Chemistry
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3.1 Physical Chemistry
3.1.1 Atomic structure
Things to remember :
- Atomic radius : lower on the table , closest to the left of the table = largest radius
- Copper and chromium electron configurations are deviations , take 1 from 4s and add it
to 3d
- When electronic configuration of ions take from 4s before 3d
- Deviations across 1st ionisation energy of period , 2-3 and 5-6 ,
- Second ionisation energy is greater than the first ionisation energy
- Second ionisation energy deviations are shifted 1 to the right ( so 2nd IE period 3
deviations = Al and Si , S and Cl )
- Atomic radius is smaller as you go across the period because higher nuclear charge so
more attraction so closer, but same in isotopes because same no protons
- Mr of a molecule = the highest m/z value on the mass spectrum
- The orbitals are the boxes and subshells are 1s , 2p etc Z is proton number and A is
mass number
- Electron impact is for small Mr and can fragment
- In mass spectrum - lower m/z peaks are due to fragmentation , molecular ion peak is
100% and m+1 ion peak is due to 13C or 2H
- If it asks for the species responsible for a peak then use X+ rather than X to show that it
is a 1+ ion.
- The Pauli exclusion principle – electrons in the same orbital have opposite spins
- The Aufbau principle – Electrons fill orbitals in order of increasing energy (4s before 3d
etc)
- Hund’s rule – electrons fill orbitals in the same subshell singly before paring
Notes
3.1.1 Atomic Structure :
- Atomic structure : JJ Thomson - Plum pudding Model ( cloud of +ve charge, electrons
scattered ) → Ernest Rutherford - Nuclear Model ( fired alpha particles (+ve) at gold foil,
most passed through, some deflected at large angles > mostly empty space, positive
charge concentrated in nucleus ) → Niels Bohr - (electrons orbit around the nucleus
in energy levels) → Erwin Schrodinger - Quantum model ( electrons move in waves )
- Isotope : Atoms of the same element with the same number of protons but different
numbers of neutrons. They have the same chemical properties ( same electronic
configuration ) and different physical properties ( different masses )
- Relative Atomic Mass : The average mass of the naturally occurring isotopes of an
element compared to 1/12th the mass of an atom of carbon-12.
,- RAM = Sum of ( isotopic mass * relative abundance ) / 100
- Mass spectrometry is used for : determination of mass and abundance of each
isotope of an element to find the RAM ( and to find the RMM of substances made of
molecules )
- Electron Impact : For Low Formula Mass ( inorganic/organic ) Vapourised → High
energy electrons are fired at the sample from electron gun ( hot wire filament with a
current which emits electrons ) → Knocks off 1 electron → X(g) + e- → X+(g) + 2e- , X+(g) (
molecular ion ) - can fragment → the 1+ ions are attracted to the -ve electric plate
where they are accelerated
- Electrospray : High Molecular Mass ( biological molecules ), Dissolved in volatile
solvents (water/methanol)→ Injected into a fine hypodermic needle to give a fine mist
→ needle is attached to the positive end of a high voltage power supply → particles gain
a proton. X(g) + H+ → XH+(g) , solvent evaporates while XH+ ions are attracted to a
negative plate
- Acceleration : All have the same KE due to the electric field KE = ½ mv2 , the velocity
depends on the mass ( lighter particles will have a faster velocity )
- Detection : positive ion hits -ve plate → they are discharged by gaining an electron
from the plate → generates a movement of electrons and hence a current → size of
current is proportional to number of ions hitting the plate
- Electrospray mass spectrum - m/z = mass of each ion ( -1 from m/z to find RMM )
Electron impact mass spectrum - greatest m/z is the molecular ion (m/z = molecular
mass) , small peaks due to molecular ions containing different isotopes ( C13 and H2 )
and there may be fragmentation
- Mass of one ion (kg) = Relative isotopic mass * 10-3 / (6.022 x 1023 )
- KE = ½ mv2( where m is in kg) , v = d/t - (m/s)
- When 1+ isotopes are accelerated they have the same KE and distance
- Energy Level 1 → 2 e- → 1s ( s- 1 orbital )
- Energy Level 2 → 8 e-. → 2s , 2p ( p - 3 orbitals )
- Energy Level 3 → 18 e- → 3s , 3p , 3d ( d - 5 orbitals )
