Atoms, Molecules, and Subatomic Particles

• Atoms and molecules form the basic building blocks of matter.
• Different types of matter arise because they contain different atoms.

Subatomic Particles and Atomic Models

• Scientists in the late 19th century aimed to determine the structure of atoms and explain their properties.
• Elucidating atomic structure relies on experimental evidence.
• Static electricity experiments showed that atoms are divisible.

Discovery of Charged Particles

• Activity 4.1 showcases that rubbing objects together can electrically charge them.
• Atoms are divisible and consist of charged particles.
• Around 1900, it became clear that atoms contain subatomic particles like the electron.

Electron and Proton

• J.J. Thomson identified the electron.
• E. Goldstein discovered canal rays in 1886, indicating positively charged radiation.
• This radiation led to the discovery of another subatomic particle.
• Protons have a charge equal in magnitude but opposite in sign to electrons.
• Protons are ~2000 times more massive.
• Electrons are represented as 'e⁻' and protons as 'p⁺'.
• An electron’s mass is considered negligible.
• An electron possesses a charge of minus one, while a proton has a charge of plus one.
• The discovery of electrons and protons suggested that atoms might consist of both, with their charges balanced.
• The protons reside in the atom's interior, while electrons can be removed more easily.

Thomson's Atomic Model (1904)

• J.J. Thomson was the first to propose a model for the structure of an atom.
• Thomson's model pictured the atom as a positively charged sphere with electrons embedded in it.
• Postulates of Thomson model
• An atom consists of a sphere of positive charge with electrons embedded in it.
• Negative and positive charges are equal in magnitude, making the atom electrically neutral.
• Example: plum pudding. A Christmas pudding with electrons like currants, or a watermelon with seeds.
• Although Thomson's model explained electrical neutrality, it could not explain the results of later experiments.

Rutherford's Atomic Model

• Ernest Rutherford aimed to understand electron arrangement within the atom.
• Rutherford designed an experiment involving fast-moving alpha (α)-particles.
• Alpha particles were directed at a thin gold foil about 1000 atoms thick.
• Alpha particles are doubly-charged helium ions, mass of 4 u, and considerable energy.
• It was expected the alpha particles would deflect slightly due to subatomic particles.
• The alpha particles are much heavier than the protons.
• Observations from the alpha-particle scattering experiment were unexpected:
• Most alpha-particles passed straight through the gold foil without deflection.
• Some alpha-particles were deflected by the foil at small angles.
• Approximately 1 in 12,000 alpha particles rebounded.
• Rutherford concluded:
• Most of the space inside the atom is empty.
• The positive charge occupies very little space.
• the positive charge and mass are concentrated in a small volume within the atom.

Rutherford's Nuclear Model of the Atom

• Based on these scattering results, also calculated the radius of the nucleus is about 10^5 times less than the radius of the atom.
• There is a positively charged center in an atom called the nucleus.
• Nearly all the mass of an atom resides in the nucleus.
• Electrons revolve around the nucleus in circular paths forming the bulk of the atom.
• The size of the nucleus is very small compared to the size of the atom.
• The model suffered from a drawback.
• Electrons revolving in circular orbits would undergo acceleration and radiate energy.
• Revolving electrons would lose energy and eventually fall into the nucleus.
• If this occurred, atoms would be unstable, contradicting the observed stability of matter.

Bohr's Model of Atom

• Niels Bohr addressed the issues with Rutherford's model and made the following postulates:
• Only specific orbits for electrons exist within the atom.
• Electrons do not radiate energy while revolving in these orbits.
• These orbits, or shells, represent energy levels within the atom.

Neutrons and Atomic Mass

• James Chadwick discovered neutrons in 1932.
• Neutrons are subatomic particles with no charge (neutral).
• Neutrons have a mass nearly equal to that of a proton.
• Present in the nucleus of all atoms except hydrogen.
• The mass of an atom is determined by the sum of the masses of the protons and neutrons in the nucleus.

Electron Distribution in Orbits

• Bohr and Bury proposed rules for how electrons fill energy levels or shells:
• The maximum number of electrons in a shell is given by 2n², where n is the orbit number.
• n = 1,2,3..
• The first orbit (K-shell) can hold up to 2 electrons = 2 × 1² = 2
• The second orbit (L-shell) can hold up to 8 electrons = 2 × 2² = 8
• The third orbit (M-shell) can hold up to 18 electrons = 2 × 3² = 18
• The fourth orbit (N-shell) can hold up to 32 electrons = 2 × 4² = 32, and so on.
• The outermost orbit can accommodate a maximum of 8 electrons.
• Electrons are not added to a shell until the inner shells are filled.

Valency and Chemical Behavior

• Electrons in the outermost shell determine the atom's chemical behavior.
• The outermost shell known as the valence shell.
• Atoms with a complete outermost shell (8 electrons, or 2 for hydrogen) exhibit little chemical activity.
• Atoms combine to achieve a full outermost shell, known as an octet (or duplet for hydrogen).
• Atoms share, gain, or lose electrons to achieve an octet.

Valency

• The valency is the number of electrons gained, lost, or shared to achieve an octet.
• Elements with one electron in their outermost shell have a valency of one.
• Magnesium has valency of two, aluminium has a valency of three
• Number of electrons close to a full capacity, valency is determined differently.
• Flourine, with seven electrons is easier to gain one electron rather than lose seven, hence a valency of one.

Atomic Number (Z)

• The atomic number (Z) is the number of protons in an atom's nucleus.
• All atoms of an element have the same atomic number.
• Hydrogen, Z = 1
• Carbon, Z = 6

Mass Number (A)

• The mass number (A) is the total number of protons and neutrons (nucleons) in an atom's nucleus
• For example:
• Carbon is 12 u because it has 6 protons and 6 neutrons (6 u + 6 u = 12 u)
• Aluminium is 27 u (13 protons + 14 neutrons)
• Notation for an atom includes the symbol, atomic number, and mass number
• For example , nitrogen (N), ^14N(7)

Isotopes

• Isotopes are atoms of the same element with the same atomic number but different mass numbers.
• Hydrogen has three isotopes: protium (¹H), deuterium (²H or D), and tritium (³H or T).
• 12C6 and 14C6

Average Atomic Mass

• Calculated based on the relative abundance of each isotope in nature.
• If the element has no isotopes, then the mass of an element is the same as the sum of the mass of protons and neutrons.

Applications of Isotopes

• Uranium-235: used as fuel in nuclear reactors.
• Cobalt-60: used in cancer treatment.
• Iodine-131: used in treating goiters.

Isobars

• Isobars are atoms of different elements (different atomic numbers) that have the same mass number.
• Calcium (atomic number 20) and argon (atomic number 18) both have a mass number of 40.

Explore the fundamental building blocks of matter with atoms and molecules. Investigate the structure of atoms, their divisibility, and the discovery of charged particles like electrons and protons through experiments and radiation.