Introduction to Chemistry: Nature of Matter and Atomic Structure

These flashcards cover core chemistry concepts including the nature of matter, atomic theory, subatomic particles, naming conventions, quantum mechanics, and periodic trends based on the lecture notes.

Chemistry

The science of matter and the transformations it can undergo.

Hypothesis

A tentative idea to explain observations that suggests further experiments to check if it is correct.

Quantitative data

Information that contains numbers or numerical values.

Law

A summary of a large number of experiments.

Theory

A unifying principle used to explain a body of facts and the laws based on them.

Nanoscale

Objects with dimensions approximately the size of an atom, utilizing the SI prefix for $10^{-9}$, where $1\text{ nm} = 1 \times 10^{-9}\text{ m}$.

Kinetic-Molecular Theory

The theory stating that matter consists of tiny particles in constant motion.

Pure Substances

Matter that can only be separated through chemical methods and contains only one component.

Heterogeneous mixture

A mixture that is different throughout its composition.

Homogeneous mixture

A mixture that is the same throughout its composition.

Physical Properties

Attributes that can be measured without changing the composition of a substance, such as temperature, density, or color.

Chemical property

A description of a chemical reaction that a substance can undergo.

Density

A physical property calculated as mass divided by volume ($d = \frac{m}{V}$).

Significant digit

The digits in a measurement including all certain digits and one estimated (uncertain) digit.

Law of Conservation of Mass

The law stating that mass is neither created nor destroyed in a chemical reaction.

Law of Definite Proportions

The principle that a given compound will always contain the same proportion of elements by mass.

Radioactivity

A term coined by Marie Curie for the phenomenon where ores (like Uranium) emit rays that can fog photographic plates.

Electrons ($e^-$)

Subatomic particles discovered by Thomson using cathode rays, with a negative charge and a mass of approximately $9.109 \times 10^{-28}\text{ g}$.

Protons ($p^+$)

Positive fundamental particles located in the nucleus with a mass of approximately $1.672 \times 10^{-24}\text{ g}$.

Neutrons ($n^0$)

Neutral subatomic particles discovered by Chadwick in 1932, present in all atoms except normal Hydrogen.

Nucleus

A small, dense core of an atom containing protons and neutrons, concentrating most of the atom's mass and all positive charge.

Atomic number (Z)

The number of protons in an atom, which defines the element.

Mass number (A)

The total number of protons plus the number of neutrons in an atom.

Isotopes

Atoms of the same element that have equal numbers of protons but different numbers of neutrons.

Cation

A positively charged ion formed when an element (typically a metal) loses electrons.

Anion

A negatively charged ion formed when an element (typically a nonmetal) gains electrons.

Polyatomic Ions

Units composed of multiple atoms that carry a net electrical charge.

Crystal Lattice

A structure in which ionic compounds exist where each ion is surrounded by many others in a regular array.

Electromagnetic Radiation

Oscillating traveling waves consisting of electric and magnetic fields that move through a vacuum at the speed of light ($c \text{ approximately } 2.998 \times 10^8\text{ m/s}$).

Wavelength ($\text{λ}$)

The distance between two consecutive crests of a wave.

Frequency ($\text{ν}$)

The number of crests passing a fixed point per unit of time, measured in Hertz ($1\text{ Hz} = 1\text{ s}^{-1}$).

Quantum

The smallest packet of energy, as described by Planck, where $E = h\nu$.

Photons

Massless particles that represent quantized electromagnetic radiation.

Wave-Particle Duality

The concept that electromagnetic radiation exhibits both wave and particle properties.

Ground state

The lowest energy state of an atom (where $n = 1$ in the Bohr model).

Heisenberg Uncertainty Principle

The principle stating it is impossible to know both the exact position and exact momentum of an electron.

Orbital

An H-atom wavefunction ($\text{ψ}$) representing a probability map for finding an electron in space.

Pauli exclusion principle

The rule that every electron in an atom must have a unique set of four quantum numbers ($n, \text{ℓ}, m_{\text{ℓ}}, m_s$).

Hund’s rule

The rule stating that the most stable electron arrangement in a subshell is the one with the maximum number of unpaired electrons with the same spin.

Valence Electrons

Electrons in the outer shell, including those in incomplete shells and partially-filled $d$ and $f$ orbitals.

Isoelectronic

Species that possess an equal number of electrons.

Paramagnetic

An atom that is attracted to magnetic fields because it contains unpaired electrons.

Ionization energy (IE)

The energy required to remove an electron from a gas-phase atom.

Electron Affinity (EA)

The energy released when an electron is added to a gas-phase atom.

Born-Haber cycle

A series of steps used to calculate the lattice energy of an ionic compound.