Flashcards · 9701 · AS Topic 1

Atomic structure (AS Level) flashcards

27 cards on AS Topic 1, subtopics 1.1 to 1.4. Pick a subtopic or study them all.

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1.1 Particles in the atom and atomic radius

Relative charge and relative mass of a proton, a neutron and an electron: Proton: charge +1, mass 1. Neutron: charge 0, mass 1. Electron: charge −1, mass 1/1840 (negligible).
Describe the distribution of mass and of charge in an atom: Almost all the mass, and all the positive charge, is in the tiny, dense nucleus (protons and neutrons). The negative charge (electrons) is spread through the large, mostly empty space around it.
Beams of protons, neutrons and electrons at the same velocity pass between charged plates. What happens to each?: Protons curve towards the negative plate (small angle). Electrons curve towards the positive plate through a much larger angle. Neutrons pass straight through, undeflected (no charge).
What decides how much a charged particle beam is deflected in an electric field (at the same velocity)?: Charge ÷ mass. The sign of the charge decides the direction; charge ÷ mass decides the size. ¹H⁺ (1 ÷ 1 = 1) is deflected more than ⁴He²⁺ (2 ÷ 4 = 0.5).
Atomic radius across a period and down a group: trends and reasons: Across a period: decreases. Nuclear charge increases, electrons go into the same shell, shielding roughly constant. Down a group: increases. An extra shell each time; greater distance and shielding outweigh the extra nuclear charge.
Why is a cation smaller than its atom, and an anion larger?: Cation: usually loses its whole outer shell, and the remaining electrons are attracted by more protons than there are electrons. Anion: extra electrons enter the same outer shell, increasing electron–electron repulsion with the same nuclear charge.
Order Na⁺, Mg²⁺ and Al³⁺ by ionic radius and explain: Al³⁺ < Mg²⁺ < Na⁺. All are isoelectronic (2,8), but the nuclear charge rises from 11 to 13, so the electrons are pulled in more strongly.

1.2 Isotopes

Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.
Atomic (proton) number and mass (nucleon) number: Atomic number, Z: the number of protons in the nucleus (identifies the element). Mass number, A: the total number of protons + neutrons in the nucleus.
Why do isotopes of an element have the same chemical properties?: They have the same number of electrons and the same electronic configuration, and chemical reactions involve the electrons.
How do the physical properties of isotopes differ (9701 1.2)?: Only in mass and density: different numbers of neutrons give atoms of different mass, so the heavier isotope gives a denser sample (e.g. water made with ²H is denser than ordinary water).
Protons, neutrons and electrons in ⁵⁶₂₆Fe³⁺: 26 protons, 56 − 26 = 30 neutrons, 26 − 3 = 23 electrons.
Chlorine is 75% ³⁵Cl and 25% ³⁷Cl (illustrative values). Calculate Ar.: Ar = (35 × 75 + 37 × 25) ÷ 100 = 3550 ÷ 100 = 35.5

1.3 Electrons, energy levels and atomic orbitals

Number of orbitals and maximum electrons in s, p and d sub-shells: s: 1 orbital, 2 electrons. p: 3 orbitals, 6 electrons. d: 5 orbitals, 10 electrons. Each orbital holds at most two electrons, with opposite spins.
Shapes of s and p orbitals: s: spherical, centred on the nucleus. p: two lobes either side of the nucleus (dumbbell); a p sub-shell has three, pointing along different axes.
Order of filling up to Kr, and the point about 4s and 3d: 1s 2s 2p 3s 3p 4s 3d 4p. 4s fills before 3d because it is slightly lower in energy, but 4s electrons are removed first on ionisation.
Electron configurations of Fe, Fe²⁺ and Fe³⁺: Fe: [Ar] 3d⁶4s² Fe²⁺: [Ar] 3d⁶ (the 4s electrons leave first) Fe³⁺: [Ar] 3d⁵
The two exceptions in the first row of the d block: Chromium: [Ar] 3d⁵4s¹. Copper: [Ar] 3d¹⁰4s¹. A half-filled or full d sub-shell is more stable.
Hund's rule: Electrons occupy separate orbitals within a sub-shell, with unpaired spins, before any orbital is doubly occupied. This minimises repulsion between electrons.
Free radical (in terms of electrons): A species with one or more unpaired electrons.

1.4 Ionisation energy

First ionisation energy (with equation): The energy required to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous 1+ ions. X(g) → X⁺(g) + e⁻
Three factors that control ionisation energy: Nuclear charge (more protons, higher IE), atomic radius (further out, lower IE) and shielding by inner shells (more shielding, lower IE).
Trends in first ionisation energy across a period and down a group: Across a period: increases (nuclear charge rises, shielding roughly constant, radius decreases). Down a group: decreases (radius and shielding increase, outweighing the greater nuclear charge).
Why is aluminium's first ionisation energy lower than magnesium's?: Al's outer electron is in a 3p orbital, higher in energy than Mg's 3s and partly shielded by the 3s electrons, so it is removed more easily. Not 'more shielding by inner shells'.
Why is sulfur's first ionisation energy lower than phosphorus's?: In sulfur two electrons share one 3p orbital and repel each other, so one of them is more easily removed.
Successive IEs (kJ mol⁻¹): 738, 1451, 7733, 10 540. Which group?: Group 2. The big jump is between the 2nd and 3rd: the 3rd electron comes from a shell closer to the nucleus, so there are two outer electrons.
Is ionisation energy exothermic or endothermic?: Always endothermic (positive): energy is required to remove an electron. Successive ionisation energies always increase.