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.