Chapter 1: Atoms — Structure and Properties (Notes)

Matter from the Particulate Point of View

  • Matter is composed of particles. Examples include subatomic particles (neutrons, protons, electrons) that make up elemental atoms and molecules composed of elemental atoms.
  • The way particles come together determines the physical properties of matter.
  • Matter is defined as anything that has mass and occupies space (i.e., has volume).
  • Chemistry is the discipline that seeks to understand matter and its properties.

The Classification of Matter

  • Matter can be classified by state (physical form): solid, liquid, or gas, based on properties exhibited.
  • State changes from solid to liquid to gas with increasing temperature.

Classification of Matter by Components

  • Matter can be classified by composition into: elements, compounds, and mixtures.

Early Ideas about the Building Blocks of Matter

  • Leucippus (5th century BCE) and Democritus (460–370 BCE) proposed matter is made of small, indestructible particles (atoms).
  • Democritus asserted: “Nothing exists except atoms and empty space; everything else is opinion.” Atoms were imagined in different shapes/sizes and moving randomly through empty space.
  • Plato and Aristotle did not embrace atomic ideas: they held matter had no smallest parts and proposed fire, air, earth, and water as fundamental substances in varying proportions.

Early Ideas about the Building Blocks of Matter (continued)

  • John Dalton (1766–1844) offered convincing evidence for atomic ideas through his atomic theory of matter.
  • Modern Atomic Theory and Its Laws:
    • Law of conservation of mass
    • Law of definite proportions
    • Law of multiple proportions

John Dalton and the Atomic Theory

Dalton’s atomic theory explained the laws as follows:

  • 1. Each element is composed of tiny, indestructible particles called atoms.
  • 2. All atoms of a given element have the same mass and other properties that distinguish them from atoms of other elements.
  • 3. Atoms combine in simple, whole-number ratios to form compounds.
  • 4. Atoms of one element cannot change into atoms of another element. In a chemical reaction, atoms change only the way they are bound together with other atoms.

The Law of Conservation of Mass

  • In a chemical reaction, matter is neither created nor destroyed.
  • When a chemical reaction occurs, the total mass of the substances involved in the reaction does not change.
  • This law is consistent with the idea that matter is composed of small, indestructible particles.

The Law of Definite Proportions

  • In 1797, Joseph Proust observed that all samples of a given compound have the same proportions of their constituent elements.
  • Hence, the law of definite proportions is also called the law of constant composition.

An Example of the Law of Definite Proportions

  • Decomposition of 18.0 g of water (H₂O) yields 16.0 g of oxygen (O₂) and 2.0 g of hydrogen (H₂), giving an oxygen-to-hydrogen mass ratio of 8:1.

The Law of Multiple Proportions

  • In 1804, Dalton published the law of multiple proportions: when two elements (A and B) form two different compounds, the masses of element B that combine with 1 g of element A can be expressed as a ratio of small whole numbers.
  • When an atom of A combines with 1, 2, 3, or more atoms of B, possible molecular formulas are AB₁, AB₂, AB₃, etc.

The Law of Multiple Proportions (Example: CO and CO₂)

  • Carbon monoxide (CO) and carbon dioxide (CO₂) both contain carbon and oxygen.
  • In CO₂, the mass ratio of oxygen to carbon is 2.67:1, meaning 2.67 g O reacts with 1 g C.
  • In CO, the mass ratio of oxygen to carbon is 1.33:1, or 1.33 g O per 1 g C.
  • The ratio of these two masses is a small whole number, consistent with the law of multiple proportions.

The Discovery of the Electron: J. J. Thomson's Cathode Ray Experiment

  • Thomson used a partially evacuated glass tube with a cathode and an anode, applying high voltage to create cathode rays.
  • Cathode rays were deflected by electric and magnetic fields, indicating they carried charge.
  • Thomson’s experiments demonstrated the existence of the electron, a negatively charged, low-mass particle present within atoms.
  • The apparatus and observations established the charge-to-mass ratio of the electron.

