Atomic Structure and Subatomic Particles
Rutherford's Gold Foil Experiment and Atomic Structure:
Rutherford used alpha particles, which expose photographic paper (or modern mirror-based detectors), to probe the structure of atoms.
The experiment involved a very thin material, allowing alpha particles to pass through a few atoms.
Observations:
Most alpha particles passed through undeflected, indicating that atoms are mostly empty space.
Some particles were deflected at large angles (disturbing deflections).
A significant finding was that 1 out of every 20,000 particles actually came all the way back to the emitter.
Conclusions:
This proved that atoms are not a homogeneous, light substance.
For an alpha particle (which is 8,000 times heavier than an electron) to be reflected back, it must hit something significantly heavier than itself.
The core of the atom is a very dense, heavy portion called the nucleus.
The nucleus contains all of the positive charge, concentrated in a very small area.
A direct hit on this positively charged, dense nucleus would cause the alpha particle to reflect straight back.
Components of the Atom:
Nucleus:
The dense core of the atom, containing all protons and neutrons.
It is very, very small, often depicted much larger in diagrams for clarity, but in reality, it's a speck at the atom's center.
Contains all the positive charge of the atom.
Protons:
Located in the nucleus.
Carry a positive charge, represented as +1 in relative units.
Typically colored red in diagrams.
Neutrons:
Located in the nucleus.
Are electrically neutral, represented as 0 in relative units.
Typically colored white in diagrams.
Along with protons, they are collectively called nucleons.
Electron Cloud:
The region where electrons exist, orbiting around the nucleus.
Most things that do not hit the nucleus should be able to pass through the electron cloud.
Electrons are easily accessible and can be pulled out of the atom by an electric field because they orbit around the outside.
Subatomic Particle Properties (Charge and Mass):
Electron:
Relative charge: -1
Relative mass: 1 (base unit of mass)
Actual mass: 9.11 imes 10^{-31} kg
Proton:
Relative charge: +1
Relative mass: Approximately 2,000 times heavier than an electron.
Neutron:
Relative charge: 0
Relative mass: Approximately 2,000 times heavier than an electron.
Important Point: The neutron is slightly heavier than the proton (a difference of about 5 units in relative mass, which has a specific reason for being slightly heavier).
Atomic Notation and Isotopes:
Atomic Number (Z):
Defined as the number of protons in an atom.
This number uniquely identifies an element and will never change for a given element.
It is typically found on the periodic table.
Mass Number (A):
Defined as the sum of protons and neutrons in an atom (A = ext{protons} + ext{neutrons}).
Represents the total number of nucleons.
Isotope Notation:
An element is typically represented as _Z^A ext{X}, where (X) is the element symbol.
Isotopes:
Atoms of the same element (meaning they have the same atomic number, Z) but have different numbers of neutrons.
Since they have different numbers of neutrons, they will have different mass numbers (A).
Example: Iron can have different isotopes, all with the same number of protons (characteristic of iron) but varying numbers of neutrons.
The number of electrons usually balances the number of protons to maintain neutrality in an atom.
Atomic Mass Unit (AMU):
Defined as 1/12 the mass of a carbon-12 atom.
This unit provides a standard for measuring the mass of atoms, taking into account the varying weights of different atoms and their isotopes in a sample.