Chemistry: Matter and Compounds

Matter

• Definition: Anything that occupies space.
• Examples: Pure substances and mixtures.
• Mass: The total amount of matter.
• Weight: Similar to mass, but influenced by gravity.
• Atom: The basic unit of matter.
• States of Matter: Solid, liquid, gas, plasma, Bose-Einstein condensate.
• Elements: Mercury, antimony, uranium.
• Compounds: Glucose, salt, ammonia.

Naming Compounds

• Prefixes for Numbers:
  • di - 2
  • tri - 3
  • tetra - 4
  • penta - 5
  • hexa - 6
  • hepta - 7
  • octa - 8
  • nona - 9
  • deca - 10
• Example: SO₄ (Sulfur Tetroxide)
• Example: CCl₅ (Carbon Pentachloride).
• The second element suffix changes to ide/ite/ate.

Subatomic Particles

• Proton: Positive charge.
• Electron: Negative charge.
• Neutron: No charge.
• Finding Proton Value: The atomic number.
• Finding Electron Value: Same as the proton value unless otherwise specified with a charge.
  • Subtract charge from electron value for positive ion.
  • Add charge to electron value for negative ion.
• Finding Neutron Value: Subtract the atomic number from the atomic mass.

Nuclide Symbol

• Parts: Mass number (A), charge (N), atomic symbol (X), atomic number (Z).
• Location of Parts:
  • Atomic mass (A): Upper right of the symbol.
  • Charge (N): Upper left of the symbol.
  • Element symbol (X): Center of the symbol.
  • Atomic number (Z): Lower left of the symbol.
• Example (Bromine): ⁸⁰₋₁Br³⁵

Historical Development of Atomic Theory

• 400 BC:
  • Leucippus: Atomos (uncut).
  • Democritus: Developed Leucippus' idea.
  • Aristotle: Rejected atoms, proposed four elements (fire, earth, water, air).
• 1774: Antoine Lavoisier—Law of Conservation of Mass (matter can't be created or destroyed).
• 1779: Joseph Louis Proust—Law of Definite Proportions (same compound, same mass ratio).
• 1803: John Dalton—Dalton's Atomic Theory (5 points):
  • Matter consists of atoms. (Partially false)
  • Atoms are indestructible. (Partially false)
  • Atoms of the same element have identical properties. (True- law of definite proportions)
  • Compounds are composed of atoms in whole number ratios. (True - law of definite composition)
  • Chemical reactions involve combination/separation/rearrangement of atoms. (True- law of conservation of mass)
• 1879: Sir William Crookes—Cathode ray tube.
• 1897: J.J. Thomson—Plum pudding model.
• 1909: Ernest Rutherford—Alpha scattering experiment.
• 1923: Niels Bohr—Electron orbits/energy levels.
• 1932: James Chadwick—Neutron discovery.

Isotopes

• Atoms of the same element with different numbers of neutrons.
• Examples: Carbon-12, Carbon-13, Carbon-14.

Periodic Table

• Historical Development:

  • Dobereiner: Law of Triads.
  • De Chancourtois: Helical graph.
  • Newlands: Law of Octaves.
  • Lothar Meyer: Arranged elements by atomic weight.
  • Mendeleev: Developed a periodic table with increasing atomic mass, organized elements into 17 columns.
  • Rayleigh/Ramsay: Noble gases.
  • Moseley: Rearranged based on atomic number.
  • Seaborg: Actinides/Lanthanides.

• Arrangement:

  • Groups (vertical): 1A-8A, transition metals (1B-8B) , labeling the families
  • Periods (horizontal): Shows the highest energy level.

• Group Names (1A-8A):

  • 1A: Alkali metals
  • 2A: Alkaline earth metals
  • 3A: Boron family
  • 4A: Carbon family
  • 5A: Nitrogen family
  • 6A: Oxygen family
  • 7A: Halogens
  • 8A: Noble/inert gases

Electronic Configuration

• Arrangement of electrons in orbitals around the nucleus.
• Subshells: s (max 2), p (max 6), d (max 10), f (max 14).
• Example (Neon): 1s²2s²2p⁶ (atomic number 10).
• Noble Gas Notation:
  • Use noble gas before element in brackets and write the remaining electrons
  • Example : Silicon (Si) [Ne]3s²3p²

Valence Electrons

• Same number as the group number (for A groups).

This quiz covers the fundamental concepts of matter, including definitions, states, and examples. It also explores the naming of compounds using prefixes and the details of subatomic particles. Test your knowledge on these essential chemistry topics!