CH1010 Exam 1

Law of Conservation of Matter

Matter occupies space and has mass; while it can change forms or state, total matter is conserved in chemical and physical processes.

Pure Substance

A form of matter that has a constant chemical composition and identical properties throughout (e.g., pure water, aluminum).

Mixture

A combination of two or more types of matter present in variable amounts that can be separated by physical means.

Homogeneous Mixture (Solution)

A mixture that possesses a uniform composition and appearance throughout (e.g., sugar completely dissolved in water).

Heterogeneous Mixture

A mixture with a non-uniform composition where individual phases or components remain distinct (e.g., iron nails in a cup of water).

Element

A pure substance consisting of only one type of atom that cannot be decomposed into simpler substances by chemical reactions.

Compound

A pure substance composed of two or more different elements chemically combined in fixed proportions; can be broken down into elements via chemical change.

Molecule

Two or more atoms bonded together by covalent chemical bonds acting as a discrete unit.

Physical Change

A change in the state or properties of matter without any accompanying change in its chemical composition or atomic structure.

Chemical Change

A process where atoms are rearranged, altering chemical composition and forming new substances with distinct chemical properties.

SI Base Units

The standard units of measurement: meter (m) for length, kilogram (kg) for mass, second (s) for time, Kelvin (K) for temperature, Ampere (A) for current, and mole (mol) for amount.

Intensive Property

A physical property that remains independent of the amount of matter present (e.g., temperature, density, pressure, boiling point).

Extensive Property

A physical property that depends directly on the quantity or mass of matter being measured (e.g., volume, mass, total internal energy).

Density

An intensive property defined as mass per unit volume (d = m/V).

Accuracy vs. Precision

Accuracy is how close a measured value is to the true or accepted value. Precision is how close repeated measurements are to one another.

Exact Numbers

Numbers with no uncertainty, having infinitely many significant figures (e.g., explicit counting values, metric conversion factors like 1 m = 100 cm).

Inexact Numbers

Numbers obtained through measurement that contain a degree of uncertainty and a limited count of significant figures.

Fundamental Atomic Particles

Protons (m_p ~ 1.0073 amu, q = +1), Neutrons (m_n ~ 1.0087 amu, q = 0), and Electrons (m_e ~ 0.00055 amu, q = -1).

Fundamental Unit of Charge

The absolute electric charge carried by a single proton or electron, equal to 1.602 x 10^-19 C.

Atomic Mass Unit (amu) Definition

Precisely defined as 1/12th of the mass of a single carbon-12 (_¹²C) atom.

Atomic Number (Z)

The number of protons in an atom's nucleus; uniquely defines the chemical identity of an element.

Mass Number (A)

The total sum of protons and neutrons present in an atom's nucleus (A = Z + neutrons).

Atomic Charge

Calculated as the total number of protons minus the total number of electrons.

Isotopes

Atoms of the same element containing equal numbers of protons but different numbers of neutrons, resulting in different mass numbers but identical chemical properties.

Average Atomic Mass Formula

The weighted average of isotope masses: AM = (M1 * A1) + (M2 * A2) + …, where M is isotopic mass and A is fractional abundance.

Structural Formula

A chemical representation displaying how atoms are bonded and connected together in space within a molecule.

Molecular Formula

A chemical formula showing the exact numbers of each type of atom present in a single molecule.

Empirical Formula

A formula providing the simplest whole-number ratio of atoms of each element in a chemical compound.

Isomers

Molecules that share identical molecular formulas but possess different structural arrangements or spatial orientations of their atoms.

Properties of Metals

Shiny/lustrous, malleable, ductile, good conductors of heat/electricity, and tend to lose electrons to form cations.

Properties of Nonmetals

Dull, brittle in solid state, poor conductors of heat/electricity, and tend to gain electrons (forming anions) or share electrons in covalent bonds.

Metalloids

Elements along the staircase dividing line possessing intermediate physical and chemical properties between metals and nonmetals.

Main Group/Representative Elements

Elements residing in Groups 1A-8A (s and p blocks) whose outer electrons fill s or p subshells and share strong vertical periodic similarities.

Transition Metals

Elements occupying the d-block (Groups 3-12) characterized by the progressive filling of d subshells.

Inner Transition Metals

Lanthanides and actinides occupying the f-block at the bottom of the periodic table.

Group 1A

Alkali metals

Group 2A

Alkaline earth metals

Group 7A

Halogens

Group 8A

Noble/Inert gases

Group 5A

Pnictogens

Group 6A

Chalcogens

Wavelength

Distance between consecutive crests

Frequency

Cycles per second (Hz)

Amplitude

Wave height from center

Intensity

Proportional to amplitude squared

Speed of Light Relation

In a vacuum, all electromagnetic radiation travels at c = 2.998 x 10^8 m/s, defined by c = wavelength * frequency.

Visible Spectrum Range

The narrow region of the electromagnetic spectrum visible to human eyes, bounded between 400 nm (violet) and 750 nm (red).

Standing Wave Nodes

Points along a stationary wave where the wave amplitude deviation is strictly zero (excluding the fixed endpoints).

