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the nucleus.","correct":false,"explanation":"Inner-shell electrons are held strongly by the nucleus and are generally not involved in bonding. The outer-shell electrons, called valence electrons, participate in chemical reactions."},{"text":"Protons - positively charged nuclear particles.","correct":false,"explanation":"Protons are positively charged particles in the nucleus and cannot be transferred or shared between atoms in ordinary chemical bonding. Valence electrons are the bonding electrons."},{"text":"Neutrons - uncharged nuclear particles.","correct":false,"explanation":"Neutrons are uncharged particles in the nucleus and do not determine how atoms bond. Chemical bonding mainly involves the outer-shell, or valence, electrons."}]},{"id":"dda4c8e8-685c-4ff5-8878-4cc87f214a31","cardType":"cloze","points":null,"front":"Metals tend to lose electrons to form positive [ ]; non-metals tend to gain electrons to form negative [ ].","back":"Metals tend to lose electrons to form positive cations; non-metals tend to gain electrons to form negative anions.","mcqOptions":[{"text":"Metals tend to lose electrons to form positive cations; non-metals tend to gain electrons to form negative anions.","correct":true,"explanation":"Metals lose electrons and become positively charged cations, whereas non-metals gain electrons and become negatively charged anions. The charges result from the electron transfer."},{"text":"Metals tend to lose electrons to form positive anions; non-metals tend to gain electrons to form negative cations.","correct":false,"explanation":"Anions are negatively charged ions formed when atoms gain electrons, so they cannot be the products of electron loss by metals. Metals form cations, while non-metals form anions."},{"text":"Metals tend to lose electrons to form positive atoms; non-metals tend to gain electrons to form negative atoms.","correct":false,"explanation":"Electron loss produces positive ions, not neutral atoms, and electron gain produces negative ions. The correct terms are cations for metals and anions for non-metals."},{"text":"Metals tend to lose electrons to form positive anions; non-metals tend to gain electrons to form negative ions.","correct":false,"explanation":"Metals that lose electrons become positively charged, whereas non-metals that gain electrons become negatively charged. Calling these ions positive anions and negative cations reverses both definitions."}]},{"id":"47a1314e-f33c-4e4c-bc6e-a5e4c312c1ec","cardType":"cloze","points":null,"front":"Ionic bonding is the [ ] between [ ] in a giant lattice.","back":"Ionic bonding is the strong electrostatic attraction between oppositely charged ions in a giant lattice.","mcqOptions":[{"text":"Ionic bonding is the strong electrostatic attraction between oppositely charged ions in a giant lattice.","correct":true,"explanation":"Ionic bonding is the strong electrostatic attraction between oppositely charged ions. This attraction acts throughout the giant lattice and holds the ions together."},{"text":"Ionic bonding is the sharing of electrons between oppositely charged ions in a giant lattice.","correct":false,"explanation":"Ionic bonding is an electrostatic attraction, not a covalent sharing of electrons. The ions involved have opposite charges, rather than the same charge."},{"text":"Ionic bonding is the strong electrostatic attraction between similarly charged ions in a giant lattice.","correct":false,"explanation":"Like charges repel, so similarly charged ions would not produce the attractive force that defines ionic bonding. Ionic bonding involves strong attraction between oppositely charged ions."},{"text":"Ionic bonding is the weak intermolecular force between neutral atoms in a giant lattice.","correct":false,"explanation":"Ionic bonding is not a weak force between neutral atoms. It is a strong electrostatic attraction caused by the opposite charges on ions in the giant lattice."}]},{"id":"d5914b6f-900b-415e-87f9-c3613a43447f","cardType":"basic","points":2,"front":"In MgCl₂, why are two Cl^- ions needed for each Mg²⁺ ion?","back":"Two Cl^- ions balance one Mg²⁺ ion.","mcqOptions":[{"text":"Two Cl^- ions balance one Mg²⁺ ion.","correct":true,"explanation":"Each Mg²⁺ ion has a charge of +2, while each Cl^- ion has a charge of -1. Two chloride ions therefore give a total charge of -2 and balance one magnesium ion."},{"text":"One Cl^- ion balances one Mg²⁺ ion.","correct":false,"explanation":"One Cl^- ion contributes only -1, which does not cancel the +2 charge of Mg²⁺. Two chloride ions are required to produce a total charge of -2."},{"text":"Two Mg²⁺ ions balance one Cl^- ion.","correct":false,"explanation":"Two Mg²⁺ ions would give a total charge of +4, while one Cl^- gives only -1. Charge balance requires two chloride ions for each magnesium ion."},{"text":"Two Cl^- ions balance two Mg²⁺ ions.","correct":false,"explanation":"The charges do not balance in this ratio: two Cl^- ions give -2, while two Mg²⁺ ions give +4. The correct ratio is one magnesium ion to two chloride ions."