Chemistry Notes: Foundations of Matter and Atomic Structure
Core Concepts of Matter and Atomic Structure
Types of Matter
Substances: Matter with a fixed composition.
Elements: Pure substances that cannot be broken down into simpler substances by any physical or chemical means.
Compounds: Pure substances formed when two or more elements are chemically combined in fixed proportions.
Mixtures: A group of two or more substances that are physically intermingled but not chemically bonded.
Atoms and Molecules
Atoms:
The smallest quantity of matter that still retains all the properties of an element.
All atoms of a given element are identical.
Atoms of one element are different from atoms of any other element.
Composed of tiny particles.
Molecules: Two or more atoms joined together and acting as a single unit.
The smallest part of a compound that still retains the properties of that compound.
Historical Development of Atomic Theory
Democritus (460-370 BC): Often called "The Father of Atomism." Proposed that all matter is made of indivisible particles called atoms.
Robert Boyle (17th Century): Suggested that elements are made up of one type of simple body.
Antoine Lavoisier (18th Century): Established the Law of Conservation of Mass.
John Dalton (19th Century): Proposed the atomic theory, building upon earlier observations and laws.
Dalton's Atomic Theory (1808)
Elements are composed of tiny, indivisible particles called atoms.
Each element is characterized by the mass of its atoms. Atoms of the same element have identical masses, but atoms of different elements have different masses.
Chemical compounds are formed when atoms of different elements combine in small, whole-number ratios.
Chemical reactions only rearrange how atoms are combined in chemical compounds; the atoms themselves do not change (i.e., they are neither created nor destroyed, nor converted into atoms of another element).
Fundamental Laws Governing Chemical Reactions
Law of Conservation of Mass
Statement: Mass is neither created nor destroyed in a chemical reaction.
Proponent: Antoine Lavoisier (18th Century).
Implication: The total mass of reactants in a chemical reaction must equal the total mass of the products.
Example Reaction: Hg(NO3)2(aq) + 2 KI(aq) \rightarrow HgI2(s) + 2 KNO3(aq)
Aqueous solutions of mercury(II) nitrate and potassium iodide react to form a precipitate of mercury(II) iodide and aqueous potassium nitrate. If known amounts of solid KI and solid Hg(NO3)2 are weighed, dissolved, mixed, and then the products (HgI2 and KNO3) are isolated and weighed, the combined masses of the products will equal the combined masses of the reactants.
Law of Definite Proportions
Statement: Different samples of a pure chemical compound always contain the same proportion of elements by mass.
Implication: No matter its source, a particular compound is always composed of the same elements in the same parts (fractions) by mass.
Example: Calcium Carbonate (CaCO_3) always contains approximately 40 \text{ Mass % Ca}, 12 \text{ Mass % C}, and 48 \text{ Mass % O}.
Law of Multiple Proportions
Statement: When two elements form a series of compounds, the ratio of the masses of one element that combine with a fixed mass of the other element can always be reduced to small whole numbers.
Proponent: John Dalton.
Example (Compounds between N and O):
Compound A: 1.750 \text{ g} nitrogen combines with 1 \text{ g} oxygen.
Compound B: 0.8750 \text{ g} nitrogen combines with 1 \text{ g} oxygen.
Compound C: 0.4375 \text{ g} nitrogen combines with 1 \text{ g} oxygen.
Ratios to demonstrate the law:
Ratio A:B = \frac{1.750 \text{ g N}}{0.8750 \text{ g N}} = 2/1
Ratio B:C = \frac{0.8750 \text{ g N}}{0.4375 \text{ g N}} = 2/1
Ratio A:C = \frac{1.750 \text{ g N}}{0.4375 \text{ g N}} = 4/1
These ratios are small whole numbers, confirming the law.
Characterization of Atoms: Discoveries of Subatomic Particles
J. J. Thomson and the Electron (1897)
Experiment: Conducted experiments using cathode-ray tubes. Applied high voltage to partially evacuated glass tubes, observing a stream of negative particles (cathode rays).