- Energy Level 4 → 32 e- → 4s , 4p , 4d , 4f (f - 7 orbitals )
- Orbital - region of space where there is a high probability of finding an electron []
- Period number for S & P : Last Subshell (i.e. 2p , 3s)
- Period number for D: Last Subshell - 1 (i.e.(2-1)d)
- Group number / Charge : number of electrons (i.e. 3p3)
- The Pauli exclusion principle – electrons in the same orbital have opposite spins
- The Aufbau principle – Electrons fill orbitals in order of increasing energy (4s before 3d
etc)
- Hund’s rule – electrons fill orbitals in the same subshell singly before paring
- Copper and Chromium take one from 4s and give it to 3d ( Cr = [Ar] 4s1 3d4 )
- 4s electrons are lost first in ions because 3d becomes lower energy when it contains
electrons
- Isoelectronic - same electronic configuration
- Enthalpy change of ionisation - The energy required to remove one mole of electrons
from one mole of atoms to form one mole of ions in the gaseous state
, - Enthalpy change of ionisation ( First ionisation energy ) - X(g) → X+(g) + e-
- Ionisation energy is affected by : Nuclear charge , Atomic radius , Shielding
- Down the group : nuclear charge increases , atomic radius increases , shielding
increases
- Across the period : nuclear charge increases , atomic radius decreases , shielding is the
same
- Deviations across 1st ionisation energy across a period, 2-3 (p is higher energy, further
away, more shielded than s) and 5-6 (paired electron in 3p which repel)
- SIE > FIE ( electron being removed from +ve ion , closer to nucleus )
- Same atomic radius of isotopes because there are the same number of protons
- 2+ ion : m/z is halved , abundance is same
Exam Questions
Definitions:
- Define relative atomic mass: The average mass of an atom of an element compared to
1/12th the mass of an atom of carbon-12.
- Define first ionisation energy. The energy needed to remove one mole of electrons
from one mole of atoms, in the gaseous state to form one mole of ions.
TOF Mass Spectrometry:
- Why is it necessary to ionise the isotopes before they are analysed in a TOF mass
spectrometer? The ions will interact with and be accelerated by an electric field, the
ions will create a current when hitting the detector
- Describe how molecules are ionised using electrospray ionisation. The sample is
dissolved in a volatile solvent. This is then injected through a needle at high voltage so
that each particle gains a proton (H+). (X + H+ → XH+)
- Describe how molecules are ionised using electron impact ionisation. The sample
is bombarded by high energy electrons, using an electron gun so that the sample loses
an electron (as one electron is knocked out) forming X+.
- Outline how the TOF MS is able to separate two species. Positive ions are
accelerated by an electric field, to a constant kinetic energy. Lighter ions will move faster
, and heavier ions will move slower, so lighter ions will arrive at the detector first.
- Suggest what might cause the RAM of this sample to be different from the RAM
given in the periodic table. There may be other isotopes present.
- Why does the mass spectrum of 128 Te also have a small peak at m/z = 127? A 2+
ion is formed from the 128Te.
- Why do isotopes have the same chemical properties? They have the same number
of electrons.
- Explain how a current is generated in the TOF MS. Electrons are transferred and flow
from the detector plate to the positive ion. The current generated is proportional to the
abundance of the isotope.
Trends in Ionisation Energy:
- Explain the pattern in the first ionisation energies of the elements from lithium to
neon. The general trend is that across period 2, the first ionisation energy increases.
This is because there are more protons (/increased nuclear charge), and all of the
electrons are in the same energy level and there is similar shielding so there is a
, stronger attraction between the nucleus and the outer e-. The first ionisation energy of B
is lower than Be (which is a trend deviation), this is because the outer electron is in the
2p orbital, which has a higher energy than the 2s orbital, so is easier to remove. The first
ionisation energy of O is lower than N (which is a trend deviation), this is because there
are 2 electrons in the 2p orbital that need to pair, which causes repulsion.