The Discovery of the Electron (Summary)

  • J. J. Thomson discovered the electron as a component of atoms.
  • The electron is negatively charged and has a very small mass relative to protons and neutrons.

Millikan’s Oil Drop Experiment: Determining the Charge of an Electron

  • Millikan measured the strength of the electric field needed to halt the free fall of oil drops and inferred their charges.
  • From the radii and density, the mass of each drop was determined.
  • The charge on any oil drop was always a whole-number multiple of –1.60 × 10^{-19} C, establishing the fundamental charge of a single electron.

Millikan’s Oil Drop Experiment: The Charge-to-Mass Ratio for an Electron

  • Using Millikan’s charge data and Thomson’s electron mass-to-charge ratio, the electron mass can be deduced.
  • Resulting electron mass is on the order of 9.109 × 10^{-31} kg (not shown here, but implied by the combination of charge and m/e ratio).

The Structure of the Atom: The Early Models

  • J. J. Thomson’s Plum-Pudding Model:
    • Positively charged sphere with embedded negatively charged electrons.
    • Model widely accepted before nucleus discovery.

Rutherford’s Model and the Gold Foil Experiment

  • In 1909, Rutherford performed the Gold Foil Experiment to test Thomson’s model.
  • He directed positively charged particles at a thin sheet of gold foil.

Rutherford’s Gold Foil Experiment Results

  • Most particles passed through the foil, but some were deflected, and a few (about 1 in 20,000) were bounced back.
  • Conclusion: Matter is not uniform; it contains large regions of empty space, with a very small, dense nucleus at the center.

Building on the Rutherford Atomic Model: The Nuclear Atom Model

  • Three basic parts:
    1) Most of the atom’s mass and all of its positive charge reside in the nucleus.
    2) The atom is mostly empty space; electrons are dispersed in the empty space surrounding the nucleus.
    3) For electrical neutrality, the number of negatively charged electrons outside the nucleus equals the number of positively charged particles (protons) inside the nucleus.

The Neutral Particles: Neutrons

  • Rutherford’s model could not account for the neutron’s mass in the nucleus.
  • 1932: Rutherford and James Chadwick demonstrated the presence of neutrons.
  • Neutron properties:
    • Mass is similar to that of a proton.
    • No electrical charge.

Subatomic Particles (Table 1.1)

  • Proton: mass ≈ 1.67262 × 10^{-27} kg; mass in amu ≈ 1.00727; charge +1; charge in C ≈ +1.60218 × 10^{-19}
  • Neutron: mass ≈ 1.67493 × 10^{-27} kg; mass in amu ≈ 1.00866; charge 0; charge in C ≈ 0
  • Electron: mass ≈ 0.00091 × 10^{-27} kg; mass in amu ≈ 0.00055; charge −1; charge in C ≈ −1.60218 × 10^{-19}

Elements: Defined by Their Numbers of Protons

  • The number of protons in the nucleus defines the element.
  • Atomic number Z represents the number of protons.

Elements & the Periodic Table

  • Elements are arranged in the periodic table by their atomic number Z.
  • Elements in the same column (group) have very similar physical and chemical properties.
  • Elements are represented in the periodic table by their symbol and atomic number.

Isotopes: Elements with Varied Number of Neutrons

  • All atoms of a given element have the same number of protons (Z) but may have different numbers of neutrons.
  • Example: Neon atoms all have 10 protons, but may have 10, 11, or 12 neutrons.
  • All three types of neon isotopes exist and have slightly different masses.

Isotopes: Representation

  • The mass number notation: A = number of protons + number of neutrons, i.e., A = p^+ + n.
  • X denotes the chemical symbol; A is the mass number; Z is the atomic number.

Isotopes: Representation (Alternative Notation)

  • An isotope is also represented as the chemical symbol (or name) followed by a dash and the mass number, e.g., ^A_ZX or X-A.

