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Bonding Flashcards

1
Q

What is ionic bonding?

A

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

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2
Q

Formation of cation

A

A positive ion (cation) is formed when an atom loses an electron

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3
Q

Formation of anion

A

A negative ion (anion) is formed when an atom gains an electron

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4
Q

Migration of ions as evidence for existence of charged particles (ions):

A

Electrolysis of copper (III) chromate (VI) on a wet filter paper shows evidence for charged particles

When a drop of green copper (III) chromate (VI) is placed on wet filter paper and electricity passed through it, the ions start to separate

The positive Cu2- ions move towards the negative cathode (opposite charges attract). You can see the blue solution move

The negative CrO42- ions move towards the positive anode (opposite charges attract). You can see the yellow solution move

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5
Q

Physical properties of ionic compounds:

A

Ionic compounds, like sodium chloride, have a giant ionic structure

They conduct electricity when molten or dissolved in solution as the ions are free to move around

Most ionic compounds dissolve in water as water molecules are polar and so they can attract the positive and negative ions and break up the structure

Have high melting points as there are many strong electrostatic forces between oppositely charged ions.

Lots of energy is needed to overcome these forces

They are brittle. When struck with a hammer the layers slide and you get positive ions aligned with positives and negative ions aligned with negatives. They repel, and the structure breaks apart

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6
Q

What are covalent bonds?

A

Covalent bond is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them:

Forms between elements that have high electronegativity values (non metals)

There are single, double and triple covalent bonds. More electrons are being shared

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7
Q

How can covalent bonds be represented?

A

Covalent bonds can be represented by lines too

An arrow represents a coordinate bond

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

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8
Q

Reasons for the shapes offend bond angles in, simple molecules

A

Electron pairs repel and try to get as far apart as possible (or to a position of minimum repulsion.)

If there are no lone pairs the electron pairs repel equally

If there are lone pairs of electrons, then state that lone pairs repel more than bonding pairs.

Lone pairs repel more than bonding pairs and so reduce bond angles (by about 2.5o per lone pair

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9
Q

2 bonding pairs

0 lone pairs

A

Linear shape

Bond angle - 180

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10
Q

3 bonding pairs

0 lone pairs

A

Trigonal planar shape

Bond angle - 120

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11
Q

4 bonding pairs

0 lone pairs

A

Tetrahedral shape

Bond angle - 109.5

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12
Q

5 bonding pairs

0 lone pairs

A

Trigonal bipyramidal shape

Bond angle - 120 and 90

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13
Q

6 bonding pairs

0 lone pairs

A

Octahedral shape

Bond angle - 90

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14
Q

3 bonding pairs

1 lone pairs

A

Trigonal pyramidal shape

Bond angle - 107

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15
Q

2 bonding pairs

2 lone pairs

A

Brent shape

Bond angle 104.5

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16
Q

Factors affecting electronegativity:

A

Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more.

It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases

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17
Q

Ionic and covalent bonding

A

Ionic and covalent bonding are the extremes of a continuum of bonding type.

Differences in electronegativity between elements can determine where a compound lies on this scale

A compound containing elements of similar electronegativity and hence a small electronegativity difference will be purely covalent

A compound containing elements of very different electronegativity and hence a very large electronegativity difference(> 1.7) will be ionic

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18
Q

Why may molecules with polar bonds may not be polar molecules?

A

If the polar bonds in a molecule are arranged symmetrically, the partial charges cancel out and the molecule is non-polar.

Due to directly opposing pulls by the electronegative element.

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19
Q

London forces

A

London forces occur between all molecular substances and noble gases. They do not occur in ionic substances

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. This leads to instantaneous dipoles

Instantaneous dipole induces a dipole in nearby molecules. Induced dipoles attract one another

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20
Q

Main factor affecting size of London Forces

A

The more electrons there are in the molecule the higher the chance that temporary dipoles will form.

This makes the London forces stronger between the molecules and more energy is needed to break them so boiling points will be greater.

