Atomic Structure

Democritus -- 400 B.C.

l Student of Arostotle

l Basic particles make-up all matter

l The smallest, indivisible particles of matter are called atomos

Robert Boyle

l First to study chemistry as a separate intellectual discipline

l First to carry out rigorous chemical exmperiments

l First to clearly define an element

Joseph Priestley - 1774

l Isolated the gas oxygen by heating mercury oxide (HgO)

2HgO ® 2Hg + O₂

l Chemical formula

l Chemical equation

Law of Conservation of Mass: Mass is neither created nor destroyed in chemical reactions

John Dalton -- 1807

l Studied chemical reactions which investigate the conservation of mass

l Developed Dalton’s Atomic Theory

l All matter is composed of tiny, indivisible particles called atoms which cannot be created, destroyed, or interconverted

l Atoms of any particular element are identical; whereas, atoms of one element differ from atoms of other elements

l Chemical change is a union, separation, or rearrangement of atoms

l If the experimental conditions of a chemical reaction are changed, the combining ratio of one element with another element may also change

Sir William Crookes 1879

• Developed the cathode ray tube

• cathode rays: negatively charged particles

The Physics of Waves

l l - (lamda) Represents wavelength of a wave

l n - (nu) Represents the frequency of a wave

l v=ln

l E=hn

J.J. Balmer -- 1885

l Swiss scientist

l Developed an equation to calculate the wavelength of lines in the hydrogen atom spectrum

Eugen Goldstein -- 1886

l Used a Crookes tube with holes in the cathode

l Observed another kind of ray which originated near the anode and passed through the holes in the cathode

l Canal rays

J.R. Rydberg -- 1890

l German

l Developed an equation based on Balmer’s work, which describes the energies associated with various energy levels in the hydrogen atom

l E=hn=2.179 X 10^-18J(1/n₁^{2 -} 1/n₂²)

J.J. Thomson -- 1897

l English physicist

l Measured the deflection of cathode ray particles in both a magnetic and an electric field

l Determined the charge (e) to mass (m) ratio and found them to be identical for all particles regardless of the metal used as an electrode or the type of gas within the tube

Max Planck -- 1900

l German physicist

l Proposed a quantum theory that described the light emitted from a hot object as composed of discrete unit called quanta or photons

l E=hn

J.J. Thomson -- 1904

l Proposed a model of the atom with electrons embedded in a sea of positive charges

l Called the “plum-pudding model”

Albert Einstein -- 1905

l Published an explanation of the photoelectric effect

• today the photoelectric effect has resulted in such technology as automatic doors

l Electrons are emitted from metals when these metals are exposed to light of the proper frequency

l E=hn=hc/l

The Photoelectric Effect: Radiation causes electrons to move in a metal

J.J. Thomson -- 1907

l Determined that Goldstein’s rays are positively charged particles called protons

l The mass was determined to be

• 1.0073 amu

• 1.673 X 10^-27 Kg

Robert Millikan -- 1909

l Millikan Experiment

l Determined the charge of the electron

l From the value of e/m, he found the mass to be 1/1837 of the mass of a hydrogen atom

• 0.00055 amu

• 9.11 X 10^-31 Kg

Rutherford’s Experiment occurred in 1909. Rutherford bombarded very thin sheet of gold foil with radioactive a particles. He was then able to detect the scattering pattern of the particles using photographic plates.

Ernest Rutherford -- 1911

l Published his 1909 work in 1911

l Projected a beam of a particles onto a very thin gold foil

l From experimental results, Rutherford concluded:

• the volume occupied by an atom is largely empty space

• each atom contains a massive, positively charged nucleus

• electrons move about the nucleus giving the atom its volume

Niels Bohr -- 1913

l Proposed that the electron’s energy is quantized

l Developed the Bohr model of the atom

• Often called the planetary model

l The electron of hydrogen moves about the nucleus in a circular orbit

l The centrifugal force due to this motion counterbalances the electrostatic attraction between the nucleus and the electron

l The energy of the electron is restricted to certain values, each corresponding to an orbit with a different radius

l Quantum number

Henry Moseley -- 1913

l English physicist

l Killed during World War I

l Used X-rays striking an element to determine the number of protons (atomic number) in the nucleus

l X-rays produced by various elements were measured

l The X-ray energies were dependent on the atomic number of each element

F.W. Aston -- 1913-14

l Developed the mass spectrometer

• Gaseous substance are bombarded by high-energy electrons, thus knocking off electrons so as to produce positively charged ions

• Ions are directed through a magnetic or electrical field, which deflects their paths depending on their mass/charge ratio

Mass Spectrometer

l Atoms of the same element have differing masses

l Identified the isotopes of various elements

Masson, Harkins, & Rutherford
1920

l Orme Masson (Australian)

l William Harkins ( American)

l Ernest Rutherford (New Zealander)

l Electrostatic attraction of positive nucleus and negative electrons.

l Centrifugal force holding the nucleus and electrons apart.

Failures with the Bohr model lead to a need for a new atomic model

Bohr’s Model Provides Insight into Atom’s Behavior

l Energies of electrons (energy levels) are quantized.

l Quantum numbers describe such electron properties as energy and location.

l An electron’s energy changes with distance from the nucleus.

l Spectral lines of the elements are due to quantized electronic energies.

Nucleons: The total of all nuclear particles, includes both protons & neutrons.

Atomic Mass: The mass in atomic mass units of an element

Molecular Mass: The mass in atomic mass unitsof a molecule

Formula Mass or Empirical Mass: The mass in atomic mass units of an ionic compound

Na - 23.0 X 1 = 23.0

Cl - 35.5 X 1 = 35.5

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58.5 amu

C - 12 X 1 = 12

O - 16 X 2 = 32

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44 amu

Moles

l The quantity of matter containing Avogadro’s number of particles

l 6.022 X 10²³ particles

l Particles may include:

• subatomic particles

• ions

• atoms

• molecules

Molar Mass: The mass of one mole of a substance in grams

Calculating moles from grams: Divide grams by molar mass

Calculating grams from moles: Multiply moles by molar mass

Nuclear Chemistry

l Nuclear equation: Elemental symbols represent only the nuclei of atoms.

l The subscript represents only the number of nuclear charges (protons or atomic number).

Alpha (a) Radiation (Helium nuclei)

Beta (b) Radiation (electron)

Gamma (g) Radiation (Electromagnetic radiation)

Positron Emission: A proton changes into a neutron plus an ejected positron.

Electron Capture: The nucleus captures one of the surrounding electrons in an atom, thereby converting a proton into a neutron.

Nuclear Stability – The ratio of protons to neutrons within the nucleus which allows for nuclear stability. A ratio that is not in the correct proportions will result in an atom that is radioactive and will go through a nuclear change to move into the band of stability.