Chemistry: Atoms, Molecules, and Ions

Chapter 2: Atoms, Molecules, and Ions

2.1 Atomic Theory of Matter

  • Early Philosophical Ideas:

    • Some Greek philosophers, like Democritus, believed in a smallest, indivisible particle called "atomos" (meaning "uncuttable") that constituted all of nature.

  • John Dalton's Organized Atomic Theory (early 1800s):

    • Developed based on experiments conducted in the eighteenth and nineteenth centuries.
    • Foundation for modern understanding of atomic structure.
    • Based on three fundamental laws:
      • The Law of Constant Composition
      • The Law of Conservation of Mass
      • The Law of Multiple Proportions

Law of Constant Composition

  • Definition: Compounds have a definite composition, meaning the relative number of atoms of each element in a given compound is always the same, regardless of the sample size or source.
  • Discoverer: Joseph Proust.
  • Significance: This law was one of the foundations for Dalton's atomic theory (specifically Postulate 4).

Law of Conservation of Mass

  • Definition: The total mass of substances present at the end of a chemical process is precisely the same as the total mass of substances present before the process took place.
  • Discoverer: Antoine Lavoisier.
  • Significance: This law was one of the foundations for Dalton's atomic theory (specifically Postulate 3). Further details are explained in Chapter 3.

Postulates of Dalton's Atomic Theory

  1. Composition of Elements: Each element is composed of extremely small, fundamental particles called atoms.
    • Visual representation: An atom of oxygen appears distinct from an atom of nitrogen.
  2. Identity of Atoms: All atoms of a given element are identical to one another in mass and other properties.
    • However, atoms of one element are distinctly different from atoms of all other elements.
    • Visual representation: All oxygen atoms are alike, but different from nitrogen atoms.
  3. Chemical Reactions: Atoms of one element cannot be changed into atoms of a different element through chemical reactions.
    • Atoms are neither created nor destroyed during chemical reactions.
    • This postulate supports the Law of Conservation of Mass.
  4. Formation of Compounds: Compounds are formed when atoms of more than one element combine.
    • A given compound will always have the same relative number and kind of atoms.
    • This postulate supports the Law of Constant Composition.
    • Visual representation: Oxygen atoms combining with nitrogen atoms to form a compound (e.g., NO).

Discovery of Subatomic Particles

Cathode Rays and the Electron

  • Experiment: Streams of negatively charged particles were observed to emanate from cathode tubes, causing fluorescence on a screen.
  • Cathode Ray Properties:
    • They move from the negative cathode to the positive anode.
    • They are deflected by a magnet, indicating they possess charge.
  • Discoverer: J. J. Thomson (credited in 1897) discovered these particles, which he called electrons.
  • Thomson's Measurements: His experiments allowed for the calculation of the charge-to-mass ratio (e/m) of the electron by balancing electric and magnetic field strengths to undeflect the electron beam.

Millikan Oil-Drop Experiment

  • Objective: To determine the fundamental charge of an electron.
  • Method:
    • Oil drops were sprayed into a chamber.
    • X-ray irradiation caused the drops to pick up electrons and become negatively charged.
    • The force of gravity pulling the drops downward was balanced by an electric field pushing the negatively charged drops upward.
  • Discoverer: Robert Millikan (in 1909) successfully determined the charge on the electron.
  • Significance: With the electron's charge known and Thomson's charge-to-mass ratio established, the mass of an electron could then be calculated.

Radioactivity

  • Discoverer: Ernest Rutherford identified three types of radiation emitted by radioactive substances:
    • \alpha particles: Positively charged particles that bend toward the negatively charged plate.
    • \beta particles: Negatively charged particles (like electrons) that bend toward the positively charged plate.
    • \gamma rays: Uncharged (neutral) radiation that is unaffected by electrically charged plates.

The Atom, Circa 1900: Thomson's "Plum Pudding" Model

  • Prevailing Theory: J. J. Thomson proposed that the atom consisted of a diffuse positive sphere of matter, with tiny negative electrons embedded within it, much like plums in a pudding.

Discovery of the Nucleus: Rutherford's Gold Foil Experiment

  • Experiment: Ernest Rutherford directed a beam of positively charged \alpha particles at an extremely thin sheet of gold foil and observed their scattering pattern using a circular fluorescent screen.
  • Observations:
    • Most \alpha particles passed straight through the gold foil with little or no deflection.
    • A tiny fraction of the \alpha particles were scattered at large angles, and some were even deflected backward.
  • Interpretation and the Nuclear Atom Model:
    • Thomson's "plum pudding" model could not explain the large-angle deflections.
    • Rutherford concluded that the atom's positive charge and most of its mass are concentrated in an extremely small, dense region at its center, which he called the nucleus.
    • The rest of the atom is mostly empty space.
    • The large-angle deflections occurred when \alpha particles closely approached or directly hit this highly charged nucleus.

Atomic Structure

  • Atomic Number (Z):

    • Defined as the number of protons in the nucleus of an atom.
    • This number uniquely identifies an element.
    • For a neutral atom, the number of protons is equal to the number of electrons (since atoms have no overall net charge).

