Physical chemistry

Explanation

• Atomic models changed when new experimental evidence could not be explained by an earlier model.
• The accepted model has a very small nucleus containing protons and neutrons, with electrons occupying the surrounding space.
• A proton has relative charge $+1$ and relative mass $1$; a neutron has charge $0$ and relative mass $1$; an electron has charge $-1$ and relative mass about $1/1836$.
• A neutral atom has equal numbers of protons and electrons.
• An ion forms only by losing or gaining electrons, so its nucleus and element identity remain unchanged.

Explanation

• Atomic number $Z$ is the number of protons; mass number $A$ is the total number of protons and neutrons, so neutrons $=A-Z$.
• Isotopes are atoms of one element with equal proton numbers but different neutron numbers.
• In time-of-flight mass spectrometry, gaseous particles are ionised, accelerated so all ions have the same kinetic energy, separated by flight time and detected.
• For mononuclear $1+$ ions, the mass-to-charge value identifies relative isotopic mass and signal size gives abundance.
• The weighted mean $A_r$ can identify an element, while a molecular-ion peak can give $M_r$.

Explanation

• For atoms up to $Z=36$, electrons fill $1s$, $2s$, $2p$, $3s$, $3p$, $4s$, $3d$, then $4p$. Chromium is $\mathrm{[Ar]3d^5\,4s^1}$ and copper is $\mathrm{[Ar]3d^{10}\,4s^1}$; transition-metal ions lose $4s$ electrons before $3d$.
• First ionisation energy removes one electron from each atom in one mole of gaseous atoms.
• Across Period 3, increasing nuclear charge generally raises first ionisation energy, with sub-shell and electron-pairing anomalies.
• Down Group 2, increasing distance and shielding lower it.
• A large successive-ionisation jump shows removal from a shell closer to the nucleus and provides evidence for shell structure.

Explanation

• Relative atomic mass, $A_r$, is the weighted mean mass of an atom of an element relative to one twelfth of the mass of a carbon-12 atom. Relative molecular mass, $M_r$, is the mean mass of a molecule on the same scale.
• Both are ratios and therefore have no units.
• Calculate $M_r$ by multiplying each element's $A_r$ by the number of its atoms and summing.
• For an ionic substance, the equivalent quantity is relative formula mass because the solid contains a lattice rather than separate molecules.
• Brackets and waters of crystallisation must be counted fully.

Explanation

• One mole contains the Avogadro constant, $N_A$, of specified entities, such as atoms, molecules, ions, electrons or formula units.
• Amount of substance links measurable quantities to particle number: $n=m/M$, $N=nN_A$, and for a solution $n=cV$ when $V$ is in $\mathrm{dm^3}$.
• A secure calculation first converts the given quantity to moles, applies the chemical formula or equation ratio, then converts to the requested quantity.
• The named entity matters: one mole of $\mathrm{MgCl_2}$ formula units contains two moles of chloride ions.

Explanation

• The ideal gas equation is $pV=nRT$.
• The gas constant $R$ has the value $8.31$ in units of $\mathrm{J\,K^{-1}\,mol^{-1}}$, so pressure must be in $\mathrm{Pa}$, volume in $\mathrm{m^3}$, temperature in $\mathrm{K}$ and amount in $\mathrm{mol}$.
• Convert $\mathrm{kPa}$ to $\mathrm{Pa}$ by multiplying by $10^3$, $\mathrm{cm^3}$ to $\mathrm{m^3}$ by multiplying by $10^{-6}$, and $\mathrm{dm^3}$ to $\mathrm{m^3}$ by multiplying by $10^{-3}$.
• Convert Celsius using $T/\mathrm{K}=\theta/^{\circ}\mathrm{C}+273$.
• Rearranging the equation before substitution reduces unit and algebra errors in multi-step calculations.

Explanation

• An empirical formula is the simplest whole-number ratio of atoms of each element; a molecular formula gives the actual numbers of atoms in one molecule.
• From masses or percentages, divide each value by the element's $A_r$ to obtain relative moles, divide every result by the smallest, then multiply all ratios by the same small integer if necessary.
• A ratio near $1.5$ requires doubling, not rounding.
• To obtain a molecular formula, calculate the empirical formula mass and find the integer $M_r/(\text{empirical formula mass})$.
• Combustion data must be converted from product moles to constituent atom moles.

Explanation

• Balanced full and ionic equations conserve every atom and the total charge; only coefficients may be changed.
• Quantitative work follows a mole route: convert the known quantity to moles, apply the balanced-equation ratio, identify any limiting reactant, then convert to the requested mass, volume or concentration.
• Percentage yield is $100\times\text{actual yield}/\text{theoretical yield}$.
• Atom economy is $100\times M_r(\text{desired product})/\sum M_r(\text{reactants})$, including stoichiometric coefficients.
• A high atom economy reduces waste and raw-material use, with economic, environmental and ethical benefits, but it does not guarantee a high experimental yield.

