Comprehensive Final Review: Units 1, 2, & 3 Notes

Unit 1: Matter

• All matter is composed entirely of atoms. • Matter can exist in various forms; for example, water can be found as ice, liquid, or steam. • The specific arrangement and movement of atoms are the factors that determine an object’s state of matter. • Kinetic Energy ($KE$) is the primary form of energy that determines the state of matter.     • Kinetic energy is defined as the energy of movement an object possesses.     • To determine the state of matter, scientific observation focuses on the kinetic energy of the object’s individual particles. • Solids:     • Represent the state of matter with the lowest kinetic energy.     • Characterized by having a definite size, volume, and shape.     • Particles are packed close together with insufficient room for significant movement. • Liquids:     • Possess more kinetic energy than solids.     • Characterized by having a definite volume but no definite shape; they take the shape of their container.     • Particles can move past each other and are held together by intermolecular bonds. • Gas:     • Possess more kinetic energy than liquids.     • Characterized by having no defined volume or shape.     • Particles move at high speeds, and there is a significant amount of space between them. • Phase Changes:     • A change in the amount of kinetic energy within matter can result in a change in its state.     • Changes in kinetic energy directly affect the speed of the matter’s molecules (speeding them up or slowing them down).     • Sublimation: involves an increase in kinetic energy ($+KE$).     • Deposition: involves a decrease in kinetic energy ($-KE$).     • Melting: involves an increase in kinetic energy ($+KE$).     • Solidification: involves a significant decrease in kinetic energy ($-KE$).     • Evaporation: involves an increase in kinetic energy ($+KE$).     • Condensation: involves a decrease in kinetic energy ($-KE$).

The Periodic Table Structure and Properties

• Groups:     • These are the vertical columns on the periodic table, read from top to bottom.     • They are also referred to as families.     • There are a total of $18$ groups.     • Elements within the same group tend to react in a similar fashion. • Periods:     • These are the horizontal rows on the periodic table, read from left to right.     • There are a total of $7$ periods.     • An element’s row number (period) is equal to the number of electron orbits it possesses. • Metals:     • Located to the left of the zig-zag line on the periodic table.     • Properties: Lustrous (shiny), malleable (can be hammered into sheets), ductile (can be drawn into wires), and typically silver in color.     • They are good conductors of heat and electricity and possess low specific heat. • Nonmetals:     • Properties: Brittle (breaking easily with a smooth fracture), dull in appearance, and vary in color.     • They have lower densities than metals and usually gain or share electrons during the bonding process. • Metalloids:     • Located directly on the zig-zag line.     • Properties: Solids that are brittle and hard.     • They are somewhat reactive and function as semiconductors. • Specific Families Identified on the Table:     • Group 1: Alkali metals.     • Group 2: Alkaline earth metals.     • Groups 3–12: Transition metals.     • Group 15: Pnictogens.     • Group 16: Chalcogens.     • Group 17: Halogens.     • Group 18: Noble gases.     • Lanthanoids and Actinoids (located at the bottom).

Unit 2: Elements and Atomic Theory

• Historical Development of Atomic Theory:     • Kanad: An Indian philosopher who originated the first idea of the atom.     • Democritus & Leucippus: Greek philosophers who developed similar early concepts of atoms.     • John Dalton: Known as the founder of atomic theory; he developed the formal atomic theory and the first scientific model of the atom.     • J.J. Thomson: Discovered the electron and proposed the "plum pudding" model.     • Ernest Rutherford: Discovered the nucleus and concluded that atoms consist mostly of empty space.     • Niels Bohr: Proposed that electrons move between specific orbits and that these orbits have limits on the number of electrons they can hold.     • Werner Heisenberg: Developed the Uncertainty Principle and the quantum model of the atom. • Parts of an Atom:     • Protons: Carry a positive charge ($+$), have a mass of $1$ atomic mass unit ($amu$), and are located in the nucleus. The number of protons determines the identity of the element.     • Neutrons: Carry a neutral charge ($0$), have a mass of $1$ atomic mass unit ($amu$), and are located in the nucleus.     • Electrons: Carry a negative charge ($-$), have a mass of approximately $0\,amu$, and are located in the electron cloud outside the nucleus. • Atomic Neutrality:     • The overall charge of a standard atom is neutral.     • In a neutral atom, the number of protons must equal the number of electrons ($p^+ = e^-$). • Calculating Subatomic Particles:     • Protons ($p^+$) = Atomic Number.     • Neutrons ($n^0$) = Atomic Mass $-$ Atomic Number.     • Electrons ($e^-$) = For a neutral atom, this is equal to the number of protons. • Electron Shell Capacities:     • Shell 1: Maximum capacity of $2$ electrons.     • Shell 2: Maximum capacity of $8$ electrons.     • Shell 3: Maximum capacity of $18$ electrons.     • Shell 4: Maximum capacity of $32$ electrons.

Valence Electrons and Dot Diagrams

• Valence Electron ($ve^-$) Patterns:     • Group 1: $1\,ve^-$     • Group 2: $2\,ve^-$     • Group 13: $3\,ve^-$     • Group 14: $4\,ve^-$     • Group 15: $5\,ve^-$     • Group 16: $6\,ve^-$     • Group 17: $7\,ve^-$     • Group 18: $8\,ve^-$ • Rules for Electron Dot Diagrams:     • These diagrams include the element’s chemical symbol surrounded by dots representing the valence electrons.     • No more than two electrons may be placed on any single side of the symbol.     • The dots must be large enough to be clearly visible.     • The element symbol must be written exactly as it appears on the Periodic Table.

Unit 3: Isotopes and Ions

• Isotopes:     • Atoms of the same element that have the same number of protons but a different number of neutrons.     • Different isotopes of an element share the same chemical and physical properties as the base element (e.g., melting or boiling points).     • Naming convention: Element Name $-$ Mass Number (e.g., Carbon-$14$). • Ions:     • An atom that has acquired an electrical charge.     • Formed when an atom gains or loses valence electrons.     • In an ion, the number of protons does NOT equal the number of electrons.     • The mass number is not affected by ion formation.     • Anion: Formed by gaining an electron; the atom becomes a negative ion.     • Cation: Formed by losing an electron; the atom becomes a positive ion. • Ion Calculation Data:     • Oxygen $-1$ ion (Group $16$): Gaining electron. Calculation: $8 + 1 = 9$ total electrons.     • Calcium $+1$ ion (Group $2$): Losing electron. Calculation: $20 - 1 = 19$ total electrons.     • Germanium $-3$ ion (Group $14$): Gaining electron. Calculation: $32 + 3 = 35$ total electrons.     • Carbon $+3$ ion (Group $14$): Losing electron. Calculation: $6 - 3 = 3$ total electrons.     • Boron $+2$ ion (Group $13$): Losing electron. Calculation: $5 - 2 = 3$ total electrons. • Orbital Changes in Ions:     • Elements can gain or lose valence electrons, but a valence orbit always exists.     • If an element loses all existing valence electrons, that orbit is considered empty and no longer there; this results in the creation of a new valence orbit from the shell beneath.     • Example: Neutral Magnesium ($12\,p^+$, $12\,n^0$) has $2\,ve^-$. Upon losing those $2\,ve^-$, it becomes a Magnesium $+2$ ion with $8\,ve^-$ in its new valence orbit.     • Adding valence electrons can also lead to the formation of new valence orbits. • Ion Dot Diagrams:     • Because a valence orbit always exists, no electron dot diagram should ever feature zero dots.     • When drawing an ion, the charge must be clearly indicated at the top left (e.g., $2+$ or $2-$).