Chemistry Grade 9 - Vocabulary Flashcards (Unit 1-5)

Vocabulary flashcards covering key terms from structure of the atom, periodic table, bonding, reactions, and states of matter. Each card defines a core concept or term that appears in the notes.

Atom

The basic unit of matter that composes all substances, consisting of a nucleus (protons and neutrons) surrounded by electrons in an electron cloud.

Nucleus

The tiny, dense region at the center of an atom that contains protons and neutrons and carries most of the atom’s mass.

Electron

A negatively charged subatomic particle with a very small mass that orbits the nucleus in electron clouds.

Proton

A positively charged subatomic particle located in the nucleus with mass about 1 amu.

Neutron

A neutrally charged subatomic particle in the nucleus with mass about 1 amu.

Atomic Number (Z)

The number of protons in the nucleus, which identifies the element and equals the number of electrons in a neutral atom.

Mass Number (A)

The total number of protons and neutrons in the nucleus.

Atomic Mass Unit (amu)

A unit used to express atomic and molecular weights, defined as 1/12 the mass of a carbon-12 atom.

Isotope

Atoms of the same element (same Z) that have different numbers of neutrons, hence different mass numbers.

Relative Atomic Mass

A mass comparison scale where carbon-12 is assigned 12 amu; other atoms are measured relative to this standard.

Isotopic Designation (Hyphen notation)

A method to designate isotopes, e.g., Hydrogen-2 (deuterium) or Oxygen-18.

Nuclear Symbol (A/Z X)

A notation for isotopes showing mass number A and atomic number Z around the element symbol, e.g., 14/7N.

Dalton’s Atomic Theory

Early theory proposing that elements are made of atoms, atoms are indivisible, all atoms of an element are identical, and atoms combine to form compounds.

Modern Atomic Theory

An updated view acknowledging subatomic particles, isotopes, and electrons arranged in energy levels; atoms are divisible and some atoms of the same element have different masses.

Dalton’s Law of Conservation of Mass

Matter is neither created nor destroyed in a chemical reaction; mass of reactants equals mass of products.

Dalton’s Postulates

Key assumptions of Dalton’s theory (e.g., atoms exist, are indivisible, atoms of a given element are identical, etc.).

Cathode Ray

A beam of electrons produced in a vacuum tube; used to study electron properties.

Electron Charge-to-Mass Ratio (e/m)

A measure determined by Thomson’s cathode-ray experiments; the ratio of the electron’s charge to its mass.

Millikan Oil-Drop Experiment

Experiment that measured the charge of the electron (e−).

Discovery of the Electron

Identification of the first subatomic particle with negative charge, shown to be part of all atoms.

Nucleus Model (Rutherford)

A tiny, dense, positively charged nucleus surrounded by mostly empty space; atoms are mostly empty space.

Bohr Model

Electrons orbit the nucleus in quantized circular orbits (energy levels) with fixed energies.

Quantum Mechanical Model

Electrons are in orbitals—probability regions around the nucleus—describing electron locations without fixed paths.

Orbital

A region of space around the nucleus where an electron is likely to be found.

Sublevel (s, p, d, f)

Divisions within main energy levels where electrons reside; each has a limited number of orbitals.

Principal Quantum Number (n)

A number that indicates the main energy level (shell) occupied by electrons.

Aufbau Principle

Electrons fill the lowest-energy orbitals first before filling higher-energy ones.

Ground-State Electron Configuration

The arrangement of electrons in the lowest available energy levels for an atom.

Noble Gas Core

Using a noble gas configuration as a core and appending outer electrons for a compact electron configuration.

Valence Electrons

Electrons in the outermost shell involved in chemical bonding.

Electron Configuration Notation

A shorthand method to show electrons in each sublevel (e.g., 1s2 2s2 2p6 3s2).

Atomic Radius

Half the distance between the centers of two identical atoms in a diatomic molecule; a measure of atomic size.

Nuclear Charge (Z)

Total positive charge of the nucleus (number of protons).

Shielding (Screening) Effect

Inner electrons shield valence electrons from the full positive nuclear charge.

Effective Nuclear Charge (Zeff)

Net positive charge experienced by valence electrons after shielding; Zeff = Z − S.

Period

A horizontal row in the periodic table; elements in a period have the same number of electron shells.

Group

A vertical column in the periodic table; elements in a group have similar chemical properties and the same number of valence electrons.

Representative Elements (Main Group Elements)

Elements in groups IA–VIIIA whose outermost electrons fill s- or p-orbitals.

Transition Elements

Elements in the d-block; their outer electrons fill the d-sublevel.

Inner-Transition Elements (Lanthanides/Actinides)

Elements where outer electrons fill f-orbitals (lanthanides and actinides).

Modern Periodic Law

The properties of elements are periodic functions of their atomic numbers.

