Exploring Matter: Properties, Substances, and Mixtures

• Matter occupies space and has mass.
• Mass is the total amount of matter.
• Weight is affected by gravity.
• Inertia is a matter property that keeps it at rest or in motion unless acted upon by an external force.
• Atoms are matter's basic unit.
• Solids have definite shape and volume.
• Liquids have definite volume, taking the container's shape.
• Gases lack definite shape and volume.
• The triple point is where all matter phases meet.

Properties of Matter

• Physical properties are observed without changing the matter's composition, including color, state, odor, boiling point, and melting point.
• Intensive properties don't depend on the amount of material like melting/boiling point, color, hardness, taste, density, and conductivity.
• Extensive properties depend on the amount of material like length, surface area, volume, mass, weight, width and shape.
• Chemical properties are observed during a chemical change.

Pure Substances and Mixtures

• Pure substances have fixed composition and include elements and compounds.
• Elements contain one type of atom.
• Compounds have 2+ atoms of different elements held by chemical bonds.
• Mixtures vary in composition.
• Homogeneous mixtures have a uniform phase.
• Heterogeneous mixtures are non-uniform.

Atoms and Molecules

• Molecules form when a chemical bond exists.
• Hydrogen, Oxygen, Nitrogen, Chlorine, Bromine, Fluorine and Iodine exist as diatomic molecules.

Atomic Structure

Development of the Atomic Theory

• 400 BC: Leucippus proposed that matter consists of indivisible particles called "atomos".
• 384-322 BC: Aristotle theorized that all matter consisted of water, fire, earth, and air.
• 1774: Antoine Lavoisier's experiments led to the Law of Conservation of Mass: matter is neither created nor destroyed.
• 1779: Joseph Louis Proust developed the Law of Definite Proportion: compounds have constant element proportions by mass.
• 1803: John Dalton's theory stated that matter consists of indivisible atoms that cannot be created nor destroyed.
• 1879: William Crookes' cathode ray tube discovery showed atoms contain charged particles.
• 1897: J.J Thomson discovered electrons and proposed the Plum Pudding model.
• 1909: Ernest Rutherford discovered the nucleus, containing positively charged protons.
• 1922: Niels Bohr proposed the Planetary model of the atom.
• 1932: James Chadwick discovered neutrons through experiments with beryllium and alpha particles.

Subatomic Particles

• The nucleus is at the atom's center, tiny, dense, and positively charged.
• Protons are positively charged subatomic particles.
• Electrons are negatively charged subatomic particles.

Atomic Number and Mass Number

• Atomic number equals the number of protons.
• Mass number equals the number of protons + neutrons.
• Number of neutrons = mass number - atomic number.
• For a neutral atom, the number of protons equals the number of electrons.
• For non-neutral atoms, negatively charged atoms have extra electrons, positively charged atoms have fewer electrons.

Historical Development of the Periodic Table

• 1862: Johann Wolfgang Dobereiner grouped elements in triads, known as the Law of Triads.
• Alexander Beguyer de Chancourtois prepared a helical graph of known elements wrapped around a cylinder.
• 1863: John Newlands arranged elements by atomic weight, noting properties repeated every 8th element, known as the Law of Octaves.
• 1864: Lothar Meyer arranged 26 elements by atomic weights.
• 1869: Dmitri Mendeleev arranged elements by atomic mass, with his periodic table originally having 17 columns.
• 1894: Lord Rayleigh discovered noble/inert gases, like argon.
• Period indicates the highest energy level of outermost electrons.
• The vertical column is a group/family.

Groups

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

Trends in the Periodic Table

• Atomic Radius refers to the size of an atom.
• Ionization Energy increases from left to right and decreases from top to bottom.
• Electron Affinity releases energy.
• Electronegativity measures an atom's attraction for shared electrons in a chemical bond.
• Metallic Property decreases from left to right.