  • Atomic Mass Unit (amu):

    • Atoms have extremely small masses, necessitating a specialized mass scale.
    • Historically, hydrogen was arbitrarily assigned a mass of 1, and other atomic masses were assigned relative to it (e.g., oxygen = 16).
    • Today, an atomic mass unit (amu) is the base unit for atomic-level masses.
    • The masses of atoms are compared to the carbon-$12$ isotope (6 protons and 6 neutrons), which is defined as having a mass of exactly 12 amu.
    • Modern techniques allow for highly accurate and precise determination of atomic masses.

  • Atomic Weight:

    • In practical applications, large quantities of atoms and molecules are used, so an average mass is necessary for calculations.
    • Atomic weight is the average mass of all naturally occurring isotopes of an element, weighted by their relative abundances.
    • Measured using a mass spectrometer, which separates ions based on mass-to-charge ratio (m/z).
    • Mass spectrometry also allows for the determination of isotope abundances (e.g., the spectrum of chlorine shows distinct peaks for ^{35}\text{Cl} and ^{37}\text{Cl}).

2.5 The Periodic Table

  • Definition: A systematic organization of all known elements.

  • Organization:

    • Elements are arranged in order of increasing atomic number.
    • Periods: Horizontal rows of the periodic table.
      • Elements within a period exhibit a repeating pattern of properties and reactivity, known as periodicity.
      • Example: Neon (Ne) is a nonreactive gas, Sodium (Na) is a soft, reactive metal, Chlorine (Cl) is listed as a nonreactive gas (note: chemically, Cl is a reactive nonmetal), Argon (Ar) is a nonreactive gas, Potassium (K) is a soft, reactive metal, Calcium (Ca) is a soft, reactive metal.
    • Groups: Vertical columns of the periodic table.
      • Elements within the same group tend to have similar chemical properties.

  • Reading the Periodic Table:

    • Atomic Number: Listed above the element symbol (e.g., 19 for K).
    • Atomic Symbol: The abbreviation for the element (e.g., K for Potassium).
    • Atomic Weight: Listed below the element symbol (e.g., 39.0983 for K).

  • Classification of Elements:

    • A steplike line on the periodic table divides metals from nonmetals.
    • Metals: Located on the left side of the periodic table.
      • Properties: Shiny luster, good conductors of heat and electricity, generally solid at room temperature (except mercury, Hg, which is liquid).
      • Examples: Iron (Fe), Copper (Cu), Aluminum (Al), Silver (Ag), Lead (Pb), Gold (Au).
    • Nonmetals: Located on the right side of the periodic table (hydrogen, H, is also a nonmetal).
      • Properties: Can exist as solids (e.g., carbon, sulfur), liquids (e.g., bromine), or gases (e.g., neon) at room temperature.
      • Examples: Carbon (C), Sulfur (S), Bromine (Br), Phosphorus (P).
    • Metalloids: Elements situated along the steplike line, exhibiting properties intermediate between metals and nonmetals.

2.6 Chemical Formulas

  • Empirical Formulas:
    • Represent the lowest whole-number ratio of atoms of each element in a compound.
  • Molecular Formulas:
    • Represent the exact number of atoms of each element in a compound.
    • If the molecular formula is known, the empirical formula can be determined (by simplifying ratios).
    • The reverse is not true without additional information (e.g., molecular mass).

Picturing Molecules (Representing 3D Structure)

  • Structural Formulas (2D):
    • Show the order in which atoms are connected or attached to one another.
    • Do NOT depict the three-dimensional (3D) shape of molecules.
  • Three-Dimensional Models:
    • Perspective Drawings: Use wedges to indicate bonds pointing out of the page (solid wedge) or behind the page (dashed wedge), and solid lines for bonds in the plane of the page.
    • Ball-and-Stick Models: Represent atoms as spheres and bonds as sticks, showing connectivity and relative angles.
    • Space-Filling Models: Show the relative sizes of atoms and their overall spatial arrangement by depicting overlapping spheres.
    • Example: Methane (CH4) can be shown by molecular formula, structural formula (H-C-H with H-C above/below), perspective drawing, ball-and-stick model, and space-filling model.

2.7 Ions and Ionic Compounds

  • Formation of Ions: An atom or group of atoms becomes an ion when it gains or loses one or more electrons.

  • Cations:

    • Formed when an atom (or group of atoms) loses at least one electron.
    • Possess a net positive charge.
    • Monatomic cations are typically formed by metals (e.g., Na+, Mg2+, K+, Ca2+).

  • Anions:

    • Formed when an atom (or group of atoms) gains at least one electron.
    • Possess a net negative charge.
    • Monatomic anions are typically formed by nonmetals (except noble gases) (e.g., H-, N3-, O2-, F-, S2-, Cl-, Br-, I-, Te2-).

