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Chemistry year 12 - speedrun > Topic 1 - Atomic Structure and the Periodic Table > Flashcards

Topic 1 - Atomic Structure and the Periodic Table Flashcards

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1
Q

what are atoms?

A

atoms of the same element have the same number of protons.

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2
Q

what are isotopes?

A

isotopes are atoms of the same element with the same number of protons and electrons, but a different number of neutrons.

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3
Q

what are ions?

A

ions are charged particles formed when atoms lose or gain electrons to gain stability.

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4
Q

what is relative isotopic mass?

A

the mass of an atom of an isotope compared with 1/12th the mass of an atom of carbon-12.

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5
Q

what is relative atomic mass?

A

the weighted average mass of an atom of an element compared with 1/12th the mass of an atom of carbon-12.

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6
Q

what does a mass spectrometer measure?

A

the mass/charge ratio and abundance of each isotope of an element.

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7
Q

How do you write a species of a mass spectrometer?

A

eg. 24Mg+ (DONT FORGET THE +). Include + and mass number.

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8
Q

How do you calculate relative atomic mass?

A

sum of (isotopic mass x abundance) / 100.

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9
Q

What is the rightmost peak in a mass spectrum of a molecule?

A

molecular ion. E.g, molecular ion for butane is C4H10+.

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10
Q

Examples of uses of mass spectrometers?

A
  • planetary space probes
  • drug testing in sports
  • quality control in pharmaceutical industry.
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11
Q

first ionisation energy?

A

the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

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12
Q

second ionisation energy?

A

the energy required to remove one electron from each ion in one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.

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13
Q

factors affecting ionisation energy?

A
  • attraction of the nucleus
  • atomic radius (distance of outermost electrons from nucleus)
  • shielding effect.
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14
Q

ionisation energy trend down a group?

A
  • although nuclear charge increases
  • atomic radius increases (number of shells increase)
  • electron shielding effect increases
  • force of attraction between outermost electrons and nucleus decreases
  • first ionisation energy decreases.
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15
Q

ionisation energy trend across a period?

A
  • nuclear charge increases
  • atomic radius decreases (number of shells stays the same, number of protons and electrons increases, force of attraction between them increases).
  • electron shielding effect stays the same (same number of shells)
  • force of attraction between outermost electrons and nucleus increases
  • first ionisation energy increases.
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16
Q

trend in first ionisation energy from end of period to start of new period?

A
  • although nuclear charge increases
  • atomic radius increases (1 more shell)
  • electron shielding affect increases
  • force of attraction between outermost electrons and nucleus decreases.
  • first ionisation energy decreases.
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17
Q

why is second ionisation energy always greater than the first?

A
  • when the first electron is removed, a positive ion is formed.
  • nuclear charge remains the same, but there is 1 less electron, so the force of attraction between the nucleus and the remaining electrons increases.
  • energy required to remove another electron increases.
18
Q

how can you tell what group an element is in by looking at ionisation energies?

A
  • you would see a big jump in ionisation energy, which indicates the number of electrons in the outer shell, hence the group.
19
Q

Why is there a small drop in first ionisation energy from Mg to Al? (group 2 and 3)

A
  • Al starts to fill in the 3p subshell, whereas Mg has its outer electrons in the 3s subshell.
  • electrons in the 3p subshell are higher in energy than 3s electrons.
  • electrons in 3p sub shell are also slightly shielded by 3s electrons.
  • therefore first ionisation energy is lower for group 3 than 2.
20
Q

why is there a small drop from phosphorus to sulfur? (group 5 and 6).

A
  • sulfur has 4 electrons in the 3p subshell.
  • 4th electron fully fills in the first 3p orbital.
  • the electron pair in the orbital slightly repel each other due to opposite spin, making it easier to remove an electron.
  • hence first ionisation energy decreases.
21
Q

what is an atomic orbital?

A

a region of space around the nucleus of an atom that can hold up to 2 electrons with opposite spin.