- State and explain the trend in first ionisation energies of the elements in group 2
from magnesium to barium. The trend is that down the group, the first ionisation
energy decreases. This is because the ions get bigger (there are more energy shells) so
there is a weaker attraction of the ion to the lost electron.
- Explain why the SIE of boron is higher than the FIE of boron. The electron is being
removed from a positive ion, which is closer to the nucleus, therefore requires more
energy.
- Predict the element with the highest second ionisation energy. Sodium, as the
electron is removed from the 2p orbital.
- Explain why the ionisation energy of every element is endothermic. Heat/energy is
needed to overcome the attraction between the negative electron and the positive
nucleus.
- Predict whether an atom of 88Sr will have an atomic radius that is larger than,
smaller than or the same as the atomic radius of 87Rb. Smaller, since it has a bigger
nuclear charge and similar shielding.
- Explain why the SIE of calcium is lower than the SIE of potassium. In Ca+ the outer
electron is further from the nucleus and there is more shielding.
3.1.2 Amount of substance
Things to remember :
- Add Mr inside brackets and then multiply
- 1000 cm^3 = 1 dm^3
- Cm3 → m3 = /1000000
- Cm3 is smaller than dm3
- When calculating atom economy use molar ratio (aka big numbers)
- pV = nRT , v in m3 and p in Pa and T in K
- If you have 2HCl, you should use moles/2 for the pV=nRT equation if it is limiting when
n/2
- When writing oxygen in balanced equation write O2
- Oxygen gas = O2
- When finding the moles of something do the whole thing (use the molar ratio , e.g. if
asks for moles of 2O2 , times molar ratio by 2 )
- Water of crystallisation , the ‘.’ means that there are water molecules are in the
crystalline structure but not part of the compound
- Mg → g = / 1000
Notes
3.1.2 Amount of substance (practical - prep of standard solution and titration) :
- Avogadro's Constant = 6.02 x 1023 , Mr is 6.02 x 1023 units of something
https://chemrevise.org/wp-content/uploads/2021/02/3.7-revision-guide-naming-and-isomerism.p
df
3.1 Physical Chemistry
3.1.1 Atomic structure
Things to remember :
- Atomic radius : lower on the table , closest to the left of the table = largest radius
- Copper and chromium electron configurations are deviations , take 1 from 4s and add it
to 3d
- When electronic configuration of ions take from 4s before 3d
- Deviations across 1st ionisation energy of period , 2-3 and 5-6 ,
- Second ionisation energy is greater than the first ionisation energy
- Second ionisation energy deviations are shifted 1 to the right ( so 2nd IE period 3
deviations = Al and Si , S and Cl )
- Atomic radius is smaller as you go across the period because higher nuclear charge so
more attraction so closer, but same in isotopes because same no protons
- Mr of a molecule = the highest m/z value on the mass spectrum
- The orbitals are the boxes and subshells are 1s , 2p etc Z is proton number and A is
mass number
- Electron impact is for small Mr and can fragment
- In mass spectrum - lower m/z peaks are due to fragmentation , molecular ion peak is
100% and m+1 ion peak is due to 13C or 2H
- If it asks for the species responsible for a peak then use X+ rather than X to show that it
is a 1+ ion.
- The Pauli exclusion principle – electrons in the same orbital have opposite spins
- The Aufbau principle – Electrons fill orbitals in order of increasing energy (4s before 3d
etc)
- Hund’s rule – electrons fill orbitals in the same subshell singly before paring
Notes
3.1.1 Atomic Structure :
- Atomic structure : JJ Thomson - Plum pudding Model ( cloud of +ve charge, electrons
scattered ) → Ernest Rutherford - Nuclear Model ( fired alpha particles (+ve) at gold foil,
most passed through, some deflected at large angles > mostly empty space, positive
charge concentrated in nucleus ) → Niels Bohr - (electrons orbit around the nucleus
in energy levels) → Erwin Schrodinger - Quantum model ( electrons move in waves )
- Isotope : Atoms of the same element with the same number of protons but different
numbers of neutrons. They have the same chemical properties ( same electronic
configuration ) and different physical properties ( different masses )
- Relative Atomic Mass : The average mass of the naturally occurring isotopes of an
element compared to 1/12th the mass of an atom of carbon-12.
