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

the structure of atoms—subatomic particles

protons

• positively charged (+)

• located inside the nucleus

• relatively massive

neutrons

• no charge—neutral (/)

• located inside the nucleus

• relatively massive

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

• 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

atomic mass

• 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