Chemistry: Thermochemistry, Bonding, Structure, and Molecular Geometry

Vocabulary practice flashcards generated directly from lecture material covering thermochemistry concepts, activation energy, catalysis, bonding types, solid structures, VSEPR theory, and molecular geometry.

Exothermic Reaction

A chemical reaction where energy is released to the surroundings, causing the surroundings to warm up ("feels hot") as temperature increases, where reactants have more energy than products, and the enthalpy change is negative ($-\Delta H$). Examples include respiration, combustion, and the reaction of $\text{Mg}$ and $\text{HCl}$.

Endothermic Reaction

A chemical reaction where energy is absorbed from the surroundings, causing the surroundings to cool down ("feels cold") as temperature decreases, where reactants have less energy than products, and the enthalpy change is positive ($+\Delta H$). Examples include dissolving $\text{NH}_4\text{Cl}$ and photosynthesis.

Bond Breaking

A chemical process that requires energy to be absorbed from the surroundings, meaning it is always endothermic.

Bond Making

A chemical process that releases energy to the surroundings, meaning it is always exothermic.

Activation Energy ($E_a$)

The minimum energy required for colliding reactant particles to undergo a successful chemical reaction.

Catalyst

A substance that increases the rate of a chemical reaction by providing an alternative pathway of lower activation energy ($E_a$) so a greater proportion of collisions are successful, without altering the enthalpy change ($\Delta H$).

Units for Enthalpy Changes

The standard metric units used to express enthalpy change ($\Delta H$), written as $kJ$ or $kJ\,mol^{-1}$.

Fusion and Vaporisation

State changes from solid to liquid ($\text{S} \rightarrow \text{L}$, fusion) and liquid to gas ($\text{L} \rightarrow \text{G}$, vaporisation), which are both endothermic processes requiring energy absorption.

Condensation and Freezing

State changes from gas to liquid ($\text{G} \rightarrow \text{L}$) and liquid to solid ($\text{L} \rightarrow \text{S}$), which are both exothermic processes releasing energy.

Reaction Energy Profile (Exothermic)

An energy level diagram for an exothermic reaction where the reactants are depicted at a higher energy level than the products, reflecting a negative enthalpy change ($-\Delta H$).

Reaction Energy Profile (Endothermic)

An energy level diagram for an endothermic reaction where the reactants are depicted at a lower energy level than the products, reflecting a positive enthalpy change ($+\Delta H$).

Catalysed vs Uncatalysed Energy Pathway

A reaction coordinate profile comparing an uncatalysed reaction pathway with higher activation energy ($E_a$) to a catalysed alternative pathway showing a reduced activation energy barrier.

Lewis Structure

A diagram showing the bonding between atoms in a molecule as well as any lone pairs of valence electrons.

Molecule

A neutral group of two or more atoms held together by covalent bonds.

Multiple Bond

A covalent bond formed between two atoms containing more than one shared pair of electrons, such as a double or triple bond.

Covalent Molecular Solid

A solid consisting of discrete, small covalently bonded molecules held together by weak intermolecular attractions, characterised by low melting and boiling points.

Ionic Solid

A 3D crystal lattice composed of alternating positively and negatively charged ions held together by strong electrostatic attractions.

Metallic Solid

A solid structure composed of metal cations bound together by non-directional attractions to a delocalised sea of valence electrons.

Covalent Network Solid

A giant structure composed of many atoms organized in a regular 3D lattice linked together by strong covalent bonds (e.g., diamond, graphite, and $\text{SiO}_2$).

Ion

A charged particle formed when an atom or group of atoms loses or gains electrons.

Cation

A positively charged ion formed when an atom loses one or more electrons.

Anion

A negatively charged ion formed when an atom gains one or more electrons.

Atom

The smallest unit of an element that can participate in a chemical reaction.

Ionic Bond

The electrostatic attraction present between oppositely charged ions in a crystal lattice.

Metallic Bond

The electrostatic attraction between the positively charged nuclei of metal atoms and loosely held delocalised valence electrons.

Covalent Bond

A chemical bond formed between two non-metal atoms by the sharing of a pair of electrons.

Double Bond

A covalent bond formed between two atoms involving 4 shared electrons (2 electron pairs).

Triple Bond

A covalent bond formed between two atoms involving 6 shared electrons (3 electron pairs).

