Atomic Structure
Smallest particle of an element w/ properties of that element About 90 natural elements Combine to form compounds ◦ Atoms elements compounds
460 – 370 BC: Democritus ◦ “atomos” (indivisible & indestructible) ◦ Greeks believed in 4 basic elements ◦ Accepted for 2000 years
Law of Conservation of Matter ◦ Matter is neither created nor destroyed
Law of Definite Proportions ◦ A compound always contains elements in a certain, definite proportion & no other combinations ◦ Ex. H2OH2O CO HCl
Law of Multiple Proportions ◦ Elements may combine in more than one proportion forming different compounds ◦ Led to the atomic theory 1. all matter is composed of atoms which are indivisible 2. all atoms of the same element are alike, atoms of different elements are different
3. Compounds are formed when atoms of different elements combined in fixed proportions 4. Chemical reactions involve a rearrangement of atoms Model:
Discovers proton using cathode ray tube & noticing that rays traveled in opposite direction of cathode (-) Mass is 1840x’s that of an electron
Discovered electron using cathode ray tube where flow goes from cathode (-) to anode (+) Millikan in 1916 found mass(1/1840 x’s mass of H atom) & charge(-1) Model:
Discovers a condensed core w/ gold foil experiment Called it the ‘nucleus’ & composed of protons & neutrons Model known as “Nuclear Atom” Model:
Discovers subatomic particle w/ no charge aka Neutron Mass nearly equal to proton (slightly larger)
Protons ◦ Positive ◦ In nucleus ◦ Mass = 1840 x’s an electron ◦ Same as atomic # Electrons ◦ Negative ◦ Outside nucleus ◦ 1/1840 mass of proton Neutrons ◦ Neutral ◦ In nucleus ◦ Almost the same as proton, slightly larger
Mass # = #protons + # neutrons Mass # for element & # neutrons may vary Ex. ◦ 14 C 6 p +, 8n o ◦ 12 C 6 p +, 6n o
Increase/decrease in # of electrons In atom: ◦ Protons = electrons Ex. Na ◦ 11p +, 11e - Na + ◦ 11p +, 10e -
Found using mass spectrometer Based on average of relative abundance of isotopes of the element Measured in amu’s (1/12 of a carbon atom)
Number of stable isotopes % of each isotope Mass of each isotope