Atom = the smallest part of an element that can exist and still have the properties of that element.
The diameter of atoms is in the range of  10^-18cm.

Range of atomic mass = 10^-24 to 10^-22 grams
        1 lb Au contains = 1.4 x 10²⁴ atoms

Molecule = two or more atoms tightly bound together that act as a single unit

• diatomic molecule H₂
• diatomic molecule CO
• triatomic molecule CO₂

Heteroatomic molecules:     NH₃ ammonia,     C₆H₁₂O₆ glucose

Pure Substances:   compounds and  elements

Compounds exist in two types: molecular and ionic.

• Molecular compounds consist of two or more atoms bound together into a neutral molecule.
• Ionic compounds contain + and - electrically charged particles called ions.

Chemical Formula, a symbolic notation of elements present in a compound.
     examples:    C₆H₁₂O₆,     C₁₂H₂₂O_11,     C8H18,       Al2(Cr2O7)3

 

Name	Symbol `	Relative Mass	Relative Charge
proton	p+1	1  amu	+1
neutron	n	1  amu	none
electron	e-1	1/1840	-1

The protons and neutrons (nucleons) are in the central nucleus of the atom. They occupy <1% of the volume and contain >99% of the mass. The nucleus is surrounded by a cloud of electrons.

Gas discharge tube experiments by J.J. Thompson and Eugene Goldstein around the turn of this century lead of the characterization of electrons and protons.

The atomic number (Z) is the distinguishing characteristic of an element. It equals the numbers of protons in the nucleus of an atom of that element.

In a neutral atom (no net electric charge) the number of protons equal the number of electrons. The mass number (A) of an atom equals the number of protons + neutrons.

             A                                           Note:  A is the mass # = p + n
                X
             Z                                            Note Z is the atomic number = # p
 

              131                                       Note this represents 53p + 78n = 131 neculeons.
                   I                                       Note that Th. "I" is the symbol for the element iodine.
              53                                         Note this represents 53 p.

A-Z = # neutrons

Isotopes = atoms of the same element with different numbers of neutrons. Isotopes have different masses but exhibit no chemical differences.

Isobars are atoms of different elements with the same mass numbers.

⁵⁸Ni and ⁵⁸Fe  are isobars of each other.
^{28              26}

23 elements have only one naturally occurring isotope. Of the ~1900 isotopes that have been observed ~270 occur naturally. Synthetic isotopes are radioactive.

Atomic mass (atomic weight) is the weighted average mass of all the isotopes of an element. The weighing is based on the % abundance's found on earth. The reference standard is one atom of
¹²C = 12 amu.
⁶

Given that copper has two naturally occurring isotopes; ⁶³Cu (isotopic mass 62.93 amu /atom, percent abundance 69.09%) and ⁶⁵Cu ( isotopic mass 64.93 amu / atom, percent abundance 30.91%) calculate the atomic mass of copper.

Note: amu = atomic mass unit. It is a unit of mass based on the standard of one ¹²C atom being defined as 12 amu.                                                                                            ⁶

6.02 x 10²³ amu = 1.00 g

23 of the known 111 elements are "synthetic". Atomic masses for these elements, based on % abundance on earth, can not be calculated. The periodic chart lists the mass number of the most stable isotope in place of atomic mass.