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A-level Chemistry > Atomic Chemistry > Flashcards

Atomic Chemistry Flashcards

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1
Q

What is mass spectroscopy?

A

Powerful Instrumental technique to find the relative mass of elements and compunds

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2
Q

What are the 4 stages of time of flight mass spectroscopy?

A
  • ionisation
  • acceleration of ions
  • separation of charged ions
  • detection
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3
Q

What are the 2 ionisation techniques in mass spectroscopy?

A
  • electron impact
  • electrospray ionisation
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4
Q

What is electron impact? What is it used for?
What is the equation?

A
  • elements and low Mr compounds
  • high energy electrons fired at sample from electron gun
  • knocks ff one electron from each atom/molecule to form 1+ ion
    X(g) –> X+(g) + e–
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5
Q

What is electrospray ionisation? What is it used for?
What is the equation?

A
  • high Mr compounds (proteins)
  • sample dissolved in volatile solvent (methanol/water), injected through fine hypodermic needle as fine spray into vacuum in ionisation chamber
  • high voltage applied to needle where spray emerges (positively charged)
  • particles gain proton and become ions as a fine mist
    X(g) + H+(g) –> XH+(g)
  • solvent evaporates leaving 1+ ions
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6
Q

What happens in the acceleration of ions in mass spectroscopy?

A

ions accelerated using electric field so all have same KE

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7
Q

What happens in the separation of charges ion in mass spectroscopy?
What is the equation?

A
  • ion drift - ions then enter flight tube
  • ions with different masses had different time of flight
  • light ions = faster = less time to reach detector
    t = d√m/2KE
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8
Q

What happens in detection in mass spectroscopy?

A
  • current produced when ions hit detector (negatively charged plate)
    more ions hit = bigger current
  • mass ions calculated from time of flight
  • mass spectrum shows number of particles (abundance) of each mass that hit detector
  • horizontal axis mass:charge ratio (m/z) but as charge usually 1+, m/z ratio effectively mass
  • in an electrospray ionisation mass spectrum, main peak at Mr 1+, so Mr is 1 less than molecular ion peak
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9
Q

How do you find the Mr from a mass spectra?

A
  • main peak = Mr.
  • often small peak at the Mr+1 due to molecules containing 13C/2H atoms
    e.g. butane Mr 58, w/ small peak at 59 due to isotopes
  • in molecules/diatomic molecules containing atoms w/ significant no. isotopes, peaks more significant
    e.g. Cl2 (w/ isotopes 35Cl + 37Cl in ratio 3:1)
  • if the m/z value is half of the isotopes mass then the z value has gone up to 2 so the ion would be 2+ but the peak would be small as it is unlikely to happen
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10
Q

What are the equations used in mass spectroscopy?

A

t = d/v
v = √2KE/m
d/t = √2KE/m
d^2/t^2 = 2KE/m
t^2/d^2 = m/2KE
t^2 = d^2 x m/2KE
t = d x √m/2KE

t = √2KEt^2/m

t = √md^2/2KE

m = 2KEt^2/d^2

m/t^2 = m/t^2

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11
Q

The element Mg (Ar 24.3) has 3 isotopes 24Mg, 25Mg, 26Mg. If the % of the heaviest isotope is 11.0% what is the % of the lightest?

A

((26x11)+(24x𝑥)+(25x89-𝑥))/100 = 24.3
286+24𝑥+2225-25𝑥=2430
2511-𝑥=2430
81=𝑥
24Mg = 81%
25Mg = 8%
26Mg = 11%

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12
Q

What do you do to find the mass no. in a mass spectroscopy calculation?

A
  • find the mass of one mole
  • multiply by Avogadro’s constant (6.022 × 10²³)
  • convert to g
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13
Q

What is a shell?

A
  • Group of orbitals whose radial distributions from the nucleus is approximately equal
  • the energy level of the orbitals in a shell are not necessarily equal
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14
Q

What is an orbital?

A
  • volume of space in which there is a high probability of finding an electron
  • an orbital has a fixed energy level
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15
Q

How many s orbitals are there?
How many p orbitals are there?
How many d orbitals are there?

A

1 (max no. electrons in sublevel :2 )
3 (max no. electrons in sublevel :6 )
5 (max no. electrons in sublevel :10 )

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16
Q

Shells and orbitals - Pauli: exclusion principle

A

maximum number of electrons that can be put into an orbital is 2 (if there are 2 electrons in an orbital they will move in opposite directions)

17
Q

Shells and orbitals - Hund’s rule

A

Electrons fill orbitals of equal energy singly before pairing up

Electrons prefer to occupy orbital son their own and only pair up when no empty orbitals of the smae energy as available

18
Q

Shells and orbitals - Aufbau principle

A

each electron within an atom occupies the orbital whose energy is the lowest available to it

Electrons enter the lowest energy orbital available

19
Q

In what order are the orbitals filled?

A

1 - 1s
2 - 2s
3 - 2p
4 - 3s
5 - 3p
6 - 4s
7 - 3d
8 - 4p

20
Q

What is the electron configuration for the first 20 elements?

A

H 1s1
He 1s2
Li 1s2 2s1
Be 1s2 2s2
B 1s2 2s2 2p1
C 1s2 2s2 2p2
N 1s2 2s2 2p3
O 1s2 2s2 2p4
F 1s2 2s2 2p5
Ne 1s2 2s2 2p6
Na 1s2 2s2 2p6 3s1
Mg 1s2 2s2 2p6 3s2
Al 1s2 2s2 2p6 3s2 3p1
Si 1s2 2s2 2p6 3s2 3p2
P 1s2 2s2 2p6 3s2 3p3
S 1s2 2s2 2p6 3s2 3p4
Cl 1s2 2s2 2p6 3s2 3p5
Ar 1s2 2s2 2p6 3s2 3p6
K 1s2 2s2 2p6 3s2 3p6 4s1
Ca 1s2 2s2 2p6 3s2 3p6 4s2

21
Q

What is the electron configuration of the transition metals?

