Chemistry: Atomic Theory and Early Atomic Structure Notes

Methanol and Ethanol: Formulas and mass concepts

Methanol (a simple alcohol): formula options

$ext{CH}<em>4 ext{O}$ or $ext{CH}</em>3 ext{OH}$

Both representations describe methanol; one emphasizes the empirical formula (CH4O) and the other the structural form (CH3OH).

Ethanol: two-carbon alcohol

Formula: $ext{C}<em>2 ext{H}</em>6 ext{O}$

Often written as $ext{C}<em>2 ext{H}</em>5 ext{OH}$ to show the -OH group attached to a two-carbon chain.

Other example discussed: a five-carbon alcohol with a triple bond and an OH group

Pentynol (pent-yn-ol) scenario from the lecture: formula $ext{C}<em>5 ext{H}</em>8 ext{O}$

Rationale: five carbons, one triple bond (yn) and an alcohol (-OH) group; hydrogens adjusted accordingly in a saturated triple-bond context.

Mass and atomic weights basics used in the example

For methanol, mass (molar mass) is calculated from atomic weights: carbon ≈ $12.01, ext{H} ext{(each)} <br>≈ 1.008, ext{ O} ext{≈ } 15.999$

Formula for methanol: $ext{CH}<em>4 ext{O} ext{ or } ext{CH}</em>3 ext{OH}$

Molar mass of methanol: $M( ext{CH}_4 ext{O}) = 12.01 + 4(1.008) + 15.999 \approx 32.042 ext{ g/mol}$

Short-hand mass reference sometimes given in class: ~32 g/mol (rough estimate)

Quick notes on significant figures (sig figs)

If carbon is given as 12.01, hydrogen as 1.008, oxygen as 15.999, the result is typically reported with the appropriate sig figs, e.g. $M ext{ with 4 s.f.} \approx 32.04 ext{ g/mol}$

If a problem specifies only 2 sig figs, report as oxed{32 ext{ g/mol}}; with 3 sig figs, oxed{32.0 ext{ g/mol}}; with 4 sig figs, oxed{32.04 ext{ g/mol}}

Practice takeaway: from the formula, count atoms to confirm composition and relate to mass via atomic weights; the mass (32) is tied to identity (methanol).

Teacher tips mentioned in the lecture

Bring a calculator you know how to use; avoid getting bogged down with unfamiliar calculators.

Stoichiometry, mass, and identity of compounds

• When a compound has the atoms arranged in a fixed ratio, that ratio defines the compound’s composition

• Methanol’s atom ratio (per molecule) is consistent: CH4O or CH3OH → C:H:O = 1:4:1

• Concept of mass conservation in reactions

• The total mass before and after a chemical reaction stays the same (conservation of mass)

• Rearranging atoms to form new substances cannot change the total number of atoms or total mass

• Empirical vs molecular formulas (illustrated by examples)

• For methanol, the empirical formula CH4O is the same as the molecular formula CH3OH

Atomic theory concepts covered in the lecture

• Dalton’s atomic theory (early 1800s) summarized

• Elements are composed of indivisible units called atoms

• Atoms of the same element are identical in mass and properties (at that time)

• In chemical reactions, atoms are neither created nor destroyed; they are rearranged to form new substances

• Compounds are formed when atoms combine in fixed ratios; the total number of atoms is conserved in reactions

• Law of constant composition (definite proportions)

• A given compound always contains its elements in the same fixed ratio by mass or by number of atoms

• Example: water always has H:O in a fixed ratio (2:1 for atoms; 2 H atoms to 1 O atom in the common representation H$_2$O)

• Law of conservation of mass

• In any chemical reaction, mass is conserved; total mass of reactants equals total mass of products

• Law of multiple proportions (and related idea of definite proportions)

• When two elements form more than one compound, the ratio of the masses of the second element that combine with a fixed mass of the first element is a simple whole-number ratio

• Example implications discussed: nitrogen and oxygen can form NO and NO$_2$ with distinct simple mass relationships; you cannot have fractional or non-integer ratios in these simple compounds from fixed masses

