General Chemistry: Chapters 1-9 Review

Flashcards covering atomic history, quantum mechanics, stoichiometry, thermodynamics, gas laws, and properties of liquids and solids based on lecture topics.

Plum Pudding Model

A model proposed by J.J. Thomson after discovering electrons, visualizing the atom as a positive space ('pudding') with negative electrons ('plums') stuck inside.

Gold Foil Experiment

An experiment by Lord Rutherford where alpha particles were shot at gold foil; most went through, but a few bounced back, proving the existence of a tiny, dense, positive nucleus.

Law of Conservation of Mass

A principle from Dalton’s atomic theory stating that matter cannot be created or destroyed in a chemical reaction.

Law of Definite Proportion

States that a specific compound always contains exactly the same proportion of elements by mass.

Radioactivity

The discovery by Henri Becquerel that atoms can spontaneously break apart and release energy.

Qualitative Observation

Descriptions or observations made using the senses without numerical data.

Quantitative Observation

Observations involving numbers and measurements.

Cation

A positively charged ion, typically formed by metals.

Anion

A negatively charged ion, typically formed by nonmetals.

Wavelength ($\lambda$)

The distance between two peaks of a wave; longer wavelengths correspond to lower energy.

Frequency ($\nu$)

The number of waves that pass a point per second; higher frequency corresponds to more energy.

Planck’s Constant ($h$)

A fundamental constant approximately equal to $6.626 \times 10^{-34}\,J\,s$ used in the equation $E = h \times \nu$.

Aufbau Principle

The rule stating that electrons must fill the lowest energy orbitals first (e.g., $1s$ before $2s$).

Hund’s Rule

Often called the 'Bus Seat Rule,' it states that electrons occupy empty orbitals in a subshell singly before pairing up.

Pauli Exclusion Principle

States that an orbital can hold a maximum of two electrons, and they must have opposite spins ($m_s = +1/2$ and $-1/2$).

Principal Quantum Number ($n$)

Indicates the shell, size, and energy level of an orbital.

Angular Momentum Quantum Number ($l$)

Defines the subshell and the shape of the orbital ($s, p, d, f$).

Nodal Surface

A region within an atomic orbital where the probability of finding an electron is zero.

Ionic Bond

A bond formed between a metal and a non-metal where electrons are transferred, resulting in an electronegativity difference ($\Delta EN$) greater than $1.7$.

Polar Covalent Bond

A bond where electrons are shared unequally between non-metals, typically with a $\Delta EN$ between $0.4$ and $1.7$.

Isoelectronic

Refers to different elements or ions that have the exact same number of electrons.

Resonance

Occurs when more than one valid Lewis structure can be drawn for a single molecule; the actual molecule is a resonance hybrid.

Delocalized Electrons

Electrons that are shared by three or more atoms simultaneously rather than being localized between two atoms.

Formal Charge

The charge assigned to an atom in a molecule, calculated by assuming electrons in bonds are shared equally to identify the most stable Lewis structure.

VSEPR Model

The Valence Shell Electron-Pair Repulsion model, which postulates that electron pairs stay as far apart as possible to minimize repulsion.

Sigma ($\sigma$) Molecular Orbital

A molecular orbital formed by the head-on overlap of atomic orbitals, where electron density is concentrated between the nuclei.

Bond Order

A measure of molecular stability calculated as $\frac{(\text{no. of bonding } e^-) - (\text{no. of antibonding } e^-)}{2}$.

Avogadro’s Number

The number of entities in one mole, equal to $6.022 \times 10^{23}$, used to convert between atomic mass units (amu) and grams.

Empirical Formula

The simplest whole-number ratio of the elements present in a compound.

Limiting Reactant

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

Molarity ($M$)

A unit of concentration defined as the moles of solute per liter of solution.

Strong Electrolyte

A substance that undergoes $100\%$ dissociation into ions when dissolved in water, resulting in high conductivity.

Reducing Agent

The species in a redox reaction that loses electrons and is itself oxidized.

Potential Energy ($E_p$)

Energy stored due to position or composition, such as chemical energy stored in bonds.

First Law of Thermodynamics

States that energy cannot be created or destroyed, expressed by the equation $\Delta E = q + w$.

Enthalpy ($H$)

A state function defined as the sum of internal energy ($E$) and the product of pressure and volume ($PV$); at constant pressure, $\Delta H = q_p$.

Hess’s Law

States that the total enthalpy change for a reaction is the same regardless of whether it occurs in one step or several, because enthalpy is a state function.

Boyle’s Law

States that for a fixed amount of gas at constant temperature, volume is inversely proportional to pressure ($P_1 \times V_1 = P_2 \times V_2$).

Charles’s Law

States that for a fixed amount of gas at constant pressure, volume is directly proportional to absolute temperature ($\frac{V_1}{T_1} = \frac{V_2}{T_2}$).

Ideal Gas Law

The combined gas law expressed as $PV = nRT$, where $R = 0.08206\,L\times atm / K\times mol$.

Root-Mean-Square Velocity ($u_{rms}$)

The average velocity of gas particles, calculated as $\sqrt{\frac{3RT}{M}}$.

Viscosity

A liquid’s resistance to flow; it is directly proportional to the strength of intermolecular forces.

Vapor Pressure

The pressure exerted by a gas in equilibrium with its liquid phase in a closed container; it is inversely proportional to the strength of intermolecular forces.

Unit Cell

The smallest repeating unit of a crystal lattice that reproduces the entire lattice when stacked in three dimensions.

Sublimation

The direct phase transition from a solid to a gas without passing through the liquid phase.