A

Sc 1s2 2s2 2p6 3s2 3p6 3d1 4s2
Ti 1s2 2s2 2p6 3s2 3p6 3d2 4s2
V 1s2 2s2 2p6 3s2 3p6 3d3 4s2
Cr 1s2 2s2 2p6 3s2 3p6 3d5 4s1 *
Mn 1s2 2s2 2p6 3s2 3p6 3d5 4s2
Fe 1s2 2s2 2p6 3s2 3p6 3d6 4s2
Co 1s2 2s2 2p6 3s2 3p6 3d7 4s2
Ni 1s2 2s2 2p6 3s2 3p6 3d8 4s2
Cu 1s2 2s2 2p6 3s2 3p6 3d10 4s1 * (this is a slightly lower energy arrangement as the reduced e- - e- repulsion makes up for the fact one electron is in a slightly higher energy level)
Zn 1s2 2s2 2p6 3s2 3p6 3d10 4s2

22
Q

What does isoelectronic mean?

A

different ions with the same electron configuration

23
Q

Explain electron orbitals:

A
  • electron arranged in electrons shells (energy levels)
  • which themselves have sub-shells
  • each sub-shell consists of electron orbitals (region of space the electrons spends most of its time)
  • each orbitals can hold 2 electrons with opposite spins (1cw 1acw)
  • In ions electrons in highest energy levels are lost first (electrons are lost from 4s before 3d)
24
Q

What is ionisation energy?

A

Energy required in order to remove a single electron from each of one mole of gaseous atoms or positive ions

energy needed to remove one mole of electrons from one mole of an element in its gaseous state

25
Q

What are the 1st and 2nd I.E.?

A

1st I.E. = energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of 1+ gaseous ions
X(g) –> X+ (g) + e-

2nd I.E. = energy required to remove second electron (not both electrons)
X+(g) + e- –> X2+ (g) +2e-

successive ionisation energies

26
Q

Describe I.E. down a group:

A
  • Atoms get bigger
  • More shielding
    = weaker attraction from nucleus to electron in outer shell
27
Q

Describe I.E. across a period:

A

I.E. generally increases
General Trend:
- Increased nuclear charge (more protons)
- Atoms get smaller
= stronger attraction from nucleus to electron in outer shell
Group 2 → 3
- G3 electron lost from p orbital / G2 electron lost from s orbital
- p orbital higher energy than s orbital = easier to lose electron
Group 5 → 6
- G6 element loses electron from orbital w/ 2 electrons
- G5 element loses electron from orbital w/ 1 electron
- Extra electron-electron repulsions = easier to lose electron from p4 than p3

28
Q

Describe 2 exceptions to I.E. trends across a period:

A
  • magnesium -> aluminium ionisation energy goes down: as outer electron in aluminium is in a 3p orbital which is of a slightly higher energy than the 3s orbital, therefore needing less energy to remove it
  • phosphorus -> sulfur ionisation energy goes down: in phosphorus each of the 3p orbitals contains just one electron, while in sulfur one of the 3p orbitals must contain 2 electrons. The repulsion between these paired electrons makes it easier to remove one of them despite the increase in nuclear charge
  • Both cases evidence that confirms existence of s- and p-sub levels
29
Q

Which of these elements have the higher first ionisation energy? Why?
argon v potassium
phosphorus v sulfur
magnesium v calcium
magnesium v aluminium
oxygen v fluorine

A

argon - outermost electron closest to nucleus, smaller atomic radius and less shielding
phosphorus - P orbital 1 electron, S orbital 2 electrons=less electron-electron repulsion
magnesium - Mg electron shell3, Ca shell4, Mg smaller atomic radius, less shielding
magnesium - Mg electron from 3s, Al 3p (3p higher energy than 3s)
fluorine - smaller atomic radius + more protons

30
Q

Why is the first I.E. of oxygen less than that of nitrogen?

A

Electron being added to half full orbital in oxygen, which results in electron-electron repulsion, which will lower the ionisation energy

31
Q

Why does helium have the highest first I.E. of all the elements?

A

higher nuclear charge and same atomic radius as hydrogen = magnitude of force acting on each electron higher

32
Q

Why is the second ionisation of an atom always greater than the first?

A

requires even more energy to remove an electron from a cation than from a neutral atom

33
Q

Why does atomic size decrease across a period?

A

Higher nuclear charge causes greater attractions to electrons, pulling electron cloud closer to nucleus = smaller atomic radius

34
Q

Why does atomic size increase down a group?

A

number of energy levels increases = greater distance between nucleus and outermost orbital

35
Q

Why are cations always smaller than corresponding atoms?

A

Electrons are lost further reducing radius of ion

36
Q

Why are anions always larger than the corresponding atoms?

A

Adding electrons increases no. of electron-electron repulsion interactions

37
Q

What do I.E. graphs show?

A

Gives evidence shells exist
- outermost electron = lowest I.E. = easiest to remove
- grouped/similar as in the second shell
- innermost shell = highest I.E. = closest to nucleus

38
Q

What is relative molecular mass? (Mr)

A

average mass of a molecule relative to 1/12 of the mass of an atom of carbon-12

39
Q

What is relative atomic mass? (Ar)

A

average mass of an elements atoms compared to 1/12 the mass of a carbon-12 atom

A-level Chemistry (2 decks)

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