topic 2 - bonding Flashcards
Ammonia and boron trifluoride react to form a compound NH3BF3 which contains a dative
covalent bond. Each of the molecules, NH3 and BF3, has a different feature of its electronic
structure that allows this to happen. Use these two different features to explain how a dative
covalent bond is formed.
(2)
- donation of lone pair of electrons from nitrogen / lone pair from ammonia
- to the boron atom which is electron deficient
State what is meant by the term electronegativity and hence explain the polarity, if any, of
the bonds in chlorine trifluoride, ClF3.
(3)
- electronegativity is the relative ability of an atom to attract the bonding electrons in a covalent bond
- flouring is more electronegative than chlorine/flourine is the most electronegative
- so flourine is delta negative and chlorine is delta positive
Explain why the melting temperature of silicon(IV) oxide is much higher than that of iodine,
even though the bonding in both is covalent.
(3)
- silicon(IV) oxide is a giant structure therefore contains many strong covalent bonds
- iodine only has weak intermolecular/ london forces
- more energy is required to break the stronger bonds in silicon(IV) oxide hence higher melting temperature
This is a question about water.
Explain why both water and carbon dioxide molecules have polar bonds but only
water is a polar molecule.
- oxygen is more electronegative than hydrogen and carbon
- which results in a polar bond with oxygen delta negative so carbon and hydrogen hydrogen delta positive
- carbon dioxide is a symmetrical/linear molecule and so the dipole moments cancel
- the lone pairs of electrons of oxygen and the shape of the water molecules mean that the dipole moments do not cancel
Explain, in terms of the structure and bonding of each element, the difference between these
values.
(3)
- silicon is a giant covalent structure and contains covalent bonds
- chlorine is a simple molecular structure and contains london forces
- covalent bonds in silicon are stronger than london forces in chlorine and covalent bonds take more energy to break than london forces
This question is about crystalline solids.
Iodine and diamond are crystalline solids at room temperature.
Explain why diamond has a much higher melting temperature than iodine.
(5)
- iodine is a simple molecular
- diamond is a giant covalent lattice structure with 4 covalent bonds per carbon atom
- iodine molecules are held together by weak london forces
- carbon atoms in diamond are held together by strong covalent bonds
- strong covalent bonds require more energy to break them than intermolecular forces
Boron and aluminium are in the same group of the Periodic Table. Both form compounds
with chlorine and with fluorine.
Aluminium fluoride and aluminium chloride are both crystalline solids at room temperature.
Aluminium fluoride sublimes to form a gas at 1291°C (1564 K), whilst aluminium chloride
sublimes at 178°C (451 K).
Use the Pauling electronegativity values in the Data Booklet to explain these differences in
sublimation temperature.
(6)
- aluminium and chlorine electronegativity difference 1.5 AND aluminimum and flourine electronegativity difference 2.5
- aluminium chloride mostly covalent small molecule
- aluminium flouride bonds are more polar
- aluminium chloride is molecular so weaker london forces
- aluminium flouride is a giant structure which has stronger electrostatic forces of attraction between the ions
- more energy needed to break the stronger bonds to cause sublimation in aluminium flouride
This is a question about water.
Water might be expected to have a lower boiling temperature than hydrogen sulfide
but it actually has a higher boiling temperature.
Comment on this statement by referring to the intermolecular forces in both these
substances.
A detailed description of how the intermolecular forces arise is not required.
(4)
- a lower boiling point is expected because water has fewer electrons than hydrogen sulfide
- water has weaker, less london forces
- a higher boiling point occurs because water has hydrogen bonding
- hydrogen bonding is stronger than london forces and requires more energy to break and results in a higher boiling temperature
This question is about compounds of Group 5 elements.
Nitrogen trichloride, NCl3, has a boiling temperature of 344 K, and
nitrogen trifluoride, NF3, has a boiling temperature of 144 K.
Explain this difference in boiling temperatures, by referring to all the
intermolecular forces present. (5)
- london forces are greater in NCl3
- as NCl3 has more electrons
- permanent dipole-dipole forces stronger in NF3 than NCl3
- as F is more electronegative than Cl
- london forces are more significant / more energy required to break the intermolecular forces between the NCl3 molecules than NF3
Explain why the boiling temperatures increase from chlorine to iodine.