Hypothesis: Assumed the cathode ray comprised electrons.
Measurement: Determined the charge-to-mass ratio (e/m) of an electron.
e is the charge of the electron in coulombs (C).
m is the mass in grams (g).
Observations from Cathode Rays:
Rays bend in magnetic fields.
Rays bend towards a positive plate in electric fields.
Rays are identical for any type of cathode material.
Conclusions:
Rays consist of charged particles.
Particles are negatively charged.
These particles (electrons) are found in all matter.
The Plum Pudding Model
Hypothesis (Thomson): Atoms were envisioned as a diffused, positively charged cloud with negative electrons embedded randomly within it, like plums in a pudding.
Robert Millikan and the Electron's Charge (1909)
Experiment: Performed the oil drop experiment, involving electrically charged plates and oil droplets, some of which were ionized by X-rays.
Measurement: Determined the magnitude of the charge on a single electron.
Charge of electron = -1.602 \times 10^{-19} \text{ C} (The negative sign indicates its charge).
Using Thomson's charge-to-mass ratio and this determined charge, he calculated the mass of the electron.
Calculated Mass of Electron: 9.11 \times 10^{-31} \text{ kg}.
Radioactivity (Turn of the 20th Century)
Discovery: Accidentally discovered by Henri Becquerel, who observed that a uranium-containing mineral produced an image on a photographic plate in the dark.
Definition: The spontaneous emission of particles or radiation from the atoms of certain elements.
Types of Radioactive Emissions:
Gamma rays ($\gamma$): High-energy electromagnetic radiation (often referred to as "light").
Beta particles ($\beta$): High-speed electrons.
Alpha particles ($\alpha$): Possess a positive charge twice that of the electron's charge magnitude (i.e., +2 charge) and are approximately 7,000 times more massive than an electron. They have a velocity of approximately 1.4 \times 10^7 \text{ m/s} (about 5\% the speed of light).
Ernest Rutherford and the Atomic Nucleus (1908 Nobel Prize in Chemistry)
Experiment (Gold Foil Experiment): Directed a beam of positively charged alpha ($\alpha$) particles at a very thin gold foil.
Results (Observed vs. Expected):
Expected (based on Plum Pudding Model): Alpha particles, being massive and fast, were expected to pass through the diffuse positive charge of the gold atoms with only minor deflections, if any.
Actual:
Most alpha particles passed straight through the foil undeflected.
A small number (about 1 in every 20,000) were deflected at large angles.
A very few actually bounced back toward the alpha particle source.
Conclusion:
Because the majority of particles passed through undeflected, the gold atoms must be almost entirely empty space.
The atom's mass and all of its positive charge must be concentrated in a tiny, dense central core, which Rutherford called the nucleus.
Only alpha particles that happened to strike or pass very close to this positively charged nucleus were deflected at large angles or bounced back.
Modern View of Atomic Structure
Composition of an Atom
Nucleus:
Small compared to the overall size of the atom.
Extremely dense.
Accounts for almost all of the atom’s mass (99.97\%), but occupies a negligible volume.
Contains protons and neutrons.
Protons:
Found in the nucleus.
Possess a positive charge (+1) equal in magnitude to the electron’s negative charge.
Mass is approximately 1836 times the mass of an electron.
Neutrons:
Found in the nucleus.
Have no electrical charge (neutral).
Virtually the same mass as a proton.
Electrons:
Found outside the nucleus, moving around it.
Possess a negative charge (-1).
Account for most of the atom's volume but contribute very little to its mass.
Comparison of Subatomic Particles
Particle | Mass (grams) | Mass (u)* | Charge (coulombs) | Charge (e) | Location |
|---|---|---|---|---|---|
Electron | 9.109383 \times 10^{-28} | 5.485798 \times 10^{-4} | -1.602177 \times 10^{-19} | -1 | Outside nucleus |
Proton | 1.672622 \times 10^{-24} | 1.007276 | +1.602177 \times 10^{-19} | +1 | Nucleus |
Neutron | 1.674927 \times 10^{-24} | 1.008665 | 0 | 0 | Nucleus |
*The unified atomic mass unit (u) is defined as exactly 1/12 the mass of an atom of carbon--12 (1 \text{ u} = 1.660538783 \times 10^{-24} \text{ g}).
