Atomic Structure

Atomic Theory Democritus ( B.C.) –Greek philosopher –Democritus proposed that the world is made up of empty space and tiny particles called atomos (or atoms)

Development of Atomic Theory John Dalton ( ) – developed Dalton’s Atomic Theory –All elements are made up of atoms –Atoms of the same element are identical to each other but are different than the atoms of other elements –Atoms of different elements can combine to form compounds –Atoms are indestructible and cannot be divided into smaller particles. Which one is NOT TRUE???

J.J. Thompson –Discovered electrons –Cathode Ray ExperimentCathode Ray Experiment –Plum Pudding Model: electrons stuck into a lump of positive charge, like raisins stuck in dough (or chocolate chips stuck in cookie dough)

Ernest Rutherford –Proposed that the atom is mostly empty space with all of the mass and positive charge concentrated within the nucleus –Conducted the “Gold Foil Experiment”Gold Foil Experiment –Concluded that the atom is mostly empty space but contains a small, dense, positively charged central core called an atom. –Nuclear Model

Atomic Structure ParticleLocationChargeRelative Mass Proton (p + ) Nucleus+1 Relatively large Neutron (n 0 ) Nucleus01 Same size as protons Electron (e - ) Shells-0 Very small

The Atom Nucleus – center core of an atom, very dense, contains protons and neutrons, overall positive charge Electron shell/orbit – surrounds the nucleus, where electrons are located

Atomic Number Tells the number of protons in the nucleus Because an atom is electrically neutral, the atomic number also indicates the number of electrons (#p + = #e - ) C Carbon

Mass Number Tells the number of protons PLUS neutrons Mass number - atomic number number of neutrons C Carbon

Protons – 6 Electrons – 6 Neutrons – = C Carbon 6p + 6n 0 2e - 4e -

Isotopes Atoms that have the same number of protons but a different number of neutrons The atomic number stays the same in isotopes (the number of protons identifies the element) The mass number changes

Isotope Examples 6p + 6n 0 2e - 4e - 6p + 7n 0 2e - 4e- C-12: 6p + 8n 0 4e - C-13: C-14:

Isotope Examples H-1 (Protium): 1p + 0n 0 1e - H-2 (Deuterium): 1p + 1n 0 1e- H-3 (Tritium): 1p + 2n 0 1e -

Ions Cations – an atom with a positive charge because electrons have been lost Anion – an atom with a negative charge because electrons have been gained A Negative I O N

Average Atomic Mass Weighted average of all naturally occurring isotopes of an element Explains why the mass number on the periodic table is a decimal number

Chlorine-35 has a percent abundance of and an amu of Chlorine-37 has a percent abundance of and an amu of Calculate the atomic mass of chlorine. Cl-35: X = Cl-37: X = % abundance must be converted to relative abundance by dividing it by 100 The number after the symbol represent the atomic mass of the isotope

Electrons in Atoms Electrons move very quickly around the nucleus. The can also move up or down energy levels. Electrons move up an energy level when they are excited or given energy. Electrons move down an energy level when they lose energy.

Valence Electrons Electrons located in the outermost energy level (the last shell) Number of valence electrons = group number

Lewis Dot Diagrams Use dots to represent the valence electrons Steps to drawing dot diagrams: –Write the chemical symbol –Determine the number of valence electrons –Draw out the dots in the following configuration Cl

The Bohr Model Electrons are found in specific paths (orbits) located around the nucleus. An electron must be ON an energy level. A “quantum” of energy will allow an electron to move from one energy level to another.

Higher energy levels are located farther away from the nucleus Drawing Bohr models: HIGHER ENERGY LEVEL LOWER ENERGY LEVEL Nucleus 6p + 6n 0 2e - 4e - 1 st energy level – 2 electrons 2 nd energy level – 8 electrons 3 rd energy level – 18 electrons 4 th energy level – 32 electrons

Quantum Mechanical Model Developed by Erwin Schrodinger Describes the probability of finding an electron at various locations around the nucleus

Quantum Numbers Describes the location of the outermost electrons The electrons “zip code” Each element on the periodic table has a unique four digit quantum number Pauli Exclusion Principle – no two elements can have the same set of quantum numbers

Practice Problems K 4, 0, 0, -1/2 W 5, 2, 1, -1/2 Cu 3, 2, 1, 1/2 5, 0, 0, ½ Sr 5, 3, -2, -1/2 Th 4, 2, 2, ½ Cd

Practice Problems Nitrogen 1s 2 2s 2 2p 3 Silicon 1s 2 2s 2 2p 6 3s 2 3p 2 Helium 1s 2 Chromium 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Gold 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 1 4f 14 5d 10

Short-Hand Electron Configuration Write the symbol for the noble gas (column 18) from the row before the element in brackets Write out the electron configuration for the entire row the element is located in

Silicon –[Ne]3s 2 3p 2 Chromium –[Ar]4s 1 3d 5 Copper –[Ar]4s 1 3d 10 Bismuth –[Xe]6s 2 4f 14 5d 10 6p 3 Practice Problems

Orbital Diagrams Series of lines and arrows used to represent the order in which the electrons fill the orbitals Write out the electron configuration Draw lines to represent the orbitals (s gets 1, p gets 3, d gets 5, and f gets 7) Draw arrows to represent the electrons (the superscript)

Atomic Spectra –Atoms that have absorbed energy have electrons that move to a higher energy level –When these electrons lose energy, they emit light as they drop back to a lower energy level (their ground state)

–The light that is emitted by these atoms have very specific frequencies that appear as discrete lines when viewed through a diffraction grate –The atomic spectrum of each element is unique because each element has a unique electron configuration –The light emitted by the electron is directly proportional to the energy change of the electrons; the greater the energy change, the greater the frequency of light emitted