ATOMIC STRUCTURE SC Science Standards Interpret Daltons atomic

ATOMIC STRUCTURE

SC Science Standards • Interpret Dalton’s atomic theory in terms of the Laws of Conservation of Mass, Constant Composition, and Multiple Proportions. • Compare and contrast the contributions of Dalton, Thomson, Rutherford, Bohr, Planck and Schodinger to the development of the current atomic model. • Based on the quantum theory, write electron configurations and orbital notation for the representative elements. • Use Bohr’s model of the atom to explain the bright line spectrum in terms of electrons moving between energy levels. • Describe and identify the regions of the electromagnetic spectrum in terms of frequency, wavelength and energy.

Atomic Theory • Democritus…(>2000 years ago) • Greek philosopher • First suggested the idea of atoms • Matter is composed of tiny indivisible particles • Named these particles “atomos” • Now called atoms • Ideas lacked experimental support

Atomic Theory • John Dalton (1776 – 1844) • English school teacher • Studied chemistry • Particularly interested in meteorology • Performed experiments • Studied the ratios that chemicals combine to form compounds • Formulated the Atomic Theory

Dalton’s Atomic Theory • All elements are composed of submicroscopic indivisible particles called atoms • Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

Dalton’s Theory Cont… • Atoms of different elements can physically mix together or can chemically combine with one another in simple wholenumber ratios to form compounds. • Chemical reactions occur when atoms are separated, joined, or rearranged. However, atoms of one element are never changed into atoms of another element as a result of a chemical reaction.

What is an atom? • The smallest particle of an element that retains the properties of that element • Individual atoms are visible with the proper instrument

Subatomic Particles • Particles that are smaller than atoms • Three main subatomic particles • Protons • Neutrons • Electrons

Electrons • Negatively charged • Discovered by JJ Thomson in 1897 • Experimented with the flow of electric current through gases in cathode ray tubes • Found that the cathode rays were attracted to the metal plates with a positive charge and repelled by metal plates with a negative charge

Electrons cont… • Thomson • Concluded that cathode rays are composed of negatively charged particles • Called these negatively charged particles electrons • Concluded that electrons are a part of the atoms of every element • Electron has 1 unit of negative charge • Electron has mass of about 1/2000 of a hydrogen atom

Thomson’s Plum Pudding Model

Protons • Positively charged subatomic particle • Discovered by E. Goldstein in 1886 • Has one unit of positive charge

Neutrons • Discovered by Sir James Chadwick in 1932 • Subatomic particle with no electric charge • Mass is equal to the mass of a proton

Structure of the atom • Ernest Rutherford (1871 – 1937) • Performed famous gold foil experiment • Tested popular theory that atoms were composed of evenly distributed protons and electrons • Experimented with alpha particles (+ charges) aimed at a thin sheet of gold foil • Most particles went straight through • Some (a very few) were bounced back

Rutherford’s experiment • Rutherford proposed • Most of the mass and all of the positive charge of the atom is concentrated in a small region at the center of the atom • Called the center region the nucleus • The nucleus is the center core of the atom and is composed of protons and neutrons

The Nucleus • Very small and dense • If the nucleus were the size of a pea, its mass would be 250 tons! • Has a positive charge • Occupies a very small volume of the atom • Electrons occupy the largest volume of the atom outside of the nucleus

Atomic Number • Different numbers of protons make atoms different • Protons determine the identity of an element • Atomic number is the number of protons in the nucleus of the atom • Each element has a unique atomic number • Reported on the periodic table

Atoms • Atoms are electrically neutral • Number of protons must be equal to the number of electrons

Mass Number • Most of an atom’s mass is concentrated within the nucleus • Protons and neutrons contribute to the mass • Mass number = # protons + # neutrons

Mass number 197 79 Au Atomic number

Isotopes • Atoms of the same element may have different nuclear structures • Number of neutrons may vary within atoms of the same element • Isotopes are atoms that have the same number of protons, but different numbers of neutrons

Average Atomic Mass • Masses of atoms are measured in units called atomic mass units (amu) • An atomic mass unit is defined as 1/12 the mass of a carbon 12 atom • The mass of Carbon-12 is 12. 000000 amu • Mass of a single proton or neutron is approximately 1 amu

Atomic Mass • In nature, most elements exist as a mixture of 2 or more isotopes • Each isotope has a fixed mass and a natural percent abundance • Atomic mass is the weighted average mass of the atoms in a naturally occurring sample of the element

Calculating Atomic Masses • You need to know: • The number of stable isotopes of the element • The mass of each isotope • The natural percent abundance of each isotope • Masses and relative abundances are values that can be looked up in chemical reference books

Atomic Mass of Element X • Element X has two natural isotopes. The isotope with mass of 10. 013 amu has a relative abundance of 19. 90%. The isotope with mass 11. 0093 has a relative abundance of 80. 10%. Calculate the atomic mass of this element and name it.

Nitrogen mass number 14 15 exact weight percent abundance 14. 003074 15. 000108 99. 63 0. 37

Chlorine mass number exact weight 35 37 34. 968852 36. 965903 percent abundance 75. 77 24. 23

Silicon mass number exact weight percent abundance 28 27. 976927 92. 23 29 28. 976495 4. 67 30 29. 973770 3. 10

mass number exact weight percent abundance 24 23. 985042 78. 99 25 24. 985837 10. 00 26 25. 982593 11. 01

mass number exact weight 92 94 95 96 97 98 100 91. 906808 93. 905085 94. 905840 95. 904678 96. 906020 97. 905406 99. 907477 percent abundance 14. 84 9. 25 15. 92 16. 68 9. 55 24. 13 9. 63

mass number 112 114 115 116 117 118 119 120 122 124 exact weight percent abundance 111. 904826 0. 97 113. 902784 0. 65 114. 903348 0. 36 115. 901747 14. 53 116. 902956 7. 68 117. 901609 24. 22 118. 903310 8. 58 119. 902200 32. 59 121. 903440 4. 63 123. 905274 5. 79