General Chemistry Exam Review Flashcards

Vocabulary practice flashcards generated from lecture notes covering fundamental chemistry principles, including matter classification, measurement units, atomic theory, quantum numbers, periodic trends, and chemical nomenclature.

Pure Elements

Elements that cannot be broken down chemically.

Compounds

Substances composed of two or more elements combined in a fixed ratio, such as $\text{H}_2\text{O}$.

Homogeneous Mixture

A mixture characterized by uniformity throughout the sample, such as salt water.

Heterogeneous Mixture

A mixture with non-uniform composition throughout, such as lemonade.

Solid

A state of matter with a fixed shape and fixed volume, in which particles are packed tightly together, ordered, and vibrate in place while moving as a unit together.

Liquid

A state of matter with a fixed volume that takes the shape of its container.

Gas

A state of matter that expands to fill the entire container.

Plasma

A state of matter described as a combination of two states, with examples including lightning, stars, and blood.

Physical Properties

Properties observable without changing the chemical composition of a substance, such as color, melting point, and boiling point.

Chemical Properties

Behaviors exhibited by a substance during chemical changes, such as flammability and rusting.

Intrinsic Property

A physical property that is independent of the amount of substance present and remains constant regardless of sample size.

Extrinsic Property

A physical property that depends on the amount of substance present, such as weight and volume.

Percy Julian

A chemist noted for his work on soybeans, which led to synthesis of steroids and hormone therapy.

Scientific Method

The systematic process of scientific investigation that involves making observations, forming a hypothesis, performing experiments, and refining hypotheses.

SI Units

The international system of units established to unify measurement standards across countries.

Kelvin Conversion Formula

The temperature conversion formula given by $K = ^\text{o}\text{C} + 273.15$.

Derived Units

Units created through mathematical operations on simple base units, such as density (calculated as $\text{mass} / \text{volume}$ with units $\text{kg}/\text{cm}^3$).

Accuracy

How close a measured value is to the true or target value.

Precision

How close repeated measurements are to each other.

Captive Zeros

Zeros located between non-zero digits, which are always significant.

Leading Zeros

Zeros that precede all non-zero digits, which are never significant.

Trailing Zeros

Zeros at the end of a number, which are significant only if a decimal point is explicitly present; otherwise, their significance is ambiguous.

John Dalton's Atomic Theory

An 1807 theory stating that matter consists of small particles called atoms, each element has one type of atom, atoms of different elements have different properties, compounds contain whole-number ratios of atoms, and atoms are conserved in chemical reactions.

Electron

A negatively charged subatomic particle located outside the nucleus ("orbiting" outside rings), with a charge of $-1$, a mass of $9.1 \times 10^{-28}\,g$ ($0.0005\,\text{amu}$), discovered by J.J. Thomson in 1897 and confirmed by Milliken in 1909.

Proton

A positively charged subatomic particle located inside the nucleus, with a charge of $+1$, a mass of $1.67 \times 10^{-24}\,g$ ($1.007\,\text{amu}$), discovered by Ernest Rutherford in 1911 using alpha particles.

Neutron

A neutral subatomic particle located inside the nucleus, with a charge of $0$ and a mass of $1.68 \times 10^{-24}\,g$ ($1.009\,\text{amu}$).

Atomic Number ($Z$)

The total number of protons present in the nucleus of an atom, a concept coined by Henry Moseley in 1913.

Isotopes

Atoms of the same element that have the same atomic number but different numbers of neutrons, resulting in different mass numbers.

Cations

Positively charged ions formed when an atom loses electrons, which are smaller in size than their neutral counterparts.

Anions

Negatively charged ions formed when an atom gains electrons, which are larger in size than their neutral counterparts.

Periodicity

The phenomenon where chemical and physical properties of elements display a repeating pattern across the periodic table.

Dmitri Mendeleev

The chemist who developed an early periodic table in the mid-1800s based on atomic mass and overlapping elemental properties.

Valence Electrons

Electrons located in the outermost shell of an atom, which are the most mobile and govern chemical reactivity.

Octet Rule

The principle stating that elements tend to gain or lose electrons to obtain $8$ valence electrons (or $2$ for small atoms), mimicking noble gas electron configurations.

Transition Metals

The largest group of elements located in the middle of the periodic table (Groups 3–12), possessing variable numbers of valence electrons and variable oxidation charges.

Noble Gases

Group 18 elements that are inert (unreactive) because they possess a full outer shell of $8$ valence electrons.

Atomic Radius

The size of an atom, which decreases across a period from left to right and increases down a group from top to bottom.

Ionization Energy

The amount of energy required to remove an electron from an atom; it increases across a period from left to right and decreases down a group from top to bottom.

Chemical Bond

The force of attraction holding two or more atoms together in a compound.

Single Bond

A covalent chemical bond formed by sharing $2$ electrons ($1$ electron pair).

Double Bond

A covalent chemical bond formed by sharing $4$ electrons ($2$ electron pairs).

Triple Bond

A covalent chemical bond formed by sharing $6$ electrons ($3$ electron pairs).

Ionic Compounds

Compounds composed of metals and nonmetals ($\text{M} + \text{Nm}$) formed via the direct transfer of electrons from the metal to the nonmetal.

Covalent Compounds

Compounds composed of nonmetals ($\text{Nm} + \text{Nm}$) formed by the sharing of electrons between atoms.

Principal Quantum Number ($n$)

The quantum number (symbol $n$, allowed values $1, 2, 3\dots$) that describes the energy level or "shell" of electrons, where larger $n$ indicates greater distance from the nucleus.

Azimuthal Quantum Number ($l$)

The quantum number (symbol $l = n - 1$; allowed values $0, 1, 2\dots$) that describes the shape of the electron orbital ($l=0$ for s, $l=1$ for p, $l=2$ for d, $l=3$ for f).

Magnetic Quantum Number ($m_l$)

The quantum number (symbol $m_l$, allowed values $-l$ to $+l$) that describes the number and 3D spatial orientation of each orbital type.

Magnetic Spin Quantum Number ($m_s$)

The quantum number (symbol $m_s$, allowed values $+\frac{1}{2}$ or $-\frac{1}{2}$) that describes the spin direction of an electron.

Aufbau Principle

The rule stating that electrons fill the lowest energy orbitals available before occupying higher energy levels.

Hund's Rule

The rule stating that when filling orbitals of equal energy, electrons singly occupy each degenerate orbital before pairing up.

Pauli's Exclusion Principle

The principle stating that no two electrons in an atom can have an identical set of electronic quantum numbers because negatively charged electrons always repel each other.