Chemistry

I. The Periodic Table

• Organization:

  • Periods: Horizontal rows. Elements in the same period have the same number of electron shells.

  • Groups (Families): Vertical columns. Elements in the same group have similar chemical properties because they have the same number of valence electrons.

  • Metals, Nonmetals, Metalloids: Be able to identify these categories on the periodic table.

    • Metals: Generally shiny, conductive, malleable, and ductile. Found on the left and middle of the table.

    • Nonmetals: Generally dull, non-conductive, and brittle. Found on the right side of the table.

    • Metalloids (Semimetals): Have properties of both metals and nonmetals. Located along the "staircase" separating metals and nonmetals.

• Key Groups (Families):

  • Alkali Metals (Group 1): Highly reactive metals.

  • Alkaline Earth Metals (Group 2): Reactive metals.

  • Halogens (Group 17): Highly reactive nonmetals.

  • Noble Gases (Group 18): Generally inert (unreactive) gases.

• Periodic Trends: Be able to explain and predict trends in:

  • Atomic Radius: Generally increases as you go down a group, decreases as you go left to right across a period.

  • Ionization Energy: Generally decreases as you go down a group, increases as you go left to right across a period.

  • Electronegativity: Generally decreases as you go down a group, increases as you go left to right across a period.

• Practice:

  • Identify elements based on their position on the periodic table.

  • Explain how an elements' position on the table is related to its properties.

  • Predict and explain trends in atomic radius, ionization energy, and electronegativity.

II. Ions

• Definition: An atom or molecule that has gained or lost electrons, thus having a net electric charge.

• Types:

  • Cations: Positively charged ions formed when an atom loses electrons (usually metals).

  • Anions: Negatively charged ions formed when an atom gains electrons (usually nonmetals).

• Formation: Ions are formed to achieve a stable electron configuration (typically a full outer electron shell). This is related to the octet rule (wanting to have 8 valence electrons)

• Ionic Charge:

  • Metals tend to lose electrons to form positive ions with a charge that is based on what it needs to gain an octet.

  • Nonmetals tend to gain electrons to form negative ions with a charge that is based on what it needs to gain an octet.

• Naming:

  • Cations: The element name followed by "ion" (e.g. sodium ion)

  • Anions: The element name with the ending replaced by "-ide" (e.g., chloride)

• Practice:

  • Predict the ionic charge of common elements based on their position on the periodic table.

  • Name ions using correct terminology.

  • Write chemical formulas for ionic compounds based on ionic charge.

III. Atomic Orbitals

• Electron Arrangement:

  • Electrons are arranged in shells or energy levels.

  • Within each energy level, electrons occupy sublevels: s, p, d, and f.

  • Each sublevel contains atomic orbitals.

  • s: spherical (1 orbital, holds 2 electrons)

  • p: dumbbell shaped (3 orbitals, hold 6 electrons)

  • d: complex shapes (5 orbitals, hold 10 electrons)

  • f: very complex shapes (7 orbitals, hold 14 electrons)

• Filling Order: Electrons fill orbitals starting from the lowest energy level (Aufbau principle).

• Hund's Rule: Within a sublevel, electrons fill orbitals singly before pairing.

• Electron Configurations: Shorthand notation of electron arrangements within an atom (e.g., 1s2 2s2 2p4 for Oxygen)

  • Be able to derive electron configurations of neutral atoms, and then use that to determine electron configuration of ions.

• Orbital Diagrams: Shows electron configurations and spin (up and down arrows) in the order electrons fill.

• Practice:

  • Draw electron configurations and orbital diagrams of given elements and ions.

  • Relate electronic configurations and atomic numbers.

  • Determine the number of valence electrons for different elements.

IV. The Atomic Nucleus

• Components:

  • Protons: Positively charged particles. Number of protons determines the element's identity.

  • Neutrons: Neutral particles. Help stabilize the nucleus by reducing repulsions between protons.

• Atomic Number: Number of protons in the nucleus. (Z) Identifies the element.

• Mass Number: Total number of protons and neutrons in the nucleus. (A)

• Nuclear Forces:

  • Strong Nuclear Force: Attractive force that holds protons and neutrons together in the nucleus; strongest of the fundamental forces

  • Electrostatic Force: Repulsive force between the positively charged protons in the nucleus.

• Nuclear Stability: Balance between attractive strong nuclear force and repulsive electrostatic force; determines if a nucleus is stable or radioactive

• Practice:

  • Explain the difference between the roles of protons and neutrons.

  • Use atomic number and mass number to determine numbers of protons, neutrons, and electrons.

  • Explain why stable nuclei exist despite the repulsion of positive charges.

V. Isotopes

• Definition: Atoms of the same element (same number of protons) with different numbers of neutrons.

• Notation: Use of the elements name, chemical symbol and number of mass numbers with atomic numbers.

• Atomic Mass: Weighted average mass of all isotopes of an element.

• Relative Abundance: Percentage of each isotope that occurs naturally.

• Practice:

  • Given the number of protons, neutrons, and electrons, determine the isotope.

  • Use abundance to calculate the average atomic mass of an element.

  • Explain why an element will have multiple isotopes with different stability.

VI. Density Modeling

• Definition: A substance's density is its mass per unit volume

• Formula: Density (d) = Mass (m) / Volume (V)

• Density Modeling:

  • Use mathematical models and relationships to predict the density of different substances.

  • Manipulate variables in density to model what it looks like on different graphs

• Applications:

  • Understanding the layers in a stratified fluid or geological system.

  • Designing materials with a desired density.

• Practice:

  • Solve numerical problems based on calculating and interpreting density.

  • Use density to predict mass or volume of a substance.

VII. Nature of Mass

• Mass

  • A measure of the quantity of matter in an object

  • Distinction from Weight: weight measures the pull of gravity on a mass

  • Units of mass are grams and kilograms

• The Law of Conservation of Mass

  • Mass is neither created nor destroyed in a chemical or physical change.

  • Important for reaction calculations.

• Mass vs. Number of Particles

  • Understanding how atoms can have different masses even though we are accounting for every single particle. (neutrons and isotopes)

  • Understanding why we use moles to account for a group of atoms in chemistry, rather than a number of atoms.

• Practice:

  • Solve problems that determine how mass is conserved in a reaction.

  • Understand what a mole is, and how to calculate how much matter is being accounted for with moles.

  • Understand how atoms have mass by understanding subatomic particles.