General Chemistry I - Unit 1: Introduction to Atoms, Energy, Forces, and Stoichiometry

Vocabulary flashcards covering key terms, historical atomic models, fundamental forces, potential energy diagrams, intermolecular vs intramolecular forces, and stoichiometry concepts from General Chemistry I (Unit 1).

Scientific Question

A question that can be answered by conducting experiments, making observations, and taking measurements.

Scientific Model

A drawing, graph, diagram, or equation used to make sense of phenomena by making testable predictions and explaining how something happens.

Scientific Theory

The best available explanation of existing evidence, data, and observations that explains why phenomena occur, makes testable predictions, is falsifiable, and can change over time as new evidence arises.

Brownian Motion

The random, continuous movement of microscopic particles suspended in a fluid or gas, caused by collisions with surrounding atoms or molecules.

Dalton's Atomic Theory

An atomic theory proposed by John Dalton in 1803 stating that elements consist of indivisible, indestructible solid sphere atoms, all atoms of a given element are identical, compounds are combinations of different types of atoms, and chemical reactions rearrange atoms without creating or destroying them.

Cathode Ray Tube Experiment

An experiment conducted by J.J. Thomson in 1897 where high voltage applied across electrodes in a glass tube under low pressure produced a beam of negative particles that deflected toward a positive plate, leading to the discovery of the electron.

Plum Pudding Model

An atomic model proposed by J.J. Thomson in 1904 depicting the atom as negatively charged electrons embedded within a uniform spherical cloud of positive charge.

Alpha ($\alpha$) Particle

A positively charged subatomic particle consisting of two protons and two neutrons (identical to a helium-4 nucleus) ejected with high energy during radioactive decay.

Rutherford Scattering Experiment

An experiment conducted by Ernest Rutherford in 1911 in which positive alpha particles were fired at thin gold foil; most passed straight through while a few deflected at large angles, proving that the atom is mostly empty space with a small, dense, positively charged nucleus.

Nuclear Model

An atomic model proposed by Ernest Rutherford in 1911 featuring a small, heavy, positively charged central nucleus surrounded mostly by empty space in which electrons reside.

Planetary Model

An atomic model proposed by Niels Bohr in 1913 stating that electrons move around the nucleus in orbits of fixed sizes and quantised energy levels.

Quantum Model

An atomic model developed by Erwin Schrödinger in 1926 stating that electrons do not move in set orbits, but exist in 3D wave-like probability clouds called orbitals.

Proton

A subatomic particle located in the atomic nucleus with a $+1$ charge and a mass of approximately $1\,amu$ ($1.67 \times 10^{-27}\,kg$).

Neutron

An uncharged subatomic particle located in the atomic nucleus with a mass of approximately $1\,amu$ ($1.675 \times 10^{-27}\,kg$), discovered by James Chadwick in 1932.

Electron

A subatomic particle with a $-1$ charge and a mass of approximately $0.0005\,amu$ ($\sim \frac{1}{2000}\,amu$) occupying the electron cloud surrounding the nucleus.

Atomic Number ($Z$)

The number of protons in the nucleus of an atom, which defines the identity of a chemical element.

Mass Number ($A$)

The total number of protons and neutrons in the nucleus of a specific atom.

Isotopes

Atoms of the same element that have the same atomic number (number of protons) but different mass numbers due to differing numbers of neutrons.

Average Atomic Mass

The weighted average mass of all naturally occurring isotopes of an element, calculated from the mass and relative fractional abundance of each isotope.

Coulomb's Law

A mathematical relationship stating that the electrostatic force ($F$) between two charged particles is directly proportional to the product of their charges ($q_1 q_2$) and inversely proportional to the square of the distance ($r$) between them: $F \propto \frac{q_1 q_2}{r^2}$.

Gravitational Force

An attractive fundamental force between any two objects with mass modeled by $F \propto \frac{M_1 M_2}{r^2}$, which is negligible at the atomic/submicroscopic level compared to electrostatic forces.

Potential Energy ($E_p$)

The energy associated with the relative position of two or more objects in a field; for two point charges, it is modeled as $E_p \propto \frac{q_1 q_2}{r}$.

Kinetic Energy ($E_k$)

The energy an object possesses due to its motion, defined by $E_k = \frac{1}{2} m v^2$.

First Law of Thermodynamics

The law of conservation of energy stating that energy cannot be created or destroyed, only transferred between system and surroundings or transformed from one form to another.

System and Surroundings

The system is the specific part of the universe defined for study, while the surroundings comprise everything else in the universe outside the system.

Instantaneous Dipole

A temporary, asymmetric fluctuation in the electron density of an atom or nonpolar molecule that creates brief partial positive ($\delta+$) and partial negative ($\delta-$) charges.

Induced Dipole

A temporary separation of charge created in a neighboring atom or molecule when its electron cloud is distorted by an adjacent dipole.

London Dispersion Forces (LDFs)

Weak intermolecular attractive forces present between all atoms and molecules caused by momentary fluctuations in electron density inducing dipoles in neighboring particles.

Potential Energy Minimum

The internuclear distance on a potential energy curve where the system achieves maximum stability because electrostatic attractive forces equal repulsive forces ($F_{attractive} = F_{repulsive}$).

Thermal Energy

The total internal kinetic energy of all atoms or molecules in a sample, represented as the sum $\sum \frac{1}{2} m v^2$.

Temperature

A physical property that is directly proportional to the average kinetic energy ($\frac{1}{2} m v^2$) of the individual particles in a sample.

Covalent Bond

A strong intramolecular interaction formed when nonmetal atoms share electrons, requiring significant thermal energy ($>6000\,K$ for $\text{H}_2$) to break.

Intermolecular Forces (IMFs)

Forces of attraction that exist between separate molecules or atoms (e.g., LDFs), which determine phase change behavior and physical properties like melting and boiling points.

Intramolecular Forces

Strong attractive forces within a single molecule that hold its constituent atoms together, such as covalent bonds.

Avogadro's Number ($N_A$)

The number of representative particles (atoms, molecules, or ions) present in one mole of any substance, equal to $6.022 \times 10^{23}\,mol^{-1}$.

Mole

The SI base unit for the amount of substance, defined as containing exactly $6.022 \times 10^{23}$ particles.

Molar Mass

The mass in grams of one mole of a chemical substance, expressed in units of $g\,mol^{-1}$.

Solution

A homogeneous mixture of two or more pure substances that has a uniform composition throughout.

Solvent

The component of a solution present in the greatest abundance in which solutes are dissolved.

Solute

Any substance dissolved in the solvent to form a solution.

Molarity ($M$)

A unit of solution concentration defined as the number of moles of solute per liter of solution ($mol\,L^{-1}$).

Dilution

The process of reducing a solution's concentration by adding solvent, increasing total volume while keeping the total moles of solute constant.

Stoichiometry

The quantitative relationship between relative amounts of reactants and products consumed and produced in a balanced chemical reaction.

Limiting Reactant

The reactant in a chemical reaction that is completely consumed first, thereby limiting the maximum amount of product that can be formed.

Theoretical Yield

The maximum amount of product calculated to form in a chemical reaction assuming complete consumption of the limiting reactant.

Actual Yield

The measured mass or quantity of product actually collected from a laboratory reaction.

Percent Yield

The measure of a reaction's efficiency, calculated as $\text{Percent Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100\%$.