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ALEVEL Physical Chemistry > Bonding > Flashcards

Bonding Flashcards

1
Q

Definition of ionic bonding

A

Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by electron transfer.

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2
Q

Strength and melting points of ionic bonding

A

Ionic bonding is stronger and the melting points higher when the ions are smaller and/ or have higher charges

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3
Q

Ionic radii of atoms

A

Positive ions are smaller compared to their atoms because it has one less shell of electrons
the ratio of protons to electrons has increased
so there is greater net force on remaining electrons holding them more closely.

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4
Q

Trends of ionic radii down the group

A

Within a group the size of the ionic radii increases going down the group. This is because as one goes down the group the ions have more shells of electrons.

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5
Q

Why are negative ions larger than corresponding atoms

A

The negative ions formed from groups five to seven are larger than the corresponding atoms.
The negative ion has more electrons than the corresponding atom but the same number of protons. So the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger.

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6
Q

Covalent bond definition

A

A covalent bond is a shared pair of electrons

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7
Q

Covalent bond definition

A

A covalent bond is a shared pair of electrons

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8
Q

Dative covalent bonding

A

A dative covalent bond forms when the shared pair of electrons in the covalent bond come from only one of the bonding atoms.

A dative covalent bond is also called co-ordinate bonding.

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9
Q

Arrows in covalent diagrams

A

The direction of the arrow goes from the atom that is providing the lone pair to the atom that is deficien

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10
Q

Metallic bonding definition

A

Metallic bonding is the electrostatic force of attraction between positive metal ions and the delocalised electrons

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11
Q

What are the three main factors that affect the strength of metallic bond?

A

Number of protons/strength of nuclear attraction (the more protons the stronger the bonds)

Number of delocalised electrons per atom (the outer shell electrons are localised) - the more electrons the stronger the bond

Size of ion (the smaller the ion the stronger the bonds)

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12
Q

Linear

A

Number of bonding pairs = 2

Number of lone pairs = 0

Bond angle = 180

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13
Q

Trigonal planar

A

Number of bonding pairs = 3

Number of lone pairs = 0

Bond angle = 120

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14
Q

Tetrahedral

A

Number of bonding pairs = 4

Number of lone pairs = 0

Bond angle = 109.5

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15
Q

Trigonal pyramidal

A

Number of bonding pairs = 3

Number of lone pairs = 1

Bond angle = 107

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16
Q

Bent

A

Number of bonding pairs = 2

Number of lone pairs = 2

Bond angle = 104.5

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17
Q

Trigonal bipyramidal

A

Number of bonding pairs = 5

Number of lone pairs =0

Bond angle = 120 and 90

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18
Q

Octahedral

A

Number of bonding pairs = 6

Number of lone pairs = 0

Bond angle = 90

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19
Q

How to explain shape

A
  • State the number of bonding pairs and lone pairs of electrons
  • state that electron pairs repel and try to get as far apart as possible (or toa position of minimum repulsion)
  • if there are no lone hairs state that the electrons repel equally
  • if there are lone pairs then that that the lone pairs repel more than bonding pairs
  • state actual shape and bond angles
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20
Q

What do lone pairs do

A

Lone pairs repel more than bonding pairs so reduce bond angles by about 2.5 degrees

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21
Q

when are more complex shapes seen

A

Occasionally more complex shapes are seen that are variations of octahedral and trigonal bipyramidal where some of the bonds are replaced with lone pairs.

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22
Q

when are more complex shapes seen

A

Occasionally more complex shapes are seen that are variations of octahedral and trigonal bipyramidal where some of the bonds are replaced with lone pairs.

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23
Q

Examples of complex shapes - XeF4

A

Xe has 8 electrons in its outer shell. 4 F’s add 4 more electrons. This makes a total of 12 electrons made up of 4 bond pairs and 2 lone pairs. The means it is a variation of the 6 bond pair shape (octahedral)

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24
Q

Examples of complex shapes ClF3

A

Cl has 7 electrons in its outer shell. 3 F’s add 3 more electrons. This makes a total of 10 electrons made up of 3 bond pairs and 2 lone pairs. The means it is a variation of the 5 bond pair shape (trigonal bipyramidal)

25
Q

Examples of complex shapes - IF4+

A

I has 7 electrons in its outer shell. 4 F’s add 4 more electrons. Remove one electron as positively charged. This makes a total of 10 electrons made up of 4 bond pairs and 1 lone pair. The means it is a variation of the 5 bond pair shape (trigonal bipyramidal)

26
Q

What is the definition of electronegativity

A

ability or tendency of an atom to attract a pair of electrons in a covalent bond.

