1.1 Atomic Structure: The Nucleus

Atomic Structure and the Nucleus

Atoms are the fundamental units of matter, and their structure determines how they behave in chemical reactions. For organic chemistry, you need a solid grasp of what's inside an atom, because bonding, reactivity, and molecular shape all trace back to atomic structure.

Structure of Atoms

An atom consists of a dense, positively charged nucleus surrounded by negatively charged electrons. The nucleus is tiny compared to the overall atom, but it contains nearly all of the atom's mass.

The nucleus holds two types of particles:

• Protons carry a positive charge and have a mass of approximately 1 atomic mass unit (amu)
• Neutrons carry no charge and also have a mass of approximately 1 amu

Electrons occupy shells (energy levels) and orbitals around the nucleus. They carry a negative charge, but their mass is negligible compared to protons and neutrons. The number and arrangement of electrons determine an element's chemical properties, including how it bonds and reacts.

In a neutral atom, the number of protons equals the number of electrons, so the charges balance out. When an atom gains or loses electrons, it becomes an ion. Losing electrons produces a positively charged cation, while gaining electrons produces a negatively charged anion.

Isotopes and Atomic Properties

Isotopes are atoms of the same element that have different numbers of neutrons. They share the same number of protons and electrons, so their chemical behavior is nearly identical, but they differ in mass.

Carbon is a good example. All carbon atoms have 6 protons, but:

• $^{12}C$ has 6 neutrons (mass number = 12)
• $^{13}C$ has 7 neutrons (mass number = 13)
• $^{14}C$ has 8 neutrons (mass number = 14)

Two key numbers define any atom:

• Atomic number (Z): the number of protons in the nucleus. This defines the element. Every carbon atom has $Z = 6$, no exceptions. All isotopes of an element share the same atomic number.
• Mass number (A): the total number of protons plus neutrons. This varies among isotopes and is written as a superscript to the left of the element symbol ($^{12}C$, $^{13}C$).

While isotopes of an element behave the same chemically, their physical properties differ. Some isotopes are stable, while others are radioactive, meaning they spontaneously decay by emitting particles or energy over time. $^{14}C$ is radioactive; $^{12}C$ and $^{13}C$ are stable.

Nuclear Properties and Stability

Protons are all positively charged, so they repel each other inside the nucleus. What holds the nucleus together is the strong nuclear force, a powerful attractive force between nucleons (protons and neutrons) that operates at very short range and overcomes electrostatic repulsion.

Nuclear stability depends on the ratio of neutrons to protons. Stable nuclei fall within a specific range of neutron-to-proton ratios. When a nucleus has too many or too few neutrons relative to its protons, it becomes unstable and undergoes radioactive decay.

Binding energy is the energy required to completely break apart a nucleus into individual protons and neutrons. A higher binding energy per nucleon generally means a more stable nucleus.

Calculation of Atomic Weight

Atomic weight is the weighted average mass of an element's naturally occurring isotopes. It accounts for both the mass and the relative abundance of each isotope.

To calculate atomic weight:

1. Convert each isotope's percent abundance to a decimal (divide by 100)
2. Multiply each isotope's mass by its fractional abundance
3. Add the products together

Example: Chlorine has two naturally occurring isotopes:

• $^{35}Cl$: mass = 34.97 amu, abundance = 75.77%
• $^{37}Cl$: mass = 36.97 amu, abundance = 24.23%

$\text{Atomic weight} = (34.97 \times 0.7577) + (36.97 \times 0.2423) = 35.45 \text{ amu}$

Notice that the atomic weight (35.45) is closer to 35 than to 37, which makes sense because $^{35}Cl$ is far more abundant.

Atomic weights feed directly into stoichiometry. The molar mass of a compound equals the sum of the atomic weights of all its atoms, expressed in g/mol. You'll use this constantly when calculating reaction yields and molecular formulas.