,- RAM = Sum of ( isotopic mass * relative abundance ) / 100
- Mass spectrometry is used for : determination of mass and abundance of each
isotope of an element to find the RAM ( and to find the RMM of substances made of
molecules )
- Electron Impact : For Low Formula Mass ( inorganic/organic ) Vapourised → High
energy electrons are fired at the sample from electron gun ( hot wire filament with a
current which emits electrons ) → Knocks off 1 electron → X(g) + e- → X+(g) + 2e- , X+(g) (
molecular ion ) - can fragment → the 1+ ions are attracted to the -ve electric plate
where they are accelerated
- Electrospray : High Molecular Mass ( biological molecules ), Dissolved in volatile
solvents (water/methanol)→ Injected into a fine hypodermic needle to give a fine mist
→ needle is attached to the positive end of a high voltage power supply → particles gain
a proton. X(g) + H+ → XH+(g) , solvent evaporates while XH+ ions are attracted to a
negative plate
- Acceleration : All have the same KE due to the electric field KE = ½ mv2 , the velocity
depends on the mass ( lighter particles will have a faster velocity )
- Detection : positive ion hits -ve plate → they are discharged by gaining an electron
from the plate → generates a movement of electrons and hence a current → size of
current is proportional to number of ions hitting the plate
- Electrospray mass spectrum - m/z = mass of each ion ( -1 from m/z to find RMM )
Electron impact mass spectrum - greatest m/z is the molecular ion (m/z = molecular
mass) , small peaks due to molecular ions containing different isotopes ( C13 and H2 )
and there may be fragmentation
- Mass of one ion (kg) = Relative isotopic mass * 10-3 / (6.022 x 1023 )
- KE = ½ mv2( where m is in kg) , v = d/t - (m/s)
- When 1+ isotopes are accelerated they have the same KE and distance
- Energy Level 1 → 2 e- → 1s ( s- 1 orbital )
- Energy Level 2 → 8 e-. → 2s , 2p ( p - 3 orbitals )
- Energy Level 3 → 18 e- → 3s , 3p , 3d ( d - 5 orbitals )
- Energy Level 4 → 32 e- → 4s , 4p , 4d , 4f (f - 7 orbitals )
- Orbital - region of space where there is a high probability of finding an electron []
- Period number for S & P : Last Subshell (i.e. 2p , 3s)
- Period number for D: Last Subshell - 1 (i.e.(2-1)d)
- Group number / Charge : number of electrons (i.e. 3p3)
- The Pauli exclusion principle – electrons in the same orbital have opposite spins
- The Aufbau principle – Electrons fill orbitals in order of increasing energy (4s before 3d
etc)
- Hund’s rule – electrons fill orbitals in the same subshell singly before paring
- Copper and Chromium take one from 4s and give it to 3d ( Cr = [Ar] 4s1 3d4 )
- 4s electrons are lost first in ions because 3d becomes lower energy when it contains
electrons
- Isoelectronic - same electronic configuration
- Enthalpy change of ionisation - The energy required to remove one mole of electrons
from one mole of atoms to form one mole of ions in the gaseous state
, - Enthalpy change of ionisation ( First ionisation energy ) - X(g) → X+(g) + e-
- Ionisation energy is affected by : Nuclear charge , Atomic radius , Shielding
- Down the group : nuclear charge increases , atomic radius increases , shielding
increases
- Across the period : nuclear charge increases , atomic radius decreases , shielding is the
same
- Deviations across 1st ionisation energy across a period, 2-3 (p is higher energy, further
away, more shielded than s) and 5-6 (paired electron in 3p which repel)
- SIE > FIE ( electron being removed from +ve ion , closer to nucleus )
- Same atomic radius of isotopes because there are the same number of protons
- 2+ ion : m/z is halved , abundance is same
Exam Questions
Definitions:
- Define relative atomic mass: The average mass of an atom of an element compared to
1/12th the mass of an atom of carbon-12.
