Atoms and Elements

Chapter 2: Atoms and Elements

Atomic Theory of Matter

• The concept of atoms dates back to Democritus, who used "atomos" (Greek for indivisible) to describe the smallest unit of matter.
• Plato and Aristotle believed substances could be broken down into smaller parts without limit.
• The atomic view diminished until its re-emergence in the early 19th century, championed by John Dalton.

Dalton's Postulates

• Postulate 1: Each element consists of extremely small particles called atoms.
• Postulate 2: All atoms of a given element are identical in mass and other properties, but atoms of one element differ from atoms of all other elements.
• Postulate 3: Atoms of an element are not changed into atoms of a different element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
• Postulate 4: Compounds form when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.

Laws Arising from Atomic Theory

• Law of Constant Composition (Law of Definite Proportions):
  • Proposed by Joseph Proust (1754–1826).
  • The elemental composition of a pure substance is always the same.
• Law of Conservation of Mass:
  • The total mass of substances present at the end of a chemical process is the same as the mass of substances present before the process.
• Law of Multiple Proportions:
  • Two elements, A & B, can combine in different ratios as long as those ratios are whole numbers.
  • Example: Water (H2O) vs. Hydrogen Peroxide (H2O_2)

Discovery of the Electron

• Streams of negatively charged particles (electrons) were found to emanate from cathode tubes.
• J.J. Thomson is credited with the discovery of the electron (1897).
• Thomson measured the charge/mass ratio of the electron:
  • 1.76 \,x\, 10^8 \,\frac{coulombs}{g}

Millikan Oil Drop Experiment

• Robert Millikan (University of Chicago) determined the charge on the electron in 1909.
• Knowing the charge/mass ratio allowed determining the mass of the electron.

Radioactivity

• The spontaneous emission of radiation by an atom.
• First observed by Henri Becquerel.
• Also studied by Marie and Pierre Curie.
• Ernest Rutherford discovered three types of radiation:
  • \alpha particles
  • \beta particles
  • \gamma rays

Models of the Atom

• Plum Pudding Model (circa 1900):
  • Proposed by Thompson.
  • A positive sphere of matter with negative electrons embedded in it.
• Discovery of the Nucleus:
  • Ernest Rutherford shot \alpha particles at a thin sheet of gold foil.
  • Observed the scattering pattern of the particles.
• The Nuclear Atom:
  • Rutherford proposed a very small, dense nucleus containing most of the atom's mass.
  • Electrons reside around the outside of the nucleus.
  • Most of the atom's volume is empty space.

Subatomic Particles

• Protons:
  • Discovered by Rutherford in 1919.
• Neutrons:
  • Discovered by James Chadwick in 1932.
• Protons and electrons have a charge; neutrons are neutral.
• Protons and neutrons have approximately the same mass.
• The mass of an electron is considered negligible.

Symbols of Elements

• Elements are symbolized by one or two letters.
• Atomic Number (Z):
  • Number of protons in an atom's nucleus.
  • All atoms of the same element have the same number of protons.
• Mass Number:
  • Number of protons plus neutrons in the nucleus.
• Representation:
  • ^{Mass\, Number}{Atomic\, Number}Symbol (e.g., ^{12}6C)

Atomic Mass

• The mass of an atom in atomic mass units (amu) is the total number of protons and neutrons in the atom.
• Isotopes:
  • Atoms of the same element with different masses.
  • Isotopes have different numbers of neutrons.
  • Examples: ^{11}6C, ^{12}6C, ^{13}6C, ^{14}6C
• Mass Spectrometer:
  • Used to measure atomic and molecular masses with great accuracy.
  • Separates ions based on mass differences.
• Average Atomic Mass:
  • Calculated from the isotopes of an element weighted by their relative abundances.
  • Used in calculations due to dealing with large amounts of atoms/molecules.
  • Example: Cl & Br

Periodic Table

• Rows are called periods.
• Columns are called groups.
• Elements in the same group have similar chemical properties.
• Group Names:
  • 1A: Alkali Metals (Li, Na, K, Rb, Cs, Fr)
  • 2A: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)
  • 6A: Chalcogens (O, S, Se, Te, Po)
  • 7A: Halogens (F, Cl, Br, I, At)
  • 8A: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)
• Metals:
  • Located on the left side of the periodic table, generally.
• Nonmetals:
  • Located on the right side of the periodic table, generally.
• Metalloids:
  • Border the stair-step line (with exceptions like Al and Po).

Chemical Formulas

• The subscript to the right of an element's symbol indicates the number of atoms of that element in one molecule of the compound.

