Chemistry: Matter and Particle Theory

Particle Theory

• All matter is composed of tiny particles.
• Particles are constantly moving.
• Particle movement increases with temperature.
• Particles attract each other.
• Each pure substance has unique particles.

Matter Categories

• Matter is categorized as either pure substances or mixtures.
• Pure substances are further categorized as elements or compounds.
• Mixtures can be heterogeneous or homogeneous.

Physical Properties

• Clarity
• Temperature
• Density
• Malleability
• Luster
• Ductility
• Color
• Conductivity
• State (solid, liquid, gas)
• Hardness
• Melting/boiling point
• Taste
• Odor
• Solubility
• Texture

Chemical Properties

• How a substance reacts with other substances.
• Reaction with acid
• Combustibility (reaction with oxygen)

Chemical and Physical Changes

• Physical change: The substance remains the same.
  • Examples: changes of state, dissolving.
• Chemical change: A new substance is formed.

Signs of a Chemical Change

• Precipitate (solid formed from two solutions)
• Bubbles/gas/odor
• Change in color
• Heat and/or light
• Usually hard to reverse.

Gas Tests

• Physical properties often insufficient to identify gases.
• Specific gas tests are used.

Atomic Structure

• Proton: Found in the nucleus, +1 charge.
• Electron: Found in orbitals, -1 charge.
• Neutron: Found in the nucleus, 0 charge.
• Atomic number (Z): Number of protons, defines the element.
• Mass number: Number of protons and neutrons.
• Isotopes: Atoms of the same element with different numbers of neutrons.
• Normally, number of protons equals number of electrons.
• Neutrons: Hold the nucleus together with the strong nuclear force.
• Unstable nuclei result from too many or too few neutrons.
• Radioactive atoms result from unstable nuclei.

Bohr-Rutherford Model

• Orbitals are rings.
• First orbital holds 2 electrons.
• Second orbital holds 8 electrons.
• Third orbital holds 8 electrons (but can hold 18).

Valence Electrons

• Valence electrons are in the outer orbital.
• Valence electrons determine chemical reactivity.
• Atoms react to achieve a full valence orbital (octet rule).
• Atoms with <4 valence electrons tend to lose electrons.
• Atoms with >4 valence electrons tend to gain electrons.
• Atoms with 4 valence electrons can either gain or lose electrons.
• Group numbers correspond to the number of valence electrons (except transition metals).

Periodic Table Groups

• Group 1: Alkali Metals
• Group 2: Alkaline Earth Metals
• Groups 3-12: Transition Metals
• Group 17: Halogens
• Group 18: Noble Gases

Lewis Diagrams

• Element symbol in the center, dots for valence electrons.

Ions

• An ion is a charged atom (gained/lost electrons).
• Cation: Positive ion (lost electrons).
• Anion: Negative ion (gained electrons).
• Elements tend to gain/lose electrons to achieve full valence shells.
• Ion charge prediction based on group number (transition metals are exceptions).

Naming Ions

• Cations: Element name + "ion".
  • Example: Magnesium ion (Mg²⁺)
• Anions: Element stem + "ide" + "ion".
  • Example: Chloride ion (Cl⁻)

Ionic Compounds

• Formed between metals and nonmetals.
• Metals form cations, nonmetals form anions.
• Ionic bond: Attraction between oppositely charged ions.
• Ionic compounds tend to be hard, brittle, with high melting/boiling points.
• Crystal structure.
• Dissolving: Ions separate in water, water molecules surround them.
• Electrolytes: Dissolve in water to form a solution that conducts electricity.

Naming Ionic Compounds

• Metal name followed by nonmetal name (with anion ending)
• Transition metals use Roman numerals to show ion charge.

Polyatomic Ions

• Stable group of atoms acting as a single charged particle.
• Example: Carbonate (CO³⁻²).
• Many anions - use brackets for multiple polyatomics.

Naming and Formulas of Ionic Compounds w/ Polyatomics

• Element + polyatomic name
• Polyatomics that have subscripts must use brackets - Example: Calcium Nitrate = Ca(NO₃)₂

Molecular Compounds

• Formed between two or more nonmetals.
• Covalent bonds: Electrons are shared.
• Lower melting/boiling points than ionic compounds.
• Usually do not dissolve in water.
• Generally do not conduct electricity.

Diatomic Molecules

• 7 diatomic elements: HOFBrINCl (Hydrogen, Oxygen, Fluorine, Bromine, Iodine, Nitrogen, Chlorine).

Naming Molecular Compounds

• Use prefixes (mono, di, tri, tetra, penta, hexa, hepta, octa) to indicate numbers of atoms.
• "Mono" not used for first element.
• Examples: CO (Carbon monoxide), N₂O₅ (dinitrogen pentoxide).

Chemical Reactions

• Reactants → Products
• Word equations vs. chemical equations (with coefficients).

Balancing Chemical Equations

• Use coefficients to ensure equal numbers of each atom on both sides of the equation.
• Never change subscripts.
• Follow steps to balance.

Types of Chemical Reactions

• Synthesis (combination): A + B → AB
• Decomposition: AB → A + B
• Single displacement: AB + C → AC + B
• Double displacement: AB + CD → AD + CB

Combustion & Corrosion

• Combustion: Rapid reaction with oxygen, produces oxides & energy.
• Incomplete combustion: Limited oxygen supply, produces soot (carbon) and carbon monoxide (CO).
• Corrosion: Breakdown of a metal due to chemical reactions with environment.
• Rusting: Specific type of corrosion affecting iron. Preventing corrosion: coatings, materials, galvanizing.
  • Coatings: Protecting the metal.

Acids and Bases

• Acids: Substance that releases H+ ions in water.

  • Taste: sour
  • Reaction with metals: Produce H₂ gas.
  • Reaction with carbonates: Produce CO₂ gas.
  • Often start with hydrogen.

• Bases: Substances that release OH⁻ or CO₃²⁻ ions in water.

• Taste: bitter

• Conduct electricity when dissolved in water

• Often contain hydroxide (OH⁻), or carbonate (CO₃²⁻)

Naming Acids

• Binary acids: Hydro + stem + "ic" acid (Ex: HCl → hydrochloric acid).
• Oxyacids: Stem end + "ic" or "ous" (Ex: HNO₃ → nitric acid).

pH Scale

• pH measures acidity/basicity of a solution (concentration of H⁺ ions).
• Low pH = strong acid, high pH = strong base.
• pH 7 = neutral.
• pH scale is logarithmic.