Mendeleev’s Periodic Law

Properties of elements recur periodically with increasing atomic mass; he predicted undiscovered elements.

Mosley’s Contribution

Showed that elements are better organized by atomic number, not atomic mass.

Periodic Table Blocks (s, p, d, f)

Regions in the periodic table corresponding to which sublevel is being filled.

Ionization Energy

Energy required to remove the outermost electron from a gaseous atom; tends to increase across a period and decrease down a group.

Electron Affinity

Energy change when an electron is added to a neutral atom to form an anion; often releases energy (negative value).

Electronegativity

Ability of an atom to attract electrons in a chemical bond (Pauling scale).

Shielding vs Zeff

Shielding reduces Zeff; higher shielding lowers effective attraction of nucleus on valence electrons.

Diagonal Relationship

Similarities between certain pairs of elements across periods (e.g., Li–Mg, Be–Al, B–Si).

Ionic Bond

Electrostatic attraction between positively charged cations and negatively charged anions formed by electron transfer.

Covalent Bond

Bond formed by sharing electron pairs between atoms.

Polar Covalent Bond

Covalent bond with unequal electron sharing leading to partial charges; e.g., HCl.

Non-Polar Covalent Bond

Covalent bond where electrons are shared equally between identical atoms; e.g., H2.

Coordinate (Dative) Bond

Covalent bond where one atom donates both electrons to the bond.

Metallic Bond

Bonding in metals arising from a ‘sea’ of delocalized electrons surrounding a lattice of positive ions.

Lewis Structure

Electron-dot diagram showing bonding and lone pairs around atoms in a molecule.

Octet Rule

Atoms tend to have eight electrons in their valence shell to achieve stability.

Dipole

A separation of electrical charges within a molecule creating a polar bond.

Dipole–Dipole Forces

Intermolecular forces between polar molecules.

London Dispersion Forces

Weak intermolecular forces arising from temporary dipoles; present in all molecules, especially nonpolar ones.

Hydrogen Bond

A strong dipole–dipole interaction between a hydrogen atom bonded to N, O, or F and a lone pair on another electronegative atom.

Intermolecular Forces

Forces of attraction between molecules (Van der Waals forces, including dipole-dipole, London dispersion, and hydrogen bonding).

Ionic Compound

A compound composed of cations and anions held together by ionic bonds; typically solid with high melting points.

Covalent Compound

A compound formed by sharing electrons; may be liquids or gases at room temperature.

Molar Mass (Molar Masses)

Mass of one mole of a substance, in g/mol; e.g., 44.01 g/mol for CO2.

Aqueous Solution

A solution in which water is the solvent.

Reaction

A process in which reactants are transformed into products.

Reactants vs. Products

Substances that react vs. substances produced by a chemical reaction.

Balanced Chemical Equation

An equation with the same number of each type of atom on both sides.

Law of Conservation of Mass

Matter cannot be created or destroyed in a chemical reaction; mass is conserved.

Law of Definite Proportions

A compound always has the same proportion by mass of its elements.

Law of Multiple Proportions

When two elements form more than one compound, the ratios of the masses of the one element that combine with a fixed mass of the other are small whole numbers.

Stoichiometry

Quantitative relationships between reactants and products in a chemical reaction.

Limiting Reactant

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

Theoretical Yield

The maximum amount of product that could be formed from given amounts of reactants.

Actual Yield

The amount of product actually obtained from a reaction.

Percentage Yield

Actual yield divided by theoretical yield, multiplied by 100%.

Avogadro’s Law

Equal volumes of gases at the same temperature and pressure contain the same number of molecules; volume is proportional to moles.

Ideal Gas Law

PV = nRT; relates pressure, volume, temperature, and amount of gas.

Graham’s Law of Diffusion

Rate of diffusion is inversely proportional to the square root of molar mass.

Diffusion

Spread of gas molecules from high to low concentration.

Vapor Pressure

Pressure exerted by a vapor in equilibrium with its liquid at a given temperature.

Boiling Point

Temperature at which a liquid’s vapor pressure equals external pressure; bubbles form throughout the liquid.

Heat of Fusion

Energy required to melt one mole of a solid at its melting point.

Heat of Vaporization

Energy required to vaporize one mole of a liquid at its boiling point.

Heat of Sublimation

Energy required to sublime one mole of a solid.

Absolute Zero

Lowest possible temperature, 0 K (−273.15°C); molecular motion is minimal.

STP

Standard Temperature and Pressure: 0°C (273.15 K) and 1 atm.

Phase Diagram

Graph showing phases of a substance under different temperatures and pressures; heating curves illustrate phase changes.

Sublimation

Phase transition from solid directly to gas.

Deposition

Phase transition from gas directly to solid.

Phase Change

Transition between solid, liquid, and gas (fusion/melting, freezing, vaporization, condensation, sublimation, deposition).