Molecular Compounds

• Composed of molecules and typically contain only nonmetals.
• Examples: Water (H2O), Carbon Dioxide (CO2), Carbon Monoxide (CO), Methane (CH4), Oxygen (O2), Hydrogen Peroxide (H2O2)

Diatomic Molecules

• Seven elements occur naturally as molecules containing two atoms:
  • H2, N2, O2, F2, Cl2, Br2, I_2

Types of Formulas

• Empirical Formulas:
  • Give the lowest whole-number ratio of atoms of each element in a compound.
• Molecular Formulas:
  • Give the exact number of atoms of each element in a compound.
  • Example: Hydrogen Peroxide (HO vs. H2O2)
• Structural Formulas:
  • Show the order in which atoms are bonded.
• Perspective Drawings:
  • Show the three-dimensional arrangement of atoms in a compound.

Ionic Charges and Formulas

• Ions:
  • Atoms or groups of atoms with a net electrical charge.
• Cations:
  • Positively charged ions.
  • Formed by losing electrons.
• Anions:
  • Negatively charged ions.
  • Formed by gaining electrons.
• Predicting Ionic Charges:
  • Group 1A (1): Lose 1 electron, 1+ charge (e.g., Li^+, Na^+, K^+)
  • Group 2A (2): Lose 2 electrons, 2+ charge (e.g., Mg^{2+}, Ca^{2+})
  • Group 3A (13): Lose 3 electrons, 3+ charge (e.g., Al^{3+})
  • Group 5A (15): Gain 3 electrons, 3- charge (e.g., N^{3-}, P^{3-})
  • Group 6A (16): Gain 2 electrons, 2- charge (e.g., O^{2-}, S^{2-})
  • Group 7A (17): Gain 1 electron, 1- charge (e.g., F^{-}$, Cl^{-}$, Br^{-}$, I^{-})

Writing Formulas for Ionic Compounds

• Compounds are electrically neutral.
• The charge on the cation becomes the subscript on the anion.
• The charge on the anion becomes the subscript on the cation.
• Reduce subscripts to the lowest whole-number ratio.

Naming Ionic Compounds

• Cation:
  • Positive charge, usually a metal.
  • Named as the element (e.g., Sodium, Aluminum).
• Anion:
  • Negative charge, usually a nonmetal.
  • Named as the element with the suffix -ide (e.g., Chloride, Oxide).
• Examples:
  • NaCl: Sodium Chloride
  • K_2S: Potassium Sulfide
  • MgO: Magnesium Oxide
  • CaI_2: Calcium Iodide
  • Al2O3: Aluminum Oxide

Naming Compounds with Polyatomic Ions

• The positive ion is named first, followed by the name of the polyatomic ion.
• Examples:
  • NaNO_3: Sodium Nitrate
  • K2SO4: Potassium Sulfate
  • Fe(HCO3)3: Iron(III) Bicarbonate (or Iron(III) Hydrogen Carbonate)
  • (NH4)3PO_3: Ammonium Phosphite

Transition Metals and Variable Charges

• Most transition metals and Group 4A (14) metals form 2 or more positive ions.
• Zn^{2+}, Ag^+, and Cd^{2+} form only one ion.
• Use Roman numerals to indicate the ionic charge for metals with multiple possible charges.
• Examples:
  • Lead (Pb): Pb^{2+} Lead(II), Pb^{4+} Lead(IV)

Older Naming System for Variable Charge Metals

• Metal names are derived from their Latin root.
  • Copper becomes cuprum
  • Iron becomes ferrum
• Higher charge ends in -ic.
• Lower charge ends in -ous.
• Examples:
  • Fe^{2+}: Iron(II) or Ferrous
  • Fe^{3+}: Iron(III) or Ferric
  • Cu^+: Copper(I) or Cuprous
  • Cu^{2+}: Copper(II) or Cupric

Naming Binary Molecular Compounds

• The less electronegative atom is typically listed first.
• Prefixes indicate the number of atoms (mono- is not used for the first element).
• The more electronegative element ends in -ide.
• If the prefix ends in a or o and the element name starts with a vowel, the vowels may be elided.
• Examples:
  • CO_2: Carbon Dioxide
  • CCl_4: Carbon Tetrachloride
  • N2O5: Dinitrogen Pentoxide

Acid Nomenclature

• If the anion ends in -ide, the acid name starts with hydro- and ends in -ic acid.
  • HCl: Hydrochloric Acid
  • HBr: Hydrobromic Acid
  • HI: Hydroiodic Acid
• If the anion ends in -ate, change the ending to -ic acid.
  • HClO_3: Chloric Acid
  • HClO_4: Perchloric Acid
• If the anion ends in -ite, change the ending to -ous acid.
  • HClO: Hypochlorous Acid
  • HClO_2: Chlorous Acid