Explanation

• Ionic bonding is the strong electrostatic attraction between oppositely charged ions throughout a giant lattice. Electron transfer forms the ions but is not itself the bond.
• Simple-ion charges can be predicted from Periodic Table groups, and the ionic formula uses the smallest whole-number ratio giving zero overall charge.
• Compound ions remain intact; for example, two $\mathrm{Al^{3+}}$ ions balance three $\mathrm{SO_4^{2-}}$ ions to give $\mathrm{Al_2(SO_4)_3}$.
• Brackets are required when more than one compound ion is present.
• Explanations of high melting point must refer to strong attractions acting in many directions through the lattice.

Explanation

• A covalent bond is a shared pair of electrons. Double and triple bonds contain two and three shared pairs.
• In an ordinary covalent bond, each bonded atom contributes one electron to the shared pair. In a co-ordinate, or dative covalent, bond, both electrons in the shared pair are supplied by one atom.
• The displayed-formula arrow starts at a lone pair on the electron-pair donor and points towards the electron-deficient acceptor.
• Once formed, the dative bond behaves like any other covalent bond.
• Dot-and-cross diagrams must show outer-shell electrons, charges and the origin of the bonding pair accurately.

Explanation

• Metallic bonding is the strong electrostatic attraction between a lattice of positive metal ions and delocalised electrons.
• The electrons are not attached to a particular ion and can move through the structure, so a metal conducts electricity when solid or liquid.
• Metallic attraction is non-directional, allowing layers of ions to slide while remaining attracted to the delocalised electrons; this explains malleability and ductility.
• Bond strength generally increases with greater ionic charge, more delocalised electrons and smaller ionic radius because attraction between the ions and electron density is stronger.
• Metals do not consist of separate molecules.

Explanation

• The four crystal types are ionic, metallic, macromolecular and molecular. Sodium chloride, magnesium, diamond or graphite, and iodine or ice are respective examples.
• Ionic and metallic solids contain giant lattices; macromolecular substances have many covalent bonds; molecular solids contain discrete molecules held by intermolecular forces. Melting a molecular solid overcomes intermolecular attractions, not covalent bonds inside molecules.
• Ionic substances conduct only when molten or dissolved because ions can move. Metals and graphite conduct through delocalised electrons.
• Structure is inferred from combined melting-point, solubility and conductivity evidence.
• A requested diagram must show the specified number and type of particles.

Explanation

• Electron charge clouds repel and arrange as far apart as possible around a central atom. Count each bonding region, including a multiple bond as one region, and each lone pair.
• Repulsion decreases in the order lone pair–lone pair $>$ lone pair–bond pair $>$ bond pair–bond pair, so lone pairs compress adjacent bond angles.
• Base geometries are linear $180^{\circ}$, trigonal planar $120^{\circ}$, tetrahedral $109.5^{\circ}$, trigonal bipyramidal $90^{\circ}$ and $120^{\circ}$, and octahedral $90^{\circ}$.
• In a trigonal bipyramid, lone pairs occupy equatorial positions first.
• Shape names describe atom positions, not lone pairs.

Explanation

• Electronegativity is the power of an atom to attract the bonding pair of electrons in a covalent bond.
• A difference in electronegativity produces an unequal electron distribution: the more electronegative atom is $\delta^-$ and the other is $\delta^+$.
• A bond dipole has both magnitude and direction.
• To decide whether a whole molecule has a permanent dipole, identify every polar bond and use the molecular shape to combine their dipoles.
• Equal bond dipoles arranged symmetrically cancel, so a molecule can contain polar bonds but have no overall permanent dipole.

Explanation

• Induced dipole–dipole forces act between all atoms and molecules. Their strength generally increases with electron number and with molecular surface contact.
• Permanent dipole–dipole attractions also act between polar molecules.
• Hydrogen bonding requires hydrogen covalently bonded to nitrogen, oxygen or fluorine and a lone pair on nitrogen, oxygen or fluorine in another molecule.
• Boiling-point comparisons must identify the forces present and compare their overall strength, not simply molecular mass.
• In ice, hydrogen bonds hold water molecules in an open arrangement, making ice less dense than liquid water; the bonds are not described as inherently stronger.

Explanation

• Enthalpy change, $\Delta H$, is the heat-energy change at constant pressure.
• An exothermic process transfers heat to the surroundings and has $\Delta H<0$; an endothermic process absorbs heat and has $\Delta H>0$.
• A standard enthalpy change uses substances in their standard states at $100\,\mathrm{kPa}$ and a stated temperature.
• Standard enthalpy of combustion is the enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions.
• Standard enthalpy of formation is the change when one mole of a compound forms from its constituent elements in their standard states.

Explanation

• Calorimetry estimates heat transfer using $q=mc\Delta T$, where $m$ is the mass undergoing the temperature change, $c$ is its specific heat capacity and $\Delta T$ is the measured temperature change.
• In solution calorimetry, total solution volume and density commonly estimate mass, while the limiting reagent supplies the reacting amount.
• Convert heat to a molar enthalpy using $\Delta H=-q/n$ when a temperature rise shows that the reaction released heat.
• Required practical 2 measures an enthalpy change, for example by dissolution, neutralisation, displacement or alcohol combustion.
• Heat loss, incomplete combustion, evaporation and heating the apparatus create systematic differences that must be evaluated.