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Physical Chemistry > Atomic Structure 1 > Flashcards

Atomic Structure 1 Flashcards

1
Q

What did Robert Boyle (1661) propose about atomic structure

A

That some substances could not be made simpler- these were chemical elements

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2
Q

What did Dalton (1803) say about atomic structure

A
  • That elements were composed of indivisible atoms.
  • Atoms of the same element all had the same mass.
  • Atoms of different elements had different masses.
  • Atoms could not be broken down
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3
Q

What did Becquerel (1896) discover?

A

Radioactivity. This proved that particles could come from inside the atom and therefore atoms are not indivisible. This led to JJ.Thompsons discovery only a year later.

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4
Q

What did JJ.Thomson develop?

A

He discovered electrons- the first sub-atomic particle and proposed the plum pudding model of the atom. He said that electrons were negatively charged and electrons are always the same.

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5
Q

What did Rutherford (1911) discover about atomic structure

A

Rutherford carried out the alpha particle scattering experiment and discovered that most of an atoms mass is concentrated in a small, positively charged nucleus and most of an atom is empty space.

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6
Q

What is the relative mass and relative charge of an electron

A

Relative mass= 1/1840
Relative charge= -1

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7
Q

What is the relative mass and relative charge of a proton

A

Relative mass=1
Relative charge= +1

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8
Q

What is the relative charge and relative mass of a neutron

A

Relative mass= 1
Relative charge= 0/neutral

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9
Q

Name the force that holds electrons and protons together in an atom

A

Electrostatic attraction

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10
Q

Name the force that holds protons and neutrons together in the nucleus of an atom

A

Nuclear force

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11
Q

Which of these is the stronger force within an atom- electrostatic attraction nuclear force

A

Nuclear force

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12
Q

Define isotope

A

Atoms of an element with the same number of protons but a different number of neutrons. They differ in mass number due to the different number of neutrons.

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13
Q

Define relative isotopic mass

A

The relative mass of an atom of an isotope of an element on a scale where an atom of carbon-12 is exactly 12

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14
Q

Define relative atomic mass

A

The average mass of an atom of an element on a scale where an atom of carbon-12 is exactly 12

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15
Q

Define relative molecular mass

A

The average mass of a molecule on a scale where an atom of carbon-12 is exactly 12

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16
Q

What is the formula for calculating relative atomic mass

A

Ar= sum of (isotopic masses x abundances)/ total abundance/percentage

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17
Q

What does a mass spectrometer do?

A

A mass spectrometer is used to determine the mass numbers and abundances of all the different isotopes in a sample of an element.

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18
Q

Name the five main stages in mass spectrometry

A
  1. Vaporisation
  2. Ionisation
  3. Acceleration
  4. Deflection
  5. Detection
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19
Q

Why do different isotopes of an element have the same chemical properties

A

Because they have the same electron configuration and this is what determines the chemical properties of an element.

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20
Q

Why is the interior of a mass spectrometer a vacuum/near-vacuum

A

To stop the ions from colliding with molecules from the air.

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21
Q

Describe the first stage of mass spectrometry

A

Vaporisation:
The sample is turned into a gas.

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22
Q

Describe the second stage of mass spectrometry

A

Ionisation.
The atoms undergo ionisation to from positive ions. This can be done via electrospray ionisation or electron impact ionisation.

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23
Q

Describe the third stage of mass spectroscopy

A
  • Acceleration.
  • The positive ions are accelerated by an electric field.
  • This gives all ions a constant kinetic energy.
  • They pass through a velocity selector which ensures that they are all travelling at the same speed.
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24
Q

Describe the fourth stage of mass spectrometry

A
  • Deflection
  • The ions pass through a hole in the negatively charged plate into a tube and they are deflected.
  • Lighter ions have less momentum and are deflected more than heavier ones by the magnetic field.
  • Heavier ions have more momentum and so are deflected less by the magnetic field.
  • This separates out the different isotopes or chemicals in a sample as this stage is mass dependant.
25
Q

Describe the fifth (final) stage of mass spectrometry

A
  • Detection
  • Lighter ions have a greater velocity and so arrive at the detector first.
  • Heavier ions will have a lower velocity and so will reach the detector later.
  • Magnetic field strength is slowly increased to ensure that all the ions reach the detector.
  • When the ions reach the detector, each ion gains electrons from the detector which converts them back into atoms and this transfer of electrons causes a current to flow.
26
Q

Describe how the information gained from the mass spectrometer is used/what’s it tells us

A

The time taken for the ions to move down the drift chamber to the detector is used by the machine to determine the mass number of the isotope (m/z ratio)
The size of the current produced when each isotope or chemical hits the detector is used to determine abundances (of each isotope)

27
Q

What are the two methods of ionisation possible for mass spectrometry

A

Electrospray ionisation and Electron impact ionisation

28
Q

Describe electron impact ionisation

A
  • the sample is vaporised
  • high energy electrons are fired at the sample from an electron gun
  • this knocks off one electron from each particle, forming a positive +1 ion
29
Q

What is the general equation for electron impact ionisation

A

X (g) + e- —> X+ (g) + 2e-

30
Q

Describe electrospray ionisation

A
  • the sample is dissolved in a volatile solvent
  • the solvent is injected into the mass spectrometer using a hypodermic needle attached to a high voltage power supply
  • as the sample is injected, the particles are ionised by gaining a proton from the solvent.
31
Q

What is the general equation for electrospray ionisation

A

X (g) + H+ —> XH+ (g)

32
Q

Why is it necessary to ionise molecules when measuring their mass in a TOF mass spectrometer

A
  • only ions, not atoms/molecules will interact with and be accelerated by an electric field
  • only ions, not atoms/molecules will create a current when hitting the detector
33
Q

What did Bohr (1913) suggest about atomic structure

A

That electrons orbit the nucleus of an atom in fixed energy levels/shells. The movement of electrons from one shell to the next explains how atoms give out light.

