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Edexcel Chemistry A level > Bonding and structure > Flashcards

Bonding and structure Flashcards

1
Q

what is ionic bonding?

A

The electrostatic attraction between a positively charged ion and a negatively charged ion (between a non-metal and metal), where electrons are transferred

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2
Q

when is ionic bonding between a certain compound stronger?

A

when ions are smaller or have higher charges- eg a 3+ ion will have greater attraction than a 1+ ion.

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3
Q

Why do positive ions usually have a stronger net attraction?

A
  • because they are smaller ions than their atoms, meaning they have a larger number of protons than electrons, so, there is a greater attraction between the nucleus and the outer energy levels.
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4
Q

Why are negative ions from groups 5-7 LARGER than their atoms?

A

because they gain electrons, the ratio of electrons is larger than the number of protons. Therefore, the attraction of the nucleus is shared amongst more electrons making it a larger atom.

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5
Q

How does ionic radius change as one goes down the group? Why?

A

it increases because the ions gain more energy levels, so the attraction between the nucleus and the energy levels will be reduced.

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6
Q

How can we evidence charged particles using copper chromate Solution?

A

electrolysis of copper(ii) Chromate (vi) - originally green. When electricity is passed through, it separates into blue copper ions that move toward the cathode, and yellow chromate ions which move toward the anode.

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7
Q

How can we evidence charged particles using Potassium Manganate solution?

A

When placed on moist filter paper and passed through a PD, the purple color that it is starts to move to the anode due to negative MnO4- ions being attracted.

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8
Q

What are some physical properties of ionic compounds

A
  • high melting and boiling point
  • can only conduct electricity when molten or aqueous
  • can’t conduct electricity when solid
  • brittle
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9
Q

why are ionic compounds brittle?

A
  • when a force is applied, it causes layers to slide
  • when like ions meet however, there will be a repulsion which forces layers apart
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10
Q

why do ionic compounds have high melting and boiling points?

A
  • the bigger the charge on an ion, the stronger the electrostatic forces,
  • because of the strong electrostatic attraction between ions, it requires a lot of energy to overcome these forces
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11
Q

How does the ionic radius of ISOELECTRIC ions change with an increasing atomic number? Why?

A

it decreases because there is a greater number of protons while the number of electrons stays the same.

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12
Q

define isoelectric

A

having the same number of electrons or the same electronic strucutre

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13
Q

What is a covalent bond?

A

caused by the electrostatic attraction between the bonding shared pair of electrons and the 2 nuclei.

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14
Q

what is meant by electron density?

A

the probability of an electron being present at a specific location.

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15
Q

what do multiple covalent bonds having a greater electron density cause?

A

causes a greater force of attraction between the nuclei and electrons between them, resulting in a shorter bond length and stronger bond strength.

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16
Q

what is a dative/coordinate covalent bond?

A

forms when a shared pair of electrons comes only from one of the bonding atoms, needs an empty energy level, and a lone pair of electrons

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17
Q

What are some examples of molecules/ions that contain a coordinate bond?

A

NH4+, H3O+, NH3BF3

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18
Q

What shape contains 2BP, 0LP and has a bond angle of 180?

A

Linear

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19
Q

What shape contains 3BP, 0LP and has an angle of 120?

A

Triagonal planar

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20
Q

What shape contains 4BP, 0LP, and an angle of 109.5?

A

Tetrahedral (draw with one triangle looking thing and one dashed line)

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21
Q

What shape contains 3BP, 1LP and has an angle of 107 degrees?

A

Triagonal pyramidal

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22
Q

What shape contains 2BP, 2LP and an angle of 104.5?

A

Bent

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23
Q

what shape contains 5BP, 0LP and has both angles of 90 and 120?

A

Triagonal Bipyramidal

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24
Q

What shape has 6BP, 0LP and has angles of 90?

A

Octahedral

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25
Q

What shape has 4BP, 2LP and has angles of 90?

A

square planer

26
Q

define electronegativity

A

the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself.

27
Q

What are the factors affecting electronegativity?

A
  • number of unshielded protons
  • atomic radius
  • number of electron shells
28
Q

How does electronegativity vary across a group?

A

decreases

29
Q

How does electronegativity vary across a period?

A

Increases

30
Q

WHY does electronegativity vary across a group?

A

Decreases because atoms get larger and shielding increases, increasing the distance between the shells and nucleus.

31
Q

WHY does electronegativity vary across a period?

A

Because the number of protons increases, causing the atomic radius to decrease, causing a stronger force of attraction between electrons and the nucleus.

32
Q

why is an element/compound not purely ionic or covalent? What determines this?

A

because ionic and covalent bonds are extremes of bonding. Differences in electronegativity can determine this.

33
Q

What will a compound be if there is a very large difference in electronegativity within elements?

A

ionic

34
Q

What will a compound be if there is a small difference in electronegativity within elements?

A

covalent

35
Q

What is a polar bond?

A

Forms when elements have different electronegativities, forming an unequal distribution of electrons, producing charge separations.

36
Q

What structure of molecule will be polar?

A

non-symetrical

37
Q

What structure of molecule will be non-polar?

A

Symmetrical

38
Q

Why is it that even though a molecule could have polar bonds, itself might not be polar?

A

due to the individual dipoles canceling each other out, so there is no net moment, causing the molecule to be non-polar.

39
Q

What experiment can you do to show that a liquid is Polar?

A

a charged rod is brought close to a flowing stream of liquid from a burette.

If the liquid is polar, the stream will be deflected- attracted/repelled.

40
Q

What are the 3 types of intermolecular forces?

