# Lecture 2: Atoms, Molecules, and Ions
## General Chemistry I
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## Learning Objectives
By the end of this lecture, students will be able to:
1. Describe the key experiments that led to the modern atomic model (Dalton, Thomson, Millikan, Rutherford)
2. Identify the subatomic particles and their properties (charge, mass, location)
3. Use atomic number, mass number, and isotope notation to describe atoms
4. Calculate average atomic mass from isotopic abundances
5. Distinguish between molecular compounds and ionic compounds
6. Name and write formulas for common ionic and molecular compounds
7. Recognize common polyatomic ions and acids
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## Lecture Content
### I. Historical Development of Atomic Theory
#### A. Dalton's Atomic Theory (1803)
John Dalton proposed that all matter is composed of indivisible atoms, and that atoms of a given element are identical in mass and properties. He further stated that compounds form through the combination of atoms of different elements in fixed, whole-number ratios, and that chemical reactions simply rearrange atoms without creating or destroying them -- a statement of the conservation of mass. Modern science has since refined Dalton's ideas: we now know that atoms are divisible into subatomic particles, and that isotopes mean not all atoms of a given element have identical mass. Nevertheless, Dalton's framework remains a cornerstone of chemistry.
#### B. Discovery of the Electron -- Thomson (1897)
J. J. Thomson's cathode ray tube experiments demonstrated that rays produced in the tube were deflected by electric and magnetic fields. Because the rays traveled from the cathode (negative terminal) to the anode (positive terminal), Thomson concluded they carried a negative charge. The charge-to-mass ratio (e/m) was the same regardless of the cathode material, revealing that this particle was universal. Thomson proposed the "plum pudding" model, in which negatively charged electrons were embedded in a diffuse sphere of positive charge.
#### C. Millikan's Oil Drop Experiment (1909)
Robert Millikan measured the charge on a single electron to be e = 1.602 x 10^-19 C. By combining this value with Thomson's e/m ratio, he calculated the electron's mass as 9.109 x 10^-31 kg. Crucially, Millikan's experiment demonstrated that electric charge is quantized -- it always appears as a whole-number multiple of the fundamental charge e.
#### D. Rutherford's Gold Foil Experiment (1911)
Ernest Rutherford directed alpha particles at a thin gold foil and observed that most passed straight through, indicating that the atom is mostly empty space. A small fraction were deflected at large angles, and very few bounced directly back. This led Rutherford to conclude that all of the atom's positive charge and nearly all of its mass are concentrated in a tiny, dense nucleus. His nuclear model placed a small positive nucleus at the center of the atom, with electrons occupying the relatively vast space surrounding it.
A four-panel timeline illustration of key atomic experiments. Panel A (Dalton, 1803): simple spherical atoms of different elements shown as different-colored solid spheres combining in fixed ratios. Panel B (Thomson, 1897): a cathode ray tube diagram showing the cathode, anode, deflection plates, and fluorescent screen with the beam path curving toward the positive plate; inset shows the plum pudding model. Panel C (Millikan, 1909): oil drop apparatus with charged plates, microscope, and X-ray source, showing a suspended oil droplet between the plates. Panel D (Rutherford, 1911): alpha particle source aimed at gold foil, with most particles passing through, some deflected at angles, and a few bouncing back; inset shows the nuclear model with a tiny dense nucleus and orbiting electrons.
### II. Subatomic Particles
| Particle | Symbol | Charge | Relative Mass (amu) | Location |
|----------|--------|--------|---------------------|----------|
| Proton | p+ | +1 | 1.0073 | Nucleus |
| Neutron | n0 | 0 | 1.0087 | Nucleus |
| Electron | e- | -1 | 0.00055 | Electron cloud |
The nucleus contains protons and neutrons, collectively known as nucleons. Although it accounts for more than 99.97% of the atom's mass, the nucleus occupies a remarkably tiny volume with a diameter on the order of 10^-15 m. Electrons occupy the much larger space surrounding the nucleus, giving the atom an overall radius of roughly 10^-10 m. In a neutral atom, the number of protons equals the number of electrons, ensuring electrical neutrality.
### III. Atomic Number, Mass Number, and Isotopes
The atomic number (Z) is the number of protons in the nucleus and defines the element's identity. The mass number (A) is the total count of protons plus neutrons (A = Z + N). Isotope notation is written as ^A_Z X -- for example, ^12_6 C, ^13_6 C, and ^14_6 C all represent isotopes of carbon. Isotopes are atoms of the same element that differ in their number of neutrons. Because they share the same electron configuration, isotopes exhibit the same chemical properties, but they differ in mass and nuclear stability. The three isotopes of hydrogen illustrate this well: protium (^1H) has no neutrons, deuterium (^2H) has one, and tritium (^3H) has two.
### IV. Average Atomic Mass
The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes of an element. It is calculated by summing the product of each isotope's fractional abundance and its isotopic mass. Consider chlorine as an example: ^35Cl has a mass of 34.969 amu and an abundance of 75.76%, while ^37Cl has a mass of 36.966 amu and an abundance of 24.24%. The average atomic mass is therefore (0.7576)(34.969) + (0.2424)(36.966) = 35.45 amu. The atomic mass unit (amu) itself is defined as one-twelfth the mass of a ^12C atom.
