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Atoms and Elements

Atoms and Elements. Chapter 2. Atomic theory. John Dalton, 1808: matter is made of tiny indestructible particles called atoms What evidence persuaded John Dalton that matter was made of atoms?. Evidence for Atoms. Boyle’s Law (1660s) Gases can be compressed (PV = constant)

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Atoms and Elements

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  1. AtomsandElements Chapter 2

  2. Atomic theory • John Dalton, 1808: matter is made of tiny indestructible particles called atoms • What evidence persuaded John Dalton that matter was made of atoms?

  3. Evidence for Atoms • Boyle’s Law (1660s) • Gases can be compressed (PV = constant) • Suggests that a gas is made of particles with space between • Conservation of mass (Lavoisier, 1789) • In a chemical reaction, matter is neither created nor destroyed

  4. Evidence for Atoms • Law of Definite Proportions (Proust, 1797) • All samples of a compound, regardless of source or how prepared, have the same proportions of their constituent elements (constant composition) • Law of Multiple Proportions (Dalton, 1803) • When two elements (A and B) form two different compounds, the masses of B that combine with 1 g of A can be expressed as a ratio of small whole numbers.

  5. Law of Multiple Proportions

  6. Evidence for Atoms • Combining Volumes of Gases • When two gases combine to form a new compound, the volumes that combine will be in a ratio of small whole numbers.

  7. A New System of Chemical Philosophy: John Dalton, 1808 • Each element is composed of tiny indestructible particles called atoms. • All atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements. • Atoms combine in simple, whole-number ratios to form compounds. • Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms change the way that they are bound together with other atoms to form a new substance.

  8. Cathode ray tube

  9. Discovery of the electron • J.J. Thompson (1897) • Cathode rays are deflected by an electric or magnetic field • Cathode ray = beam of negatively charged particles (electrons) coming out of cathode metal atoms • Determined electronmass-to-charge ratio = –5.6857 x 10–9 g/C

  10. Millikan oil-drop experiment

  11. Charge of the electron • Robert Millikan (1909) • Drop charges are integer multiples of 1.60 x 10–19 Coulomb

  12. Radioactivity • Some elements spontaneously emit radiation called radioactivity • Three types of radioactivity: • Alpha (a) = heavy, +2 charge • Beta (b) = light, –1 charge • Gamma (g= high-energy photon

  13. The gold foil experiment • Ernest Rutherford (1910) • Bombarded gold foil with alpha particles • Most went straight through • Some slightly deflected • A few strongly deflected • “about as credible as if you had fired a 15-inch shell at a piece of tissue paper, and it came back and hit you”

  14. Alpha particle scattering explained

  15. 10-10 m 10-15 m Rutherford’s nuclear model • Mass & positive charge concentrated in nucleus • Most  particles miss the nucleus & are not deflected: most of the atom is empty space • Some  particles come near the nucleus & are deflected: nucleus is positively charged and very dense • Tiny, lightweight electrons circle the nucleus, like planets around the sun

  16. Composition of the atom • Protons • +1 charge • 1.67262 x 10–24 g = 1.0073 amu • Neutrons • 0 charge • 1.67493 x 10–24 g = 1.0087 amu • Electrons • –1 charge • 0.00091 x 10–24 g = 0.00055 amu

  17. Atomic number and mass number • Atomic number (Z) = protons • Mass number (A) = protons + neutrons • A 35Cl atom has 17 protons and 18 neutrons • Also shown as chlorine-35

  18. 23 11 Na1+ Ion charge • When an atom loses or gains an electron, it becomes anion • Ion charge is shown in the upper right corner of the atomic symbol • If no charge is shown, the charge is zero mass number, A protons + neutrons ion charge protons – electrons atomic number, Z protons

  19. Give the structure (p+, n0, e–)of these atoms • 9 p+, 10 n0, 9 e– • 19 p+, 21 n0, 18 e– • 92 p+, 143 n0, 92 e– • 16 p+, 18 n0, 18 e– 19F 40K1+ 235U 34S2–

  20. Give the symbol of each atom or ion • 35 p+, 44 n0, 36 e– • 79Br1– • 47 p+, 62 n0, 46 e– • 109Ag1+ • 26 p+, 28 n0, 26 e– • 54Fe

  21. Isotopes • Atoms with the same number of protons but different numbers of neutrons are isotopes of the same element • 6Li and 7Li are isotopes of lithium • Both are the element lithium • 6Li has 3 protons, 3 neutrons • 7Li has 3 protons, 4 neutrons

  22. A Mass Spectrometer

  23. Mass Spectrum of Neon 20Ne 90.48% 21Ne 0.27% 22Ne 9.26%

  24. Average atomic mass • Element’s atomic mass = weighted average of masses of all naturally-occurring isotopes of that element

  25. Average Atomic Mass • Calculate the average atomic mass of Ne: • ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE 20Ne 19.99244 amu 90.48% 21Ne 20.99395 amu 0.27% 22Ne 21.99138 amu 9.26%

  26. Average Atomic Mass • Calculate the average atomic mass of Ne: • ISOTOPE ISOTOPIC MASS (amu) ABUNDANCE 20Ne 19.99244 amu 90.48% 21Ne 20.99395 amu 0.27% 22Ne 21.99138 amu 9.26% (0.9048)(19.99244 amu) = 18.09 amu (0.0027)(20.99395 amu) = 0.057 amu (0.0926)(21.99138 amu) = 2.04 amu 20.18 amu

  27. Atomic mass • Two natural isotopes of antimony exist. 57.3% exists as 121Sb (mass 120.9038 amu), and the rest is 123Sb (mass 122.9041 amu). What is the atomic mass of antimony? • The abundances must total 100%,so abundance of 123Sb = 100 – 57.3 = 42.7% • (0.573 x 120.9038) + (0.427 x 122.9041) = 69.3 + 52.5 = 121.8 amu

  28. Percent abundance • Two natural isotopes of copper exist, 63Cu (62.9296 amu) and 65Cu (64.9278 amu). What is the abundance of each isotope? • The abundances must total 100% • If x = 63Cu abundance, then 65Cu = 1–x • Average mass (from periodic table) = 63.546 amu

  29. Percent abundance • x = 63Cu (62.9296 amu) & 1–x = 65Cu (64.9278 amu) • Average mass = 63.546 amu • 63Cu = 69.15% and 65Cu = 30.85%

  30. A unit for counting atoms • We have units for describing mass, volume, temperature, and so forth, but • we need a unit for counting the number of items • For large items like eggs or donuts, we can use dozen: one dozen = 12 items • For tiny items like atoms or molecules, we need . . . the mole!

  31. What is a mole? • One mole is defined as the number of atoms in exactly 12 grams of 12C • One mole contains 6.0221421 x 1023 particles • This number is called Avogadro’s number (NA) There is Avogadro’s number of particles in a mole of any substance

  32. Calculations with Avogadro • How many atoms of gold are present in 0.0507 mol Au? • How many moles of Pb atoms are 8.27 x 1022 atoms of Pb?

  33. Mass on the periodic table One atom of Cl weighs 35.45 amu OR One mole of Cl atoms weighs 35.45 grams The molar mass of Cl is 35.45 g/mol

  34. Calculations with molar mass • What is the mass of 1.38 mol Al? • How many moles are in 35 g of Zn?

  35. Avogadro and molar mass • What is the mass of 2.35 x 1024 atoms of Cu?

  36. Avogadro and molar mass • How many He atoms are present in a 22.6 g sample of He gas?

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