Chemistry CLEP Chapter 1: The Structure of Matter

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Knowing everything mentioned in the notes provided is necessary in order to pass the CLEP Chemistry exam. Our summarized notes only contain material relevant to the CLEP Chemistry exam.

Chapter 1: The Structure of Matter

1.1: Evidence for Early Atomic Theory

John Dalton
•Claimed that all matter is composed of extremely small particles called atoms.
•Claimed that atoms cannot be created, destroyed, or divided.
•Suggested that elements combined in fixed ratios to form compounds.
 
J.J. Thompson
•Experimented with cathode ray tubes and discovered electrons
•Proposed the ”plum-pudding” model of the atom.

-Random (+) and (-) charge spread out throughout the atom

Ernest Rutherford
•Fired alpha particles through gold foil and noticed that most of these alpha particles, which are positively charged, went through the foil without any trouble whereas a few particles repelled. Because positive charges repel each other and this phenomenon was rare, Rutherford concluded that positive charge takes up only a small volume of the atom.
   –Positive charge is contained in a small, dense nucleus.
 
Max Planck
•Concluded that electromagnetic energy is quantized. E = h f
   –Where E = energy, f = frequency, and h = Planck’s constant.
      »In S.I. Units, h = 6.626 × 10^-34 m^2 kg / s
 
Niels Bohr
•Showed electrons exist around nucleus at fixed radius.
•Electrons with higher energy exist farther from the nucleus.
•The Bohr Model works well for hydrogen, not for other elements.
 
Werner Heisenberg
•Uncertainty Principle: Cannot know both momentum and position of a particle accurately. The more you know about position, the less you know about momentum and vice-verse.
 
Erwin Schrodinger
•His work on wavefunctions led to the idea of orbitals.
 
1.2: Atomic Number – Atomic Mass – Isotopes
 
Atomic Mass
•The mass of an atom is the sum of all the particles inside the atom including all protons, neutrons, and electrons. Because the mass of the electron is negligible compared to the mass of the neutron and proton, we can find the atomic mass by adding up the masses of the neutrons and protons that are in the atom.
 
•Example: A  helium atom consists of two protons and two neutrons. Its atomic mass is 4 amu (atomic mass units or the mass of one proton or neutron since their masses are nearly identical). It is referred to as helium-4.
 
Atomic  Number
•Number of protons in the nucleus of the atom.
•Number of electrons surrounding the nucleus for neutral atoms.
•Example: An atom of helium-4 has an atomic number of 2. This means that when the atom has neutral charge, it has 2 protons, 2 neutrons, and 2 electrons.
– Mass of electrons is negligible compared to the mass of neutrons and protons.
 
Isotopes
•Isotopes are atoms with the same number of protons as each other but different number of neutrons.
•For example, carbon-12 and carbon-14 are the same element, carbon, but carbon-14 has 2 more neutrons than carbon-12. Both carbon atoms must have 6 protons to be a carbon by definition.
 
Average Mass Number
•Elements naturally occur with more than one isotope.
•The average mass number on the periodic table is based on the relative frequencies of the different isotope so that on average the element weighs what the periodic table states.
•The average mass number = atomic weight.
•The average mass number of the element is also its molar mass. The molar mass is mass of one mole of atoms given in grams. One mole contains 6.02 x 10^23 atoms; that quantity is known as Avogadro’s Number.
•The atomic weight has a value close to the most frequently occurring isotope a given element.
•Sample Problem: Suppose lead exists as a combination of 4 common isotopes: Pb-204, Pb-206, Pb-207, and Pb-208. The relative abundance of each of the 4 isotopes is known.
–Pb-204  1.40%
–Pb-206  24.10%
–Pb-207  22.10%
–Pb-208  52.40%
 
 

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