Lecture Notes: Matter, Atomic Theory, Nuclear Chemistry, and Periodic Trends (Vocabulary Flashcards)
Matter and Its Properties
Matter is anything that has mass and volume.
Three states of matter:
Solid: rigid; fixed shape and fixed volume.
Liquid: definite volume; takes the shape of its container (shape adapts to container).
Gas: no fixed volume or shape; takes the shape and volume of its container.
Physical properties and changes:
Physical properties: properties that can be measured or observed without a substance turning into an entirely different substance.
Examples of physical properties include texture, size, shape, color, etc., as well as melting point, boiling point, and freezing point.
Melting point: the temperature at which a solid changes to liquid.
Boiling point: the temperature at which a liquid evaporates.
Freezing point: the temperature at which a liquid becomes a solid.
Solubility: the ability of a solute to dissolve in a solvent.
Texture-related properties include:
Size
Color
Shape
Chemical properties and changes:
A chemical property is a characteristic of a substance observed during a chemical reaction that can lead to the formation of new substances.
Examples include biodegradability and combustibility.
Biodegradability: the capacity of a material to decompose through microorganisms.
Combustion: a chemical reaction between a substance and oxygen that releases heat and light (flame).
Elements, Compounds, and Mixtures
Elements: cannot be broken down into other substances by chemical means. Examples: Iron (Fe), Aluminum (Al), Oxygen (O).
Compounds: contain atoms of different elements in a fixed ratio. Example: Water –
Water: H
two atoms of Hydrogen and one atom of Oxygen; chemical formula: ext{H}_2 ext{O}
Mixtures: classification of matter made up of two or more pure substances with varying compositions. Examples: Wine, Coffee.
Pure substances: have the same composition throughout; they are either elements or compounds.
History and Fundamentals: Atomic Theory
Democritus (around 440 B.C.E.): proposed that everything is made of tiny particles called "atomos" (indivisible).
John Dalton (1808): Dalton’s Atomic Theory:
1) Each chemical element is composed of extremely small particles called atoms, which are indivisible and not visible by naked eye.
2) All atoms of an element are alike in mass and properties; atoms of one element differ from those of other elements.
3) For each compound, different elements combine in simple numerical ratios.
J.J. Thomson (1897): discovered electrons via cathode ray experiment; proposed the Plum Pudding model with electrons embedded in a positively charged sphere.
Ernest Rutherford (1909): Gold foil experiment led to the Nuclear Model; atoms have a small, dense, positively charged nucleus containing protons; electrons orbit outside.
James Chadwick (1932): discovered neutrons.
Niels Bohr (1913): refined the nuclear model with electrons occupying fixed-energy orbits around the nucleus (Planetary Model).
Erwin Schrödinger (1926): developed the Quantum Mechanical Model by treating electrons as waves; Schrödinger equation yields wavefunctions and orbitals.
Contributions of Planck and Einstein as groundwork for quantum ideas.
Atomic Models and Key Concepts
Dalton’s Solid Sphere Model (1808): atoms as indivisible spheres.
Thomson’s Plum Pudding Model (1897): electrons within a positively charged matrix.
Rutherford’s Nuclear Model (1909): nucleus with protons; electrons outside in mostly empty space.
Bohr’s Planetary Model (1913): electrons in fixed orbits with quantized energies.
Schrödinger’s Quantum Mechanical Model (1926): electrons described by wavefunctions; orbitals define regions of probability for finding electrons.
The modern view uses wavefunctions and orbitals rather than fixed paths.
Quantum Mechanics, Orbitals, and Electron Arrangement
Atomic orbitals: regions of space where the probability of finding an electron is highest.
s-orbitals: spherical shapes.
p-orbitals: dumbbell shapes (two lobes).
d-orbitals: more complex shapes.
f-orbitals: even more complex shapes.
Sublevels and shapes:
Sublevel labels: s, p, d, f.
Each orbital can hold up to 2 electrons (with opposite spins).
Quantum numbers (to describe electrons):
Principal quantum number: n — main energy level (shell).
Azimuthal quantum number: l = 0,1,2,3 corresponding to s,p,d,f sublevels respectively.
Magnetic quantum number: m_l = -l, -l+1, …, +l — orientation of the orbital.
Spin quantum number: m_s = - frac{1}{2}, + frac{1}{2} — electron spin direction.
Rules governing electron configurations:
Aufbau Principle: electrons fill lower-energy atomic orbitals before filling higher-energy ones.
Pauli Exclusion Principle: no two electrons in an atom can have identical values for all four quantum numbers; each orbital can hold at most 2 electrons with opposite spins.
Hund’s Rule: electrons occupy degenerate orbitals singly with parallel spins before pairing up.
Shell capacities:
The maximum number of electrons in shell with principal quantum number n is 2n^2.
Examples: for n=1,2,3,4 the capacities are 2, 8, 18, 32 respectively.
Subshell capacities (per energy level):
s subshell: 2 electrons
p subshell: 6 electrons
d subshell: 10 electrons
f subshell: 14 electrons
Electron configuration notation reflects filling order and subshell capacities.
Total electrons in a given shell follow 2n^2; total electrons in a subshell follow its capacity (2, 6, 10, 14).
Electron cloud concept: a probabilistic region around the nucleus where electrons are likely to be found.
Isotopes, Nuclear Chemistry, and Nuclear Reactions
Isotopes: atoms of the same element with different mass numbers (same Z, different A).
Atomic mass unit (amu): standard unit for atomic and molecular masses.
Four fundamental forces: gravity, electromagnetic, strong nuclear force, weak nuclear force.
Gravity: weakest, long-range.
Electromagnetic: holds the atom together and governs chemical interactions.
