Comprehensive Notes on Atomic Structure

Early model: Atom as a solid ball with negative charges (electrons) scattered throughout (plum pudding model).
Rutherford's discovery: Atom has a small, dense nucleus with electrons orbiting around it. This is the nuclear atom model.

Nucleus: Contains protons and neutrons.
Electron cloud: Region surrounding the nucleus where electrons are located. Electrons move all around the nucleus, making their exact location difficult to pinpoint.
Most of the atom is empty space.

The weight of an atom is determined by the number of protons and neutrons in the nucleus.
Mass number: Total number of protons and neutrons in the nucleus.
- $Mass\ Number = Number\ of\ Protons + Number\ of\ Neutrons$ or $A = Z + N$
Atomic Mass Unit (AMU): Unit of mass for atomic particles.
- Each proton and each neutron weigh approximately 1 AMU.
- Electrons have negligible weight compared to protons and neutrons.
Example: Carbon atom. If a carbon atom has six protons and six neutrons, its mass number is 12 AMU. Electrons do not contribute to this mass. It represents the mass of the atom's nucleus.

Lithium example: Visual representation shows electrons surrounding a nucleus with protons and neutrons.
Electron density: Darker areas indicate a higher probability of finding electrons. Electrons are found outside the nucleus.
Lithium has three protons and four neutrons, giving it a mass of seven.
To balance the charge of three protons, lithium has three electrons.

Atomic number: Distinguishes elements from one another; the number above each element symbol on the periodic table. It represents the number of protons in the atom.
Carbon's atomic number is six, signifying six protons.
In a neutral atom, for example carbon, the number of protons (6) equals the number of electrons (6).
For lithium, the atomic number is 3; it has three protons and, in its neutral state, three electrons.

To determine the number of neutrons, subtract the atomic number (number of protons) from the mass number.
- $Number\ of\ Neutrons = Mass\ Number - Atomic\ Number$ or $N = A - Z$
Example: given an isotope $^{56}_{26}Fe$ . This element has 26 protons and 30 neutrons.

Isotopes: Atoms with the same number of protons but a different number of neutrons.
Examples: Carbon-12, Carbon-13, Carbon-14.
The number following the element name indicates the mass number.
All carbon isotopes have an atomic number of 6 (6 protons).
- Carbon-12 has 6 neutrons ( $12 - 6 = 6$ ).
Hydrogen Isotopes: Hydrogen, Deuterium, Tritium (Hydrogen-1, Hydrogen-2, Hydrogen-3).
Deuterium and Tritium react in the sun.

Given Carbon-14 ( $_{6}^{14}C$ ), determine the number of protons, neutrons, and electrons.
- Protons: 6 (same as atomic number).
- Electrons: 6 (same as number of protons in a neutral atom).
- Neutrons: 8 ( $14 - 6 = 8$ ).

Given an element with a mass number of 79 and an atomic number of 35, find the number of neutrons, protons and electrons.
- Number of Protons: 35 (equal to the atomic number).
- Number of Electrons: 35 (equal to the number of protons for a neutral atom).
- Number of Neutrons: 44 ( $79 - 35 = 44$ ).

Mass Number: The sum of protons and neutrons in an atom's nucleus. Always a whole number.
Atomic Mass: A weighted average of the masses of all naturally occurring isotopes of an element. This is why atomic masses are not whole numbers and have decimals.

Average atomic mass is calculated based on the abundance (percentage) of each isotope.
Mass Spectrometer: A device used to determine the mass numbers and abundances of different isotopes, allowing accurate experimental measurements.
Weighted Average Calculation:
1. Multiply the mass of each isotope by its percentage abundance (expressed as a decimal).
2. Sum the results to get the weighted average atomic mass.
Example: Silicon (Si) has three isotopes: Silicon-28, Silicon-29, Silicon-30.
- Isotope Masses: 27.98 AMU, 28.98 AMU, 29.97 AMU.
- Abundances: 92.2%, 4.7%, 3.1%.
- Weighted Average Calculation:
  - $(27.98 \times \frac{92.2}{100}) + (28.98 \times \frac{4.7}{100}) + (29.97 \times \frac{3.1}{100})$
  - $25.8 + 1.36 + .93$
  - $\approx 28.1 AMU$
Important: This calculation involves a weighted average, not a simple average. Do not simply sum the masses and divide by the number of isotopes.

If given the total percentage and one isotope percentage, subtract to find the other percentage.
Example: If one isotope has 75.77% abundance, the other isotope has 24.23% abundance ( $100 - 75.77 = 24.23$ ).