Ch.1 Atomic Structure

Study of Matter

Anything that occupies space (volume) and has mass

Everything around us (solid, liquid, gas)

An atom

• fundamental unit of matter

• smallest unit of matter to keep its identity

• smallest identifiable unit of an element

• from ancient Greek word: Atomos: meaning “uncuttable” or “indivisible” (cannot be further divided)

subatomic particles: protons, neutrons, electrons

• Found in the nucleus of an atom

• Have a positive electrical charge (+1)

• Mass of 1 amu (atomic mass unit)

• The atomic number (Z) of an element is equal to the number of protons found in an atom of that element

• are neutral (no charge)

• mass is slightly larger than that of a proton

• found in the nucleus of an atom

• mass of 1 amu

Protons and Neutrons

Protons and Neutrons make up almost the entire mass of an atom

• represented by the letter (A)

• sum of protons & neutrons

• represented by the letter (Z)

• unique identifier of each element because elements are defined by the number of protons they contain

• equal to the number of protons

Atoms that have the same number of protons (atomic number) but different number of neutrons (mass number)

• move through the space surrounding the nucleus

• associated with varying levels of energy

• found in space surrounding nucleus

• -1 electromagnetic charge (a charge equal in magnitude to that of a proton but with opposite sign (-))

• mass of an e- is ~1/2000 of a proton

• Because subatomic particles’ masses are so small

• Electron shells

• Electrons move around nucleus at varying distances, which corresponds to varying levels of electrical potential energy

  • electrons closer to nucleus at lower energy levels

  • electrons further out have higher energy

• Electrons furthest from nucleus interact strongest with surrounding environment because weakest interaction with nucleus

• electrons FURTHEST from nucleus

  • much more likely to become involved in bonds with other atoms because they experience least electrostatic pull from their own nucleus

• equal number of protons and electrons

• an atom that has lost an electron and become POSITIVELY charged

  • cats have paws PAW-sitive

• an atom that has gained an electron and become NEGATIVELY charged

• 1/12th the mass of a Carbon-12 atom

• 1.66 X 10^-24g

• Carbon-12 has 6 protons and 6 neutrons so 1 amu is equal to the ~mass of a proton or a neutron (proton and neutrons almost the same mass, very smallll difference)

• isotopes have the same number of protons and electrons, so they generally exhibit similar chemical properties

• in nature almost all elements exist as 2 or more isotopes

• isotopes are usually present in same proportions in any sample of naturally occurring element

• the weighted average of these different isotopes is referred to as atomic weight and is the number reported on the periodic table

• No, e.g. Bromine is listed on periodic table as 79.9 amu, which is the average weight of the naturally occurring Bromine isotopes (79, 81), which occur in almost equal proportions. No bromine atom with an actual mass of 79.9 amu (naturally occurring Bromine atoms have a mass of 79 amu or 81 amu)

• weighted average of naturally occurring isotopes of an element

• mass of 1 mole of an element in grams

• a mole is a number of “things” (atoms, ions, molecules)

  • that number is equal to Avogadro’s number (6.02 ×10²³)

6.02 x 10²³

500 BCE Leucippus and Democritus propose matter is made up of indivisible particles called “atomos”

1789 Antoine Lavoisier: Law of Conservation of Mass

1794 Joseph Proust: Law of Constant Proportions

1803 John Dalton: Atomic Theory

1897 Discovery of Electron: Cathode Ray Experiment (JJ Thomson)

1898 Discovery of Proton: Anode Ray Experiment (E Goldstein)

1900 Max Planck: Quantum Theory, Planck’s Constant

1904 JJ Thomson: Plum Pudding Model

1911 Ernest Rutherford: Nuclear Model

1913 Niels Bohr: Planetary Model (Bohr Model)

1926 Erwin Schrodinger: Quantum Model