Molecules: Shapes & Forces

## Count electron domains around a central atom - VSEPR models the repulsion between regions of valence electron density around a central atom. Their arrangement minimises repulsion; the model predicts geometry rather than electron trajectories. - Each single, double or triple bond counts as one bonding region. Each central-atom lone pair adds one further region, so domain count is different from the number of shared electron pairs. - Start with a valid Lewis structure and account for the overall charge. The total number of valence electrons increases for a negative ion and decreases for a positive ion. ## Distinguish electron arrangement from molecular shape - An electron-domain arrangement includes both bonding regions and lone pairs. A molecular shape describes the positions of atoms and does not name a lone pair as if it were another atom. - Four regions have a tetrahedral parent arrangement. Four bonded atoms give a tetrahedral molecule; three bonded atoms and a lone pair give a trigonal pyramid; two bonded atoms and two lone pairs give a bent molecule. - Lone pairs occupy space and affect repulsion even though they do not appear in the molecular shape name. A flat Lewis drawing alone is not a reliable picture of the three-dimensional molecule. ## Explain linear and trigonal planar molecules - Two bonding regions and no central lone pairs give a linear molecule with a 180° bond angle, as in CO₂. Its two double bonds remain two regions, not four. - Three bonding regions and no central lone pairs give trigonal planar geometry with ideal 120° angles, as in BF₃. The three bonds lie in one plane around boron. - These predictions concern the local geometry around the named atom. Electron-deficient boron can still be treated by VSEPR; a complete octet is not a prerequisite for a shape prediction. ## Explain tetrahedral and pyramidal shapes - Methane has four C–H bonding regions and no central lone pairs. Equivalent regions point towards the corners of a tetrahedron, with angles of approximately 109.5°. - Ammonia has three N–H bonding regions and one central lone pair. Its molecular shape is trigonal pyramidal, with an H–N–H angle of about 107°. - In the AQA repulsion model, lone pair–lone pair repulsion exceeds lone pair–bond pair repulsion, which exceeds bond pair–bond pair repulsion. A lone pair can compress angles between bonds. ## Explain water and the limits of angle shortcuts - Water has two O–H bonding regions and two central lone pairs. Its molecular shape is bent, and the H–O–H angle is approximately 104.5°. - Compare water and ammonia with the 109.5° tetrahedral parent arrangement. The differing repulsions help explain their compressed bond angles. - Do not subtract a fixed 2.5° for every lone pair in every species. Actual angles also depend on the bonded atoms and bonding; 107° and 104.5° are familiar examples rather than universal constants. ## Include charge when predicting ion geometry - NH₄⁺ contains eight valence electrons: 5 + 4 − 1. Four N–H bonds use four pairs, leaving no nitrogen lone pair, so its molecular geometry is tetrahedral. - H₃O⁺ also contains eight valence electrons: 6 + 3 − 1. Three O–H bonds leave one oxygen lone pair, giving a trigonal pyramidal molecular shape. - The coordinate origin of a bond does not require a permanently different bond direction or strength after an ion forms. Use the final electron arrangement for VSEPR. ## Interpret five electron domains - Five bonding regions and no central lone pairs give trigonal bipyramidal geometry, illustrated by molecular PF₅. Three equatorial bonds lie in a plane; two axial bonds lie on an axis perpendicular to that plane. - Ideal equatorial–equatorial angles are 120°, axial–equatorial angles are 90°, and the two opposite axial bonds form 180°. - Use molecular species under stated conditions. Gas-phase PCl₅ can be treated as a trigonal bipyramidal molecule, while its solid-state structure requires a different description. ## Place lone pairs in a five-domain arrangement - In a trigonal bipyramidal parent arrangement, an equatorial position has fewer 90° interactions than an axial position. Lone pairs preferentially occupy equatorial positions in the simple model. - AX₄E therefore gives a seesaw molecular shape, AX₃E₂ gives a T-shape, and AX₂E₃ gives a linear molecule. Here X is a bonded atom and E is a central lone pair. - Non-equivalent repulsions distort some angles from their ideal parent values. Do not describe every five-domain molecule as trigonal bipyramidal simply because its domain count is five. ## Interpret six electron domains - Six bonding regions and no central lone pairs give octahedral molecular geometry, illustrated by