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“ What are ATOMS ?”

“ What are ATOMS ?”. Leucippus and Democritus (400? BC). Greek philosopher and his “student” New idea : - there is a limit to how far matter can be divided - atomos : “uncuttable” particles, different shape/size. Robert Boyle (1627-1691) English scientist

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“ What are ATOMS ?”

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  1. “What are ATOMS?”

  2. Leucippus and Democritus (400? BC) • Greek philosopher and his “student” • New idea: - there is a limit to how far matter can be divided - atomos: “uncuttable” particles, different shape/size

  3. Robert Boyle (1627-1691) English scientist and philosopher Hypothesis: everything made of corpuscles (tiny particles of various sizes/shapes).

  4. Thinking about Atoms…

  5. Some Definitions: Matter Classification • Element: 1) A pure substance that cannot be broken down into simpler substances by a chemical reaction. 2) A substance made of unique, (nearly) identical atoms. • Compound: 1) A pure substance that can be broken down into simpler substances by a chemical reaction. 2) A substance made in a chemical reaction by combining two or more different elements.

  6. Start of Modern Era of Atoms:Dalton’s Atomic Theory John Dalton (1766-1844) British chemist, lecturer, and meteorologist

  7. Dalton’s Atomic Theory (1803) - 1 • All matter is made up of indivisible and indestructiblebasicparticles called atoms. • All atoms of a givenelement are identical, both in mass and in properties. Atoms of different elements have different masses and properties. • Compounds are formed when atoms of different elementscombine in the ratio of small whole numbers.

  8. Dalton’s Atomic Theory (1803) - 2 • Elements and compounds are composed of definite arrangements of atoms. Chemical change occurs when the atomic arrays are rearranged.

  9. Significance of Dalton’s Atomic Theory • Broke down 18th-century view of “elements” • Bridged gap between lab data and hypothetical atom. - way of calculating relative atomic weights. • Explained Law of Conservation of Mass - “Initial Mass = Final Mass” - Only reorganizing of unchangeable atoms occurs in chemical reaction.

  10. Dalton: inconsistencies uncovered… • The basic state of an element is one atom? Perhaps… basic naturalstate of an element may be a molecule made of 2 or more atoms. 2) Dalton: “Thou knows…no man can split the atom.” Not so: radioactivity. 3) Atoms of given element have same mass and properties? Not so: isotopes exist…

  11. More Definitions: • Molecule: - any combination of 2 or more atoms. - atoms can be same or different elements. • Allotropes: - different molecular combinations of a single element. Ex] oxygen as O2 or ozone O3 • Isotopes: (more later…) - atoms of same element that differ in mass only

  12. Identification of Elements • Physical properties • Chemical properties • Flame test for solids/solutions • Interaction with light: line-absorption spectrum line-emission spectrum

  13. Flame Test for Element Identification (From left) Sodium, potassium, lithium; strontium, barium, potassium.

  14. Elements: Ages of Discovery

  15. Classification of the Elements:Development of the Periodic Table • Dobereiner: “Triads” • Newlands: “Octaves” • Mendeleev: first-published “Period” definition • Meyer: second-published “Period” definition

  16. Dobereiner’s “Law of Triads” (1817) • Organized a few elements by similar chemical properties • Hypotheses: - Groups of three elements (triads) existed. - In order of increasing atomic weight, middle element’s properties were the average of other 2 elements’ properties.

  17. Newlands’ “Law of Octaves” (1863?) • Every eighth element repeats properties of first “Octave” element [like middle-C to high-C on piano]

  18. Dmitri Mendeleev (1834-1907) “Creator of the Periodic Table” (probably formulated periodic idea at same time as Meyer)

  19. Mendeleev’s early notes for the Periodic Table (1869)

  20. Characteristics of Mendeleev’s Table • Organized 60+ known elements… - in families (groups) with similar properties - by valence = “combining number” (split out elements with multiple valence) - roughly by atomic weight (moved 17 elements based on properties rather than weight) • Could use to predict existence of new elements (of 10, found 7; other 3 do not exist) • NOTE: at first, no “rare gases” were classified

  21. Comparison of eka-silicon’s predicted properties and known Group 4 properties Eka: “one beyond”

  22. Discovery of Atomic Structure;Sub-atomic Particles • Thomson: electron mass-to-charge ratio • Millikan: electron charge • Rutherford: mass and charge of nucleus • Chadwick: neutron (1932). (Nobel prize in 1935) • Bohr: electron energy levels (Topic 7)

  23. Joseph John Thomson (1856-1940) British physicist and mathematician Nobel Prize in 1906 (existence of electrons) 1897: calc’d electron’s charge-to-mass ratio in cathode-ray tube expt.

  24. Robert Millikan (1868-1953) U.S. physicist Nobel Prize in 1923 (charge of electron: 1909 oildrop expt.) With Thomson’s result, this allowed calculation of electron mass. Millikan’s experimental apparatus.

  25. Ernest Rutherford (1871-1937) nuclear physicist, New Zealander teaching in Great Britain Nobel Prize in 1908 (radioactive decay) Gold foil experiments

  26. Rutherford’s Experiments (1910-11)(done by undergrad Ernest Marsden/physicist Hans Geiger) • Fired beam of alpha particles [He+2] at gold foil. • Most particles went straight through, someweredeflected, BUT a few were reflected straight back to source!

  27. Rutherford’s Experiments (cont.) • Rutherford’s description: “It was about as incredible as if you had fired a 15-inch shell at a piece of paper and it came back and hit you.” • Interpretation: gold atom has small, dense, positively-chargednucleus surrounded by “mostly empty” space in which the electrons must exist. • Calculated nuclear mass as mass of positively- charged protons. Protons about half of actual mass: suggests neutral particles of same mass as proton?

  28. Known Properties of Subatomic Particles

  29. Proton/Electron Count in Atoms • Every neutral atom has an equal number of electrons and protons. • The number of protons, or “atomic number”, is unique for each element. • If an atom has unequal numbers of protons and electrons, it is called an ion (which is a charged atom). Example] Chlorine has atomic number 17. How many protons & electrons are in one Cl atom? How many protons & electrons are in one Cl- (chloride ion)?

  30. Neutron Count in Atoms: Isotopes • For many elements, the atoms of that specific element are only “nearly identical” to each other: the number of neutrons may vary. • The “nearly identical” forms are called isotopes. Example: Chlorine atoms exist in nature as a mixture of two isotopes: Cl-35 and Cl-37. 3 in four Cl atoms are Cl-35; 1 in four is Cl-37. The “atomic mass” of Cl in Table is an average.

  31. Modern “Periodic Table” Organization Elements are NOW placed in order of increasing atomic number (# of protons). - Why? Refer to “isotopes”: natural weight irregularities exist This relationship between nuclear charge and arrangement of elements in Table was discovered by Henry Moseley in 1914. In 1860s, Mendeleev / Meyer could NOT have predicted a relationship to subatomic particles!

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