}]},{"id":"2c5551ba-2ee8-4a2b-82cb-3fcdcce60c94","cardType":"cloze","points":null,"front":"A covalent bond is a [ ] with both nuclei attracted to that pair.","back":"A covalent bond is a shared pair of electrons with both nuclei attracted to that pair.","mcqOptions":[{"text":"shared pair of electrons","correct":true,"explanation":"A covalent bond consists of a shared pair of electrons. Both atomic nuclei are attracted to this negative electron pair, holding the atoms together."},{"text":"shared pair of protons","correct":false,"explanation":"A shared pair of protons cannot form a covalent bond because protons remain in atomic nuclei and repel one another. Covalent bonding involves shared electrons."},{"text":"shared electron","correct":false,"explanation":"A covalent bond is not a single shared electron. It is a shared pair of electrons attracted by both nuclei, giving the bond its stability."},{"text":"lone pair of electrons","correct":false,"explanation":"A lone pair is located on one atom and is not necessarily shared between two nuclei. A covalent bond specifically contains a shared pair of electrons."}]},{"id":"25f69663-7eff-482f-8bc1-05a212ec103b","cardType":"basic","points":1,"front":"In a coordinate bond, where do the two shared electrons come from?","back":"The same atom donates both electrons.","mcqOptions":[{"text":"The same atom donates both electrons.","correct":true,"explanation":"In a coordinate bond, one atom supplies both electrons in the shared pair, usually from a lone pair. The other atom provides an empty orbital or space for the pair."},{"text":"Each atom donates one electron.","correct":false,"explanation":"If each atom donates one electron, the bond is an ordinary covalent bond rather than a coordinate bond. A coordinate bond has both electrons from one atom."},{"text":"The two nuclei donate one electron each.","correct":false,"explanation":"The electrons in a coordinate bond come from one bonding atom, not from the nucleus. Nuclei contain protons and neutrons, whereas bonding involves electrons."},{"text":"An external source donates both electrons.","correct":false,"explanation":"A coordinate bond is formed when one atom provides both electrons, not when the electrons originate outside the atoms. The donated pair usually comes from a lone pair."}]},{"id":"198ed59a-f80a-4575-97c9-5c8ea24077bd","cardType":"cloze","points":null,"front":"After NH₄⁺ forms, all four N-H bonds are [ ]; coordinate describes [ ].","back":"After NH₄⁺ forms, all four N-H bonds are equivalent; coordinate describes how the bond formed.","mcqOptions":[{"text":"After NH₄⁺ forms, all four N-H bonds are equivalent; coordinate describes how the bond formed.","correct":true,"explanation":"Once NH₄⁺ has formed, the four N-H bonds are equivalent. The term coordinate refers only to the way one bond initially formed, not to a permanent bond type."},{"text":"After NH₄⁺ forms, all four N-H bonds are different lengths; coordinate describes how the bond formed.","correct":false,"explanation":"The four N-H bonds in NH₄⁺ are equivalent after formation, because the electron arrangement becomes symmetrical. Coordinate describes how the bond formed, not its length."},{"text":"After NH₄⁺ forms, all four N-H bonds are different strengths; coordinate describes how the bond formed.","correct":false,"explanation":"Coordinate bonding does not make one N-H bond permanently stronger or weaker in NH₄⁺. All four bonds are equivalent, and coordinate describes their formation route."},{"text":"After NH₄⁺ forms, all four N-H bonds are equivalent; coordinate describes the permanent bond type.","correct":false,"explanation":"The bonds do not remain distinguishable by their electron origins after NH₄⁺ forms. All four are equivalent; coordinate refers to how one bond was formed initially."}]},{"id":"fedf4038-3c51-4c08-883d-6c9646a04e20","cardType":"basic","points":1,"front":"In a bond between atoms with different electronegativities, which atom becomes δ-?","back":"The more electronegative atom.","mcqOptions":[{"text":"The more electronegative atom.","correct":true,"explanation":"The more electronegative atom attracts the shared bonding electrons more strongly, giving it a partial negative charge, δ-."},{"text":"The less electronegative atom.","correct":false,"explanation":"The less electronegative atom has weaker attraction for the shared electrons and therefore becomes δ+. The more electronegative atom becomes δ-."},{"text":"The larger atom.","correct":false,"explanation":"Electronegativity, not atomic size alone, determines which atom attracts the shared pair more strongly. The more electronegative atom receives the δ- charge."},{"text":"The atom with lower ionisation energy.","correct":false,"explanation":"The atom with the lower ionisation energy is not automatically the one that becomes δ-. Partial negative charge develops on the atom with greater electronegativity."