- Define first ionisation energy. The energy needed to remove one mole of electrons
from one mole of atoms, in the gaseous state to form one mole of ions.
TOF Mass Spectrometry:
- Why is it necessary to ionise the isotopes before they are analysed in a TOF mass
spectrometer? The ions will interact with and be accelerated by an electric field, the
ions will create a current when hitting the detector
- Describe how molecules are ionised using electrospray ionisation. The sample is
dissolved in a volatile solvent. This is then injected through a needle at high voltage so
that each particle gains a proton (H+). (X + H+ → XH+)
- Describe how molecules are ionised using electron impact ionisation. The sample
is bombarded by high energy electrons, using an electron gun so that the sample loses
an electron (as one electron is knocked out) forming X+.
- Outline how the TOF MS is able to separate two species. Positive ions are
accelerated by an electric field, to a constant kinetic energy. Lighter ions will move faster
, and heavier ions will move slower, so lighter ions will arrive at the detector first.
- Suggest what might cause the RAM of this sample to be different from the RAM
given in the periodic table. There may be other isotopes present.
- Why does the mass spectrum of 128 Te also have a small peak at m/z = 127? A 2+
ion is formed from the 128Te.
- Why do isotopes have the same chemical properties? They have the same number
of electrons.
- Explain how a current is generated in the TOF MS. Electrons are transferred and flow
from the detector plate to the positive ion. The current generated is proportional to the
abundance of the isotope.
Trends in Ionisation Energy:
- Explain the pattern in the first ionisation energies of the elements from lithium to
neon. The general trend is that across period 2, the first ionisation energy increases.
This is because there are more protons (/increased nuclear charge), and all of the
electrons are in the same energy level and there is similar shielding so there is a
, stronger attraction between the nucleus and the outer e-. The first ionisation energy of B
is lower than Be (which is a trend deviation), this is because the outer electron is in the
2p orbital, which has a higher energy than the 2s orbital, so is easier to remove. The first
ionisation energy of O is lower than N (which is a trend deviation), this is because there
are 2 electrons in the 2p orbital that need to pair, which causes repulsion.
- State and explain the trend in first ionisation energies of the elements in group 2
from magnesium to barium. The trend is that down the group, the first ionisation
energy decreases. This is because the ions get bigger (there are more energy shells) so
there is a weaker attraction of the ion to the lost electron.
- Explain why the SIE of boron is higher than the FIE of boron. The electron is being
removed from a positive ion, which is closer to the nucleus, therefore requires more
energy.
- Predict the element with the highest second ionisation energy. Sodium, as the
electron is removed from the 2p orbital.
- Explain why the ionisation energy of every element is endothermic. Heat/energy is
needed to overcome the attraction between the negative electron and the positive
nucleus.
- Predict whether an atom of 88Sr will have an atomic radius that is larger than,
smaller than or the same as the atomic radius of 87Rb. Smaller, since it has a bigger
nuclear charge and similar shielding.
- Explain why the SIE of calcium is lower than the SIE of potassium. In Ca+ the outer
electron is further from the nucleus and there is more shielding.
3.1.2 Amount of substance
Things to remember :
- Add Mr inside brackets and then multiply
- 1000 cm^3 = 1 dm^3
- Cm3 → m3 = /1000000
- Cm3 is smaller than dm3
- When calculating atom economy use molar ratio (aka big numbers)
- pV = nRT , v in m3 and p in Pa and T in K
- If you have 2HCl, you should use moles/2 for the pV=nRT equation if it is limiting when
n/2
- When writing oxygen in balanced equation write O2
- Oxygen gas = O2
- When finding the moles of something do the whole thing (use the molar ratio , e.g. if
asks for moles of 2O2 , times molar ratio by 2 )
- Water of crystallisation , the ‘.’ means that there are water molecules are in the
crystalline structure but not part of the compound
- Mg → g = / 1000
Notes
3.1.2 Amount of substance (practical - prep of standard solution and titration) :
- Avogadro's Constant = 6.02 x 1023 , Mr is 6.02 x 1023 units of something
