Atom

Atomic structure

visualized by various atomic models
composed of subatomic particles: protons, neutrons, electrons
nucleus
- dense center of the atom
- contains protons and neutrons
- holds most of the atom’s mass
electron cloud (see electron cloud model)
- where electrons are likely to be found
- makes up most of the atom’s volume

historically used Atomic Mass Unit ( $am u$ )
modern version is Unified Atomic Mass Unit ( $u$ ) ¹
$u = 1.660540 \times 1 0^{- 27} k g$
the periodic table shows elements’ average atomic mass, the weighted average based on abundance of its isotopes
- ex. Iron has isotopes with % abundance: $Fe - 54$ at $5.9%$ , $Fe - 56$ at $91.8%$ , and $Fe - 57$ at $2.3%$
$average atomic mass = (54 \cdot 0.059) + (56 \cdot .918) + (57 \cdot 0.023) = 55.905 u$

different types of atoms, each with a unique set of physical and chemical properties
an element is identified by the number of protons in its atoms
an elements atomic number ( $Z$ ) is equal to the number of protons in its atoms
- helium: $Z = 2$ , 2 protons
- iron: $Z = 26$ , 26 protons
atomic number also indicates how many electrons are in a “neutral” atom of that element
in the periodic table, elements are organized by their atomic numbers
- increasing atomic number left→right and top→bottom
- elements in the same column tend to have similar physical and chemical properties
the total mass of an element is calculated based on its number of protons and neutrons
- $Mass number = # of protons + # of neutrons$
- electrons have very little mass

atoms of the same element with a different number of neutrons
has different mass from other isotopes of the same element
notation: $X_{Z}^{A} X_{2 Z}^{2 A} X$
- A = mass number (# of protons + # of neutrons)
- Z = atomic number (# of protons)
- X = chemical symbol
another common notation: $X - A$
example: carbon, atomic number 6
- in nature, mainly composed of 2 isotopes
  - 6 neutrons: $X_{6}^{12} X_{26}^{212} C$ or $C - 12$
  - 7 neutrons: $X_{6}^{13} X_{26}^{213} C$ or $C - 13$

Changed from $am u$ to $u$ in 1961 to solve conflict between physicists and chemists. $am u$ was based on oxygen, when scientists believed all oxygen was oxygen-16. Later, isotopes of oxygen were discovered, confusing the definition of $am u$ . New standard $u$ was defined using carbon-12, where $u = 1/12 \cdot mass of X_{6}^{12} X_{26}^{212} C$ . $C - 12$ is the only atom with a whole-number mass ( $12 u$ ). https://en.wikipedia.org/wiki/Carbon-12#History ↩