SF₆. Four bonds lie in an equatorial square and two lie on the perpendicular axis. - Adjacent bonds form ideal 90° angles and opposite bonds form 180°. A two-dimensional perspective view does not make the actual molecule planar. - With one central lone pair, AX₅E is square pyramidal. With two opposite central lone pairs, AX₄E₂ is square planar; these are molecular shapes derived from the octahedral parent arrangement. ## Read perspective and compare unfamiliar species - An ordinary line represents a bond in the page, a solid wedge comes towards the viewer, and a hashed wedge goes away. These symbols encode perspective rather than literal triangular bonds. - Determine domains from electron counting before interpreting the drawing. Rotate a model mentally: a rotated tetrahedron is still a tetrahedron. - For an unfamiliar species, state the central bonding and lone-pair counts, the parent domain arrangement, the molecular shape and appropriate angles. Do not infer a shape from a molecular formula without checking its bonding. ## Define electronegativity and explain trends - Electronegativity is an atom’s ability to attract the shared pair of electrons in a covalent bond. It is not the energy change when a gaseous atom gains an electron. - Across a typical main-group period, increasing nuclear charge with broadly similar inner-shell shielding strengthens attraction for bonding electrons. Down a group, increased distance and shielding usually outweigh increased nuclear charge. - Pauling values are dimensionless and scale-dependent. Fluorine is the most electronegative element on the usual scale; use supplied values rather than assigning unsupported values to all noble gases. ## Represent a polar covalent bond - Different electronegativities produce unequal sharing of electron density. The more electronegative atom carries δ− and its partner δ+ in the bond model. - Partial charges do not mean that a complete electron has transferred or that the molecule consists of separate integer ions. A covalent bond can remain covalent while being polar. - A chemical bond-dipole arrow points towards the negative end and has a cross at the positive end. Always identify the convention when comparing arrows from different scientific contexts. ## Decide whether bond dipoles cancel - A polar bond does not automatically make a polar molecule. The vector sum of the bond dipoles depends on molecular geometry and the identity of the surrounding atoms. - In linear CO₂, equal opposite C=O dipoles cancel. In trigonal planar BF₃ and tetrahedral CCl₄, the symmetric arrangements of equivalent bonds also cancel. - Water’s bent geometry prevents the two O–H dipoles from cancelling. Replacing atoms around a tetrahedral centre can also remove cancellation, as in CH₃Cl. ## Explain permanent dipole interactions - Polar molecules have an uneven charge distribution and can orient with a δ+ region near a neighbour’s δ− region. Their electrostatic attraction is called a permanent dipole–dipole interaction. - At finite temperature, molecules move and rotate, so the interaction is not a rigid head-to-tail chain of stationary particles. Orientation and separation affect the interaction. - Polar molecules also experience dispersion forces. Identifying a permanent dipole adds an interaction; it does not remove the others. ## Explain instantaneous and induced dipoles - Electron-density fluctuations can create an instantaneous dipole, which changes the electron distribution of a neighbouring particle and induces a correlated dipole. - The resulting dispersion attraction exists between all atoms and molecules, including nonpolar particles. The dipoles fluctuate; this is not the transfer of an electron to make permanent ions. - Larger, more easily distorted electron clouds are generally more polarisable. Electron number can support a comparison within related species, but is not a complete universal formula for attraction strength. ## Compare polarisability and molecular contact - Within a related series, larger electron clouds are often more polarisable, giving stronger dispersion interactions and higher boiling temperatures under the same pressure. - For isomers, molecular shape affects close contact. A less compact shape can allow more extensive contact than a compact, branched shape and often has a higher boiling point. - Keep variables and conditions explicit. Molecular mass alone does not determine boiling point, and melting points also depend strongly on crystal packing. ## Identify hydrogen-bond donors and acceptors - In the AQA model, hydrogen bonding involves H covalently bonded to N, O or F and an available lone pair on an N, O or F acceptor nearby. - Distinguish the covalent donor bond from