2.2: isotopes, isotope structures, and calculating average atomic mass

protons

neutrons

electrons

negatively charged (-)
located outside nucleus in orbitals
extremely small compared to protons and neutrons
do not appreciably contribute to the mass of the atom (which is generally calculated with protons and neutrons)

charge is determined by the combination of protons and electrons; neutrons have no effect as they have no charge.

isotopes: atoms of the same element with different masses
have different numbers of neutrons, but the same number of protons
atom: has a neutral charge (# of protons = # of electrons)
ion: has a charge (positive or negative; # of protons =/= # of electrons)
atomic number: the number of protons in the nucleus of an element
the atomic number “defines” an element and contributes to its placement on the periodic table
eg. aluminum has atomic number 13, which means that every aluminum atom has 13 protons in its nucleus
z is shorthand for atomic number
for Al, z=13
mass number: the number of protons + the number of neutrons in the nucleus of an atom
A is shorthand for mass number
in a neutral atom, the number of protons = the number of electrons
isotope symbols

atoms have extremely small masses
the heaviest naturally occurring atoms (uranium atoms) have a mass of roughly 4 x 10^-22 g
a mass scale on the atomic level is used to determine average atomic mass, where an atomic mass unit (AMU) is the base unit
1 AMU = 1.66054 x 10^-24 g

atomic mass (weight) measurement

atomic and molecular weight can be measured with great accuracy using a mass spectrometer
masses of atoms are compared to the carbon atom with 6 protons and 6 neutrons
in the real world, we use large amounts of atoms and molecules, so we use average masses in calculations
an average mass is found using all isotopes of an element weighted by their relative abundances

average atomic mass

average atomic mass: the weighted average of the masses of all the isotopes of the element