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History of Chemistry and Atomic Structure

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History of Chemistry and Atomic Structure

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  1. History of Chemistry and Atomic Structure Unit 3

  2. Early developments • Chemistry has been around ever since people started working with dyes (more than 20,000 years ago) and hardening clay for ceramics (8,000 years ago). • Later people worked with (pure) metals (Iron Age, Bronze age; 4,000 – 1,000 BC).

  3. Early ideas • In the 7th century BC Greek philosophers were the first to developed an idea about matter: it is composed of four elements: earth, water, air, and fire. • They also speculated whether one compound could be converted into another. • “Chemistry” is thought to be derived from the Greek work χημεια (khemeia) meaning “to transform one substance into another”.

  4. Early Models • The Greek philosopher, Democritus (460 B.C.-370 B.C.), was among the 1st to suggest the existence of atoms (ατομος): indivisible small particles. • His ideas were not based on scientific method, but they did agree with later scientific theory. • However, his ideas were not universally accepted.

  5. Dark Ages • The end of the Golden Age of Greece (200 BC) ended scientific development for a long time. • Alchemy (elixir of life in the East; gold out of lead and other metals in the West) was common practice up until the 16th- 17th century. • Alchemy did lead to the discovery of gunpowder (China) and better purification methods for metals and alloys (mixed metals).

  6. New developments • In the 1660’s Irishman Robert Boyle, while working with gases, revived the idea of atoms. He also argued that ideas about chemistry should be based on evidence. • In the same time German Hennig Brandt was the first to isolate a non-metallic element, phosphorus.

  7. Antoine Lavoisier (1743-1793) • Showed that hydrogen and oxygen combined to form water and that oxygen is involved in combustion. • He also showed that no mass was lost. • Law of Conservation of Matter: matter can not be destroyed nor created.

  8. Antoine Lavoisier (1743-1793) • Lavoisier demonstrated the importance of carefully making and recording all measurements. • He even wrote the first chemistry text. • For all his contributions, Lavoisier is often referred to as the Father of Modern Chemistry.

  9. John Dalton (1766-1844) • In the early 1800’s an English school master, John Dalton, used the work of Lavoisier and others, and his own research to develop the first modern atomic theory. • Part of Dalton’s work was developing the law of multiple proportions. • According to this law, different compounds made of the same elements, have mass ratios related by small whole numbers.

  10. John Dalton’s Atomic Theory • All elements are composed of atoms. • Atoms of the same element are identical. Atoms of any one element are different than atoms of another element. • Atoms of different elements can physically mix together or can chemically combine in whole number ratios to form compounds.

  11. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element. • Dalton’s theory was accepted, unchanged, for nearly a century, when the existence of the first subatomic particle is established.

  12. Subatomic particles • In the late 1800’s and early 1900’s it became clear that atoms are divisible into smaller particles: • Electrons • Protons • Neutrons

  13. Electrons • Discovered by J.J. Thomson in 1897 using a cathode ray tube. • Electrons are negatively charged subatomic particles. • Robert Millikan measured the charge of the electron in 1916.

  14. Protons and Neutrons • Eugene Goldstein observed positively charged particles (protons) in 1886. • Protons have a mass 1840 times that of an electron. • In 1932, James Chadwick discovered the existence of a subatomic particle (neutrons) with no charge, but a mass slightly larger than the mass of a proton.

  15. The Atomic Model • When the subatomic particles were discovered, scientists wondered how they were put together in an atom. • Thomson believed that the particles were all distributed evenly in what became to be known as the “plum pudding model”.

  16. Plum Pudding Model

  17. Gold foil experiment • Ernest Rutherford, a former student of Thomson, is credited with coming up with an improved model of an atom. • In 1911 he did his gold-foil experiment. Using positively charged alpha particles, he found that most of these went right through the gold foil. Only few particles were redirected.

  18. Rutherford concluded that the atom is mainly empty space with all the positive charge and almost all the mass centrally located in what he called the nucleus. The nucleus contains the protons and neutrons, with the electrons distributed around the nucleus.

  19. Subatomic Particles

  20. amu • The atomic mass unit (amu) is a relative mass value. • It is equal to 1/12 the mass of the carbon-12 nucleus. • The electron actually has a mass of 9.11 x 10-28g. • This mass is insignificant (1/1840) compared to that of the proton, or neutron, so it is given a value of 0 amu.

  21. Distinguishing among Atoms • Elements are different because they have different numbers of protons. • The atomic number of an element is the number of protons in the nucleus of an atom of that element. • Since atoms are electrically neutral, the number of electrons (-) must equal the number of protons (+).

  22. Atomic Number on the Periodic Table

  23. Mass number is the total number of protons and neutrons in an atom. • # of neutrons = mass # - atomic # Example: Oxygen Mass # of 16 Atomic # of 8 16 – 8 = 8 neutrons

  24. Atomic Symbol

  25. Write the atomic symbols for atoms with the following: A. 8 p+, 8 n, 8 e- ___________ B. 17p+, 20n, 17e- ___________ C. 47p+, 60 n, 47 e- ___________

  26. Isotopes • Atoms that have the same number of protons but different numbers of neutrons. • Because isotopes have different numbers of neutrons, they have different mass numbers. • Atomic mass is a weighted average mass of the atoms in a naturally occurring sample of the element.

  27. Isotope Symbol

  28. Example: Carbon Carbon-12 Carbon-13 Carbon-14 # p # n # e Symbol

  29. Calculating atomic mass • Multiply the atomic mass of an isotope by its percentage/100. • Do this for every isotope. • Add the atomic masses found.

  30. Isotopes • Gallium is a metallic element found in small lasers used in compact disc players. In a sample of gallium, there is 60.2% of gallium-69 (68.9 amu) atoms and 39.8% of gallium-71 (70.9 amu) atoms. What is the atomic mass of gallium?

  31. Isotopes Ga-69 68.9 amu x 60.2 = 41.5 amu for 69Ga 100 Ga-71 70.9 amu x 39.8 = 28.2 amu for 71Ga 100 Atomic mass Ga = 69.7 amu

  32. Lead has four different isotopes. 204Pb has a mass of 203.994 amu and an abundance of 1.32%. 206Pb has a mass of 205.993 and an abundance of 26.31%. 207Pb and 208Pb have masses of 206.991 and 207.899 amu and abundances of 20.78% and 51.59%, respectively. Calculate the relative (average) atomic mass of lead.

  33. 203.994 amu x .0132 = 2.69 amu • 205.993 amu x .2631 = 54.20 amu • 206.991 amu x .2078 = 43.01 amu • 207.899 amu x .5159 = 107.3 amu • 207.2 amu