  • Ionic Compounds:

    • Generally formed between metals (which tend to form cations) and nonmetals (which tend to form anions).
    • Involve the transfer of electrons from the metal to the nonmetal.
    • The resulting oppositely charged ions are held together by strong electrostatic attraction.
    • Only empirical formulas are written for ionic compounds, representing the simplest whole-number ratio of ions.
    • Example: Sodium (Na) atom loses an electron to become Na+ ion. Chlorine (Cl) atom gains an electron to become Cl- ion. These oppositely charged ions attract to form NaCl.

2.8 Naming Inorganic Compounds (Chemical Nomenclature)

  • Definition: Chemical nomenclature is the systematic process of naming chemical compounds.
  • Key Requirement: Names of common ions (both cations and anions) must be memorized.
  • Naming Rules Apply to:
    • Ionic compounds
    • Acids
    • Binary molecular compounds

Writing Inorganic Formulas (for Ionic Compounds)

  • Principle: Compounds are electrically neutral; the total positive charge from cations must balance the total negative charge from anions.
  • Method (Crisscross Rule):
    1. The magnitude of the charge on the cation becomes the subscript for the anion.
    2. The magnitude of the charge on the anion becomes the subscript for the cation.
    3. If the resulting subscripts are not in the lowest whole-number ratio, divide them by their greatest common factor to simplify the formula.
    • Example: Magnesium ion (Mg^{2+}) and Nitride ion (N^{3-}) combine to form Mg3N2.

Patterns in Oxyanion Nomenclature

  • Definition: Oxyanions are polyatomic anions containing one or more oxygen atoms bonded to a central atom.

  • General Rules:

    • Period 2 Central Atoms: Central atoms in the second row of the periodic table (e.g., C, N) typically bond to at most three oxygen atoms (e.g., Carbonate, CO3^{2-} ; Nitrate, NO3^{-}; note: slide shows CO2^{2-}, likely typo).
    • Period 3 Central Atoms: Central atoms in the third row (e.g., P, S, Cl) can bond to up to four oxygen atoms (e.g., Phosphate, PO4^{3-} ; Sulfate, SO4^{2-} ; Perchlorate, ClO4^{-}).
    • Charge Trend: Ion charges generally decrease as you move from left to right across a period for elements forming oxyanions.

  • Halogen Oxyanion Series (Example with Chlorine):

    • This system uses prefixes and suffixes to indicate the number of oxygen atoms relative to a common or representative oxyanion (usually the one ending in "-ate").
    • Common Anion (no oxygen): Ends in "-ide" (e.g., Chloride, Cl^{-}).
    • Per- -ate: Has one more oxygen atom than the "-ate" ion (e.g., Perchlorate, ClO4^{-}).
    • -ate: The common or representative oxyanion (e.g., Chlorate, ClO3^{-}).
    • -ite: Has one fewer oxygen atom than the "-ate" ion (e.g., Chlorite, ClO2^{-}).
    • Hypo- -ite: Has two fewer oxygen atoms than the "-ate" ion (e.g., Hypochlorite, ClO^{-}).

Nomenclature of Binary Molecular Compounds

  • Rules:
    1. The first element in the formula is named first.
    2. The ending of the second element's name is changed to "-ide".
    3. Prefixes are used to indicate the number of atoms of each element present:
      • mono- (1)
      • di- (2)
      • tri- (3)
      • tetra- (4)
      • penta- (5)
      • hexa- (6)
      • hepta- (7)
      • octa- (8)
      • nona- (9)
      • deca- (10)
    4. The prefix "mono-" is generally omitted for the first element, but always used for the second element (e.g., carbon monoxide, not monocarbon monoxide).
    5. When a prefix ends with "a" or "o" and the name of the element begins with a vowel, the successive vowels are often elided into one (e.g., dinitrogen pentoxide, not dinitrogen pentaoxide).
  • Examples:
    • CO_2: carbon dioxide
    • CCl_4: carbon tetrachloride
    • CO: carbon monoxide

2.9 Some Simple Organic Compounds

  • Organic Chemistry: The branch of chemistry dedicated to the study of carbon compounds.
  • Nomenclature: Organic chemistry has its own distinct system of naming compounds.

Alkanes (Simplest Hydrocarbons)

  • Definition: Hydrocarbons are compounds composed solely of carbon and hydrogen atoms.
  • Alkanes: The simplest class of hydrocarbons.
  • Naming Convention:
    • The first part of the name indicates the number of carbon atoms:
      • meth- = 1 carbon
      • eth- = 2 carbons
      • prop- = 3 carbons
      • (etc.)
    • This prefix is followed by the suffix "-ane".
  • Examples:
    • Methane (CH4)
    • Ethane (C2H6)
    • Propane (C3H8)

Alcohols

  • Definition: Alcohols are compounds derived from alkanes where one or more hydrogen atoms have been replaced by a hydroxyl functional group (-OH).
  • Naming Convention:
    • Similar to alkanes, the name is derived from the parent alkane.
    • The ending of the name denotes the type of compound and for alcohols, it is "-ol".
  • Examples:
    • Methanol (CH3OH)
    • Ethanol (C2H5OH)
    • 1-Propanol (C3H7OH)