exam questions Flashcards
The sample of chromium is analysed in a time of flight (TOF) mass spectrometer.
Give two reasons why it is necessary to ionise the isotopes of chromium before they can be analysed in a TOF mass spectrometer.
- (Ions will interact with and) be accelerated (by an electric field)
- Ions create a current when hitting the detector
The first ionisation energies of the elements in Period 2 change as the atomic
number increases.
Explain the pattern in the first ionisation energies of the elements from lithium to
neon.
Stage 1: General Trend (Li → Ne)
1a. 1st IE increases
1b. More protons/increased nuclear charge
1c. Electrons in same energy level / shell
1d. No extra/similar shielding
1e. Stronger attraction between nucleus and outer e OR outer e closer to
nucleus
Stage 2: Deviation Be → B
2a. B lower than Be
2b. Outer electron in (2)p
2c. higher in energy than (2)s
Stage 3: Deviation N → O
3a. O lower than N
3b. 2 electrons in (2)p need to pair
3c. pairing causes repulsion
State which of the elements magnesium and aluminium has the lower first
ionisation energy.
Explain your answer.
Al
(Outer) electron in (3)p sublevel / orbital
Higher in energy / further from the nucleus
so easier to remove
State how, if at all, the chemical properties of these isotopes differ.
Give a reason for your answer.
Chemical properties
No difference in chemical properties
reason
Because all have the same electronic structure (configuration)
A TOF mass spectrometer can be used to determine the relative molecular
mass of molecular substances.
Explain why it is necessary to ionise molecules when measuring their mass
in a TOF mass spectrometer.
Ions, not molecules, will interact with and be accelerated by an electric field
Only ions will create a current when hitting the detector
Outline how the TOF mass spectrometer is able to separate these two
species to give two peaks.
Positive ions are accelerated by an electric field
To a constant kinetic energy
The positive ions with m / z of 104 have the same kinetic energy as
those with m / z of 118 and move faster
Therefore, ions with m / z of 104 arrive at the detector first
The melting point of XeF4 is higher than the melting point of PF3
Explain why the melting points of these two compounds are different.
In your answer you should give the shape of each molecule, explain why each
molecule has that shape and how the shape influences the forces that affect the
melting point.
Stage 1 electron pairs
-XeF4 4BP and 2LP around Xe
-PF3 3BP and 1LP around P
Stage 2 explanation of shapes
- XeF4 is square planar
-PF3 is pyramidal (allow tetrahedral)
-Electron pairs repel as far as possible or Lone pair repels more than bonding
pairs
Stage 3 IMF
-XeF4 has vdw forces and PF3 has dipole-dipole forces (and vdw)
-Stronger/more intermolecular forces in XeF4
Due to larger Mr or more electrons or larger molecules or packs more closely
together
Sodium fluoride contains sodium ions (Na+) and fluoride ions (F–).
Na+ and F– have the same electron configuration.
Explain why a fluoride ion is larger than a sodium ion.
(Electrostatic) forces of attraction between
oppositely charged ions/Na+ and F–
Lots of energy needed to overcome/break forces
Explain, in terms of structure and bonding, why the melting point of sodium
fluoride is high
Fluoride ion has (two) fewer protons/lower nuclear charge
Weaker attraction between nucleus and (outer) electrons
Methoxymethane (CH3OCH3) is an isomer of ethanol.
The table shows the boiling points of ethanol and methoxymethane.
Compound Boiling point / °C
ethanol 78
methoxymethane −24
In terms of the intermolecular forces involved, explain the difference in boiling points.
Hydrogen bonds (between ethanol molecules)
(permanent) dipole-dipole OR van der Waals force (between methoxymethane molecules)
Hydrogen bonds are stronger/est intermolecular force
Describe the bonding in magnesium.
Attraction between (lattice of) Mg2+ ions and delocalised electrons
Explain, in terms of structure and bonding, why magnesium chloride has a
high melting point.
