Atomic Structure. SC Science Standards. Interpret Dalton’s atomic theory in terms of the Laws of Conservation of Mass, Constant Composition, and Multiple Proportions.
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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 whole-number 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
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
Au 197 79 Mass number 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.