Strong nuclear force: very short-range but strongest; holds nucleus together.
Weak nuclear force: involved in certain types of radioactive decay.
Nuclear particles: protons (p), neutrons (n), electrons (e), with nucleus containing protons and neutrons.
Nuclear stability and radioactivity:
In large nuclei, repulsive electromagnetic forces between many protons can overcome the strong force, leading to instability and radioactive decay.
Nuclei with too many neutrons or protons tend to be unstable and radioactive, especially when Z > 83 (beyond Bismuth).
Radioactivity: spontaneous decay of unstable nuclei with emission of particles and energy; results in daughter nuclei with different mass/energy.
Discovery history:
Henri Becquerel discovered natural radioactivity (co-discoveries by Marie and Pierre Curie).
Nuclear reactions and conservation:
Balanced nuclear reactions conserve total mass number and total charge:
Mass-number balance: \sum A{\text{reactants}} = \sum A{\text{products}}
Charge balance: \sum Z{\text{reactants}} = \sum Z{\text{products}}
Types of nuclear reactions:
Nuclear decay (radioactive decay): unstable nucleus emits radiation and transforms into another nucleus.
Nuclear transmutation: a nucleus reacts with another particle to become a different nuclide.
Typical radioactive emissions (radiation types):
Alpha (α): helium nucleus; symbolized as \alpha; relatively heavy and slow.
Beta (β−): electron emitted when a neutron converts to a proton.
Positron (β+): the antimatter counterpart of an electron.
Gamma (γ): high-energy photon; no mass or charge.
Neutron emission: emission of a free neutron.
Penetration abilities (relative to shielding):
Alpha cannot pass through paper.
Beta can pass through paper and skin but is stopped by dense materials.
Neutrons can pass through paper, skin, and some materials; require proper shielding.
Gamma rays can pass through paper, skin, metal, water, etc., but are attenuated by dense materials like lead.
Nuclear notation and reactions often include mass-number and atomic-number changes depending on the decay mode.
Isotopic Ratios and Nuclear Stability
Proton-neutron ratios influence stability:
Light, stable isotopes tend toward a close ratio (roughly 1:1 for very light nuclei).
Heavier nuclei require increasing numbers of neutrons to offset increasing repulsion among protons (ratio shifts as mass number grows).
Stability concepts tie to binding energy and decay paths; heavy nuclei tend to decay toward more stable configurations.
Practical Nuclear Chemistry: Applications and Examples
Nuclear transmutation can occur via particle bombardment (e.g., a nucleus reacts with a particle to become another nuclide).
Radioactive dating, medical imaging and therapy, and energy production are practical applications of nuclear chemistry.
Notation, Formulas, and Key Equations
Maximum electrons in a shell: N_{\text{shell}} = 2n^2\,, where n = 1,2,3,4,…
Subshell electron capacity: N_{\text{subshell}} = 2, 6, 10, 14 for s, p, d, f respectively.
Total electrons in an atom obeys the sum of shell capacities up to the atomic number Z.
Aufbau, Pauli, and Hund’s rules govern the order and arrangement of electrons in orbitals.
Nuclear conservation laws (mass number and charge):
\sum A{\text{reactants}} = \sum A{\text{products}}
\sum Z{\text{reactants}} = \sum Z{\text{products}}
Periodic Trends and Basic Properties (Overview)
Atomic Radius: half the distance between the nuclei of two identical atoms; trend varies across periods and groups.
Electronegativity: attraction of an atom for shared electrons in a chemical bond.
Ionization Energy: energy required to remove an electron from a neutral atom in the gaseous phase.
Electron Affinity: tendency of an atom to accept an electron.
Metallic Character: ease with which an element loses electrons; generally increases left to right across a period and decreases downward in a group (contextual notes from the material).
Laboratory Equipment and Common Glassware (Overview)
Glassware and containers:
Beaker: holds liquids; approximate volumes; used for pouring; accuracy roughly ±5%.
Florence flask: long-neck with rounded bottom; suitable for heating; stores liquids; covers hot vapors.
Petri dish: transparent; used to observe culture growth; prevents contamination.
Volumetric glassware: measures precise volumes; spout for pouring; reading meniscus carefully.
Pipette: manual or bulb/aspirator; various types; used to transfer precise volumes.
Burette: used for titrations; delivers accurate volumes.
Support equipment:
Test tube rack: holds multiple test tubes.
Ring stand with clamps: holds equipment during experiments.
Wire gauze and clay triangles for balancing beakers on stands.
Miscellaneous lab safety and functionality:
Tongs, crucibles, crucible tongs for heating solids at high temperatures.
Evaporation dishes, watch glasses, and sponges to handle liquids and moisture.
Desiccants and drying equipment to remove moisture from samples.
Safety and measurement devices:
Thermometers: measures temperature.
Balance/scale: measures mass; required for quantitative experiments.
pH meter or indicators: measures acidity/basicity (less or more accurate depending on device).
Reagents and reagents bottles: proper labeling and handling.
General lab practices:
Read and interpret menisci for accurate volume readings.
Use clamps and stands to secure glassware during procedures.
Handle hot equipment with tongs and heat-resistant gloves.
connections to Foundational Principles and Real-World Relevance
The evolution of atomic theory connects to the model of matter, radioactive decay, spectroscopy, and quantum chemistry.
Understanding physical vs chemical properties helps in material science, medicine, environmental science, and industry.
Nuclear chemistry concepts underpin energy production, medical imaging, radiotherapy, and dating techniques.
Periodic trends guide predictions about element reactivity, bonding, and material properties.
Lab equipment and safe handling practices are foundational for conducting reliable experiments and ensuring safety in scientific work.