Chemical Properties
All atoms have the same components, but their different chemical properties are caused by differences in:
The number of electrons.
The arrangement of electrons.
Electrons of different atoms intermingle to form molecules, and the degree of interaction in an atom is determined by the number of electrons it possesses.
Atomic Number (Z) and Mass Number (A)
Atomic Number (Z):
Represents the number of protons in an atom's nucleus.
It defines the element.
In a neutral atom, the number of electrons equals the number of protons.
Mass Number (A):
Represents the total number of protons and neutrons in an atom's nucleus.
Isotopes
Definition: Atoms of the same element (i.e., same number of protons, or same atomic number Z) but with different numbers of neutrons (and thus different mass numbers A).
Representation: The symbol for an isotope is typically written as ^{A}_{Z}X , where X is the element symbol.
Example: Two Isotopes of Sodium
Sodium-23 ($^{23}_{11}Na$): Contains 11 protons, 11 electrons, and 12 neutrons (23-11=12).
Sodium-24 ($^{24}_{11}Na$): Contains 11 protons, 11 electrons, and 13 neutrons (24-11=13).
Practice Problem Example:
Question: Write the symbol for the atom that has an atomic number of 9 and a mass number of 19. How many electrons and how many neutrons does this atom have?
Solution:
Atomic number 9 identifies the element as Fluorine (F).
The symbol is ^{19}_{9}F. This isotope is called fluorine-19.
Since it's a neutral atom with 9 protons, it must also have 9 electrons.
The number of neutrons is the mass number minus the atomic number: 19 - 9 = 10 neutrons.
Unified Atomic Mass Unit (u or amu)
Definition: One unified atomic mass unit (u) is defined as exactly 1/12 the mass of an atom of carbon-12 ($^{12}C$).
Value: 1 \text{ u} = 1.660538783 \times 10^{-24} \text{ g}.
Relevance: The mass of an atom is measured relative to this standard. The mass of carbon-12 is the only one that is exact (defined as 12 \text{ u}).
Atomic Mass (Atomic Weight)
Definition: The weighted average of the atomic masses of an element’s naturally occurring isotopes.
Formula: Atomic Weight = (P1\% \times M1) + (P2\% \times M2) + (P3\% \times M3) + \ldots
P\% is the natural abundance of each isotope.
M is the mass of the isotope in atomic mass units (u).
Example: Atomic Weight of Carbon
Carbon-12: 98.89\% natural abundance, mass = 12 \text{ u}.
Carbon-13: 1.11\% natural abundance, mass = 13.0034 \text{ u}.
Atomic Weight of Carbon = (12 \text{ u})(0.9889) + (13.0034 \text{ u})(0.0111) = 11.867 \text{ u} + 0.14433774 \text{ u} \approx 12.011 \text{ u} . (The 11.867 \text{ u} mentioned in the transcript is only the first part of the sum).
Ions
Definition: Electrically charged atoms formed by either removing or adding electrons to a neutral atom.
Cations:
Positively charged ions.
Formed when a neutral atom loses one or more electrons.
Example: Removing an electron from a neutral sodium atom (Na) forms a sodium ion (Na^+}).
Anions:
Negatively charged ions.
Formed when a neutral atom gains one or more electrons.
Example: Adding an electron to a neutral chlorine atom (Cl) forms a chloride ion (Cl^-).
Moles and Molar Mass
Avogadro’s Number (N$_A$):
Represents the number of particles (atoms, molecules, ions, etc.) in one mole of any substance.
N_A = 6.022 \times 10^{23} particles/mole.
Molar Mass:
The mass in grams of one mole of any element or compound.
Numerically, for an element, it is equivalent to its atomic weight (expressed in grams per mole, g/mol).
For a compound/molecule, it is equivalent to its molecular weight (expressed in grams per mole, g/mol).