Atomic Theory And Structure

Question 0

Thompson

Answer 0

J.J. Thompson observed the deflection of particles in a cathode ray tube.

Concluded that atoms are composed of positive and negative charges.

Called negative charges electrons, and he suggested that the positive charges were distributed in islands throughout the atom, like raisins in raisin bread. The ‘plum pudding’ model of the atom.

Question 1

Dalton

Answer 1

Atoms are never created or destroyed during chemical reactions.

Proposed that all matter is composed of subunits called atoms. Atoms had different identities called elements. Elements combined in definite ratios to form compounds.

John Dalton performed chemical reactions and carefully measured the masses of reactants and products.

Question 2

Millikan

Answer 2

Robert Millikan used oil drops falling in an electric field of known strength to calculate the charge-to-mass ratio of electrons and to surmise the charge contained by a single electron.

Question 3

Rutherford

Answer 3

Ernest Rutherford fired alpha particles, that he knew to be positively charged, through thin gold foil.

Measuring the resulting scatter patterns he found that most past through but some went off at odd angles as if they collided with a heavier object.

Rutherford concluded that the positive charge and the mass of the atom are concentrated at the center of the atom, and the rest is mostly empty space- countering Thompson’s ‘plum pudding’ model.

Question 4

Planck

Answer 4

Max Planck determined that electromagnetic energy is quantized, or composed of discrete bundles, expressed by the equation:

E = h v
or
E = hc/λ

E- energy of photon, J
h- Planck’s constant, 6.63•10(to the negative 34th) J•sec
v- frequency of light, sec(to the neg one power)
λ- wavelength of light, m
c- speed of light, 3.00•10(to the 8th power) m/sec

Question 5

Bohr

Answer 5

Niels Bohr applied the idea of quantized energy to show that electrons exist around the nucleus at a fixed radius.

Electrons with higher energy exist farther from the nucleus.

The Bohr model is accurate only for atoms and ions with one electron. It was clear that a more complex model was needed to explain atoms with multiple electrons.

Electrons give of energy in the form of electromagnetic radiation when they move from a higher level, or an excited state, to a lower level. The energy represented by light, using Planck’s equation, represents the difference between the two energy levels of the electron.

Question 6

DeBrogile

Answer 6

Louis deBrigile identified the wave characteristics of matter by combining Einstein’s relationship between mass and energy (E=mc-squared) and the relationship between velocity and the wavelength of light (E=hv).

This shows that all particles with momentum have a corresponding wave nature.

DeBrogile Wavelength of Particles:
λ = h/mv

λ- wavelength associated with particle, m
m- mass of particle, kg
h- Planck's constant
v- velocity of particle, m/sec
mv- momentum of particle, kg m/sec

Question 7

Heisenberg

Answer 7

Werner Heisenberg, in the early 20th century, said that it is impossible to simultaneously know both the position and the momentum of an electron.

For small particles, such as electrons, this uncertainty suggests that we need a wave model, rather that a Newtonian model, to understand their behavior.

Question 8

Schrödinger

Answer 8

Erwin Schrödinger, in the early 20th century, attributed a wave function to electrons.

The wave function describes the probability of where an electron might exist. The regions of high probability are called orbitals, even though they are more like clouds than orbits.

The orbitals of each electron is described in Schödinger’s equation. These orbitals can be described as s, p, d, or f orbitals, as used in electron configuration.