22
Q

main features of s-orbital?

A
  • each shell contains 1
  • spherical in shape
  • each s subshell holds 2 electrons.
23
Q

main features of p-orbital?

A
  • 3 in every shell from 2nd shell
  • dumbbell shape
  • each p subshell holds 6 electrons.
24
Q

d-orbital?

A
  • 5 in every shell from 3rd shell.

- each s subshell can hold 10 electrons.

25
Q

f- orbial?

A
  • 7 in every shell from 4th shell.

- each f subshell can hold 14 electrons.

26
Q

why does 4s get filled before 3d?

A

sub shells fill in order of increasing energy. 4s subshell is lower in energy than 3d subshell so 4s is filled in first. However, when electrons are lost, they are lost from 4s subshell first.

27
Q

how are electron configurations drawn?

A
  • 2 electrons of opposite spin in each orbital.

- each orbital gains 1 electron before the sub shell finishes filling up.

28
Q

why are copper and chromium exceptions to the normal electron structure?

A

in both cases, it is a more stable arrangement to move 1 electron from the 4s subshell to ‘fill out’ the 3d subshell. (so there would be 1 electron in the 4s subshell for them).

29
Q

what do you need to remember about ions when writing electron configurations?

A

4s is part of the outermost shell so electrons are lost from there first (even though it was filled up first than 3d).

30
Q

when writing electron configurations and electrons filled in 3s and 3d, what do you write first?

A

even though 4s subshell is filled in first than 3d, you write it in number order. (e.g …3p6 3d9 4s2)

31
Q

what does the term ‘periodicity’ mean?

A

the repeating pattern of physical or chemical properties going across the periods.

32
Q

classification of elements in s, p, d, f blocks?

A

elements are classified as s, p, d, f block, according to which orbitals their highest energy electrons are in.

33
Q

why does atomic radii decrease across a period?

A
  • nuclear charge increases
  • number of shells stays the same
  • electron shielding effect is the same
  • force of attraction between outermost electrons and nucleus increases.
  • outer electrons are pulled in closer towards the nucleus.
  • atomic radii decreases across a period.
34
Q

trend in melting points for Na to Al?

A
  • metallic bonding increases.
  • proton number increases, so there are more electrons on the outermost shell.
  • more electrons are released into the sea of delocalised electrons.
  • ionic radii also decreases across period.
  • more delocalised electrons per unit volume
  • (delocalised electrons are closer to the nucleus of the cations).
  • stronger metallic bonding
  • increase in melting temperature.
35
Q

Why does Si have a greater melting/boiling point than Na to Al?

A
  • Si is a giant covalent structure.
  • many strong covalent bonds.
  • a very high amount of energy is required to overcome the strong covalent bonds, hence the high mp and bp.
  • covalent bonds are stronger than metallic bonds.
36
Q

P, S, Cl, Ar boiling points are low. Why?

A

Only weak intermolecular forces (London forces) between the molecules. They require only a little amount of energy to overcome, therefore low mp and bp.

37
Q

Why does Sulfur have a higher boiling point than Chlorine?

A

Sulfur exists as molecules of 8 atoms, whereas chlorine exists as diatomic molecules of 2 atoms. Sulfur molecules contain a greater number of electrons than chlorine diatomic molecules, therefore there are greater London forces between sulfur molecules. Hence more energy is required to overcome them, so higher mp and bp.

38
Q

Why does Argon have a lower mp than chlorine?

A

argon exists as single atoms, chlorine exists as diatomic molecules. Chlorine molecules have more electrons, forms greater London forces, so more energy is required to overcome them, hence chlorine has a higher mp and bp than argon.

39
Q

what is atomic (proton) number?

A

the number of protons in the nucleus of an atom.

40
Q

what is mass number?

A

the sum of the number of protons and the number of neutrons in the nucleus of an atom.

Chemistry year 12 - speedrun (7 decks)

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