LC

Unit 1 Review

Unit 1 Review: Matter and Qualitative Analysis

Early Atomic Theories & Atomic Structure

  • Solid Sphere Model: Proposed by John Dalton (1803)

  • Plum Pudding Model: Proposed by J.J. Thomson (1904)

  • Nuclear Model: Proposed by Ernest Rutherford (1911)

  • Planetary Model: Proposed by Niels Bohr (1913)

  • Quantum Model: Proposed by Erwin Schrödinger (1926)

Dalton’s Theory (1803)

  • Described the atom as a uniformly dense, solid sphere.

  • Key Points:

    • Matter is made of indestructible particles called atoms.

    • Law of Conservation of Mass: Atoms cannot be created or destroyed.

    • Matter consists of definite particles called atoms.

    • Atoms of different elements have unique properties.

    • Atoms can combine in constant ratios to form new substances.

Thomson’s Theory (1897)

  • Model: Raisin Bun/Plum Pudding Model.

  • Discovered electrons via cathode ray tube experiments.

  • The remainder of the atom is a positively charged sphere with empty space.

  • Atoms can gain or lose electrons, forming ions.

  • Element identity is based on electron count.

Rutherford’s Theory (1912)

  • Conducted the Gold Foil Experiment:

    • Most particles passed straight through, indicating empty space.

    • Some particles bounced back, revealing a tiny positive nucleus.

  • Nucleus Details:

    • Contains positive charges (protons).

    • Electrons are located outside the nucleus in volumes.

    • Protons determine element identity, with Chadwick later identifying neutrons.

Bohr’s Theory (1913)

  • Electrons exist in specific energy levels (orbits).

  • Electrons can move to higher energy levels with energy input; emitted light when falling back.

Standard Atomic Notation

  • Mass Number (Atomic Mass): # protons + neutrons (rounded).

  • Atomic Number: # of protons (equal # of electrons in a neutral atom).

  • Neutrons: Mass Number - Atomic Number.

Isotopes

  • Atoms with the same number of protons but different numbers of neutrons.

  • Example: C-12 vs. C-14 (varying mass numbers but similar chemical properties).

Bohr-Rutherford Model

  • Protons and neutrons in the nucleus; electrons in energy levels or orbitals.

  • Orbital Capacities:

    • Orbital #1: 2 electrons

    • Orbital #2: 8 electrons

    • Orbital #3: 8 electrons.

The Octet Rule

  • Atoms tend to achieve electron configurations like the nearest noble gas.

  • Typically results in 8 outer electrons.

Lewis Dot Diagrams

  • Show only valence electrons in an atom.

  • Important for understanding atomic bonding.

Organization of the Periodic Table

  • Periodic Table: Arranged by increasing atomic number.

  • Mendeleev's initial version left gaps for undiscovered elements.

  • Periodic Law: Properties of elements recur periodically as atomic number increases.

Groups of Elements

  • Groups/Families: Vertical columns showing similar chemical properties due to similar valence electron numbers.

  • Reactivity Trends: Group 1 (alkali metals) react violently with water; Group 18 (noble gases) are non-reactive.

Group 1: Alkali Metals

  • Valence Electrons: 1

  • Properties:

    • Solids, shiny, soft, extremely reactive.

    • React with water to produce hydrogen gas.

    • Form +1 ions.

Group 2: Alkaline Earth Metals

  • Valence Electrons: 2

  • Properties:

    • Solids, denser than Group 1.

    • Form +2 ions.

Group 17: Halogens

  • Valence Electrons: 7

  • Properties: Nonmetals with distinct visual appearances.

  • Extremely reactive; form -1 ions.

Group 18: Noble Gases

  • Valence Electrons: 8 (full).

  • Properties: Nonmetals, stable, odorless, colorless.

Transition Metals (Groups 3-12)

  • Exhibit variable reactivity and higher melting points.

Lanthanides and Actinides

  • Inner transition metals, rare and found below the periodic table.

Chemical Change vs. Physical Change

  • Physical Change: Substance changes form but retains chemical identity (e.g., melting, boiling).

  • Chemical Change: Substance transforms into one or more new substances with different properties (e.g., combustion).

Recognizing Chemical Changes

  • Evidence of a chemical change:

    • Color change.

    • Formation of gas.

    • Change in odor.

    • Formation of a precipitate.

Chemical Reactions Basics

  • Reactants: Substances before the reaction.

  • Products: New substances formed by the reaction.

Balancing Chemical Equations

  • Law of Conservation of Mass applies: total mass of reactants = total mass of products.

  • Steps include determining atom counts and equalizing on both sides of the equation.

Types of Chemical Reactions

  • Synthesis, Decomposition, Single Displacement, Double Displacement, Neutralization, and Combustion.

  • Synthesis Reactions: Combine two or more reactants to form one product.

  • Decomposition Reactions: A compound breaks down into simpler substances.

  • Single Displacement Reactions: One element replaces another in a compound; can occur for metals and nonmetals.

  • Double Displacement Reactions: Exchange of cations between two different compounds, often forming a precipitate.

Neutralization Reactions

  • Acid + Base → Salt + Water.

Combustion Reactions

  • Hydrocarbons react with oxygen to produce carbon dioxide and water.

  • General formula: CxHy + O2 → CO2 + H2O.

Ionic Compounds and Bonding

  • Formed from metals and nonmetals through electron transfer, resulting in ionic bonds.

  • Ionic compounds are usually solids with high melting points, brittle, and form crystal lattices.

Molecular Compounds

  • Formed by sharing electrons between nonmetals (covalent bonds).

  • Typically have lower melting points, can exist in various states, and do not conduct electricity.