Grade 11 Chemistry Practice Flashcards

Comprehensive vocabulary flashcards covering atomic structure, bonding, states of matter, kinetics, equilibrium, and organic chemistry based on the Grade 11 textbook.

Atomos

A term suggested by Democritus meaning tiny, indestructible, or 'indivisible' particles.

Law of Conservation of Mass

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

Law of Definite Proportions

The law stating that a given compound always contains exactly the same proportion of elements by mass.

Law of Multiple Proportions

When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers.

Cathode Rays

Rays originating from the negative electrode (cathode) in an evacuated glass tube under high-voltage, consisting of a beam of negatively charged particles or electrons.

Radioactivity

The spontaneous emission of particles and/or radiation from the unstable nuclei of certain atoms.

Alpha ($\alpha$) Rays

Positively charged particles identical to helium nuclei with a mass about four times that of a hydrogen atom and a charge twice the magnitude of an electron.

Beta ($\beta$) Rays

Electrons originating from inside the nucleus that are deflected by a negatively charged plate.

Gamma ($\gamma$) Rays

High-energy rays with no charge that are not affected by external electric or magnetic fields.

Neutron

A nuclear particle discovered by James Chadwick in 1932 having a mass almost identical to that of a proton but with no electric charge.

Proton

A nuclear particle discovered by Rutherford having a positive charge equal in magnitude to that of an electron and a mass of approximately $1.67262 \times 10^{-27}\,\text{kg}$.

Atomic Number ($Z$)

The number of protons in the nucleus of each atom of an element.

Mass Number ($A$)

The total number of protons and neutrons in the nucleus of an atom.

Isotopes

Atoms of an element that have the same atomic number but different numbers of neutrons and different mass numbers.

Atomic Mass Unit (amu)

A unit of mass equal to $\frac{1}{12}$ the mass of an atom of carbon-12.

Wavelength ($\lambda$)

The distance a wave travels during one cycle, typically expressed in meters, nanometers, or angstroms.

Frequency ($\nu$)

The number of cycles a wave undergoes per second, expressed in units of $\text{1/s}$ or hertz (Hz).

Quantum

The smallest discrete quantity of energy that can be emitted or absorbed by atoms and molecules, proportional to the frequency of radiation.

Planck’s Constant ($h$)

A fundamental constant with a value of $6.63 \times 10^{-34}\,\text{J}\cdot\text{s}$.

Photoelectric Effect

A phenomenon in which electrons are ejected from the surface of certain metals exposed to light of at least a certain minimum threshold frequency ($\nu_o$).

Photons

Particles of light or discrete energy packets.

Ground State

The lowest energy state of an atom, corresponding to $n = 1$ in the Bohr model.

Excited State

Any energy level higher than the ground state, occurring when an electron is in a level where $n > 1$.

Heisenberg Uncertainty Principle

The principle stating it is not possible to know with great certainty both an electron’s position and its momentum at the same time.

Orbital

The region in space around the nucleus where an electron is most likely to be found.

Principal Quantum Number ($n$)

Describes the main energy level or shell an electron occupies; can be any positive integer.

Angular Momentum Quantum Number ($\ell$)

Designates the shape of atomic orbitals and identifies subshells; takes values from $0$ to $n-1$.

Magnetic Quantum Number ($m_l$)

Relates to the orientation of the orbital in space relative to other orbitals; has integral values between $-\ell$ and $\ell$.

Electron Spin Quantum Number ($m_s$)

Refers to the spin of an electron and the orientation of the magnetic field produced by this spin, taking values of $+\frac{1}{2}$ or $-\frac{1}{2}$.

Aufbau Principle

The scheme used to reproduce electronic configurations by filling orbitals in order of increasing energy.

Hund’s Principle

States that equal energy (degenerate) orbitals are each occupied by a single electron before a second electron with opposite spin enters.

Pauli’s Exclusion Principle

No two electrons in an atom can have the same four quantum numbers.

