Download Chapter 1 Notes: The Science of Chemistry

Chapter 3 Notes: Atoms and Moles
Dalton’s Atomic Theory
 John Dalton (British Teacher)
 Composed in early 1800’s
 Mostly still true today
 5 parts
o All matter is composed of atoms
o All atoms of the same element have identical properties
o Atoms of different elements have different properties
o Atoms of different elements combine in definite whole number ratios
to form compounds
o In chemical reactions, atoms are combined, separated, or rearranged
but NEVER created or destroyed
Law of Definite Proportions
 Compounds always contain the same elements in the same ratios
o Water is always H2O, nothing else.
o H2O2 exists but it is a different compound
Law of Conservation of Mass
 Mass cannot be created or destroyed
 Reactants  Products
Law of Multiple Proportions
 It is possible to have 2 different compounds made from the same elements but
just in different ratios
 For example CO2 (carbon dioxide) CO (carbon monoxide)
Atomic Structure
 Atoms are composed of 3 basic subatomic particles
 Proton
o Positively charged
o Found in nucleus
o p+
 Neutron
o No charge
o Found in nucleus
o n0
o same size as proton
 Electron
o Negatively charged
o Found in the electron cloud (surrounding nucleus)
o e-
o approximately 1800 times smaller than a p+ or n0
o discovered by JJ Thomson- cathode ray experiment
Rutherford’s Experiment
 Ernest Rutherford
 Proved the existence of a small dense positively charged nucleus
 Originally thought atoms were solid; experiment ended up proving the
opposite
 See diagram of Rutherford’s exp. in Chapter
Atomic Number
 Represents the number of protons in the nucleus of an atom
 Unique number for each element
 Cannot be changed
Mass Number
 Total number of protons and neutrons in the nucleus of an atom
 It can possibly change since the number of neutrons in an atom is subject to
chance occasionally (isotopes)
Atomic Models
 Theories of what we think the atom looks like
 These have changed over time
 Evidence is mostly indirect (wrapped present)
 4 majors historic models
o Thomson’s model
 Plum pudding
 Solid shell
 Electrons move around in positively charged pudding like
material
o Rutherford’s model
 Solid dense positive nucleus
 Electrons found on the outside (nothing specific about e-)
 Atoms consist mostly of empty space
o Bohr’s model
 Just like Rutherford’s except Bohr believed that the electrons
orbited the nucleus in set paths similar to the planets around the
Sun
o Quantum model
 Current model
 Only change is that we know electrons DO NOT follow an
exact orbit like the planets
 However, electrons do stay at a set distance from the nucleus
(energy levels, orbitals, etc)
 Electrons act as both particles and waves (dual nature)
Electromagnetic Spectrum
 Range of wavelengths of EMR
 EMR- electromagnetic radiation
 Electrons thought to be behaving as waves NOT particles in EMR
Characteristics of Waves
 Wavelength- length of wave before it repeats
o All EMR travels at the speed of light (3 x 108 m/s)
 Frequency (v)- number of wavelengths that pass a certain point in space per
unit of time
o Hertz (Hz)- SI unit of frequency
o Number of waves (cycles) which pass a given point in one second
Electron Configuration in the Atom
 Ground state- electrons in their lowest possible energy level around the atom
(normal)
 Excited state- electrons that have absorbed energy and jumped to a higher
energy level (will eventually drop back to the ground state and release energy
as a result)
 See diagram discussing sublevels, orbitals, orbital shape, and energy levels.
 Aufbau principle- electrons fill orbitals of lowest energy first
 Pauli exclusion principle- only 2 electrons fit in each orbital
 Hund’s rule- for sublevels that contain more than 1 orbital (p,d,f), each orbital
must accept 1 electron before any orbital of equal energy receives a 2nd
electron in same orbital
o 2 electrons in the same orbital must have opposite spins
 WORK ON e- configuration problems
Atomic Mass
 Weight of an individual atom
 Unit used is the amu (atomic mass unit)
 1 amu= 1/12 the mass of a C-12 atom
 Protons and neutrons each have a mass of approximately 1amu
 Mass of electrons is so small to be negligible
 An amu is so small that we don’t have a tool to measure; therefore, a way to
group atoms together had to be developed
Mole (mol)
 SI unit used to discuss the number of atoms
 A mole is a group of atom
 1 mol= 6.02 x 1023 particles
 Analogous to 1 dozen = 12 particles
o 1 C atom = 12.011 amu
o 1 mol C = 12.011 grams = 6.02 x 1023 C atoms
o 6.02 x 1023 is known as Avogadro’s Number

Do problems involving conversions between mass-moles-atoms pages 102103