}]},{"id":"c90ae8e3-e066-470d-86e2-75126e92d257","cardType":"cloze","points":null,"front":"In H-Cl, chlorine is more electronegative, so Cl is [ ] and H is [ ].","back":"In H-Cl, chlorine is more electronegative, so Cl is δ- and H is δ+.","mcqOptions":[{"text":"In H-Cl, chlorine is more electronegative, so Cl is δ- and H is δ+.","correct":true,"explanation":"Chlorine attracts the bonding pair more strongly because it is more electronegative. It therefore gains a partial negative charge, leaving hydrogen with a partial positive charge."},{"text":"In H-Cl, chlorine is more electronegative, so Cl is δ+ and H is δ-.","correct":false,"explanation":"This reverses the charge distribution. Chlorine attracts the shared electrons more strongly, so Cl is δ- and the resulting electron deficiency on H makes it δ+."},{"text":"In H-Cl, chlorine is more electronegative, so Cl is δ- and H is δ-.","correct":false,"explanation":"Both atoms cannot gain electron density from the same bonding pair. Chlorine attracts the pair more strongly, making Cl δ- and H δ+."},{"text":"In H-Cl, chlorine is more electronegative, so Cl is - and H is +.","correct":false,"explanation":"Full charges would imply complete electron transfer and ionic bonding. H-Cl is polar covalent, so unequal sharing produces partial charges: Cl is δ- and H is δ+."}]},{"id":"20223e81-20df-4c14-b7c1-ae24953bae3a","cardType":"basic","points":1,"front":"What does a very large electronegativity difference often make bonding?","back":"Mostly ionic.","mcqOptions":[{"text":"Mostly ionic.","correct":true,"explanation":"A very large electronegativity difference strongly favours electron transfer rather than equal sharing, so the resulting attraction between oppositely charged ions makes the bonding mostly ionic."},{"text":"Mostly covalent.","correct":false,"explanation":"Covalent bonding involves sharing electrons and is favoured by a smaller electronegativity difference. A very large difference generally makes the bonding mostly ionic."},{"text":"Mostly non-polar.","correct":false,"explanation":"Non-polar bonding requires little or no electronegativity difference so that electrons are shared nearly equally. A very large difference instead produces mostly ionic bonding."},{"text":"Mostly metallic.","correct":false,"explanation":"Metallic bonding involves delocalised electrons among a lattice of metal ions and is not determined by a large difference between two atoms. Such a difference usually indicates mostly ionic bonding."}]},{"id":"d025d71a-c7c8-4b51-a08f-23ff67fe9665","cardType":"cloze","points":null,"front":"In VSEPR, electron pairs around a central atom [ ] and arrange [ ].","back":"In VSEPR, electron pairs around a central atom repel and arrange as far apart as possible.","mcqOptions":[{"text":"In VSEPR, electron pairs around a central atom repel and arrange as far apart as possible.","correct":true,"explanation":"All electron pairs carry negative charge, so electrostatic repulsion makes them adopt positions that maximise their separation. This arrangement minimises repulsion and determines molecular shape."},{"text":"In VSEPR, electron pairs around a central atom attract and arrange as close together as possible.","correct":false,"explanation":"Electron pairs are negatively charged and therefore repel rather than attract one another. They minimise this repulsion by arranging as far apart as possible around the central atom."},{"text":"In VSEPR, electron pairs around a central atom repel and arrange as close together as possible.","correct":false,"explanation":"Although electron pairs do repel, moving close together would increase that repulsion. VSEPR predicts that they arrange as far apart as possible to reach the lowest-energy geometry."},{"text":"In VSEPR, electron pairs around a central atom attract and arrange as far apart as possible.","correct":false,"explanation":"Electron pairs do not attract one another merely because they surround the same nucleus. Their mutual repulsion causes them to arrange as far apart as possible."}]},{"id":"2ad47374-effc-4a30-8abf-a79ab589c803","cardType":"basic","points":1,"front":"In VSEPR, how does a multiple bond count for shape prediction?","back":"As one region of electron density.","mcqOptions":[{"text":"As one region of electron density.","correct":true,"explanation":"A double or triple bond contains several shared electron pairs but occupies one direction from the central atom. VSEPR therefore treats it as one region of electron density."},{"text":"As two regions of electron density.","correct":false,"explanation":"Counting a multiple bond as two regions confuses the number of shared pairs with spatial directions. All its bonding electrons occupy one direction, so it counts as one region."},{"text":"As three regions of electron density.","correct":false,"explanation":"Only a triple bond contains three shared pairs, but even it does not count as three spatial regions. Any multiple bond counts as one region of electron density."},{"text":"As no region of electron density.","correct":false,"explanation":"A multiple bond still influences molecular shape and cannot be ignored. Because all its shared pairs lie in the same direction, it counts as one region of electron density."}]},{"id":"ca620054-7bab-4821-b38e-89fc887e0248","cardType":"basic","points":2,"front":"What shape and bond angle come from two bonding regions?","back":"