the intermolecular H···acceptor contact. An oxygen atom without an O–H bond can accept a hydrogen bond from water even if its own pure liquid lacks a donor. - Check availability of the lone pair, not just the presence of a nitrogen symbol. For example, the nitrogen in NH₄⁺ has no lone pair available to accept in this simple model. ## Count water’s hydrogen-bond contacts - A water molecule can donate through two O–H groups and accept through its two oxygen lone pairs. In an ideal ice-like network, it can therefore have four hydrogen-bonded neighbours. - The local fragment illustrates two donated and two accepted contacts around the central molecule. Its O–H covalent bonds remain inside the water molecules. - When counting network links, each intermolecular contact belongs to two neighbouring waters. Multiplying molecule count by four and stopping there counts every link twice. ## Compare the forces present in a pure liquid - First establish molecular geometry and polarity, then identify donors and acceptors. All molecules have dispersion; polar molecules add permanent dipole interactions; suitable donor–acceptor pairs permit hydrogen bonding. - Ethanol can hydrogen-bond between its own molecules. Propanone accepts hydrogen bonds from water but lacks the required donor for the same self-association pattern in the AQA model. - Avoid an absolute strength ranking across arbitrary substances. A large nonpolar molecule can have stronger total intermolecular attraction than a small polar molecule; compare the entire molecular system. ## Explain boiling using intact molecules - Boiling separates molecules against intermolecular attractions. It does not normally break the covalent bonds within every molecule or convert the sample into isolated atoms. - At the same external pressure, stronger overall intermolecular attraction generally requires a higher temperature for boiling. Always state the pressure when comparing boiling temperatures. - Butane and propanone each have 34 electrons, yet their normal boiling points differ markedly. Propanone’s permanent dipole contributes to stronger overall cohesion; both liquids still have dispersion interactions. ## Explain water’s anomalous boiling behaviour - H₂S, H₂Se and H₂Te show a broad increase in normal boiling point as electron clouds become more polarisable. These bent polar molecules have permanent-dipole interactions as well as dispersion. - Water has a much higher normal boiling point than a simple extrapolation from the heavier hydrides would suggest. Extensive hydrogen bonding provides an additional important contribution. - Do not attribute the heavier hydrides exclusively to dispersion or calculate an exact hypothetical water boiling point without a justified model. Use the supplied observations to support a qualitative explanation. ## Relate ice structure to density - Near the freezing point at ordinary pressure, ice has a more open hydrogen-bonded arrangement than liquid water. Four-neighbour local coordination helps form this extended structure. - On melting, the network partly reorganises and molecules can pack more closely on average. The same mass then occupies less volume, so liquid water is denser under these conditions. - Ice floating does not imply that every solid is less dense than its liquid or that water molecules themselves become larger. Retain the temperature and pressure conditions in the comparison. ## Explain surface tension without a rigid-skin model - Surface tension can be expressed as force per unit length or surface free energy per unit area. Creating more liquid surface requires energy because surface molecules have a different neighbour environment. - Water’s cohesive hydrogen-bond interactions contribute to its substantial surface tension. Inward attraction at a surface is a useful qualitative model, not a literal solid membrane or a guarantee that an object will be supported. - A force calculation must identify how many contact lines contribute and whether a contact-angle factor applies. Use the supplied simplified model rather than inventing extra factors. ## Evaluate what physical-property evidence establishes - Melting points depend on both interactions and crystal packing. A higher melting point does not by itself prove a specific intermolecular mechanism or a particular molecular geometry. - Use a valid structure, polarity analysis and donor–acceptor assessment alongside controlled physical-property data. Keep observations separate from the explanation proposed for them. - For an unfamiliar substance, identify what is known, what is inferred and which extra evidence could distinguish alternatives. A single boiling value cannot uniquely determine its molecular formula.