Historical experiments and what they revealed

• Cathode ray tube experiments and discovery of the electron

• A cathode ray is deflected by electric and magnetic fields, revealing a negatively charged particle: the electron

• J. J. Thomson’s work led to the concept of electrons and introduced the plum pudding model (electrons embedded in a diffuse positive matrix)

• Millikan oil-drop experiment (determining electron charge)

• Oil droplets become charged (negative) and are suspended in an electric field; balancing gravity and electric forces allows calculation of the charge on a single droplet

• From the quantified charges, the fundamental charge e is determined (approximately $-1.602 imes 10^{-19} ext{ C}$)

• By knowing the mass of the droplets (from volume and density) and their charge, Millikan and successors established the electron’s charge-to-mass ratio, $rac{e}{m_e} ext{ (≈ } 1.758 imes 10^{11} ext{ C/kg)}$

• Rutherford’s gold foil experiment and the nuclear model of the atom

• Experimental setup: alpha particles fired at thin gold foil; a detector screen observed scattering

• Classic plum pudding model (Thomson) predicted most alpha particles would pass through with slight deflections

• Observations: most particles passed through, but some deflected at large angles; a few even bounced back

• Conclusion: atoms are mostly empty space, with a very dense, small nucleus containing protons and (later known) neutrons; electrons occupy the surrounding space

• This experiment replaced the plum pudding model with the nuclear model of the atom

• Early models and interpretations mentioned in class

• Plum pudding model (Thomson): diffuse positive matrix with embedded electrons

• Nuclear model (Rutherford): dense nucleus in the center with electrons orbiting in the surrounding space

Key formulas and concepts to remember (in LaTeX)

• Molar mass of a compound: $M = ext{sum of atomic masses in the formula}$

• Methanol molar mass example: $M( ext{CH}_4 ext{O}) = 12.01 + 4(1.008) + 15.999 \approx 32.042 ext{ g/mol}$

• Volume of a sphere (relevant for Millikan oil-drop analysis of droplet volume): $V = rac{4}{3}\pi r^3$

• Mass from volume and density: $m = <br>ho V$

• Electron charge: $e \,( ext{elementary charge}) \approx -1.602\times 10^{-19} ext{ C}$

• Electron mass: $m_e \approx 9.109\times 10^{-31} ext{ kg}$

• Electron-to-mass ratio: $\frac{e}{m_e} \approx 1.758\times 10^{11} ext{ C kg}^{-1}$

Connections to broader chemistry concepts and real-world relevance

• Understanding mass, atoms, and bonding explains why materials have fixed compositions and how reactions proceed

• Electron behavior underpins covalent bonding (sharing electrons) and the stability of atoms

• The balance of forces in small systems (e.g., Millikan’s droplet) links macroscopic measurements to subatomic properties

• The shift from plum pudding to nuclear model illustrates how hypotheses are revised with new experimental data; science advances through testable predictions and revisions

Foundational takeaways and implications

• Matter is composed of atoms; atoms combine in fixed ratios to form compounds; mass is conserved in chemical changes

• The discovery of charge and mass of fundamental particles (electrons, protons, neutrons) explains chemical bonding and reactivity

• Early models (plum pudding) were superseded by models informed by precise measurements (Rutherford), highlighting the iterative nature of science

• Practical lab skills highlighted: using appropriate tools (calculators) and interpreting data with respect to significant figures and measurement precision

Quick recap of major terms

• Atom, element, compound, molecule, ion

• Dalton’s atomic theory, law of definite proportions, law of conservation of mass, law of multiple proportions

• Cathode ray, electron, proton, neutron, nucleus, plum pudding model, nuclear model

• Molar mass, empirical formula, molecular formula, sig figs

Final thought from the lecture

• The instructor emphasizes that these classic experiments laid the groundwork for modern chemistry and that many of the key ideas came from careful observation, measurement, and willingness to revise models when data disagreed with predictions