Atomic Theory and Structure Quiz

Atomic Theory and Structure

• Dalton's Atomic Theory: Proposed that elements are composed of indivisible particles called atoms; atoms of the same element are identical; atoms of different elements are different; atoms combine to form new substances; atoms cannot be created or destroyed.
• JJ Thomson's Discovery: Discovered electrons, negatively charged particles within the atom, using a cathode ray tube.
• Niels Bohr's Theory: Proposed that electrons orbit the nucleus in specific energy levels; electrons absorb or emit energy when transitioning between levels.

Atomic Structure - Energy Levels

• Bohr's Energy Levels: Electrons orbit the nucleus at defined energy levels (lowest energy in level 1); electrons can transition between levels by gaining or losing energy; filling order: 1st, then 2nd and 3rd
• Light Wavelength: Color of light emitted or absorbed depends on the specific energy levels electrons jump between (electron transitions)

Atomic Mass Unit (amu)

• Atomic Mass Unit (amu): A unit of mass equal to 1/12 the mass of a carbon-12 atom.

Elements and Properties

• Alkali Metals: Shiny, soft, solids at room temperature; highly reactive; always bonded with another element in nature
• Alkaline Earth Metals: Light, highly reactive metals; react with oxygen to form oxides; solid at room temperature.
• Transition Metals: Form multivalent ions; variable reactivity; good conductors of heat and electricity; many react with oxygen to make oxides.
• Halogens: Can be solid, liquid or gas; dull; non-conductors; highly reactive.
• Noble Gases: Colorless gases; extremely unreactive; low melting & boiling points; found in Earth's atmosphere.

Effective Nuclear Charge (Zeff)

• Effective Nuclear Charge (Zeff): The net positive charge experienced by an electron in an atom, taking into account the shielding effect of inner electrons. Formula: Zeff = Z - S (Z = number of protons, S = number of core electrons)

Electron Affinity

• Electron Affinity: The energy change associated with adding an electron to a gaseous atom. Multiple electron affinities exist for adding multiple electrons.
• Trend Down a Group: Decreases as the attraction between the nucleus and an added electron decreases and shielding increases.
• Trend Across a Period: Increases as the attraction between the nucleus and an added electron increases.

Electrolytes

• Electrolytes: Substances that form an electrically conductive solution when dissolved.
• Ionic Compounds as Electrolytes: Ionic compounds are electrolytes because their ions dissociate in water, allowing them to carry charge through the solution.

Covalent Compounds and Melting Points

• Low Melting Points of Covalent Compounds: Covalent compounds, despite having tightly bonded atoms, often have weak intermolecular forces, resulting in low melting points.

Electronegativity

• Electronegativity: The ability of an atom to attract bonding electrons.
• Trend Across a Period: Increases as atomic radius decreases and the attraction between valence electrons and the nucleus increases (additional protons).
• Trend Down a Group: Decreases as atomic radius increases and the attraction between valence electrons and the nucleus weakens.

Electronegativity Difference

• Electronegativity Difference (ΔEN): The difference in electronegativity between two atoms. This difference determines the type of bond:
  • Less than 0.5: Nonpolar covalent
  • 0.5 - 1.7: Polar covalent
  • Greater than 1.7: Ionic

Dipole

• Dipole: Separation of positive and negative charges within a polar molecule, creating positive and negative poles, indicated by δ+ and δ- respectively and the arrowhead points to the negative pole.

Intermolecular Forces

• Dipole-Dipole Forces: Attraction between the positive end of one polar molecule and the negative end of another. Strongest in polar molecules.
• London Dispersion Forces: Temporary, weak attractions between nonpolar molecules due to temporary shifts in electron cloud distribution. Present in all molecules, weaker than dipole-dipole forces unless the nonpolar molecule is very large.
• Hydrogen Bonding: Strong dipole-dipole force between molecules with a hydrogen atom bonded to a highly electronegative atom (N, O, or F).

Chemical Equations

• Nomenclature Rules and Patterns (Per-, -ate → -ite, Hypo-): Predicting formulas for and naming chemical compounds of different oxidation states.
• Common Chemical Reactions (Metal replacement): Including metal displacement from water, oxides, and acid reactions.
• Acid Rain: Caused by nitrogen and sulfur oxides reacting with water, creating nitric or sulfuric acid.

Stoichiometry

• Empirical Formula: Determining the simplest whole-number ratio of atoms in a compound.
• Molecular Formula: Finding the actual number of atoms in a molecule.
• Hydrate: Determining the ratio of water molecules to the ionic compound in a hydrate.
• Limiting and Excess Reactants: Determining the reactant that limits product formation.
• Percent Yield: Comparing actual yield of a product to its theoretical yield.

Reactions and Rates

• Rates of Dissolution: Affected by temperature, agitation, and particle size.
• Factors Affecting Solubility: Temperature, pressure, size of solute particle.

Solutions

• Surfactants: Compounds that reduce surface tension by breaking down hydrogen bonds at the surface, enabling polar solvents to dissolve nonpolar liquids.

Acids and Bases

• Naming Acids: Differentiating acids without and with oxygen.

Gases

• Kinetic Molecular Theory: Describing the behavior of gases at the molecular level, average kinetic energy dependent on temperature.
• Gas Laws: Boyle's law (P & V), Charles's law (V & T), Gay-Lussac's law (P & T), Ideal gas law (PV=nRT).
• Gas Stoichiometry: Applying gas laws to chemical reactions.

Periodic Table

• Periodic Table Arrangement: Mendeleev based it on atomic mass; Moseley suggested arranging it by atomic number.
• Periodic Trends: Elements in the same group have similar chemical properties due to similar numbers of valence electrons.

Solubility in Water

• NaCl: Highly soluble due to ionic nature and interaction with water molecules.
• Glucose: Soluble due to hydrogen bonding with water.
• Vegetable Oil: Insoluble due to nonpolar nature, inability to interact with polar water.
• Monatomic, Diatomic, Triatomic gases: Describing gases based on the number of atoms in a molecule.