Atomic Structure (AQA A Level Chemistry): Flashcards

A positive ion is formed when an atom loses electrons.

A negative ion is formed when an atom gains electrons.

Isotopes have different physical properties because they differ in the number of neutrons, which adds mass to the atom without affecting its charge.

Isotopes have the same number of electrons in their outer shells.

Electrons determine how an atom reacts chemically, so isotopes of the same element behave identically in chemical reactions.

False.

Isotopes have identical chemical properties. Neutrons do not take part in chemical reactions; it is the number of electrons that determines how an element reacts.

A_r = Σ(abundance × mass of isotope) ÷ 100

The percentage abundance of each isotope is multiplied by its mass number, all values are summed, and the total is divided by 100.

1. Ionisation

2. Acceleration

3. Ion drift

4. Detection

An Electrospray ionisation is a soft ionisation technique used for high-mass molecules such as proteins. The sample gains a proton (H⁺) from the solvent to form MH⁺, so the detected mass equals M_r + 1.

X (g) + e^- → X⁺ (g) + 2e^-

In TOF mass spectrometry, all ions are accelerated to the same kinetic energy. A lighter ion therefore has a higher velocity and a shorter time of flight than a heavier ion.

When an ion hits the detector, it gains an electron and generates a small electric current.

The size of this current is directly proportional to the abundance of that ion.

The molecular ion peak is the peak at the highest m/z value. It represents the ion formed when a molecule loses one electron without fragmentation, and its m/z value equals the relative molecular mass of the compound.

In electrospray ionisation, the detected mass of a molecule with relative molecular mass M_r is M_r + 1 because the ion formed is MH⁺ (the molecule has gained a proton).

False.

All ions are given the same kinetic energy during acceleration. Since KE = ½mv², lighter ions have a higher velocity and therefore a shorter time of flight, arriving at the detector first.

A higher value of n indicates a shell with greater energy that is further from the nucleus. The first shell (n = 1) is the closest to the nucleus and has the lowest energy.

An atomic orbital is a region of space in an atom that can hold a maximum of 2 electrons. Orbitals exist at specific energy levels; electrons can only be found at these levels, not in between.

• n = 1: 2 electrons

• n = 2: 8 electrons

• n = 3: 18 electrons

• n = 4: 32 electrons

Each atomic orbital can hold a maximum of 2 electrons. The p sub-shell has 3 orbitals and can therefore hold a maximum of 6 electrons.

The ground state is the most stable electron configuration of an atom, in which electrons fill the lowest available energy sub-shells first.

The three p orbitals in a sub-shell are oriented perpendicular to one another along the x, y and z axes. Each has a dumbbell shape.

False.

AQA mark schemes treat these as distinct. Writing "p shell" instead of "p sub-shell" or "p orbital" will not score the mark. Always use precise terminology.

Sub-shells are filled in order of increasing energy, starting with 1s. The 4s sub-shell is filled before the 3d sub-shell as it has lower energy at that point.

Degenerate orbitals are orbitals within the same sub-shell that have equal energy. For example, the three p orbitals (p_x, p_y and p_z) are degenerate.

1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹

The shorthand electron configuration replaces the inner filled shells with the symbol of the nearest preceding noble gas. For potassium, the shorthand configuration is [Ar] 4s¹.

Cr: [Ar] 3d⁵ 4s¹ and Cu: [Ar] 3d¹⁰ 4s¹.

These differ because half-filled (3d⁵) and fully filled (3d¹⁰) d sub-shells are energetically stable, making these configurations lower in energy than the alternatives.

Electrons occupy separate orbitals within the same sub-shell before pairing up. When pairing occurs, the two electrons have opposite spins to minimise spin-pair repulsion.

When a transition metal forms a positive ion, electrons are lost from the 4s sub-shell first, even though it filled before the 3d. This is because the 4s is higher in energy once the 3d sub-shell contains electrons.

True.

Full electron configurations are written in numerical order of principal quantum number. Sub-shells in n = 3 are written before those in n = 4, regardless of filling order.

The equation for the first ionisation energy of calcium is:

Ca (g) → Ca⁺ (g) + e^-

The first ionisation energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. It is measured in kJ mol^-1.

1. Nuclear charge

2. Distance between the outer electrons and the nucleus

3. Shielding by inner electrons

4. Spin-pair repulsion

First ionisation energy generally increases across a period because nuclear charge increases while shielding remains roughly constant, making it harder to remove an outer electron.

Going down a group, the distance between the outer electron and the nucleus increases (more shells are added) and shielding by inner electrons increases. These factors outweigh the increase in nuclear charge, so less energy is needed to remove the outer electron.

Shielding is the reduction in the attractive force between the nucleus and outer electrons caused by repulsion from inner shell electrons. Greater shielding makes outer electrons easier to remove, lowering the ionisation energy.

The first IE of boron is lower than that of beryllium because boron's outer electron is in the 2p sub-shell, which is higher in energy than beryllium's 2s sub-shell.

True.

The first element of the next period has an outer electron in a new shell that is further from the nucleus and more shielded by inner electrons. These two factors outweigh the increased nuclear charge.

Be⁺ (g) → Be²⁺ (g) + e^-

Successive ionisation energies are the energies required to remove each mole of electrons in turn from one mole of gaseous ions. Each successive IE is greater than the previous one, as electrons are removed from an increasingly positive ion.

Gaseous atoms and ions must carry (g) state symbols. Electrons do not take state symbols. Omitting state symbols from atoms or ions will lose the mark.

On a successive ionisation energy graph, a large jump indicates the next electron is being removed from a shell that is closer to the nucleus, where there is less shielding.

The element belongs to Group 13. There is a large jump between the third and fourth ionisation energies, indicating the first three electrons are in the outer shell and the fourth is in an inner shell closer to the nucleus.

After the first electron is removed, the remaining electrons experience less shielding and the ion carries a positive charge, increasing the attractive force between the nucleus and the remaining electrons. More energy is therefore needed to remove the next electron.

Sodium shows a large jump between its first and second ionisation energies. This indicates it has 1 electron in its outer shell, placing it in Group 1.

True.

The large jump signals a shell boundary — the fourth electron is in an inner shell, much closer to the nucleus. The first three electrons were therefore in the outer shell, confirming the element is in Group 13.