Quantization

The physical principle where a variable cannot take continuous values, but is restricted to discrete allowed energy levels or values (quanta).

Blackbody Radiation

Electromagnetic radiation emitted by a completely absorbing, ideal object heated above 0 K.

The Ultraviolet Catastrophe

The classical physics failure predicting classical blackbody intensity would approach infinity at shorter UV wavelengths; solved by Max Planck.

Planck's Quantization Hypothesis

Energy emitted/absorbed by oscillating atoms is restricted to discrete quanta (E = n * h * frequency, where n = 1, 2, 3…).

Photoelectric Effect Key Findings

1. No electrons are ejected below threshold frequency regardless of intensity.
2. Ejected electron KE increases linearly with frequency above threshold.
3. Intensity increases electron flux, not electron KE.

Photon Concept (Einstein)

Light behaves as localized packets or particles of energy called photons, where photon energy is E = h * v = h * c / wavelength.

Work Function

The minimum energy required to eject an electron from the surface of a specific metal.

Wave-Particle Duality

The principle that matter and light exhibit both wave-like (interference, diffraction) and particle-like (collisions, localized mass/energy) behavior.

de Broglie Hypothesis

Every moving particle possesses a wavelength inversely proportional to its linear momentum (wavelength = h / p = h / (m * v)).

Continuous spectrum

Shows an uninterrupted rainbow of all wavelengths

Line spectrum

Emits discrete, narrow spectral lines unique to specific gas atomic transitions.

Bohr Model Postulates

1. Electrons orbit the nucleus in fixed circular paths without emitting radiation.
2. Orbits are quantized in energy.
3. Transitions between orbits absorb/emit photons.

Ground State vs. Excited State

Ground state: The lowest available quantum energy state (n = 1). Excited state: Any quantum energy state higher than ground state (n > 1).

Zero-Energy Reference Level (n = infinity)

At n = infinity, an electron is completely removed from the pull of the nucleus, establishing zero potential energy.

Heisenberg Uncertainty Principle

Fundamentally asserts that it is impossible to simultaneously measure both the exact position and momentum of a quantum particle.

Trajectory

Classical deterministic path requiring precise position/velocity (invalid for quantum mechanics).

Orbital

Quantum probability density region where an electron is likely found.

Wavefunction and Orbital

A mathematical function whose square represents the spatial probability density of locating an electron.

Principal Quantum Number (n)

Indicates the main electron shell, energy level, and size of the atomic orbital (n = 1, 2, 3…).

Azimuthal / Angular Momentum Quantum Number (l)

Indicates the subshell and geometric shape of the orbital (l = 0 to n-1). l=0 (s), l=1 (p), l=2 (d), l=3 (f).

Magnetic Quantum Number (m_l)

Indicates the spatial orientation of an orbital along 3D axes (m_l = -l to +l).

Spin Quantum Number (m_s)

Describes intrinsic magnetic orientation of an electron; takes values +1/2 (up) or -1/2 (down).

Subshell Node Calculation

The total number of nodes in an atomic orbital is n - 1. The number of radial nodes is n - l - 1, and angular nodes equal l.

Pauli Exclusion Principle

No two electrons in the same atom can possess the exact same set of all four quantum numbers, limiting each orbital to max 2 electrons.

Aufbau Principle

Electrons occupy the lowest available energy atomic orbitals before filling higher energy subshells.

Hund's Rule

For degenerate orbitals, electrons enter singly with parallel spins before pairing up to maximize stability.

Degenerate Orbitals

Orbitals within the same subshell that share the exact same energy level (e.g., 2px, 2py, 2pz).

Core Electrons

Inner closed-shell electrons represented by noble gas core.

Valence electrons

Outer shell electrons participating in chemical reactivity.

Transition Metal Cation Ionization Rule

When transition metals form cations, they lose their highest outer s-orbital electrons FIRST before losing any d-electrons.

Paramagnetism

Possesses unpaired electrons and is attracted to magnetic fields.

Diamagnetic

All electrons paired and is weakly repelled by magnetic fields.

Atomic Radius Periodic Trend

Atomic size increases down a group (more n shells) and decreases left-to-right across a period (increasing Z_eff pulls electrons tighter).

Effective Nuclear Charge (Z_eff)

The net positive nuclear charge experienced by an electron, accounting for inner core electron shielding (Z_eff = Z - Shielding).

Ionic Radii Trends

Cations are always smaller than neutral parent atoms (less electron repulsion). Anions are always larger than neutral parent atoms (increased electron repulsion).

Isoelectronic Series Trend

A group of species sharing identical electron configurations; radius decreases systematically as nuclear charge Z (proton count) increases.

Ionization Energy (IE) Trend

Energy required to remove an electron from a gaseous atom. Increases left-to-right across periods and decreases down groups.

Ionization Energy Deviations

1. Group 2A to 3A: IE₁(B) < IE₁(Be) because 2p electron is higher in energy than 2s.
2. Group 5A to 6A: IE₁(O) < IE₁(N) due to repulsion in paired 2p orbital.

Electron Affinity (EA)

The energy change associated with adding an electron to a gaseous atom or ion.