(2)
- from chlorine to iodine, the number of electrons increases
- so the strength of the london / instantaneous dipole-dipole forces and more energy is needed to separate the molecules
The compounds hydrogen fluoride, water and methane, all have simple molecular
structures, but they have significantly different boiling temperatures.
Discuss the reasons for the differences in the boiling temperatures of the three compounds,
using the data in the table and the Pauling electronegativity values in the Data Booklet. (6)
- same number of electrons so similar london forces
-large electronegativity differences in HF and H2O and small in CH4 - only weak london forces in CH4
- hydrogen bonding in both HF and H2O but not CH4
- more hydrogen bonds/twice as many hydrogen bonds in H2O than in HF
- more energy needed to break stronger intermolecular forces
Explain why methanol and water are ‘soluble in all proportions’.
You must include a diagram in your answer.
(3)
- methanol hydrogen bonds to water
- at least one lone pair on an oxygen atom and an approximate 180 degrees OHO bond angle
- strength of all intermolecular forces between methanol and water is approximately the same as those in water and methanol
or strength/extent of H-bonding between methanol and water is the same
This is a question about hydrocarbons.
Explain why 2,2-dimethylpropane has a much lower boiling temperature than its isomer
pentane.
Detailed descriptions of the forces involved are not required.
(2)
- branching results in fewer/weaker london forces
- due to less surface area / points of contact
This question is about some halogens and their compounds.
The intermolecular attractions between halogen molecules are London forces.
(i) Describe how London forces form between halogen molecules.
(3)
setting up of the dipole
- uneven distribution of electrons
type of dipole
- results in an instantaneous dipole in the first molecule
induction od a second dipole
- causes/induces a second dipole on other molecule
Methanol, CH3OH, is miscible with water in all proportions.
Sodium chloride is much less soluble in methanol than in water.
Explain these statements using your knowledge of the interactions between solutes and
solvents.
You must use diagrams to illustrate your answers.
(6)
- hydrogen bonding between water/solvent and methanol/solvent
- suitable diagram
- same strength/comparable to bonding in either component on its own OR hydrogen bonding is present in methanol and in water
- hydration of Na+ and Cl-
- suitable diagram of at least one ion
- the ionic bonding is stronger than the bonding between sodium and/or chloride ions and methanol
Explain why hydrogen bonding causes ice to be less dense than liquid water.
(2)
- more open /more space between molecules making it less dense
- due to 3D lattice/ ring structure on ice
- hydrogen bonds longer than covalent bonds
In Stage 1, phenol is nitrated using dilute nitric acid.
The nitration of benzene requires concentrated nitric acid at 55°C with a catalyst of
concentrated sulfuric acid.
Both these reactions are electrophilic substitution.
(i) Explain why phenol can be nitrated using milder conditions than benzene.
(2)
- the electron density of the benzene ring is greater in phenol than in benzene
- because the lone pair of electrons on oxygen and overlaps with the pi cloud/ delocalised system
A mixture of 2-nitrophenol and 4-nitrophenol is produced in Stage 1.
They are separated by steam distillation.
The boiling temperature of 2-nitrophenol is 215°C and that of 4-nitrophenol is 279°C.
Explain, in terms of intermolecular forces, why 4-nitrophenol has a higher boiling
temperature than 2-nitrophenol.
You may include a diagram in your answer.
(2)
- they both form hydrogen bonds
- in 4-nitrophenol the hydrogen bonds join molecules in a straight chain OR 2-nitrophenol forms intermolecular hydrogen bonds so fewer intermolecular hydrogen bonds
This question is about chlorine and its compounds.
In many swimming pools, sodium chlorate(I) has replaced chlorine gas as a disinfectant.
Sodium chlorate(I) is an ionic compound. It is very soluble in water.
(i) Describe, using diagrams to illustrate your answer, the interactions between
each of the ions and the solvent when sodium chlorate(I) dissolves in water.
(2)
- diagram showing both Na+ ion and ClO- ion surrounded by water molecules/solvated
- correct orientation of the water molecules around both ions with a relevant dipole shown on at least one water molecule for each ion
Ethanol is very soluble in water whereas chloroethane is almost insoluble in water.
Explain this observation by comparing the types of intermolecular forces formed
by each of these molecules with water.