This depends on the number of protons in the nucleus, the distance between outer electrons and the nucleus and the degree of shielding.

27
Q

How does nuclear charge affect electronegativity

A

Attraction exists between the positively charged protons in the nucleus and negatively charged electrons found in the energy levels of an atom. An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells. Therefore, an increased nuclear charge results in an increased electronegativity.

28
Q

How does atomic radius affect electronegativity

A

The atomic radius is the distance between the nucleus and electrons in the outermost shell. Electrons closer to the nucleus are more strongly attracted towards its positive nucleus. Those electrons further away from the nucleus are less strongly attracted towards the nucleus. Therefore, an increased atomic radius results in a decreased electronegativity.

29
Q

How does sheilding affect electronegativity?

A

• Shielding - Filled energy levels can shield (mask) the effect of the nuclear charge causing the outer electrons to be less attracted to the nucleus. Therefore, the addition of extra shells and subshells in an atom will cause the outer electrons to experience less of the attractive force of the nucleus. Sodium (period 3, group 1) has higher electronegativity than caesium (period 6, group 1) as it has fewer shells and therefore the outer electrons experience less shielding than in caesium. Thus, an increased number of inner shells and subshells will result in a decreased electronegativity

30
Q

Slightly ionic (type of bonding)

A

when one atom has a greater tendency to attract electron pairs.
This makes them polar covalent where partial charges that result from the uneven balance of electrons.

31
Q

Example of slightly ionic bonding

A

Eg: HF

The electrons are closer to fluorine due to its high difference in electronegativity. This causes a greater electron density around fluorine due to the uneven distribution. This causes a slight negative charge around the fluorine and a slight positive around the hydrogen.

32
Q

Trends in electronegativity - Down a group

A

There is a decrease in electronegativity going down the group
The nuclear charge increases as more protons are being added to the nucleus
However, each element has an extra filled electron shell, which increases shielding
The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii

Overall, there is decrease in attraction between the nucleus and outer bonding electrons

33
Q

Trends in electronegativity - across a period

A

Electronegativity increases across a period
The nuclear charge increases with the addition of protons to the nucleus
Shielding remains relatively constant across the period as no new shells are being added to the atoms
The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the period of the periodic table

This results in smaller atomic radii

34
Q

How can dipole moments be used to give a measure of the overall polarity of a molecule

A

• Non-polar molecules have a 0 dipole moment, the more polar the molecule, the greater the dipole moment.

• For some molecules, e.g HCI, the polar bond gives rise to a permanent dipole- there is a positive and negative end to the molecule.

• However, the shape and symmetry of the molecule plays a role in determining whether the molecule is polar overall.

• For a molecule to be polar overall, it must contain polar bonds and its not in the same place shape must be such that the centres of positive and negative charge are

35
Q

What are the three types of intermolecular forces

A

• Van der Waal’s forces (london and dispersion forces)
• Dipole-dipole forces (stronger than van der Waal’s)
• Hydrogen bonding

36
Q

What is a van der waals force and how do they occur?

A

• The electron charge cloud in non-polar molecules or atoms are constantly moving. During this movement, the electron charge cloud can be more on one side of the atom or molecule than the other

• The electron movement gives rise to instantaneous dipole on the moment

• This temporary dipole has an inductive effect on the electron clouds of neighbouring molecules, the partial negative end of one neighbouring molecule resulting in an induced charge.

37
Q

What is the effect of mr on van der waals forces?

A

• the strength of the forces increase with mass

• As moles get bigger, there are more electrons so instantaneous dipole moments increases

• with halogens atom get bigger, more forces because there are more electrons so there are greater forces ( not strong).

