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A Level Chemistry > Topic 1: Atomic Structure and Periodic Table > Flashcards

Topic 1: Atomic Structure and Periodic Table Flashcards

1
Q

What is an atom?

A
  • consists of a tiny nucleus containing protons and neutrons
  • surrounded by electrons which make up the majority of the volume of the atom
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2
Q

Define isotopes

A
  • Isotopes are atoms of the same element that have the same number of protons, but different numbers of neutrons
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3
Q

Define relative isotopic mass

A

Relative isotopic mass is the mass of one atom of an isotope compared to one-twelfth of the mass of one atom of carbon-12

  • E.g. 35Cl ‘s isotopic mass is 35 (equal to the mass number)
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4
Q

Define relative atomic mass (Ar)

A
  • Relative atomic mass of an element is the weighted average mass of an atom of an element compared to one twelfth the mass of one atom of carbon-12
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5
Q

Define relative molecular mass (Mr)

A
  • Relative molecular mass is the average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12
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6
Q

Define relative formula mass (Mr)

A

Relative formula mass is the average mass of the formula units of a ionic compound compared to one twelfth of the mass of one atom of carbon-12

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7
Q

What is mass spectrometry?

A
  • measures the masses of atoms and molecules
  • produces positive ions that are deflected by a magnetic field according to their mass-to-charge ration (m/z)
  • also calculates relative abundance of each positive ion
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8
Q

Use of mass spectrometers

A
  • drug testing in sports
  • quality control in pharma
  • radioactive dating to determine age of fossils or human remains
  • space probes to identify elements on other planets
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9
Q

How do you work out the relative atomic mass from a mass spectra graph?

A
  • multiply each relative isotopic mass by its relative isotopic abundance and add up the results
  • divide by the sum of the isotopic abundances
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10
Q

What is the M+1 peak?

A
  • seen when analysing organic compounds with large masses
  • molecular ion peak + 1 due to the small percentage of carbon-13 present
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11
Q

Where do electrons move?

A
  • electron move around the nucleus in quantum shells, also known as energy levels
  • these shells are given numbers as principal quantum numbers
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12
Q

Do electron shells nearer or further away from the nucleus have greater energy?

A
  • further away
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13
Q

Describe the electron shells in more detail

A
  • shells contain different types of subshell
  • these have different number of orbitals, which can hold up to 2 electrons
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14
Q

How many orbitals are there in subshell s and therefore how many electrons can it hold max?

A
  • 1 orbital
  • 2 electrons
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15
Q

How many orbitals are there in subshell p and therefore how many electrons can it hold max?

A
  • 3 orbitals
  • 6 electrons
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16
Q

How many orbitals are there in subshell d and therefore how many electrons can it hold max?

A
  • 5 orbitals
  • 10 electrons
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17
Q

How many orbitals are there in subshell f and therefore how many electrons can it hold max?

A
  • 7 orbitals
  • 14 electrons
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18
Q

Which sub shells are in the 1st electron shell?

A
  • 1s
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19
Q

Which sub shells are there in the second shell?

A
  • 2s
  • 2p
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20
Q

Which subshells are there in the third shells?

A
  • 3s
  • 3p
  • 3d
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21
Q

Which sub shells are in the 4th shell?

A
  • 4s
  • 4p
  • 4d
  • 4f
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22
Q

What are orbitals?

A
  • orbitals is the space that an electron moves in
  • orbitals within the same subshell have the same energy
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23
Q

What does an s orbital look like?

A
  • they are spherical
24
Q

What do p orbitals look like?

A
  • they have dumbbell shapes
  • one on the x axis, one on y-axis and one on z axis
25
Q

How can you represent electrons in orbitals?

A
  • you can you arrows in boxes
  • each box represents one orbitals
26
Q

How do you figure out electronic configurations?

A
  • electrons fill up lower energy sub shells first
  • remember the 4s sub shells has lower energy than the 3d subsell, so fills up first
27
Q

Write the electronic configuration of calcium

A
  • 1s2 2s2 2p6 3s2 3p6 4s2
28
Q

How can you use the periodic table to work out electronic configurations?

A
  • the periodic table can be split into an s-block (groups 1 and 2), d-block (transition metals) and p-block (third part)
  • s-block elements have a valence shell electronic configuration of s1 or s2
  • p-block elements have an outer shell configuration of s2p1 to s2p6
  • d-block are harder to work out because 4s subshell fills before the 3d subshell
29
Q

Why do chromium and copper donate one of their 4s electrons to the 3d subshell?

A
  • they’re more stable with a full or half full d subshell
30
Q

What is the first ionisation energy?

A
  • the first ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole gaseous 1+ atoms
31
Q

Is ionising an atom or a molecule an exothermic or an endothermic process?

A
  • an endothermic process
  • energy is being put in to ionise
32
Q

What are the three factors that affect ionisation energy?