Linear
180°

","mcqOptions":[{"text":"

Linear
180°

","correct":true,"explanation":"Two bonding regions repel to opposite sides of the central atom, maximising their separation. This produces a linear molecular shape with a bond angle of 180°."},{"text":"

Bent
104.5°

","correct":false,"explanation":"A bent shape near 104.5° results when two bonding regions are accompanied by two lone pairs. With only two bonding regions, the arrangement is linear at 180°."},{"text":"

Trigonal planar
120°

","correct":false,"explanation":"Trigonal planar geometry requires three regions around the central atom. Two bonding regions instead move to opposite sides, producing a linear shape and a 180° angle."},{"text":"

Tetrahedral
109.5°

","correct":false,"explanation":"Tetrahedral geometry requires four regions of electron density. With two bonding regions and no stated lone pairs, maximum separation gives a linear shape with a 180° angle."}]},{"id":"44c02116-e2a5-494d-8002-19eb4608f201","cardType":"basic","points":2,"front":"What shape and bond angle come from three bonding regions?","back":"

Trigonal planar
120°

","mcqOptions":[{"text":"

Trigonal planar
120°

","correct":true,"explanation":"Three bonding regions spread equally within one plane to minimise repulsion. Their positions form a trigonal planar shape with equal bond angles of 120°."},{"text":"

Linear
180°

","correct":false,"explanation":"Linear geometry occurs with two regions, not three. Three bonding regions repel into a trigonal planar arrangement, where the equal bond angles are 120°."},{"text":"

Trigonal pyramidal
107°

","correct":false,"explanation":"Trigonal pyramidal geometry requires three bonds plus a lone pair, giving four total regions. Three bonding regions alone form a trigonal planar shape with 120° angles."},{"text":"

Tetrahedral
109.5°

","correct":false,"explanation":"Tetrahedral geometry and 109.5° angles require four regions around the central atom. Three bonding regions alone arrange in a plane with 120° between them."}]},{"id":"4aa6e54b-a7b8-465b-8bc9-4459e94075db","cardType":"basic","points":2,"front":"What shape and bond angle come from four bonding regions?","back":"

Tetrahedral
109.5°

","mcqOptions":[{"text":"

Tetrahedral
109.5°

","correct":true,"explanation":"Four bonding regions minimise repulsion by pointing towards the corners of a tetrahedron. The angle between each pair of bonds is approximately 109.5°."},{"text":"

Trigonal pyramidal
107°

","correct":false,"explanation":"Trigonal pyramidal geometry has three bonding regions and one lone pair, with the lone pair compressing the angles. Four bonding regions alone give tetrahedral geometry at 109.5°."},{"text":"

Square planar
90° and 180°

","correct":false,"explanation":"Square planar geometry usually derives from six electron regions when two opposite positions contain lone pairs. Four bonding regions alone adopt tetrahedral geometry with 109.5° angles."},{"text":"

Trigonal planar
120°

","correct":false,"explanation":"Trigonal planar geometry requires three regions of electron density. Four bonding regions cannot remain in one plane and instead form a tetrahedral shape with 109.5° angles."}]},{"id":"6f14cb99-1a66-4a29-84ad-84db3d72b33c","cardType":"basic","points":2,"front":"What shape and bond angles come from six bonding regions?","back":"