A

London (or Van der Waals)
permenant dipole-dipole
Hydrogen

41
Q

Describe how Van Der Waals are formed.

A
  • In many molecules, electrons are constantly moving, causing electron density to potentially fluctuate.
  • therefore, this can induce a dipole-dipole between molecules due to parts of molecules having a slightly more +/- charge.
42
Q

What are the factors affecting the size of London forces?

A

number of electrons.
Higher the number, the stronger the forces.

Shape of molecule: if it has more areas of contact, it will have stronger London Forces. (eg straight chain alkanes have stronger london forces than branched ones cus more contact)

43
Q

How are permanent dipole-dipole forces formed?

A

Happens between polar molecules.
Unlike London forces, these are permanent, making them stronger.

44
Q

How are hydrogen bonds formed?

A

It occurs when there is a hydrogen atom attached to any of the most electronegative elements: eg F, O, N which has an available lone pair of electrons.

45
Q

Explain why Water has a high boiling point.

A

This is due to the hydrogen bonding in water- it can form 2 H bonds/molecule due to its 2 lone pairs. Therefore, this is very strong and needs a lot of energy to overcome.

46
Q

Why does water expand when frozen?

A

Since Ice is a very regular structure, the hydrogen bonding in water pushes water molecules further apart creating ice which is less dense.

47
Q

Why do alcohols have relatively higher boiling points than an alkane with the same number of electrons?

A

Because alcohols can form Hydrogen bonds, which are very strong. Therefore, they are less volatile and have a higher boiling point since more energy is required to overcome the H bonds.

48
Q

Why is the melting point of XeF4 higher than PF3? explain through the shapes of their molecules

A

XeF4 has a square planar shape with bond angles of 90 degrees whereas PF3 has a tetrahedral shape with bond angles of 109.5 degrees. This is due to lone pairs of electrons repelling each other equally to try to get as further away as possible. XeF4 has London forces, but PF3 has london forces and dipole-dipole interactions. However, there are more intermolecular forces in XeF4, therefore meaning that more energy is needed to overcome them.

49
Q

Why do alcohols dissolve in water?

A

Because they can form H bonds with the lone pairs of electrons in polar ends of H2O and their OH group.

50
Q

Explain why the O–H and S–H bond lengths are different

A
  • sulfur is larger atom than oxygen
  • so sulfur has greater shielding
  • so reduces the attraction for bonding pairs of electrons
51
Q

Using the data provided, explain how changes in the cation affect the bond strength in an ionic
compound.

LiF= 1086
KF= 817
CaF2= 2948

A
  • as charge increase, the bond strength also increases as Li has a 1+ charge but Ca has a 2+ charge
  • so a smaller ionic radius increases bond strength as there will be a greater attractive force between opp charged ions since Li+ is smaller than K+
52
Q

Explain what is meant by a covalent bond

A
  • strong electrostatic attraction
  • between bonding pairs of electrons shared between 2 nuclei.
53
Q

what is covalent character?

A

the measure of the ability of an atom to attract a shared pair of electrons towards itself.

54
Q

The bonding in lithium iodide has some covalent character because

A the lithium ion polarizes the iodide ion.

B the iodide ion polarizes the lithium ion.

C there is a very large difference in electronegativity between lithium and iodine.

D there is a very small difference in electronegativity between lithium and iodine.

A

A

55
Q

During this reaction, (dative bond formed) the bond angles about the nitrogen atom and the boron atom change.
State the new H—N—H and F—B—F bond angles.

A

Both change to 109.5 since original angle was 107 for both, increases by 2.5 since one lone pair present.

56
Q

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).

A
  • iron atoms pack closer to each other
  • iron atoms have higher Ar than carbon atoms
57
Q

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.

HF: 89
H2O: 100
CH4: -141

all have 10 electrons

A
  • same electrons number so similar London forces.
  • both H2O and HF can form H bonds
  • due to large diff in electronegativity- H=2.1, F= 4 and O= 3.5
  • H2O higher boiling point cus 2 H bonds/molecule
  • CH4 though only has London forces
  • so more energy is needed to overcome H bonds than only london forces
58
Q

State what is meant by the term electronegativity and hence explain the polarity, if any, of the
bonds in chlorine trifluoride, ClF3.

A
  • the tendency of an atom to pull bonding pairs of electrons in a covalent bond towards itself.
  • polar due to F being more electronegative than Cl
  • so F forms δ- and Cl forms δ+ charge.
59
Q

Which of the following statements is evidence for the existence of ions in ionic compounds?
A Ionic compounds, in the solid state, conduct electricity.

B When any ionic compound in solution is electrolysed, the migration of ions
can be seen.

C In electron density maps for ionic compounds, there is no single line
representing electron density that surrounds both cations and anions.

D In electron density maps for ionic compounds, there are some single lines
representing electron density that surround both cations and anions.

A

C

60
Q

KBr is soluble in water, but not in hexane
2-methylpentane is soluble in hexane, but not in water.

Explain the findings of the investigation by considering the interactions between
the compounds and each of the solvents.

A
  • KBr soluble in water, as K+ and Br- ions become hydrated in water.
  • so enthalpy of hydration is larger than the lattice energy.
  • insoluble in hexane because London forces formed between them would be similar in magnitude of the K+ and Br- ions.
  • 2-methyl pentane is soluble in hexane because London forces are similar in magnitude.
  • 2-methyl pentane insoluble in water as it cannot form Hydrogen bonds.
  • so London forces between hexane and 2methyl pentane are similar in magnitude to London forces in hexane and 2-methyl pentane individually.
61
Q
A

Edexcel Chemistry A level (18 decks)

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