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21
Q

Permanent dipole-dipole forces

A
  • Permanent dipole-dipole forces occurs between polar molecules
  • It is stronger than London forces 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.
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22
Q

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

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23
Q

Understand the interactions in molecules, such as H2O, liquid NH3 and liquid HF, which give rise to hydrogen bonding:

A

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

The anomalously high boiling points of H2O, NH3 and HF are caused by the hydrogen bonding between these molecules in addition to their London forces.

The additional forces require more energy to break and so have higher boiling points

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24
Q

How does the shape of the molecule affect the size of London forces?

A

The shape of the molecule can also have an effect on the size of the London forces.

Long straight chain alkanes have a larger surface area of contact between molecules for London forces to form than compared to spherical shaped branched alkanes and so have stronger London forces

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25
Explain how hydrogen bending results in ice being more dense than water
For most substances, as the temperature drops below its freezing point the density increases as a solid forms. This is not the case for water. The density of ice changes at around 4*C. There are 4 groups around each oxygen atom, arranged tetrahedrally in three dimensions by hydrogen bonding. This leads to an 'open' structure with large spaces in it." Because of its open structure ice is less dense than water.
26
Explain how hydrogen bending results in the relatively high melting/ boiling points of water
Each water molecule can potentially form four hydrogen bonds with surrounding water molecules. There are exactly the right numbers of δ+ hydrogens and lone pairs so that every one of them can be involved in hydrogen bonding. This is why the boiling point of water is higher than that of ammonia or hydrogen fluoride
27
Trends in boiling temperatures of alkanes with increasing chain length
The only intermolecular force in alkanes are London forces These forces are very weak. With more electrons, the forces get stronger, as the instantaneous dipoles are likely to be stronger. This is why longer chains of alkanes have higher boiling and melting points.
28
Effect of branching in the carbon chain on the boiling temperatures of alkanes
The only intermolecular force that exist in alkanes are London forces Straight chain molecules have more places along its length where they can be attracted to other molecules, so there are more chances of London forces to be developed. Hence, they have stronger intermolecular forces (and higher boiling points) as compared to the branched chain molecules which have a compact shape, and therefore fewer spaces where they can be attracted to other molecules.
29
The relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons
Alcohols are not as volatile as alkanes with similar numbers of electrons, due to hydrogen bonding Alkanes have weaker London forces
30
The trends in boiling temperatures of the hydrogen halides, HF to HI
Among hydrogen halides, only hydrogen fluoride exhibits hydrogen bonding between molecules, and therefore has the highest melting and boiling points of the Hydrogen halides. From HCl to HI the boiling point rises. This trend is attributed to the increasing strength of intermolecular van der Waals forces , which increases as the numbers of electrons in the molecules increases.
31
For a substance to dissolve:
Solvent bonds must break Substance bonds must break New bonds must be formed between the solvent and substance
32
Ionic substances dissolving in water
When an ionic lattice dissolves in water it involves breaking up the bonds in the lattice and forming new bonds between the metal ions and water molecules. The negative ions are attracted to the δ+ hydrogens on the polar water molecules and the positive ions are attracted to the δ- oxygen on the polar water molecules. The water molecules surround the ions in a process called hydrogen
33
Solubility of simple alcohols
The smaller alcohols are soluble in water because they can form hydrogen bonds with water. The longer the hydrocarbon chain the less soluble the alcohol.
34
Insolubility of compounds in water
Compounds that cannot form hydrogen bonds with water molecules ,e.g. polar molecules such as halogenoalkanes or non polar substances like hexane will be insoluble in water.
35
Solubility in non-aqueous solvents