Octahedral
90° and 180°

","mcqOptions":[{"text":"

Octahedral
90° and 180°

","correct":true,"explanation":"Six bonding regions point towards the vertices of an octahedron. Adjacent bonds are separated by 90°, while bonds occupying opposite positions are separated by 180°."},{"text":"

Trigonal bipyramidal
90°, 120° and 180°

","correct":false,"explanation":"Trigonal bipyramidal geometry corresponds to five regions, producing 90°, 120° and 180° angles. Six bonding regions instead form an octahedral shape with 90° and 180° angles."},{"text":"

Square planar
90° and 180°

","correct":false,"explanation":"Square planar describes four bonds occupying one plane, commonly with two additional lone pairs. Six bonding regions occupy all octahedral positions, giving an octahedral shape."},{"text":"

Pentagonal bipyramidal
72°, 90° and 180°

","correct":false,"explanation":"Pentagonal bipyramidal geometry requires seven regions, including five equatorial positions. Six bonding regions instead arrange octahedrally, with bond angles of 90° and 180°."}]},{"id":"3f38c029-613b-4f8d-bd54-3ab30dbe0405","cardType":"basic","points":2,"front":"Why can lone pairs reduce bond angles?","back":"Lone pairs repel more strongly than bonding pairs.","mcqOptions":[{"text":"Lone pairs repel more strongly than bonding pairs.","correct":true,"explanation":"Lone pairs occupy more space around the central atom and repel more strongly than bonding pairs, pushing bonded atoms closer together and reducing the bond angle."},{"text":"Bonding pairs repel more strongly than lone pairs.","correct":false,"explanation":"This reverses the relative repulsion strengths. Lone pairs repel more strongly because their electron density is concentrated closer to the central atom, so they reduce the bond angle."},{"text":"Lone pairs attract more strongly than bonding pairs.","correct":false,"explanation":"Electron pairs repel one another rather than attract. Lone pairs have stronger repulsions than bonding pairs, which pushes bonded atoms together and reduces the bond angle."},{"text":"Lone pairs have no effect on bonding pairs.","correct":false,"explanation":"Lone pairs influence molecular shape through electron-pair repulsion. Their stronger repulsion compared with bonding pairs pushes the bonds closer together, reducing the bond angle."}]},{"id":"717b7be2-8743-4e8d-b8a8-7f66d62a7dd9","cardType":"basic","points":1,"front":"When can a molecule with polar bonds be non-polar overall?","back":"When symmetrical bond dipoles cancel.","mcqOptions":[{"text":"When symmetrical bond dipoles cancel.","correct":true,"explanation":"In a symmetrical molecule, equal bond dipoles point in directions that cancel each other, giving no overall dipole and making the molecule non-polar overall."},{"text":"When symmetrical bond dipoles reinforce.","correct":false,"explanation":"Reinforcing dipoles produce a net dipole, so the molecule remains polar. Non-polar molecules with polar bonds require symmetrical bond dipoles that cancel."},{"text":"When asymmetrical bond dipoles cancel.","correct":false,"explanation":"Asymmetrical arrangements usually prevent complete cancellation, leaving a net dipole. The required condition is symmetry, so symmetrical bond dipoles cancel overall."},{"text":"When all bonds have non-polar covalent character.","correct":false,"explanation":"A molecule can be non-polar even when its bonds are polar, provided their dipoles cancel through symmetry. Non-polar bonds are not required."}]},{"id":"34a3049a-5e49-45eb-b296-8bf3bff48e39","cardType":"cloze","points":null,"front":"Intermolecular forces are attractions [ ], much weaker than covalent bonds [ ].","back":"Intermolecular forces are attractions between molecules, much weaker than covalent bonds within molecules.","mcqOptions":[{"text":"Intermolecular forces are attractions between molecules, much weaker than covalent bonds within molecules.","correct":true,"explanation":"Intermolecular forces act between separate molecules, whereas covalent bonds hold atoms together within molecules. The latter are much stronger and require more energy to overcome."