Valence Electrons

Electrons in the outermost principal quantum level of an atom, involved in chemical bonding.

Core Electrons

The inner-level electrons of an atom not involved in chemical bonding.

Metalloid

An element with properties intermediate between those of metals and non-metals, such as boron, germanium, and silicon.

Atomic Radius

An estimate of atomic size measured as half the distance between the nuclei of two adjacent atoms.

Effective Nuclear Charge ($Z_{eff}$)

The nuclear charge an electron actually experiences after accounting for shielding by inner electrons.

Ionization Energy (IE)

The energy required for the complete removal of 1 mol of electrons from 1 mol of gaseous atoms or ions.

Electron Affinity (EA)

The energy change for the process of adding an electron to a neutral atom in the gaseous state to form a negative ion.

Electronegativity

Indicates the extent of attraction by which the electrons of a bond pair are attracted by an atom linked by that bond.

Chemical Bond

The attractive force which holds atoms, ions, and molecules together.

Octet Rule

The tendency of atoms to gain or lose electrons until they have achieved an outer shell containing eight electrons.

Ionic Bond

A bond formed by the electrostatic attraction between positive and negative ions resulting from the transfer of one or more electrons.

Lattice Energy ($U$)

The enthalpy change that occurs when 1 mol of an ionic solid separates into gaseous ions.

Lewis Symbol

Consists of a chemical symbol representing the nucleus and core electrons, surrounded by dots representing valence electrons.

Born–Haber Cycle

An indirect method using Hess's law to calculate the lattice energy of an ionic solid.

18-Electron Rule

A rule that replaces the octet rule for some transition and post-transition elements due to the involvement of d orbitals.

Covalent Bond

A bond formed when a pair of electrons is shared between two atoms.

Bond Length

The optimum distance between nuclei in a covalent bond where net attractive forces are maximized.

Coordinate Covalent Bond

A bond formed when one atom donates both electrons in a shared pair to another atom with a vacant valence orbital; also called a dative bond.

Resonance

A description of a molecule or ion for which two or more valid Lewis structures can be written; the true structure is a hybrid of these.

Dipole Moment ($\mu$)

The product of the magnitude of the charge ($\delta$) and the distance ($d$) separating the charges in a polar molecule.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; states that electron pairs around a central atom arrange themselves to minimize repulsions.

Intramolecular Forces

Chemical bonds (ionic, covalent, or metallic) that exist within a particle and affect chemical properties.

Intermolecular Forces

Attractive forces between neutral molecules that hold particles together and affect physical properties like boiling points.

Dipole-Dipole Forces

Attractive forces acting between molecules possessing permanent dipoles.

Hydrogen Bond

An especially strong dipole-dipole attraction occurring between a partially positive hydrogen atom and an electronegative atom like N, O, or F in another molecule.

London (Dispersion) Forces

Weak attractive forces resulting from small, instantaneous dipoles that occur due to electron motion; present between all particles.

Metallic Bonding

The sharing of free, delocalized electrons among a lattice of positively charged metal ions.

Sigma ($\sigma$) Bond

A localized covalent bond where the highest electron density lies along the axis between the two nuclei, resulting from head-on overlap.

Pi ($\pi$) Bond

A covalent bond formed by the lateral or sideways overlap of parallel p atomic orbitals, with electron density above and below the bond axis.

Hybridization

An imaginary mixing process in which atomic orbitals rearrange to form new, equivalent hybrid orbitals.

Molecular Orbital (MO) Theory

A theory describing bonding where atomic orbitals are simultaneously transformed into new molecular orbitals associated with the whole molecule.

Bond Order

The number of electrons in bonding MOs minus the number in antibonding MOs, divided by two.

Paramagnetic

A property of species with unpaired electrons that causes them to be attracted by external magnetic fields.

Crystalline Solid

A solid composed of one or more crystals with a well-defined, ordered three-dimensional structure.

Amorphous Solid

A solid with a disordered structure lacking a well-defined arrangement of basic units.