(2)
- ethanol forms hydrogen bonds with water
- chloroethane forms permanent dipole-dipole attractions and london forces with water
Nitrogen forms several hydrides. In addition to ammonia, NH3, it forms hydrazine, N2H4, in
which the two nitrogen atoms are covalently bonded together.
Hydrazine is very soluble in water.
Explain, using a labelled diagram and naming the relevant intermolecular interactions, why
hydrazine is very soluble in water.
(3)
- hydrogen bonds in water , hydrazine
- diagram showing hydrogen bond between the correct atoms
- lone pair on either nitrogen or oxygen and bond angle shown on diagrams as approximately 180 degrees.
This question is about dissolving different compounds.
The solubility of two compounds in different solvents was investigated.
A summary of the findings is shown.
Explain the findings of the investigation by considering the interactions between the
compounds and each of the solvents.
(6)
- 2-methylpentane is insoluble in water as it cannot hydrogen bond to water as non of the hydrogen atoms are electronegative
- 2-methylpentane is soluble in hexane as london forces in both compounds are similar in strength and size
- so resultant force in mixture are similar in magnitude to those in each liquid
- potassium bromide is soluble in water as its ions are hydrated when dissolved
- the enthalpy change of hydration is greater than the energy needed to break apart the lattice
- potassium bromide is insoluble in hexane as any london forces that form between hexane would be smaller in magnitude than the forces between the atoms
This question is about structure and bonding.
Water has two significant anomalous properties:
* it has a higher melting temperature than hydrogen sulfide, H2S, even though it has
fewer electrons in its molecules
* the density of ice at 0 °C is less than that of water at 0 °C.
Explain these properties.
You should include a labelled diagram to show the intermolecular forces between two
molecules of water.
(6)
- lone pair and dipole lone pair on oxygen in hydrogen bons and dipole shown with delta + on any one H and delta - on any one O
- shape hydrogen bond labelled / shown as a dotted line and hydrogfen bonds shows as approximately linearor O-H-O bond angle labelled as 180 degrees
- london forces
hydrogen sulfide has stronger london forces because it has more electrons - comparison
hydrogen bonding is stronger than london forces and requires more energy to overcome - ice at 0 degrees
water molecules are arranged in a lattice or hydrogen bonds are longer than covalent bonds - water at 0 degrees
water molecules get closer/have less distance between them so more molecules in the same volume
This question is about catalytic converters.
Catalytic converters contain metals such as platinum.
Describe the bonding in platinum.
You may include a diagram in your answer.
(2)
- regular arrangement / lattice of positive ions/metal cations in a sea of delocalised electrons
- metallic bonding is the strong electrostatic force of attracted between cations and electrons
Using the data provided, explain how changes in the cation affect the bond strength in an
ionic compound.
(2)
- the higher the charge on the cation the stronger the attraction between the ions and mention of a 2+ cation in CaF2 compared to 1+ cation in LiF
- the smaller the radius of the cation the stronger the attraction between ions and mention of Li+ being smaller than K+
State all the conditions under which magnesium bromide conducts electricity.
(1)
molten or dissolved in water
Explain why the electrical conductivity of solid potassium bromide is poor but an
aqueous solution of potassium bromide is a good electrical conductor.
(2)
- solid potassium bromide does not conduct because the ions are in fixed positions/ions are not free to move
- it does conduct electricity in solution because the ions are free to move and carry charge
Which diagram shows the trend in ionic radius for the isoelectronic ions N3−
to Al3+?
(1) explain your answer (2)
- increase in number of protons in the nucleus
-increases the attraction for the electrons bringing them closer to the nucleus
Explain the difference in the melting temperatures of magnesium oxide and potassium
bromide.
(3)
- comparison of ionic charges ( MgO has doubly charged ions and KBr has singly charged ions
- comparison of ionic radii ( Mg 2+ smaller than K+ and/or O2- smaller than Br-
- comparison of energy required ( more energy needed to overcome the electrostatic forces of attraction in MgO than in KBr/ion polarization
DO NOT MENTION MOLECULES/INTERMOLECULAR FORCES/HYDROGEN BONDING/COVALENT BONDING
Element X has the typical appearance of a metal.
Predict two other distinct physical properties that element X would exhibit if it is a metal.