},{"text":"Intermolecular forces are attractions within molecules, much weaker than covalent bonds between molecules.","correct":false,"explanation":"The locations are reversed here. Intermolecular forces act between molecules, while covalent bonds join atoms within molecules and are much stronger."},{"text":"Intermolecular forces are attractions between atoms, much weaker than covalent bonds within molecules.","correct":false,"explanation":"Attractions between atoms within a molecule are chemical bonds, not intermolecular forces. Intermolecular forces specifically act between separate molecules."},{"text":"Intermolecular forces are attractions between molecules, much weaker than covalent bonds between atoms.","correct":false,"explanation":"Although the first location is correct, the comparison is incomplete in the required wording. Covalent bonds are within molecules, joining their atoms, rather than between separate molecules."}]},{"id":"dd2b9f67-dce1-44e6-9344-0fa6c89e4fd7","cardType":"basic","points":1,"front":"When a simple molecular substance boils, what is usually overcome?","back":"Intermolecular forces between molecules, not covalent bonds.","mcqOptions":[{"text":"Intermolecular forces between molecules, not covalent bonds.","correct":true,"explanation":"Boiling separates molecules from one another, so energy is used to overcome intermolecular forces. The covalent bonds inside each molecule usually remain intact."},{"text":"Covalent bonds within molecules, not intermolecular forces.","correct":false,"explanation":"Boiling normally does not break the covalent bonds within molecules. It overcomes intermolecular attractions between molecules, allowing them to enter the gas phase."},{"text":"Ionic bonds between molecules, not covalent bonds.","correct":false,"explanation":"Simple molecular substances do not usually have ionic bonds between their molecules. Boiling overcomes intermolecular forces, while the covalent bonds within molecules remain intact."},{"text":"Electrons within molecules, not intermolecular forces.","correct":false,"explanation":"Boiling does not remove electrons from molecules. It separates intact molecules by overcoming intermolecular forces, rather than breaking their internal covalent bonds."}]},{"id":"be871709-0ef5-4467-b1e9-cf90a7c9880f","cardType":"basic","points":2,"front":"What causes London forces between molecules?","back":"A temporary dipole induces a dipole in a neighbour.","mcqOptions":[{"text":"A temporary dipole induces a dipole in a neighbour.","correct":true,"explanation":"Random electron movement creates a temporary uneven charge distribution. This temporary dipole repels or attracts electrons in a neighbouring molecule, inducing a dipole and causing attraction."},{"text":"A permanent dipole induces a dipole in a neighbour.","correct":false,"explanation":"Permanent dipoles explain attractions between polar molecules, not London forces specifically. London forces arise when random electron movement creates a temporary dipole."},{"text":"A temporary dipole neutralises a dipole in a neighbour.","correct":false,"explanation":"A temporary dipole polarises a neighbouring molecule rather than neutralising it. The induced dipole then creates an attraction between the two molecules."},{"text":"A lone pair induces a dipole in a neighbour.","correct":false,"explanation":"Lone pairs can contribute to polarity in particular molecules, but London forces occur in all particles through temporary dipoles caused by random electron movement."}]},{"id":"d3153c77-8434-4b5e-805e-4e335699c828","cardType":"basic","points":3,"front":"What makes London forces stronger?","back":"