Plasma

The fourth physical state of matter consisting of a gaseous mixture of positive ions and electrons, existing at extremely high temperatures.

Kinetic Theory of Matter

States that all matter is composed of small particles in a state of continuous and random motion.

Standard Temperature and Pressure (STP)

Standard conditions for gases defined as a pressure of 1 atmosphere and a temperature of $273\,\text{K}$ ($0\,^{\circ}\text{C}$).

Boyle’s Law

States that the volume of a fixed mass of gas is inversely proportional to its pressure at a constant temperature ($P_1V_1 = P_2V_2$).

Charles’ Law

States that the volume of a fixed mass of gas at constant pressure varies directly with its Kelvin temperature ($\frac{V_1}{T_1} = \frac{V_2}{T_2}$).

Gay-Lussac’s Law

States that at constant volume, the pressure of a fixed amount of gas varies directly with the temperature ($\frac{P_1}{T_1} = \frac{P_2}{T_2}$).

Avogadro’s Law

States that at the same temperature and pressure, equal volumes of gases contain equal numbers of moles ($V = kn$).

Ideal Gas Equation

The relationship $PV = nRT$ describing the behavior of a hypothetical gas that obeys all gas laws.

Diffusion

The spreading of gas molecules throughout a container or given space.

Graham’s Law of Diffusion

States that the rate of diffusion of a gas is inversely proportional to the square root of its density or molar mass.

Vaporization (Evaporation)

The process by which molecules on the surface of a liquid break away into the gas phase.

Condensation

The process where a vapor returns to the liquid state.

Vapor Pressure

The partial pressure of a vapor in dynamic equilibrium with its liquid in a closed container.

Boiling Point

The temperature at which a liquid’s vapor pressure equals the external atmospheric pressure.

Molar Heat of Vaporization ($\Delta H_{vap}$)

The amount of heat needed to convert 1 mole of a liquid at its boiling point to a gas.

Molar Heat of Fusion ($\Delta H_{fus}$)

The quantity of heat needed to convert one mole of a solid at its melting point to the liquid state.

Sublimation

The process of a solid changing directly to vapor without passing through the liquid state.

Deposition

The process in which gaseous substances directly change into the solid state.

Chemical Kinetics

The area of chemistry concerned with reaction rates and mechanisms.

Rate of Reaction

The change in concentration of a reactant or product per unit time.

Activation Energy ($E_a$)

The minimum amount of energy needed for a collision between reactants to result in a reaction.

Catalyst

A substance that changes reaction rate by providing a different reaction mechanism with a lower activation energy, without being consumed.

Chemical Equilibrium

A state in a reversible reaction when the rates of the forward and reverse reactions are equal and concentrations remain constant.

Dynamic Equilibrium

A state where reactants and products are interconverted continually at equal rates, resulting in no net change in composition.

Law of Mass Action

The rate of a chemical reaction is directly proportional to the product of the concentrations of the reactants, each raised to the power of its stoichiometric coefficient.

Reaction Quotient ($Q$)

The ratio of product concentrations to reactant concentrations at any stage of a reaction, used to predict the direction toward equilibrium.

Le Châtelier’s Principle

If a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change.

Haber Process

The industrial synthesis of ammonia ($NH_3$) from nitrogen and hydrogen gases.

Contact Process

The industrial production of sulfuric acid ($H_2SO_4$) via the catalytic oxidation of sulfur dioxide to sulfur trioxide.

Alcohols

Derivatives of hydrocarbons in which one or more hydrogen atoms are replaced by a hydroxyl group ($-OH$).

Fermentation

The slow decomposition of carbohydrates like sucrose in the presence of enzymes, resulting in ethanol and carbon dioxide.

Ethers

Organic compounds in which an oxygen atom is bonded to two alkyl substituents ($R-O-R'$).

Aldehyde

A compound containing a carbonyl group ($C=O$) bonded to at least one hydrogen atom ($RCHO$).