Compounds which have similar intermolecular forces to those in the solvent will generally dissolve Non-polar solutes will dissolve in non-polar solvents e.g. Iodine which has only London forces between its molecules will dissolve in a non polar solvent such as hexane which also only has London forces. Propanone is a useful solvent because it has both polar and nonpolar characteristics. It can form London forces with some non polar substances such as octane with its CH3 groups. Its polar C=O bond can also hydrogen bond with water
36
What is metallic bonding?
Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons
37
The three main factors that affect the strength of metallic bonding are:
1. Number of protons/ Strength of nuclear attraction. The more protons the stronger the bond 2. Number of delocalised electrons per atom (the outer shell electrons are delocalised) The more delocalised electrons the stronger the bond 3. Size of ion. The smaller the ion, the stronger the bond.
38
Properties of metals
Metals have high melting points because the strong electrostatic forces between positive ions and sea of delocalised electrons require a lot of energy to break Metals can conduct electricity well because the delocalised electrons can move through the structure Metals are malleable because the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another. The attractive forces in the lattice are the same whichever ions are adjacent
39
Ionic bonding structure:
Giant ionic lattice eg, sodium chloride, magnesium oxide
40
Covalent bonding structure: | Simple molecular
Simple molecular: With intermolecular forces (London forces, permanent dipoles, hydrogen bonds) between molecules Eg, Iodine, Ice, Carbon dioxide, Water, Methane
41
Covalent bonding structure: | Macromolecular
Macromolecular: giant molecular structures. Eg. Diamond, Graphite, Silicon dioxide, Silicon
42
Metallic bonding structure:
Giant metallic lattice Eg, Magnesium, Sodium (all metals)
43
Examples of giant covalent structures: | Graphite
Each carbon is bonded 3 times with 4th electron delocalised Lots of strong covalent bonds means graphite has a very high melting point Layers slide easily as there are weak forces between the layers Delocalised electrons between the layers allow graphite to conduct electricity as they can carry a charge Layers are far apart in comparison to covalent bond length. This means it has a low density It is insoluble as the covalent bonds are too strong to break
44
Examples of giant covalent structures: | Diamond
Each carbon is bonded 4 times in a tetrahedral shape The tightly packed rigid, arrangement allows heat to conduct well in diamonds Doesn't conduct electricity as it doesn't have delocalised electrons Very high melting point due to many strong covalent bonds It is insoluble as the covalent bonds are too strong to break
45
Examples of giant covalent structures: | Graphene
Graphene is one layer of graphite. It is one atom thick and is made up of hexagonal carbon rings Delocalised, free moving electrons makes graphene an excellent conductor of electricity as they can carry a charge
46
Ionic structures
Boiling and melting points: high-because of giant lattice of ions with strong electrostatic forces between oppositely charged ions Solubility in water: Generally good Conductivity when solid: poor: ions can’t move/ fixed in lattice Conductivity when molten: good: ions can move General description: crystalline solids
47
Molecular (simple) structures
Boiling and melting points: low- because of weak intermolecular forces between molecules (specify type e.g London forces/hydrogen bond) Solubility in water: generally poor Conductivity when solid: poor: no ions to conduct and electrons are localised (fixed in place) Conductivity when molten:poor: no ions General description:mostly gases and liquids
48
Macromolecular structures
Boiling and melting points: high because of many strong covalent bonds in macromolecular structure, takes a lot of energy to break the many strong bonds Solubility in water: insoluble Conductivity when solid: diamond and sand - poor, because electrons can’t move (localised), Graphite - good as free delocalised electrons between layers Conductivity when molten: poor General description: solids
49
Metallic structures
Boiling and melting points: high - strong electrostatic forces between positive ions and sea of delocalised electrons Solubility in water: insoluble Conductivity when solid: good -delocalised electrons can move through structure Conductivity when molten: good General description: shiny metal, malleable as the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another -attractive forces in the lattice are the same whichever ions are adjacent
50
What are molecular ions?
Covalently bonded atoms that lose or gain electrons
51
Which electrons are lost when an atom becomes a positive ion?
Electrons in the highest energy levels
52
Which are the 4 elements that don’t tend to form ions and why?
The elements are beryllium, boron, carbon and silicon Requires a lot of energy to transfer outer shell electrons
