Atoms, Molecules and Ions History Greeks Democritus 460-370 BC “atomos” Aristotle- elements. Alchemy 1660 - Robert Boyle- experimental definition of element. Lavoisier (1734-1794)- Father of modern chemistry. He wrote the book -1789.

Laws Conservation of Mass Law of Definite Proportion- compounds have a constant composition. Multiple Proportions- When two elements form more than one compound, the ratios of the masses of the second element that combine with one gram of the first can be reduced to small whole numbers.

Multiple What??? Water has 8 g of oxygen per g of hydrogen. Hydrogen peroxide has 16 g of oxygen per g of hydrogen. 16/8 = 2/1 Small whole number ratios.

Dalton’s Atomic Theory 1803-1807 1) Elements are made up of atoms 2) Atoms of each element are identical. Atoms of different elements are different. 3) Compounds are formed when atoms combine. Each compound has a specific number and kinds of atom. 4) Chemical reactions are rearrangement of atoms. Atoms are not created or destroyed.

A Helpful Observation Gay-Lussac- under the same conditions of temperature and pressure, compounds always react in whole number ratios by volume. Avogadro- interpreted that to mean that at the same temperature and pressure, equal volumes of gas contain the same number of particles. (called Avogadro’s Hypothesis)

Thomson’s Experiment (1897) Used the Cathode Ray Tube(CRT) to discover the electron.

Thomson’s Experiment Passing an electric current makes a beam appear to move from the negative (cathode) to the positive end (anode).

Proving Electrons had Mass Thomson devised a cathode ray tube with a paddle wheel built inside. When the high voltage electricity was turned on the paddle wheel began to rotate and move away from the cathode and towards the anode.

Determining Charge Thomson concluded from this evidence and from his previous experiments that tiny particles were being emitted from the atoms of the cathode. These tiny particles were negatively charged. He called these particles "electrons."

Thomson’s Model of the Atom Given these experimental results, Thomson proposed in 1897 what has since been referred to as the "plum pudding" model of the atom. This model depicts the atom as a diffuse cloud of positive charge with the negative electrons embedded randomly in it, like plums in pudding.

Millikan (1909) - Mass of the Electron “(Oil Drop Experiment)” Using this information he calculated the mass of the electron as 9.11 X 10-28 grams.

Radiation In 1896, the French scientist, Henri Becquerel, found that a piece of a mineral containing uranium could produce its image on a photographic plate in the absence of light. He attributed this phenomenon to a spontaneous emission of something which he called "radiation" which originated from the uranium.

Radioactivity Discovered by accident Bequerel Marie and Pierre Curie isolated the radioactive components at Bequerel’s suggestion Three types –alpha- helium nucleus (+2 charge, large mass) __beta- high speed electron –gamma- high energy light

Rutherford (1911) A fluorescent screen would detect radiation by flashing whenever it was struck by the radiation. Rutherford observed that the radiation was diffracted into three beams by the charged plates.

Rutherford “Gold Leaf Experiment” Rutherford set up the apparatus which would bombard thin gold foil with alpha particles. These particles would then be detected by the fluorescent screen. He anticipated that all of the alpha particle detection would occur on the screen directly behind the foil.

Rutherford got surprisingly different results "It was about as credible as if you had fired a 15-inch shell at a piece of tissue paper and it came back and hit you!"

Atomic Model Revised Rutherford suggested that the atom was mostly empty space with a highly charged center. Most of the particles pass through the atom undisturbed, but a few get too close to the center and are deflected.

"Planetary Model." To account for these results Rutherford proposed a new model of the atom in 1913. This model had the following characteristics: The atom is mostly empty space with the majority of its mass concentrated in the center of the atom which he called the "nucleus."

Mass Number and Atomic Mass By the early 1930's the major subatomic particles had been discovered and their physical properties had been described. James Chadwick (1932) - discovers Neutrons. Particles Mass (grams) Relative Mass (amu) Relative Charge Proton 1.67262 X 10-27 1.007 +1 Electron 9.10939 X 10-31 5.486 X 10-4 (~ 0) -1 Neutron 1.67493 X 10-27 1.009 0

Atomic Mass Unit amu - the unit that we use to measure atoms. 1 amu ~ mass of one mole of the following: ~hydrogen ~ 1 proton ~ 1 neutron Actual masses: Particle Charge Mass (amu) Proton positive(+1) 1.0073 Neutron none(neutral) 1.0087 Electron negative (-1) 5.486 x 10-4 Hydrogen none (neutral) 1.0079

Measuring Atoms o Angstrom (A) - a convenient non-SI unit of length used to express atomic dimensions. 1 Angstrom = 1 x 10-10 meters most atoms have diameters between 1 x 10-10 m and 5 x 10 -10 m, or between 1 - 5A. o

Isotopes All atoms have the same number of protons. The number of neutrons may vary for a given element. Isotope - atoms of a given element that differ in the number of neutrons and consequently the mass. May be written as the symbol 12C or simply carbon-12 as opposed to the isotope 14C or carbon-14. 6 6

Atomic Numbers and Mass Numbers Atomic Number - is the number of protons, which is shown as the subscript. (it is also the number of electrons in a neutral atom.) Mass Number - is the total number of protons plus neutrons in the atom. (which represent essentially all the mass of the atom.) ex. 12C has 6 Protons, 6 Electrons and 6 Neutrons Number of Neutrons = Mass # - Atomic # 6 6 = 12 - 6

Nuclides Nuclide - an atom of a specific isotope ex. 14C or carbon-14 6 protons, 6 electrons and 8 neutrons # of neutrons = atomic mass - atomic number 8 = 14 - 6 6