Classification of Matter and Atomic Theory

Classification of Matter

• Matter is defined as anything that possesses mass and volume and takes up space.

• Matter exists in three primary states: solid, liquid, or gas.

• Mixture: This is matter that can be separated into its individual components by physical means.

  • A mixture does not have a definite composition.

  • Heterogeneous Mixture (Mechanical Mixture): A mixture where the different components are visible to the naked eye. Examples include:

    • Salad dressing.

    • Orange juice.

    • Chilli.

  • Homogeneous Mixture (Solution): A mixture where the different components are not visible. Examples include:

    • Iced tea.

    • Tap water.

• Pure Substance: This refers to matter that has a definite composition, which can be represented by a chemical formula.

  • Element: A substance that cannot be chemically broken down into simpler substances. Examples include:

    • Oxygen ($O_2$).

    • Sodium ($Na$).

    • Gold ($Au$).

  • Compound: A substance formed when two or more elements are chemically combined. Examples include:

    • Water ($H_2O$).

    • Salt ($NaCl$).

Fundamental Concepts of Atoms

• All matter is composed of tiny particles known as atoms.

• An atom is the smallest particle that still characterizes and retains the properties of a chemical element.

• Atoms are made up of smaller constituent parts called subatomic particles.

• There are four major atomic theories that have been developed over time:

  • Dalton's Atomic Theory.

  • Thomson's Atomic Theory.

  • Rutherford's Atomic Theory.

  • Bohr's Atomic Theory.

Historical Development of Atomic Theory

• Dalton’s Atomic Theory:

  • All matter is constructed from atoms.

  • Atoms are the smallest possible particles of matter.

  • In this theory, the atom is modeled as a solid, uniform sphere.

  • This specific model is commonly referred to as “The Billiard Ball Model”.

• J.J. Thomson’s Atomic Theory:

  • Postulated that every atom contains electrons.

  • Electrons are described as being embedded within a cloud of positive charge.

  • This model is commonly referred to as “The Raisin-Bun Model”.

• Rutherford’s Atomic Theory:

  • Proposes that atoms have a positive nucleus at the center, which contains protons.

  • Negatively charged electrons rotate around the central nucleus.

  • The majority of the atom's mass is concentrated within its nucleus.

  • This model is commonly referred to as “The Solar System Model”.

• Bohr’s Atomic Theory:

  • Electrons are associated with specific, discrete energy levels.

  • Electrons have the ability to move from one energy level to another, but they are forbidden from existing in the spaces between energy levels.

  • This model is commonly referred to as the “Energy Level Model”.

Subatomic Particles and Atomic Structure

• The Nucleus: The central part of the atom containing nucleons.

• Nucleons: These are the particles found in the nucleus, which include:

  • Protons: Carry a positive charge.

  • Neutrons: Carry a neutral charge.

• Electrons:

  • Carry a negative charge.

  • Located in energy levels surrounding the nucleus.

Atomic Identity and Nuclear Notation

• Atomic Number:

  • Represents the number of protons found in the nucleus of an atom.

  • This number identifies the specific element.

  • The atomic number can never be changed for a specific element (e.g., Silver, $Ag$, has an atomic number of $47$).

• Mass Number:

  • This is the total number of protons plus neutrons in the nucleus.

  • Example: Helium ($He$) has an atomic number of $2$ and an atomic mass of approx $4.0026$.

• Nuclear Notation:

  • A standard way to represent an element's atomic statistics: ${}^A_Z X$

  • $X$ = Chemical symbol for the element.

  • $A$ = Mass number ($A = Z + N$).

  • $Z$ = Atomic number (number of protons).

  • $N$ = Neutron number.

  • Example: A Sodium ion noted as ${}^{23}_{11} Na^+$ signifies a mass number of $23$, an atomic number of $11$, and a positive charge.

• Calculating Neutrons:

  • The number of neutrons can be determined using the following formula:

  • $\text{Number of neutrons} = \text{mass number} - \text{atomic number}$

Isotopes

• Isotopes are defined as atoms of the same element that have the same number of protons but contain different numbers of neutrons.

• Isotopes of Hydrogen:

  1. Protium: Consists of $1$ proton and $0$ neutrons.

  2. Deuterium: Consists of $1$ proton and $1$ neutron.

  3. Tritium: Consists of $1$ proton and $2$ neutrons.

• Isotopes of Carbon:

  • Carbon exists in various isotopic forms including:

    • Carbon-$12$ ($^{12}C$).

    • Carbon-$13$ ($^{13}C$).

    • Carbon-$14$ ($^{14}C$).