History of the Atom - Tina's Science Class

History of the Atom - Tina's Science Class

History of the Ancient Greece The ancient Greek philosophers had two competing theories on what made up matter. Democritus and Epicurus believed that matter was made of indivisible and indestructible atoms. The word atom comes from the Greek Atomos, meaning uncuttable.

The competing theory was that all matter was made of a combination of air, water, earth, and fire. "Democritus of Abdera or Democrito " 1635 Peter Paul Aristotle opposed the idea of the atom, and supported the four elements view. Given the lack of evidence and experimentation, the elements

view was more widely accepted throughout Europe and the Muslim world. This lead to a 1600 year period sometimes called Dark Age Alchemy. The focus was on transmutation; the changing of one substance into another. The Alchemist, In Search of the Philosophers' Stone by Joseph Wright

John Dalton John Dalton was a college teacher who studied the movement and reaction of different gases. Daltons work measuring and identifying different elements lead him to develop the beginnings of the modern atomic theory. 1. Chemical elements are made of indivisible atoms. 2. All atoms of an element have identical masses. 3. Atoms from different elements have different

masses. 4. Atoms only combine in small, whole-number ratios such as 1:1, 1:2, 2:3, etc. 5. Chemical reactions involve the rearrangement of atoms. Daltons work gave rise to what is today called the Law of Definite Proportions. A given compound contains the same proportions of elements by mass regardless of the size of the sample. A tiny sample of sodium

chloride A massive spill from a Morton salt sodium chloride storage building. Dalton also discovered the Law of Multiple Proportions. Atoms of the same element can unite in more than one ratio with atoms of another element to form different compounds. Water, H2O 2:1 ratio of H

to O Hydrogen peroxide, H2O2 1:1 Ratio of H to O Up to this point, philosophers and scientists saw atoms as featureless balls that were indivisible. Nothing was known of

the subatomic particles that actually make types of atoms different from each other. Referred to as the billiard ball model. J.J. Thomson J.J. Thomson was studying the flow of electric current using Cathode Ray Tubes

(CRTs). Two metal plates with different electrical charges were placed on either end of a glass tube containing a vacuum. This created a beam, called a cathode ray. Thomson observed that when a small paddle wheel was placed in the path of the cathode rays, the wheel

was set into motion. This suggested that the cathode rays were made of tiny particles with mass. Thomson discovered that the particles of the cathode ray tube would respond to a magnet. A positively-charged magnet would attract the beam. A negatively-charged

magnet would repel the beam. Thomson had the first subatomic particle, negatively-charged electrons. Thomson proposed that the electrons were embedded in a positively-charged ball of matter. This is named the plum-pudding model of the atom, after the dessert it resembles.

The Nucleus Ernest Rutherford directed a beam of tiny, positively charged particles at a sheet of very thin gold foil. Most of the particles passed directly through the foil. A few were partially deflected. Rutherford hypothesized that a small, positivelycharged area within the

atoms of the gold foil deflected some of the particles. This area is now called the nucleus, and the positive particles within are protons. Since most of the particles passed through the foil, it was concluded that atoms are mostly empty space.

Rutherfords model of the atom had a positive nucleus surrounded by moving negative electrons. Referred to as the planetary model. The problem with this model was that the electrons did not collapse into the nucleus as would be

expected. Neils Bohr proposed that electrons only occupied certain energy levels, also called orbitals. Similar to rungs of a ladder. No electrons would appear between the rungs. Electrons closest to the nucleus have the lowest energy level. Electrons are able to jump between energy levels.

These quantum jumps result in the emission of light. Jumps between nearby energy levels produce low frequencies of light, which we perceive as red. Jumps between energy levels that are farther apart produce higher wavelengths of light, which we perceive as blue.

In-between distances can produce other colors, such as green. Electron Cloud Model The Bohr planetary model is inaccurate because electrons rapidly move around the atom in different-shaped orbitals, not just circles. The current, accepted model of the atom shows all the places

where electrons are the most likely to be found at any given moment. This is called the electron cloud model. The Quantum Mechanical Model Two physicists, Edwin Schrodinger and Werner Heisenberg, developed the electron cloud model. The electron cloud is

divided into 7 principal energy levels that are numbered 1 through 7. These correspond with the 7 periods on the periodic table and the different levels in the Bohr model of the atom. Think of these as the floors of the atom apartment building.

Each principal energy level is divided into sublevels, which indicate the overall shape of the orbital. s p d f

Each sublevel has a different number of possible orientations, called the magnetic quantum number. Think of each sublevel orientation as two-bedroom apartments within our atom apartment building. The s sublevel takes the shape of a sphere, so it only has one orientation and can only hold two electrons. The p sublevel takes the appearance of a dumbbell or

the infinity symbol. The p sublevel can be oriented in each of the three dimensions (x, y, z), meaning it can hold a total of 6 electrons. The d sublevel takes the appearance of a 4-leaf clover. There are five different orientations, each holding two electrons, meaning the d sublevel can accommodate a total of 1o electrons. The f sublevel has a total of 7 orientations, each holding two electrons, meaning it can hold a total of

14 electrons. The spin quantum number indicates what direction the electrons are spinning around their axis. Think of this as the specific bedroom within the magnetic quantum number apartment. Electron Orbital Notation For each element, an electron orbital notation can be written to indicate the location or address of each electron within the atom. Hydrogen has only one electron, and its notation is:

1S1 Each part of the notation relates to a specific quantum number: 1 is the principal quantum number and indicates the energy level . S is the azimuth quantum number (sublevel shape - sphere). 1 is the number of electrons occupying that sublevel. is the magnetic quantum number and indicates the spin direction. Helium has two electrons, and its notation is: 1S2

Each electron is in the first energy level and occupies a sphere-shaped orbital. Each electron spins in a different direction, shown by the arrows. Remember, each S sublevel can only fit two electrons, so this apartment is now full. At this point, it is helpful to divide the periodic table blocks to indicate which sublevel is being filled, and in which order.

Lithium has three electrons, 1S22S1 Beryllium has four, and now the 2S apartment is full. 1S22S2 Boron has five electrons. The fifth has to move into the first 2P apartment.

1S22S22P1 Unlike 1S and 2S, the 2P block has 3 apartments. By the time we get to neon, all of those are now full. 1S2 2S2 2P6 Notice that the second energy level is now completely full. The atom will have to add another floor to

accommodate more electrons. Rules of Electron Orbitals There are three rules of writing electron orbital diagrams: The Pauli Exclusion Principle states that no two electrons can occupy the same quantum state (bedroom). Incorrect: Correct: 1S22S2 1S22S2 The Aufbau principle states that electrons fill in

energy levels starting at the bottom (1st floor) and working their way up. Incorrect: 1S2 3S2 Correct: 1S22S2 Hunds rule states that electrons will occupy their own sublevel orientation (apartment) before sharing

one with another electron. Correct: 1S2 2S2 2P3 Incorrect: 1S2 2S2 2P3

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