Level 3 Chemistry 3.4 Thermochemistry and Particle Properties

Vocabulary flashcards covering atomic trends, electron configurations, molecular geometry, intermolecular forces, thermochemistry, and enthalpy principles based on the Level 3 Chemistry 3.4 test paper.

Electronegativity

A measure of the ability of an atom in a covalent bond to attract bonding electrons toward itself.

First Ionisation Energy

The minimum energy required to remove one mole of electrons from one mole of gaseous atoms in their ground state.

First Ionisation Energy Reaction Equation for Bromine

The chemical equation representing the first ionisation energy of bromine: $\text{Br}(g) \rightarrow \text{Br}^+(g) + e^-$

Electron Configuration of Bromide Ion ($\text{Br}^-$)

The complete electron configuration using $s, p, d$ notation: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2 4p^6$

Electron Configuration of Iron(II) Ion ($\text{Fe}^{2+}$)

The complete electron configuration using $s, p, d$ notation: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^6$

Electron Configuration of Calcium (Ca)

The complete electron configuration using $s, p, d$ notation: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2$

Atomic Radius vs. Ionic Radius ($\text{Br}$ vs. $\text{Br}^-$)

The radius of $\text{Br}$ ($114\,\text{pm}$) is smaller than $\text{Br}^-$ because adding an extra electron increases electron-electron repulsion within the valence shell, expanding the electron cloud while nuclear charge remains constant.

Group 17 Radius Trend

Atomic radius increases down Group 17 ($\text{F} = 71\,\text{pm}$, $\text{Cl} = 99\,\text{pm}$, $\text{Br} = 114\,\text{pm}$) due to an increasing number of occupied electron shells, which increases inner shell shielding and places outer electrons further from the nucleus.

Group 17 Electronegativity Trend

Electronegativity decreases down Group 17 ($\text{F} = 3.98$, $\text{Cl} = 3.16$, $\text{Br} = 2.96$) because the increased atomic radius and shielding reduce the effective electrostatic attraction between the nucleus and bonding electrons.

Xenon Tetrafluoride ($\text{XeF}_4$) Geometry

Square planar geometry, formed by 6 electron groups (4 bonding pairs and 2 lone pairs) around the central xenon atom to minimize electron pair repulsion.

Sulfur Pentafluoride Cation ($\text{SF}_5^+$) Geometry

Trigonal bipyramidal geometry, formed by 5 electron groups (5 bonding pairs and 0 lone pairs) around the central sulfur ion.

Lewis Structure of Bromine Pentafluoride ($\text{BrF}_5$)

A Lewis structure showing a central bromine atom bonded to 5 fluorine atoms with 1 lone pair on the bromine atom, resulting in a square pyramidal shape with bond angles of approximately $90^\circ$.

Spontaneous Reaction

A chemical reaction or physical process that occurs naturally without continuous external energy input, characterized by an overall increase in total entropy ($\Delta S_{\text{total}} > 0$).

Standard Enthalpy of Formation ($\Delta_f H^\circ$)

The enthalpy change when one mole of a substance is formed from its constituent elements in their standard states under standard conditions.

Enthalpy of Formation Reaction Equation for $\text{NO}_2(g)$

The chemical equation for the formation of one mole of nitrogen dioxide gas from elements in standard states: $\frac{1}{2}\text{N}_2(g) + \text{O}_2(g) \rightarrow \text{NO}_2(g)$

Enthalpy of Fusion ($\Delta_{\text{fus}} H$)

The heat energy required to convert one mole of a substance from solid to liquid state at its melting point.

Enthalpy of Vaporisation ($\Delta_{\text{vap}} H$)

The heat energy required to convert one mole of a substance from liquid to gas state at its boiling point.

Polystyrene Cup Calorimeter

An experimental apparatus comprising a polystyrene cup, lid, stirrer, and thermometer used to measure energy changes ($q = mc\Delta T$) during reactions in solution.

Specific Heat Capacity ($c$)

The amount of heat energy required to raise the temperature of one gram of a substance by one degree Celsius (e.g., $c(\text{solution}) = 4.18\,\text{J\,g}^{-1}{^\circ\text{C}}^{-1}$).

Temporary Dipole-Dipole Forces (Dispersion Forces)

Weak intermolecular attractive forces created by instantaneous, temporary fluctuations in electron density that induce temporary dipoles in adjacent atoms or molecules.

Permanent Dipole-Dipole Forces

Intermolecular attractive forces occurring between polar molecules due to permanent positive and negative partial charges across molecular bonds.

Hydrogen Bonding

A strong form of dipole-dipole attraction occurring when a hydrogen atom covalently bonded to a highly electronegative atom ($\text{N}$, $\text{O}$, or $\text{F}$) is attracted to a lone pair on a neighboring electronegative atom.

Intermolecular Forces in Methanol ($\text{CH}_3\text{OH}$)

Temporary dipole-dipole forces, permanent dipole-dipole forces, and hydrogen bonding.

Intermolecular Forces in Chlorine ($\text{Cl}_2$)

Temporary dipole-dipole forces (dispersion forces) only, as $\text{Cl}_2$ is a non-polar diatomic molecule.

Intermolecular Forces in Dichloromethane ($\text{CH}_2\text{Cl}_2$)

Temporary dipole-dipole forces and permanent dipole-dipole forces.