Isotopes: Varied Number of Neutrons

  • The relative abundance of isotopes in a naturally occurring sample is roughly constant.
  • These percentages are called the natural abundance of the isotopes.

Ions: Charged Atoms Losing and Gaining Electrons

  • In a neutral atom, the number of electrons equals the number of protons (Z).
  • Atoms can lose or gain electrons to form ions.
  • Cations are positively charged ions (e.g., Na⁺).
  • Anions are negatively charged ions (e.g., F⁻).

Atomic Mass: The Average Mass of an Element’s Atoms

  • Atomic mass (also called atomic weight or standard atomic weight) is shown beneath the element symbol in the periodic table.
  • It represents the weighted average mass of the element’s isotopes, based on natural abundances.

Atomic Mass: Calculation

  • Atomic mass = Σ (fraction of isotope n) × (mass of isotope n)
  • Example notation:
    • ext{Atomic mass} = ig( ext{fraction of isotope 1}ig) imes ig( ext{mass of isotope 1}ig) + ig( ext{fraction of isotope 2}ig) imes ig( ext{mass of isotope 2}ig) + \ ext{…}

Atomic Mass: Problem Example

  • Naturally occurring chlorine consists of 75.77% chlorine-35 (mass 34.97 amu) and 24.23% chlorine-37 (mass 36.97 amu).
  • Calculate chlorine’s atomic mass.
  • Formula to apply: ext{Atomic mass} = ( ext{fraction 35}) imes (m{35}) + ( ext{fraction 37}) imes (m{37})

Mass Spectrometry: Measuring the Mass of Atoms and Molecules

  • Mass spectrometry measures the masses of atoms and the percent abundances of isotopes.
  • It separates particles according to their mass.

What Is a Mole?

  • A mole (mol) is a counting unit used for very large numbers.
  • 1 dozen = 12, 1 gross = 144, etc.
  • A mole contains 6.02214 × 10^{23} pieces, known as Avogadro’s number.
  • Examples:
    • 1 mol of marbles corresponds to 6.02214 × 10^{23} marbles.
    • 1 mol of sand grains corresponds to 6.02214 × 10^{23} sand grains.

The Mole (Foundational Value)

  • The mole is defined by the number of atoms in exactly 12 grams of carbon-12 (C-12).
  • Therefore, 12 g of C contains 1 mole of C atoms: 6.022 × 10^{23} atoms of C.

Mole Conversions: Atoms to Moles or Moles to Atoms

  • The fundamental conversion: 1 mol atoms = 6.022 × 10^{23} atoms.
  • The conversion factors can be written in the form:
    • 1
    • 1
    • 6.022 × 10^{23}
  • (These references allow conversion between numbers of atoms and moles.)

Converting Between Mass and Moles

  • One mole of anything contains 6.022 × 10^{23} items but not necessarily the same mass.
  • Examples:
    • 26.98 g Al = 1 mol Al = 6.022 × 10^{23} atoms Al
    • 12.01 g C = 1 mol C = 6.022 × 10^{23} atoms C
    • 4.003 g He = 1 mol He = 6.022 × 10^{23} atoms He

Mass to Moles to Number of Particles: The Conceptual Plan

  • For an element:
    • Mass of element (grams) → Moles of element → Number of atoms
    • Multiply by Avogadro’s number to get number of atoms.
  • For a molecule (compound):
    • Mass of molecule (grams) → Number of molecules → Moles of molecules → Multiply by Avogadro’s number to get number of molecules.

Number of Particles to Moles to Mass: The Conceptual Plan

  • For an element:
    • Mass of element (grams) ← Moles of element ← Number of atoms × (molar mass)\
  • For a molecule (compound):
    • Mass of molecule (grams) ← Number of molecules ← Moles of molecules × (molar mass)
  • The conceptual plan emphasizes moving between mass, moles, and particle counts depending on the direction of the problem.