More electrons
Larger electron clouds
Greater surface contact

","mcqOptions":[{"text":"

More electrons
Larger electron clouds
Greater surface contact

","correct":true,"explanation":"More electrons and larger, more polarisable electron clouds produce stronger temporary dipoles. Greater surface contact allows more attraction between neighbouring molecules."},{"text":"

Fewer electrons
Smaller electron clouds
Greater surface contact

","correct":false,"explanation":"Fewer electrons and smaller electron clouds are less polarisable, weakening London forces. Greater surface contact strengthens them, but the first two trends are reversed."},{"text":"

More electrons
Smaller electron clouds
Less surface contact

","correct":false,"explanation":"More electrons can strengthen London forces, but smaller electron clouds and less surface contact weaken them. Stronger forces require larger clouds and greater contact."},{"text":"

More protons
Smaller electron clouds
Greater surface contact

","correct":false,"explanation":"The relevant factors are electron number, electron-cloud size and surface contact. More protons alone do not determine London-force strength, while smaller clouds reduce polarisability."}]},{"id":"9a1007ad-2caf-42f0-9265-e44a73fd3ad5","cardType":"cloze","points":null,"front":"Hydrogen bonding requires H covalently bonded to [ ], attracted to a [ ] in a neighbouring molecule.","back":"Hydrogen bonding requires H covalently bonded to N, O or F, attracted to a lone pair on N, O or F in a neighbouring molecule.","mcqOptions":[{"text":"Hydrogen bonding requires H covalently bonded to N, O or F, attracted to a lone pair on N, O or F in a neighbouring molecule.","correct":true,"explanation":"Hydrogen bonding occurs when hydrogen is bonded to highly electronegative nitrogen, oxygen or fluorine and is attracted to a lone pair on one of those atoms nearby."},{"text":"Hydrogen bonding requires H covalently bonded to C, N or O, attracted to a bonding pair on N, O or F in a neighbouring molecule.","correct":false,"explanation":"Carbon is not sufficiently electronegative for the usual hydrogen-bonding requirement, and hydrogen bonds involve attraction to a lone pair, not a bonding pair."},{"text":"Hydrogen bonding requires H covalently bonded to N, O or Cl, attracted to a lone pair on N, O or Cl in a neighbouring molecule.","correct":false,"explanation":"Chlorine is the tempting substitution, but the standard A-Level requirement is hydrogen bonded to nitrogen, oxygen or fluorine, attracted to a lone pair on one of them."},{"text":"Hydrogen bonding requires H covalently bonded to any halogen, attracted to a positive ion in a neighbouring molecule.","correct":false,"explanation":"Hydrogen bonding is not formed with any halogen or with a positive ion. It requires H bonded to N, O or F and attraction to a lone pair on N, O or F."}]},{"id":"bf5cbe76-8b74-4657-9f66-227217600cfb","cardType":"basic","points":2,"front":"Why does H₂O boil much higher than H₂S?","back":"H₂O forms hydrogen bonds; H₂S cannot.","mcqOptions":[{"text":"H₂O forms hydrogen bonds; H₂S cannot.","correct":true,"explanation":"The highly electronegative oxygen in water allows hydrogen bonding between molecules. Sulfur is not electronegative enough for H₂S to form significant hydrogen bonds, so its boiling point is lower."},{"text":"H₂S forms hydrogen bonds; H₂O cannot.","correct":false,"explanation":"This reverses the key property. Oxygen is sufficiently electronegative for water to form hydrogen bonds, whereas sulfur in H₂S is not."},{"text":"H₂O forms ionic bonds; H₂S cannot.","correct":false,"explanation":"Water is a simple molecular substance with covalent O-H bonds, not ionic bonds. Its higher boiling point results from intermolecular hydrogen bonding."},{"text":"H₂O forms stronger London forces; H₂S cannot.","correct":false,"explanation":"Both molecules experience London forces, and H₂S has more electrons. Water boils much higher because it forms hydrogen bonds, which H₂S cannot."}]},{"id":"2efaeb1f-8ccb-40e7-bd64-5a361bd0eec2","cardType":"basic","points":2,"front":"What does “like dissolves like” predict for solubility?","back":"

Polar substances dissolve in polar solvents.
Non-polar substances dissolve in non-polar solvents.

","mcqOptions":[{"text":"

Polar substances dissolve in polar solvents.
Non-polar substances dissolve in non-polar solvents.

","correct":true,"explanation":"Polar solutes form favourable intermolecular attractions with polar solvents, while non-polar solutes mix through similar dispersion forces with non-polar solvents."},{"text":"

Polar substances dissolve in non-polar solvents.
Non-polar substances dissolve in polar solvents.

","correct":false,"explanation":"This reverses the principle by pairing substances with unlike solvents. Polar substances dissolve in polar solvents, while non-polar substances dissolve in non-polar solvents."},{"text":"

Polar substances dissolve in non-polar solvents.
Non-polar substances dissolve in non-polar solvents.

","correct":false,"explanation":"Assuming all substances dissolve in non-polar solvents ignores the need for compatible intermolecular forces. Polar substances favour polar solvents, whereas non-polar substances favour non-polar solvents."},{"text":"

Polar substances dissolve in polar solvents.
Non-polar substances dissolve in polar solvents.

","correct":false,"explanation":"Assuming all substances dissolve in polar solvents overlooks the weak attractions between non-polar solutes and polar solvents. Non-polar substances instead dissolve in non-polar solvents."}]}]}],"examCountdown":{"examDate":"2027-05-28","paperLabel":"Component 1"},"heading":"Bonding","title":"Bonding Flashcards","topbar":{"backHref":"/a-level/chemistry/eduqas/flashcards"},"topicSidebar":"$L38"}]}],"topicId":"41b97916-5a4f-5738-844e-413a7e832ce6","walled":"$L39"}],"$L3a","$L3b"]