3 nitrogen trifluoride SO2 sulfur dioxide N2Cl4 dinitrogen tetrachloride NO2 nitrogen dioxide N2O dinitrogen monoxide 2.7

  50. 2.7

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