Explain your choices in terms of structure and bonding.
(4)
- high melting/boiling temperature because of strong electrostatic attraction between metal ions and delocalised electrons
- good electrical/thermal conductivity because of mobile delocalised electrons
- malleability/ductility because the layers of the ions/atoms can easily slide over eachother
- high density because the ions/atoms are tightly packed due to the strong attraction between them
The compressive strength is a measure of the energy required to break
some of the bonds within a substance.
Deduce possible reasons why there are two widely different values for the compressive
strength of graphite.
Both the values (2.3 and 15.3 GPa) are valid experimental results.
(2)
- lower value relates to weak london forces
- higher value refers to strong covalent bonds within each layer
Deduce two possible reasons why the density of iron (7.86 g cm−3
) is much greater
than the density of graphite (2.2 to 2.8 g cm−3
).
(2)
- iron atoms have greater mass than carbon atoms
- iron atoms pack closer than carbon atoms in graphite
This question is about crystalline solids.
Graphite is also a crystalline solid at room temperature.
Unlike diamond, graphite conducts electricity.
Describe the key feature of the bonding of the carbon atoms in graphite that results in it
being an electrical conductor.
(2)
- one electron free to move / delocalised within the layer to carry the charge
- each carbon is covalently bonded to three other carbons OR the carbon atoms are arranged in layers which allow the flow of electricity through them
Graphene, graphite and diamond are all forms of solid carbon.
Explain, in terms of structure and bonding, why graphene and graphite are good
electrical conductors but diamond is a poor electrical conductor.
You may include labelled diagrams in your answer.
(6)
- graphene has a single layer/ single sheet of hexagons/rings
- graphene has delocalised electrons that are free to move ( mobile )
- graphite has layers and each carbon is bonded to 3 others
- graphite has delocalalised electrons that are mobile/ free to move
- diamond has each carbon bonded to 4 other carbons/ diamond has a tetrahedral arrangement around each carbon structure
- diamonds carbon atoms have all their outer electrons involved in bonding ( no delocalised electrons )
Explain why the bond angle in water is less than the bond angle in ammonia.
(2)
- oxygen has one more lone pair of electrons than nitrogen/ oxygen has 2 lone pairs of electrons but nitrogen only has one
- so the repulsion from the oxygen lone pairs is greater and reduces the bond angle OR lone pair - lone pair repulsion is greater than lone pair - bonded pair and reduces the bond angle
Explain the shape of the PCl3 molecule. The bond angle is not required.
(3)
- PCl3 is a trigonal pyramidal
- has 3 bond pairs and 1 lone pair around the central P atom
- electron pair repels to positions of minimum repulsion /maximum separation
Explain why phosphorus forms PCl5 but nitrogen does not form NCl5.
(2)
- phosphorus can expand its octet/ can expand its outer shell/ can accommodate 10 electrons / has available 3d orbitals for promotion of electrons
- nitrogen does not have 2d orbitals can only accommodate eight electrons in its outer shell
Explain why a phosphorus(III) chloride molecule has this shape and bond angle.
(2)
- pyramidal because there is 3 bond pairs and 1 lone pair of electrons around central P atom and these are arranged to minimise repulsion
- bond angle less than 109.5 as lone-pair-bond pair repulsion is greater than bond pair-bond pair repulsion
why is the sigma bond stronger than the pi bond (2)
- p orbitals are parallel so poor overlap when pi bonds form
OR
-overlap of orbitals in sigma bond is along the line between the two nuclei - whereas in the pi bond , there is sideways overlap
Explain why but-1-ene does not exhibit E-Z isomerism.
(1)
- one C on the double bond has 2 hydrogen atoms attached to it
OR - C on one end of the double bond is not attached to two different atoms or groups
Describe the result of the test for the presence of a C C bond in E-but-2-ene
using bromine water. Give the displayed formula of the organic product.
(2)
(Bromine water goes from brown/ redbrown / yellow/ orange to) colourless
OR
(Bromine water is) decolorised
Another test for C C bonds is the reaction with acidified potassium manganate(VII).
Describe the result of this test using but-1-ene and give the displayed formula of
the organic product.