Intermolecular Forces

Weak attractive forces existing between separate individual molecules.

Intramolecular Forces

Strong bonding forces (such as covalent bonds) operating within a single molecule to hold its atoms together.

Halogen

A Group 17 element, such as $\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, or $\text{I}_2$.

Halide

A binary compound formed when a halogen reacts with another element (e.g., chlorides, bromides).

Electronegativity

The inherent ability of an atom in a covalent bond to attract the shared pair of electrons toward itself.

Most Electronegative Elements

The five most electronegative elements, led by Fluorine ($\text{F}$), Oxygen ($\text{O}$), Nitrogen ($\text{N}$), Chlorine ($\text{Cl}$), and Sulfur ($\text{S}$).

Polar Covalent Bond

A bond characterized by the uneven sharing of electrons between two atoms due to a difference in electronegativity.

Dipole

A separation of electric charge across a bond or molecule, producing positive and negative poles.

Polar Molecule

A molecule possessing an overall unequal distribution of electric charge due to polar bonds whose individual dipoles do not cancel out.

Non-polar Molecule

A molecule containing no polar bonds, or one where individual bond dipoles cancel out symmetrically so there is no net dipole moment.

Hydrocarbon

An organic compound consisting exclusively of hydrogen and carbon atoms.

Solute and Solvent

A solute is a substance that dissolves in a medium; a solvent is the dissolving liquid medium (typically water or an organic liquid).

Solubility Principle

The rule governing dissolution: polar solutes dissolve in polar solvents (e.g., $\text{NaCl}$ in water), whereas non-polar solutes dissolve in non-polar solvents (e.g., non-polar species in cyclohexane).

Electrical Conductivity

The ability of a substance to conduct electric current through the movement of free-moving charged particles, either delocalised electrons or mobile ions.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model used to predict 3D molecular shapes based on minimizing electrostatic repulsion between regions of electron density surrounding a central atom.

Regions of Electron Density

Regions of negative charge around a central atom consisting of bonding pairs, non-bonding (lone) pairs, double bonds, or triple bonds.

Melting Point

The temperature at which a substance transitions from solid to liquid ($\text{S} \rightarrow \text{L}$); its magnitude reflects the strength of attractive forces binding the particles together.

Boiling Point

The temperature at which a substance transitions from liquid to gas ($\text{L} \rightarrow \text{G}$); its magnitude reflects the strength of attractive forces between particles.

Diamond (C)

An allotrope of carbon forming a 3D covalent network lattice where each carbon atom is covalently bonded in a rigid tetrahedral arrangement.

Graphite (C)

An allotrope of carbon forming 2D covalent network layers with delocalised valence electrons between layers, enabling it to conduct electricity.

Silicon Dioxide ($\text{SiO}_2$)

A giant covalent network solid (silica) comprising silicon and oxygen atoms linked in a continuous 3D covalent structure.

Periodic Bonding Trend

The systematic shift in bonding type across a period from left to right, transitioning from ionic structures ($\text{NaCl}$), to metallic, to covalent molecular compounds ($\text{PCl}_3$, $\text{Cl}_2$).

Linear Molecular Geometry

A molecular shape characterized by atoms aligned in a straight line with a bond angle of $180^\circ$ (e.g., $\text{CO}_2$, $\text{HCl}$).

Trigonal Planar Molecular Geometry

A molecular geometry with three regions of electron density lying in one plane around a central atom with bond angles of approximately $120^\circ$ (e.g., $\text{COCl}_2$, $\text{H}_2\text{CO}$).

Tetrahedral Molecular Geometry

A 3D molecular geometry featuring four bonding electron regions spaced evenly around a central atom with bond angles of approximately $109^\circ$ (e.g., $\text{CCl}_4$, $\text{CH}_2\text{Cl}_2$).

V-Shaped (Bent) Molecular Geometry

A non-linear molecular shape resulting from lone pairs repelling bonding pairs around a central atom, giving bond angles of approximately $109^\circ$ (e.g., $\text{H}_2\text{O}$, $\text{SF}_2$, $\text{H}_2\text{S}$).

Trigonal Pyramidal Molecular Geometry

A 3D molecular shape with three bonding pairs and one non-bonding lone pair surrounding a central atom, with bond angles of approximately $109^\circ$ (e.g., $\text{NCl}_3$, $\text{PBr}_3$, $\text{PH}_3$).