38
Q

Relative strength of van der waals forces

A

For small molecules with the same number of electrons, permanent dipoles are stronger than induced dipoles

Butane and propanone have the same number of electrons

Butane is a nonpolar molecule and will have induced dipole forces
Propanone is a polar molecule and will have permanent dipole forces

Therefore, more energy is required to break the intermolecular forces between propanone molecules than between butane molecules

So, propanone has a higher boiling point than butane

39
Q

Dipole forces

A

Polar molecules have permanent dipoles
The molecule will always have a negatively and positively charged end

Forces between two molecules that have permanent dipoles are called permanent dipole - dipole forces

The δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other

40
Q

Hydrogen bonding

A

Hydrogen bonding is the strongest form of intermolecular bonding
Intermolecular bonds are bonds between molecules
Hydrogen bonding is a type of permanent dipole – permanent dipole bonding

41
Q

What is needed ofr a hydrogen bond to oocur

A

Hydrogen bonding is the strongest form of intermolecular bonding

Intermolecular bonds are bonds between molecules

Hydrogen bonding is a type of permanent dipole – permanent dipole bonding

42
Q

How does hydrogen bonding occur

A

When hydrogen is covalently bonded to an O, N or F, the bond becomes highly polarised

The H becomes so δ+ charged that it can form a bond with the lone pair of an O, N or F atom in another molecule

For example, in water
Water can form two hydrogen bonds, because the O has two lone pairs

43
Q

Properties of water

A

Hydrogen bonding in water, causes it to have anomalous properties such as high melting and boiling points, high surface tension and a higher density in the liquid than the solid

44
Q

Why does water have a high boiling and melting point

A

Water has high melting and boiling points due to the the strong intermolecular forces of hydrogen bonding between the molecules in both ice (solid H2O) and water (liquid H2O)

A lot of energy is therefore required to separate the water molecules and melt or boil them

45
Q

The enthalpy of vapourisation of different hydrides

A

The enthalpy changes increase going from H2S to H2Te due to the increased number of electrons in the Group 16 elements

This causes an increase in the instantaneous dipole - induced dipole forces as the molecules become larger

H2O is an anomalous cos of hydrogen

46
Q

High surface tension in water

A

Water has a high surface tension
Surface tension is the ability of a liquid surface to resist any external forces

The water molecules at the surface of liquid are bonded to other water molecules through hydrogen bonds

These molecules pull downwards the surface molecules causing the surface of them to become compressed and more tightly together at the surface

This increases water’s surface tension

47
Q

Density of water

A

the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form

Therefore, ice has a lower density than liquid water by about 9%

48
Q

Formation of a permanent dipole

A

A polar covalent bond forms when the elements in the bond have different electronegativities . (Of around 0.3 to 1.7)

When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) ends.

49
Q

Symmetric molecules

A

A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar

The individual dipoles on the bonds ‘cancel out due to the symmetrical shape of the molecule. There is no net dipole moment: the molecule is non-polar

50
Q

Main factor affecting size of Van der Waals

A

The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the Van der Waals stronger between the molecules and so boiling points will be greater.

51
Q

What substances do van der walls forces not occur in

A

Ionic

52
Q

What is a van der walls forces

A

These are also called induced dipole-dipole interactions.

In any molecule the electrons are moving constantly and randomly. As this happens the electron density can fluctuate and parts of the molecule become more or less negative i.e. small temporary or transient dipoles form. These instantaneous dipoles can cause dipoles to form in neighbouring molecules.

These are called induced dipoles. The induced dipole is always the opposite sign to the original one.

53
Q

Halogens boiling pint explained by van der walls forces

A

The increasing boiling points of the halogens down the group 7 series can be explained by the increasing number of electrons in the bigger molecules causing an increase in the size of the Van der Waals between the molecules. This is why I2 is a solid whereas Cl2 is a gas.

54
Q

How does the shape of the molecule can also have an effect on the size of the Van der Waals forces.

A

Long chain alkanes have a larger surface area of contact between molecules for Van der Waals to form than compared to spherical shaped branched alkanes and so have stronger Van der Waals.

55
Q

What is a permanent dipolw force

A

•Permanent dipole-dipole forces occurs between polar molecules
•It is stronger than Van der Waals and so the compounds have higher boiling points
•Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)
•Polar molecules are asymmetrical and have a bond where there is a significant difference in
electronegativity between the atoms.

56
Q

What is hydrogen bonding

A

It occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons. e.g. a –O-H -N-H F- H bond.
There is a large electronegativity difference between the H and the O,N,F

57
Q

Strongest type of bonding

A

Hydrogen bonding is stronger than the other two types of intermolecular bonding.

58
Q

Boiling point trend of h2o nh3 and hf explained

A

The anomalously high boiling points are caused by hydrogen bonding between the molecules. The general increase in boiling point from H2S to H2Te is caused by increasing Van der Waals forces between molecules due to an increasing number of electrons.

ALEVEL Physical Chemistry (3 decks)

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