(Colour change purple/ purple-pink / pink
to) colourless
OR
(KMnO4 is) decolorised
The electrical conductivity of pure silicon is very low. Explain why this is so in
terms of the bonding
-silicon’s (outer) electrons are fixed (in covalent bonds)
silicon’s (outer) electrons are involved in bonding
- silicon’s electrons are not
free (to move)/silicon has no delocalized electrons/silicon’s electrons cannot flow
Suggest why aqueous solutions of calcium chloride, CaCl2(aq), and barium
chloride, BaCl2(aq), of the same molar concentration, have different electrical
conductivities.
(1)
The cations / barium and calcium (ions) are
different sizes
Some buildings are made from limestone, which is mainly calcium carbonate. Gases
in the atmosphere such as sulfur dioxide, SO2, and nitrogen dioxide, NO2, can be
responsible for damaging these buildings.
Describe how these gases come to be present in the atmosphere and explain how
they can damage a limestone building
(Sulfur / nitrogen oxides) form when fossil fuels are burnt when petrol or diesel burn in vehicle engine or volcanoes
/ lightning (1)
They (react with water to) form sulfuric acid /nitric acid /acid rain (1)
Which reacts with limestone (to form soluble compounds) / limestone and acid take part in
neutralisation / dissolves building / corrodes building (1)
The lattice energy of calcium chloride, CaCl2, is −2258 kJ mol−1 based on an
experimental Born-Haber cycle and −2223 kJ mol−1
based on theoretical calculations.
Would you expect its bonding to match the ionic model? Justify your answer.
(1)
Yes, as the values match closely (so little
deviation from ionic model)
Or
no, as the values are (slightly) different so a
degree of covalency / not fully ionic
Explain why the boiling temperature of HF is the highest in the series.
(2)
- hydrogen bonds also present
- which are stronger than london forces
Explain why the values of the boiling temperatures for Cl2,CH3,Cl and HCl do not
follow the same trend as F2,CH3,F and HF.
(1)
- HCl doesnt have hydrogen bonds
Which of the following statements is evidence for the existence of ions in ionic compounds?
In electron density maps for ionic compounds, there is no single line
representing electron density that surrounds both cations and anions.
Describe the structure of a metal.
(2)
- lattice of positively charged ions in a sea of delocalised electrons
Describe the bonding in a metal.
(2)
(Electrostatic) attraction between positive ions and delocalized electrons
/ sea of electrons
Explain why the melting temperature of magnesium (650 °C) is much higher than that of sodium (98 °C).
(3)
- magnesium ion / Mg2+ has a larger charge density than the sodium ion ( Na+ )
- magnesium ions are smaller than sodium ions
- Mg2+ contributes 2 more electrons per atom to the sea of delocalised electrons
- Mg2+ has a greater attraction for the deloaclised sea of electrons
Suggest two forces of repulsion which exist in an ionic lattice.
(2)
- ions of the same charge repel
- nuclei of the atoms repel
bonding and structure of magnesium oxide (1)
giant ionic
predict two other distinct physical properties that element x would exhibit if it is a metal. explain choices in terms of structure and bonding (4)
- high melting/ boiling temperature (1)
- strong (electrostatic) attraction between metal ions and delocalised electrons (1)
- (good) electrical conductivity/ thermal conductivity(1)
- mobile delocalised electrons (1)
- malleability/ ductility (1)
- the layers of ions/ atoms can easily slide over each
other (1) - high density (1)
- the ions/ atoms are tightly packed due to the strong attraction between them
Explain why the O–H and S–H bond lengths are different (3)
- sulfur is a larger atom (than oxygen)/sulfur has a
larger atomic radius/sulfur has more shells of
electrons (than oxygen) (1) - so sulfur has greater (inner shell) shielding (than
oxygen of the nucleus)/ more repulsion from inner
shells (1) - which reduces the (nuclear) attraction for the bonding
electrons (resulting in a longer bond length) (1)
sublimation: explain how the apparatus could be altered to maximise the formation of the product
- cool the reaction vessel/ surround the flask with cool water
- in order to prevent sublimation of PCl5/ prevent PCl5 turning into a gas
deduce the shape of the carbamate ion around the carbon. justify your answer
- trigonal planar
- C double bond O treated as a single bond pair of electrons
- 3 areas of electron density
- electron pairs repel to positions of minimum repulsion
