General Chemistry – Foundations, Matter, Atomic Structure & Mass Spectrometry

50 vocabulary flashcards summarizing essential terms from the lecture covering matter classification, physical vs chemical changes, Dalton’s atomic theory, subatomic particles, isotopes, relative masses, and mass spectrometry.

Chemistry

The central science that studies the properties and behavior of matter by examining atoms and molecules.

Macroscopic Realm

The observable world of ordinary-sized objects; bulk properties of matter are viewed here.

Submicroscopic Realm

The level of atoms and molecules used to explain macroscopic properties of matter.

Matter

Anything that has mass and occupies space.

Atom

The basic building block of matter; the smallest unit of an element that retains its chemical identity.

Molecule

A group of two or more atoms joined in specific shapes by chemical bonds.

Solid

Physical state in which matter has fixed shape and volume.

Liquid

Physical state with definite volume but no fixed shape; it takes the shape of its container.

Gas

Physical state with neither fixed shape nor fixed volume; it expands to fill its container.

Element

A substance that cannot be decomposed into simpler substances; composed of one kind of atom.

Compound

A substance composed of two or more elements chemically combined in fixed proportions.

Mixture

A physical combination of two or more substances in which each retains its chemical identity.

Heterogeneous Mixture

A mixture with non-uniform composition where different components are visible (e.g., granite).

Homogeneous Mixture (Solution)

A mixture with uniform composition throughout (e.g., salt water).

Physical Property

A characteristic that can be observed without changing a substance’s composition (e.g., boiling point).

Chemical Property

A property observed only when a substance undergoes a chemical change (e.g., flammability).

Physical Change

A change that alters form or appearance but not composition (e.g., melting ice).

Chemical Change (Chemical Reaction)

A process in which one or more substances are transformed into different substances (e.g., combustion).

Law of Constant Composition

In a pure compound, the elemental composition is always the same (Dalton, Postulate 4).

Law of Multiple Proportions

When two elements form more than one compound, the masses combining with a fixed mass of one element are in small whole-number ratios.

Law of Conservation of Mass

Mass is neither created nor destroyed in a chemical process; total mass remains constant.

Cathode Rays

Streams of negatively charged particles (electrons) observed in discharge tubes, discovered by J. J. Thomson.

Electron

A subatomic particle with negative charge (–1) and very small mass (≈0.00055 amu).

Plum Pudding Model

Thomson’s early atomic model with electrons embedded in a diffuse positive sphere.

Gold Foil Experiment

Rutherford’s α-particle scattering experiment that led to discovery of the small, dense atomic nucleus.

Nucleus

The central, very small, positively charged core of an atom containing protons and neutrons.

Proton

Positively charged subatomic particle found in the nucleus; mass ≈1 amu.

Neutron

Electrically neutral subatomic particle in the nucleus; mass ≈1 amu.

Atomic Number (Z)

Number of protons in an atom; defines the element.

Mass Number (A)

Total number of protons plus neutrons in an atom.

Isotope

Atoms of the same element with the same proton number but different neutron numbers (different A).

Ion

An atom or group of atoms that carries an electrical charge due to gain or loss of electrons.

Cation

A positively charged ion formed by loss of electrons.

Anion

A negatively charged ion formed by gain of electrons.

Polyatomic Ion

A charged species composed of two or more covalently bonded atoms acting as a single ion.

Relative Atomic Mass (Ar)

Average mass of an element’s atoms compared to 1/12 the mass of a 12C atom; unit-less.

Relative Molecular Mass (Mr)

Mass of one molecule relative to 1/12 the mass of 12C; found by summing Ar of constituent atoms.

Relative Formula Mass

Mr applied to ionic compounds; sum of Ar values in the empirical formula.

Relative Isotopic Mass

Mass of a specific isotope relative to 1/12 the mass of 12C; no units.

Isotopic Abundance

The relative proportion (fraction or percent) of each isotope in a natural sample of an element.

Average Atomic Mass

Weighted average of isotopic masses based on their natural abundances (appears on periodic table).

Mass Spectrometer

Instrument that measures relative atomic/molecular masses by ionizing samples and separating ions by m/e ratio.

Mass-to-Charge Ratio (m/e)

The ratio of ion mass to its charge; determines the degree of deflection in a mass spectrometer’s magnetic field.

Vaporisation (MS Step 1)

Conversion of sample to gas before ionisation in a mass spectrometer.

Ionisation (MS Step 2)

Bombardment of gaseous atoms/molecules with high-energy electrons to produce positive ions.

Acceleration (MS Step 3)

Use of electric plates to speed up ions so they enter the magnetic field uniformly.

Deflection (MS Step 4)

Bending of ion paths by a magnetic field; amount depends on m/e.

Detection (MS Step 5)

Arrival of ions at the collector where relative abundances are recorded as peak heights to produce a mass spectrum.

Mass Spectrum

Graph of detector signal intensity versus ion mass that reveals isotopic masses and abundances.

Limiting Reactant

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

Percentage Yield

(Actual yield / Theoretical yield) × 100 %; measures efficiency of a chemical reaction.