Chemistry Chapter 6 Quiz

Chapter 2: Atoms, Molecules, and Ions

• This chapter covers the fundamentals of atoms, molecules, and ions, including their history and modern understanding.

Section 2.1: The Early History of Chemistry

• Greeks were the first to attempt to explain chemical changes.
• Alchemy dominated for 2000 years, with several elements being discovered and mineral acids prepared.
• Robert Boyle was a pivotal figure, performing quantitative experiments and developing the first experimental definition for an element.

Section 2.2: Fundamental Chemical Laws

• Law of Conservation of Mass (Lavoisier): Mass is neither created nor destroyed in a chemical reaction.
• Law of Definite Proportion (Proust): A given compound always contains exactly the same proportion of elements by mass.
• Law of Multiple Proportions (Dalton): When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can be reduced to small whole numbers.

Section 2.3: Dalton's Atomic Theory

• Dalton's Atomic theory (1808):
  • Each element is made up of tiny particles called atoms
  • The atoms of a given element are identical, but the atoms of different elements are different in some fundamental way.
  • Chemical compounds are formed when atoms of different elements combine with each other. A given compound always has the same relative numbers and types of atoms.
  • Chemical reactions involve reorganization of the atoms—changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

Section 2.4: Early Experiments to Characterize the Atom

• J. J. Thomson (1898-1903):
  • Postulated the existence of negatively charged particles (electrons) using cathode-ray tubes.
  • Determined the charge-to-mass ratio of an electron.
  • The atom must also contain positive particles to balance the negative charge of electrons.
• Ernest Rutherford (1911):
  • Explained the nuclear atom.
  • The atom has a dense center of positive charge called the nucleus.
  • Electrons travel around the nucleus at a large distance relative to the nucleus.

Section 2.5: The Modern View of Atomic Structure

• The atom contains:
  • Electrons: negatively charged, found outside the nucleus.
  • Protons: positively charged, found in the nucleus, with a magnitude of charge equal to the electron's negative charge.
  • Neutrons: no charge, found in the nucleus, with a mass virtually the same as a proton.
• The nucleus is:
  • Small compared to the overall size of the atom.
  • Extremely dense, accounting for almost all of the atom's mass.
• Isotopes: Atoms with the same number of protons but different numbers of neutrons. They show almost identical chemical properties because the chemistry of an atom is due to its electrons. Most elements contain mixtures of isotopes.
• Isotopes are identified by:
  • Atomic Number (Z): the number of protons.
  • Mass Number (A): the number of protons plus the number of neutrons.

Section 2.6: Molecules and Ions

• Chemical Bonds:
  • Covalent Bonds: Bonds form between atoms by sharing electrons. The resulting collection of atoms is called a molecule.
  • Ionic Bonds: Bonds form due to the attraction between oppositely charged ions.
    • Ion: an atom or group of atoms with a net positive or negative charge.
    • Cation: a positive ion; lost electron(s).
    • Anion: a negative ion; gained electron(s).

Section 2.7: An Introduction to the Periodic Table

• Metals vs. Nonmetals: A key classification of elements.
• Groups/Families: Vertical columns of elements with similar chemical properties.
• Periods: Horizontal rows of elements.

Section 2.8: Naming Simple Compounds

• Binary Compounds: Compounds composed of two elements (ionic or covalent).
• Binary Ionic Compounds (Type I): The cation is named first, followed by the anion. Monatomic cations take their name from the parent element. Monatomic anions have the element root with -ide added. Examples: KCl (Potassium chloride), MgBr₂ (Magnesium bromide), CaO (Calcium oxide).
• Binary Ionic Compounds (Type II): Metals that can form more than one positive ion. The charge of the metal cation must be specified using a Roman numeral. Examples: CuBr (Copper(I) bromide), FeS (Iron(II) sulfide), PbO₂ (Lead(IV) oxide).
• Polyatomic Ions: Ions containing multiple atoms. Their names need to be memorized. Examples: NaOH (Sodium hydroxide), Mg(NO₃)₂ (Magnesium nitrate), (NH₄)₂SO₄ (Ammonium sulfate).
• Binary Covalent Compounds (Type III): Formed between two nonmetals. The first element is named using its full element name, and the second element is named as an anion, using prefixes to indicate the number of atoms present. Examples: CO₂ (Carbon dioxide), SF₆ (Sulfur hexafluoride), N₂O₄ (Dinitrogen tetroxide).
• Acids: Molecules with one or more hydrogen ions (H⁺) attached to an anion. If the anion does not contain oxygen, the acid is named with the prefix "hydro-" and the suffix "-ic." If the anion contains oxygen:
  • The suffix "-ic" is added if the anion name ends in "-ate."
  • The suffix "-ous" is added if the anion name ends in "-ite." (e.g., HNO₃ (Nitric acid), H₂SO₄ (Sulfuric acid), HNO₂ (Nitrous acid), H₂SO₃ (Sulfurous acid).