A
  • nuclear charge: the more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons
  • electron shell: attraction falls off very rapidly with distance
  • an electron in an electron shell close to the nucleus will be much more strongly attached than one in a shell further away
  • shielding: as the number of electrons between valence electrons and the nucleus increases, they feel less attraction towards the nuclear charge
  • electron-electron repulsion occurs between two atoms of the same orbital and between electrons in different orbitals
33
Q

Describe and explain the trend of first ionisation energies as you go down a group

A
  • as you go down a group in the periodic table, ionisation energies generally fall, meaning it is easier to remove outer electrons because
  • electrons further down a group have extra electron shells compared to ones above. this means that the atomic radius is larger, so valence electrons are further away from the nucleus, which greatly reduces attraction to nucleus
  • the inner shells shield the outer electrons from the attraction of the nucleus
  • the positive charge of the nucleus does increase as you go down group due to extra protons, but it is overridden by the effect of extra shells
34
Q

Write out an equation for the first ionisation of oxygen

A

O(g) = O+(g) + e-

35
Q

Define second ionisation energy

A
  • the energy needed to remove 1 electron from each singly charged positive ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions
36
Q

Write an equation for the second ionisation of oxygen

A

O+(g) = O2++(g) + e-

37
Q

How does a graph of successive ionisation energies provide evidence for the shell structure of atoms?

A
  • within each shell, successive ionisation energies increase
  • this is because electrons are being removed from an increasingly positive ion: there’s less repulsion amongst the remaining electrons, so they’re held more strongly by the nucleus
  • the big jumps in ionisation energy happen when a new shell is broken into, an electron is being removed from a shell closer to the nucleus
38
Q

How can you tell from a successive ionisation energy graph which group the element belongs to?

A
  • count how many electrons are removed before the first big jump
39
Q

Why does helium have the largest first ionisation energy?

A
  • its first electron is in the first shell closest to the nucleus and it has no shielding effects from inner shells
  • it has a bigger first IE than H as it has one more proton
40
Q

Which elements in a period table have the same number of electron shells?

A
  • elements within a period
  • period has 1 shell, period 2 has 2 shells etc
41
Q

Why does Na have a much lower first ionisation energy than Neon?

A
  • Na’s valence electron is in a 3s shell
  • which is further from the nucleus and is more shielded
42
Q

Why is there a small drop of 1st IE from Mg to Al (and for Be and B)?

A
  • Al is starting to fill a 3p subshell, whereas Mg has its outer electrons in the 3s subshell
  • electrons in the 3p subshell are slightly easier to remove because the 3p electrons are higher in energy and also slightly shielded by the 3s electrons
43
Q

Why is there a small drop in 1st IE from P to S?

A
  • with sulphur, there are 4 electrons in the 3p subshell and the 4th is starting to doubly fill the first 3p orbital
  • this causes a slight repulsion between the two negatively charged electrons, which makes the second electron easier to remove
44
Q

Explain why there is a slight decrease in 1st IE from nitrogen to oxygen

A
  • the first electron removed from the oxygen atom is one of the two paired electrons in the 2px orbital
  • the presence of two electrons in a single orbital increases the electron-electron repulsion
  • therefore, less energy is required to remove one of these electrons, despite the larger nuclear charge
45
Q

What is periodicity?

A
  • the repeating trends in the physical and chemical properties of the elements across each period
46
Q

What do elements within a group have in common?

A
  • same number of electrons in their outer shell
47
Q

What chemical properties do s-block elements have?

A
  • s-block elements have 1 or 2 valence electrons
  • these are easily lost to form positive ions with an inert gas configuration
48
Q

What chemical properties of elements in groups 5, 6, 7 have?

A
  • they can gain 1, 2, or 3 electrons to form negative ions inert gas configuration
49
Q

Why is group 0 inert?

A
  • their s and p sub shells have completely filled and don’t need to gain or lose or share electrons
50
Q

What are the chemical properties of d-block elements?

A
  • they tend to lose s and d electrons to form positive ions
51
Q

What is the trend of the atomic radius across a period?

A
  • it decreases
  • as the number of protons increases, the positive charge of the nucleus increases
  • which creates a more positive charge attraction for electrons which are in the same shell with similar shielding
52
Q

What is the trend of ionisation energy across a period

A
  • it generally increases across a period
  • as you move across a period, the general trend is for the ionisation energies to increase because it is harder to remove outer electrons
  • this is because proton numbers increase, so stronger nuclear attraction
  • all the extra electrons are at roughly the same energy level, even if outer electrons are in different orbital types
  • this means there is generally little extra shielding effect or extra distance to lessen the attraction from the nucleus
  • but there are small drops between group 2 and 3, and 5 and 6
53
Q

What is the trend across the period for metals in melting and boiling points?

A
  • they increase across the period because the metallic bonds get stronger
  • metal ions have an increasing number of delocalised electrons and decreasing radius
  • this means there is a strong attraction between the metal ions and delocalised electrons so stronger metallic bonding
54
Q

Why do carbon and silicon have high melting and boiling points?

A
  • they are elements with giant covalent lattice structures
  • they have strong covalent bonds linking all their atoms together
  • a lot of energy is needed to break all these bonds
55
Q

Describe the trend across the period for melting and boiling points in simple molecular structures in period 2 and 3

A
  • melting points depends on the strength of the London dispersion forces between their molecules
  • these forces are weak and easily overcome, so these elements have low melting and boiling points
  • more electrons in a molecule mean stronger London forces
  • S8 has the most electrons so it has got higher melting and boiling points
56
Q

Why do noble gases have the lowest melting and boiling points in their periods?

A
  • they exist as individual atoms, monatomic
  • resulting in very weak London forces

A Level Chemistry (19 decks)

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