General Chemistry: Atoms and Elements (CHEM 1303)

The Reality of Atoms

• The idea of the reality of atoms was promoted by Ludwig Boltzmann at the end of the 19th Century.

• This idea was confirmed by the discovery and explanation of Brownian motion.

  • Observation by Robert Brown (1773–1858):

    • A Scottish botanist who observed water-suspended particles from pollen grains through his microscope.

    • He noted that these particles were in continuous, random motion.

  • Theory by Albert Einstein (1905):

    • Developed a quantitative theory that explained what was by then called Brownian motion.

  • Test by Jean Perrin (1908):

    • A French physicist who tested Einstein's model and confirmed its validity through his measurements.

    • Perrin was awarded the Nobel Prize in Physics in $1926$ for his contributions.

    • The combined work of Einstein and Perrin conclusively removed any remaining doubt about the particulate nature of matter.

  • Note: Albert Einstein was awarded the $1921$ Nobel Prize in Physics for his services to Theoretical Physics, specifically for his discovery of the law of the photoelectric effect.

Building Blocks: Modern Atomic Theory

Modern atomic theory is founded on several fundamental laws:

• Law of Conservation of Mass:

  • States that in a chemical reaction, matter is neither created nor destroyed. The total mass of reactants equals the total mass of products.

• Law of Definite Proportions (also known as the Law of Constant Composition):

  • States that all samples of a given compound, regardless of their source or how they were prepared, have the same proportions of their constituent elements.

• Law of Multiple Proportions:

  • States that elements combine with other elements to make compounds in small, whole-number ratios.

  • When an atom of element A combines with either one, two, three, or more atoms of element B, the following molecular compounds are possible: AB1, AB2, AB3, etc.

  • Example: Hydrogen and Oxygen Forming Two Compounds ($H<em>nO</em>m$):

    • Compound A: Decomposition yields $0.125 ext{ g}$ hydrogen for every $1.00 ext{ g}$ oxygen.

    • Compound B: Decomposition yields $0.0625 ext{ g}$ hydrogen for every $1.00 ext{ g}$ oxygen.

    • To find the ratio consistent with the law of multiple proportions, compare the mass of hydrogen in both compounds for a fixed mass of oxygen:
      $rac{ ext{Mass H in Compound A}}{ ext{Mass H in Compound B}} = rac{0.125 ext{ g H}}{0.0625 ext{ g H}} = 2 = rac{2}{1}$

    • This means that for the same amount of oxygen, Compound A contains twice the mass of hydrogen as Compound B. Possible compounds reflecting this $2:1$ ratio for hydrogen (given a fixed amount of oxygen) are:

      • Compound A: Water ($H_2O$)

      • Compound B: Hydrogen Peroxide ($H<em>2O</em>2$)

    • If water ($H<em>2O$) has an $H:O$ mass ratio close to $0.125$, then hydrogen peroxide ($H</em>2O_2$) having a $H:O$ mass ratio close to $0.0625$ fits the $2:1$ ratio of hydrogen masses for constant oxygen mass.

Dalton's Atomic Theory (Early 19th Century)

John Dalton's atomic theory, developed around $1808$, explained the observed laws of conservation of mass, definite proportions, and multiple proportions with the following postulates:

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 the 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 within a chemical reaction. Instead, atoms change only the way they are bound together with other atoms.

Discovery of the Electron (Subatomic Particles)

J. J. Thomson and Cathode Ray Experiments ($1898$)

• J. J. Thomson conducted experiments using a partially evacuated glass tube called a cathode ray tube.

• Observations:

  • Particles emitted traveled in straight lines.

  • These particles were independent of the composition of the material from which they originated (the cathode).

  • They were found to carry a negative electrical charge.

• Measurement: Thomson measured the charge-to-mass ratio ($e/m$) of these cathode ray particles by deflecting them using electric and magnetic fields.

  • The charge-to-mass ratio of the electron was determined to be $-1.76 imes 10^8 ext{ C/g}$.

• Thomson was awarded the $1906$ Nobel Prize in Physics for his work.

Robert Millikan and the Oil Drop Experiment

• American physicist Robert Millikan performed his famous oil drop experiment (published in $1909$) to deduce the charge of a single electron.

• He observed that the charge on the oil drops was always an integral multiple of $-1.60 imes 10^{-19} ext{ C}$.

• This fundamental charge, $-1.60 imes 10^{-19} ext{ C}$, was identified as the elementary charge of a single electron.

Calculating the Electron's Mass

Mass of electron (me) = (Charge of electron (e)) / (Charge-to-mass ratio (e/m))
me = (-1.60 x 10^-19 C) / (-1.76 x 10^8 C/g) = 9.0909… x

Atomic Models: From Plum Pudding to Nuclear Atom

The Atom at the End of the 19th Century (Thomson's Plum Pudding Model)

• Following his discovery of the electron in $1899$, J. J. Thomson proposed the plum pudding model of the atom.

• Hypothesis: The atom was thought to be a sphere of uniformly distributed positive charge, with tiny, negatively charged electrons (the