### V. Ions
A cation is an atom that has lost one or more electrons, giving it a net positive charge. Metals commonly form cations; for example, sodium loses one electron to become Na+. An anion is an atom that has gained one or more electrons, resulting in a net negative charge. Nonmetals commonly form anions, as when chlorine gains an electron to become Cl-. The charge on any ion equals its number of protons minus its number of electrons. Monatomic ions of main-group elements tend to achieve the electron configurations of the nearest noble gas.
### VI. Molecules and Molecular Compounds
A molecule consists of two or more atoms held together by covalent bonds. A molecular formula shows the exact number of each type of atom in the molecule (for example, H2O or C6H12O6), while an empirical formula gives the simplest whole-number ratio of atoms (CH2O for glucose). A structural formula goes further, showing how the atoms are connected by depicting bonds as lines. Seven elements naturally exist as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, and I2, often remembered by the mnemonic BrINClHOF.
### VII. Ionic Compounds
Ionic compounds form between metals (which become cations) and nonmetals (which become anions), held together by the electrostatic attraction known as an ionic bond. Rather than existing as discrete molecules, ionic compounds arrange themselves in crystal lattices. For this reason, they are represented by empirical (formula) units rather than molecular formulas -- we write NaCl, not Na2Cl2. In every ionic compound, the total positive charge must equal the total negative charge, ensuring overall electrical neutrality.
A side-by-side comparison of ionic vs. molecular compounds. Left panel (Ionic -- NaCl): a 3D crystal lattice showing alternating Na+ (small, labeled) and Cl- (large, labeled) ions in a face-centered cubic arrangement, with arrows indicating electrostatic attractions. Right panel (Molecular -- H2O): individual water molecules shown with covalent bonds (solid lines) between O and H atoms, lone pairs on oxygen shown as electron dots, and dashed lines between molecules indicating weak intermolecular forces. Labels highlight key differences: "Extended lattice vs. discrete molecules," "Ionic bonds vs. covalent bonds," "High melting points vs. lower melting points."
### VIII. Naming Ionic Compounds
#### A. Binary Ionic Compounds (Type I -- metals with one common charge)
For binary ionic compounds involving a metal that forms only one type of cation, the naming convention is straightforward: name the cation (metal) first, then name the anion using the nonmetal's root plus the suffix "-ide." Thus NaCl is sodium chloride, MgO is magnesium oxide, and CaF2 is calcium fluoride.
#### B. Metals with Variable Charges (Type II)
When the metal can form cations of different charges, a Roman numeral in parentheses indicates the charge on the metal. FeCl2 is iron(II) chloride, while FeCl3 is iron(III) chloride. Common metals with variable charges include iron, copper, tin, lead, cobalt, manganese, and chromium.
#### C. Ionic Compounds with Polyatomic Ions
When a compound contains a polyatomic ion, the ion is treated as a single unit. Important polyatomic ions to know include NH4+ (ammonium), OH- (hydroxide), NO3- (nitrate), NO2- (nitrite), SO4^2- (sulfate), SO3^2- (sulfite), CO3^2- (carbonate), PO4^3- (phosphate), and the chlorine oxyacid series: ClO4- (perchlorate), ClO3- (chlorate), ClO2- (chlorite), and ClO- (hypochlorite). Additional ions worth memorizing are CH3COO- (acetate), CN- (cyanide), MnO4- (permanganate), and Cr2O7^2- (dichromate). Examples of naming include Ca(NO3)2 as calcium nitrate and (NH4)2SO4 as ammonium sulfate.
### IX. Naming Molecular (Covalent) Compounds
Molecular compounds formed between two nonmetals are named using Greek prefixes to indicate the number of each type of atom: mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, and deca-. The first element is given a prefix (with "mono-" omitted) followed by the element name. The second element receives a prefix plus the element root with the suffix "-ide." For example, CO2 is carbon dioxide, N2O5 is dinitrogen pentoxide, and PCl3 is phosphorus trichloride.
### X. Naming Acids
Binary acids, which consist of hydrogen and a nonmetal dissolved in water, are named with the prefix "hydro-" followed by the nonmetal root and the suffix "-ic acid." HCl(aq) is hydrochloric acid and HBr(aq) is hydrobromic acid. Oxyacids contain hydrogen paired with a polyatomic ion that includes oxygen. When the polyatomic ion has the "-ate" suffix, the acid takes the "-ic acid" suffix (NO3- is nitrate, so HNO3 is nitric acid). When the ion has the "-ite" suffix, the acid takes the "-ous acid" suffix (NO2- is nitrite, so HNO2 is nitrous acid).
A reference chart of common polyatomic ions organized by charge. Column 1 (1- charge): nitrate NO3-, nitrite NO2-, hydroxide OH-, cyanide CN-, permanganate MnO4-, acetate C2H3O2-, and the halogen oxyacid series (perchlorate, chlorate, chlorite, hypochlorite) with their formulas. Column 2 (2- charge): sulfate SO4^2-, sulfite SO3^2-, carbonate CO3^2-, chromate CrO4^2-, dichromate Cr2O7^2-, oxalate C2O4^2-. Column 3 (3- charge): phosphate PO4^3-, phosphite PO3^3-. Column 4 (1+ charge): ammonium NH4+. Each ion's Lewis structure is shown as a small inset beside its name.