(Giant) ionic lattice / lots of Mg2+ and Cl– ions
Strong (electrostatic) forces of attraction
Between Mg2+ and Cl– ions
Describe the structure and bonding in magnesium
(giant) lattice of (Mg2+) cations / (giant) lattice of (Mg) atoms
(Electrostatic) attractions between cations / Mg2+ ions / nuclei and
delocalised electrons
State the trend in the atomic radius of the elements down Group 2 from Mg to Ba
Give a reason for this trend.
Trend
Reason
Trend: increases
Reason: the number of electron energy levels increases
This question is about the shapes of molecules.
Discuss the difference between the shapes of CF4 and XeF4
In your answer you should:
* name the shape of each molecule
* explain the shape of each molecule
* explain the bond angle(s) in each molecule.
Stage 1: Name of each shape
1a: CF4 = Tetrahedral
1b: XeF4 = Square Planar
Stage 2: Explanation of shape / bond angle
in CF4
2a: four bonding pairs (and zero lone pairs)
2b: electron pairs repel each other to be as far
apart as possible / electron pairs repel
each other equally
2c: so bond angle is 109.5 °
Stage 3: Explanation of shape / bond angle
in XeF4
3a: four bonding pairs
3b two lone pairs
3c: so bond angle is 90 °
3d: lone pairs repel more than bonding pairs,
3e: so lp as far apart as possible / so lone
pairs are opposite each other /
180 ° apart
Explain why the electronegativity of the halogens decreases down the group.
M1 Larger atoms / more electron shells / more shielding / bonding pair
of electrons further from the nucleus
M2
weaker attraction between nucleus and bonding pair of electrons /
weaker attraction between nucleus and shared pair of electrons /
weaker attraction between nucleus and electron density in the
covalent bond
Explain why the third ionisation energy of magnesium is much higher than the
second ionisation energy of magnesium
The electron is removed from 2p sub-shell / 2nd energy level / lower energy level / sub-shell that is closer to the nucleus
(Electron being removed is) less shielded (than 3s)
The student uses this method to complete a titration.
* Rinse a burette with distilled water.
* Fill the burette with sodium hydroxide solution.
* Use a measuring cylinder to transfer 25 cm3 of the citric acid solution into a
conical flask.
* Add 5 cm3 of indicator.
* Slowly add the sodium hydroxide solution from the burette into the conical flask.
* Add the sodium hydroxide solution dropwise near the end point until the indicator
just changes colour.
* Repeat the titration to get concordant results.
The method used by the student includes three practical steps that will lead to an
inaccurate final result.
For each of these three steps
* identify the mistake
* explain why it is a mistake
* suggest how the mistake can be overcome
Use best three of these four stages
Stage 1
a. Problem – using a measuring cylinder
b. Explanation – large uncertainty / not
accurate enough
c. Improvement – use a (volumetric)
pipette (Not dropping pipette)
Stage 2
a. Problem – too much indicator
b. Explanation – may react and affect the
endpoint reading
c. Improvement – use a smaller volume
(2-6 drops)
Stage 3
a. Problem – rinsing the burette with
distilled or deionised water
b. Explanation – will slightly dilute the
alkali solution
c. Improvement – rinse the burette with
alkali solution
Stage 4
a. Problem – adding alkali solution until
the indicator “just” changes colour
b. Explanation – acid may not have fully
reacted (as mixture not swirled)
c. Improvement – add alkali solution until
a permanent colour change is seen.
5 Suggest, in terms of the intermolecular forces for each compound, why CBr4 has a
higher boiling point than CHBr3
M1 CBr4 has van der Waals’ forces between molecules
M2 CHBr3 has van der Waals’ forces and dipole-dipole intermolecular
forces
M3 The van der Waals’ between CBr4 molecules are stronger than the
dipole-dipole and van der Waals’ forces between CHBr3 (because it
has a larger mass/more electrons/larger electron cloud)
