Fundamentals of Atoms, Isotopes, Atomic Mass, and the Mole

I. Isotopes: Counting Protons, Neutrons, and Electrons

Understanding Isotopes

Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons. This section focuses on identifying subatomic particles and calculating atomic and mass numbers.

• Proton: Positively charged particle found in the nucleus; defines the element.

• Neutron: Neutral particle in the nucleus; contributes to atomic mass.

• Electron: Negatively charged particle found outside the nucleus; determines chemical behavior.

• Atomic Number (Z): Number of protons in the nucleus.

• Mass Number (A): Total number of protons and neutrons in the nucleus.

Example Table:

Isotope	Protons	Neutrons	Electrons	Atomic Number	Mass Number
Na	11	12	11	11	23

II. Periodic Table

Organization and Groups

The periodic table arranges elements in order of increasing atomic number. Elements in the same group (vertical columns) share similar chemical and physical properties.

• Periods: Horizontal rows on the periodic table.

• Groups: Vertical columns; elements in a group have similar properties.

Common Groups and Specific Names

Group	Name	Properties
1A	Alkali metals	Very reactive with water, soft, not found free in nature
2A	Alkaline earth metals	Reactive, but less so than 1A; found in minerals
7A	Halogens	Very reactive nonmetals, form salts with metals
8A	Noble gases	Inert, least reactive elements

III. Atomic Mass: The Average Mass of an Element’s Atoms

Weighted Average of Isotopes

The atomic mass of an element is the weighted average mass of all naturally occurring isotopes of that element. Percent abundance is the percentage of a specific isotope in a natural sample of the element.

• Formula for Average Atomic Mass:

$\text{Average Atomic Mass} = \sum (\text{Fractional abundance} \times \text{isotope exact mass})$

• Example Calculation: If an element has two isotopes with masses 10.0129 amu (19.91%) and 11.0093 amu (80.09%), the average atomic mass is:

$\text{Average Atomic Mass} = (0.1991 \times 10.0129) + (0.8009 \times 11.0093) = 10.81 \text{ amu}$

• Practice: Calculate the average atomic mass of neon using the given isotopic masses and abundances.

Isotope	Mass (amu)	Percent Abundance (%)
Ne-20	19.9924	90.92
Ne-21	20.9930	0.257
Ne-22	21.9914	8.82

IV. Atoms and the Mole: How Many Particles?

Avogadro’s Number and the Mole

The mole is a counting unit for atoms, molecules, or ions. Avogadro’s number defines the number of particles in one mole.

• Avogadro’s Number:

$6.022 \times 10^{23}$ entities = 1 mole

• Example: Calculate the number of copper atoms in 2.45 moles of copper.

• Practice: Calculate the number of silver atoms in a sample containing $2.80 \times 10^{22}$ atoms.

Molar Mass of Elements

The molar mass is the mass in grams of 1 mole of a substance. The units are grams per mole (g/mol).

• Example: Calculate the number of grams in 0.90 moles of carbon.

• Practice: Calculate the number of moles in 8.57 grams of carbon.

Element	Atomic Mass (amu)	Molar Mass (g/mol)
H	1.008	1.008
C	12.01	12.01
O	16.00	16.00
Na	22.99	22.99
Cl	35.45	35.45

Additional info: The worksheet provides practice problems for calculating atomic mass, number of particles, and conversions between grams, moles, and atoms, which are foundational skills in general chemistry.