34
Q

What did Schrodinger (1926) suggest about atomic structure

A

Schrodinger formed an equation that showed that electrons have the properties of waves as well as particles. This led to the theory of quantum mechanics which can be used to predict the behaviour of sub-atomic particles.

35
Q

What did Chadwick (1932) discover

A

The neutron

36
Q

From largest to smallest, list the three components of electron arrangement

A

1) Main energy level
2) Subshells
3) orbitals

37
Q

Define electron orbital

A

A three-dimensional description of an area in space where you are most likely to find an electron

38
Q

What are the four types of sub shell

A

S, p, d, f

39
Q

How many orbitals and electrons are there in an S sub-shell

A

1 orbital, 2 electrons

40
Q

How many orbitals and electrons are there in a P sub-shell

A

3 orbitals, 6 electrons

41
Q

How many orbitals and electrons are there in a D sub-shell

A

5 orbitals, 10 electrons

42
Q

How many orbitals and electrons are there in an F sub-shell

A

7 orbitals, 14 electrons

43
Q

How many orbitals and electrons are there in an F sub-shell

A

7 orbitals, 14 electrons

44
Q

How many electrons are there in an orbital

A

A maximum of two electrons which have opposite spin

45
Q

List the electron configuration possible for the first four main levels

A

1S,2S,2P,3S,3P,4S,3D,4P,4D,4F

46
Q

What are the two exceptions that are key to remember when writing electron configuration for atoms/ions

A

1) That the 4S orbital fills before the 3D orbital but the electrons are also lost form this first so when completing electron arrangement for ions, the 4S electrons are lost before the 3D electrons are lost.
2) As a half full orbital is most stable, sometimes the electron configuration changes to make this happen. For example 4S(2),3D(4) becomes 4s(1),3D(5) and 4S(2),3D(9) becomes 4S(1),3D(10)

47
Q

Name the three key rules when completing electron configuration

A
  1. (Aufbau) = electrons enter the lowest available energy level
  2. (Hund) = when in orbitals of equal energy, electrons will try to remain unpaired
  3. (Pauli) = no two electrons can have the same four quantum numbers. I.e. A full orbital must contain opposite spin electrons.
48
Q

Define first ionisation energy

A

The energy needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous ions. It is measured in KJmol^-1

49
Q

Define ionisation energy generally

A

Ionisation energy is the energy needed to remove one electron from each atom or ion in one mole of gaseous atoms or ions to form one mole of gaseous ions.

50
Q

What is the general equation for first ionisation energy

A

X (g) —> X+ (g) + e-

51
Q

What is the general equation for second ionisation energy

A

X+ (g) —> X2+ (g) + e-

52
Q

What is the trend for successive ionisation energies

A

Successive ionisation energies increase each time because the electron is being removed from an ion that has a stronger positive charge, and therefore electrostatic attraction, from the nucleus each time.

53
Q

Why does the size of the jump between successive ionisation energies differ

A

Sometimes the ionisation energy does not change a lot as the electron was in the same orbital/sub-shell as the previous one that had been removed. The increase in ionisation energy on that case is only due to the ion becoming more highly charged. However, sometimes the ionisation energy increases a lot because the electron is from a different main level and so requires a lot more energy to remove due to the increase in nuclear and electrostatic force from the nucleus

54
Q

What is the general rule for the pattern in ionisation energies across a period in the periodic table

A

The ionisation energies increase as the nuclear charge within the atoms increase and this makes the electron harder to remove.

55
Q

What are the two exceptions for the general rule that as you go across a period on the periodic table, the ionisation energies of the elements increase

A

1) the ionisation energy of aluminium is lower than that of magnesium
2) the ionisation energy of sulphur is slightly lower than the ionisation energy of phosphorus

56
Q

Why does the ionisation energy decrease from magnesium to aluminium instead of increase as is the general pattern?

A

The ionisation energy of aluminium is lower than magnesium’s because the outer electron in aluminium is in a 3P orbital which is at a slightly higher energy level than the 3s orbital which means less energy is required to remove the electron.

57
Q

Why is the ionisation energy of sulphur slightly lower than that of phosphorus when the general pattern is that the energies should increase

A

In phosphorus, each of the orbitals contains only one electron but in sulphur one of them orbitals contains two electrons. The repulsion from these paired electrons makes it easy to lose and the outer 3P4 sub-shell in sulphur easily becomes stable, half-full, 3P3.

58
Q

What is the general trend in ionisation energies as you go down a group on the periodic table

A
  • The ionisation energies decrease because the atoms get bigger and so the electrons are in main levels that are further away from the nucleus each time.
  • This means they experience a weaker force of attraction from the nucleus and so the outer electrons require less energy to remove.
  • The nuclear charge within the nucleus does increase as you go down the group, however, this doesn’t make the ionisation energies higher due to the inner electrons having a shielding affect so the outer electrons do not experience a high proportion of the nuclear charge.
59
Q

Why does the range/size of the decrease in ionisation energy decrease as you go down a group on the periodic table

A

The range of the decrease gets smaller because the electrons in the first elements of the group are being removed from completely different ‘main levels’, however, when they get to main level three and four, they are being removed from different sub-shells but from the same ‘main level’ meaning the size of the decrease in ionisation energy is smaller

Physical Chemistry (10 decks)

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