53
What are the 3 main types of chemical bonds?
● Ionic ● Covalent ● Metallic
54
Give an example of a ionically bonded substance
NaCl (Sodium Chloride - salt)
55
What determines the strength of an ionic bond?
- Ionic radius and ionic charge | - Ionic bonding is stronger and the melting points higher when the ionic radius is smaller and/ or have higher charges.
56
Explain the trend in ionic radius down a group
Ionic radii increases going down the group. This is because down the group the ions have more shells of electrons and thus the outermost electron experience less pull from positive nucleus.
57
Explain the trend in ionic radius for this set of isoelectronic ions, e.g. N3- to Al3+
There are increasing numbers of protons from N to F and then Na to Al but the same number of electrons. Therefore nuclear attraction between the outermost electrons and nucleus increases and ions get smaller
58
What are isoelectronic ions?
Isoelectronic ions are ions that have the same number of electrons.
59
What are the physical properties of ionic compounds?
* high melting points * non conductor of electricity when solid * conductor of electricity when in solution or molten * brittle
60
In a solution of CuCrO4 with connected electrodes which electrode will the 2 ions migrate to?
Cu2+ - migrates to negative electrode CrO42- - migrates to positive electrode
61
Define metallic bonding
Electrostatic attraction between the positive metal ions and the sea of delocalised electrons
62
Define covalent bonding
Electrostatic attraction between a shared pair of electrons and the nuclei
63
Why does giant ionic lattices conduct electricity when liquid but not when solid?
In solid state the ions are in fixed positions and thus cannot move. When they are in liquid state the ions are mobile and thus can freely carry the charge
64
Giant ionic lattices have high or low melting and boiling point? Explain your answer
They have high melting and boiling point because a large amount of energy is required to overcome the electrostatic bonds
65
In what type of solvents do ionic lattices dissolve?
Polar solvents E.g water
66
Why are ionic compounds soluble in water?
Water has a polar bond. Hydrogen atoms have a δ+ charge and oxygen atoms have a δ- charge. These charges are able to attract charged ions When an ionic lattice dissolves in water it involves breaking up the bonds in the lattice and forming new bonds between the metal ions and water molecules.
67
How many covalent bonds does carbon form?
4
68
How many covalent bonds does oxygen form?
2
69
What is the effect of multiple covalent bonds on bond length and strength?
Double/triple bonds exert greater electron density therefore the attraction between nucleus and electron is greater resulting in a shorter and stronger bond.
70
What is a lone pair?
Electrons in the outer shell that are not involved in the bonding
71
What is a dative covalent bond?
A bond where both of the shared electrons are supplied by one atom
72
How are oxonium ions formed?
Formed when acid is added to water, H3O
73
What are the types of covalent structure?
● Simple molecular lattice | ● Giant covalent lattice
74
Describe the bonding in simple molecular structures?
Atoms within the same molecule are held by strong covalent bonds and different molecules are held by weak intermolecular forces
75
Why do simple molecular structures have low melting and boiling point?
Small amount of energy is enough to overcome the weak intermolecular forces
76
Can simple molecular structures conduct electricity? Explain
No, they are non conductors They have no free charged particles to move around
77
Simple molecular structures dissolve in what type of solvent
Non polar solvents
78
Give examples of giant covalent structures
● Diamond ● Graphite ● Silicon dioxide, SiO2
79
List some properties of giant covalent structures
● High melting and boiling point ● Non conductors of electricity, except graphite ● Insoluble in polar and non polar solvents
80
How does graphite conduct electricity?
Delocalised electrons present between the layers are able to move freely carrying the charge
81
Why do giant covalent structures have high melting and boiling point?
Strong covalent bonds within the molecules need to be broken which requires a lot of energy
82
What does the shape of a molecule depend on?
Number of electron pairs in the outer shell Number of these electrons which are bonded and lone pairs
83
By how many degrees does each lone pair reduce the bond angle?
2.5°
84
What is electronegativity?
Electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond The further up and right you go in the periodic table (excluding the noble gases) the more electronegative an element is. Fluorine is the most electronegative Electronegativity is measured on the Pauling scale(ranges from 0 to 4)
85
How is a polar bond formed?
Bonding atoms have different electronegativities
86
What are the 3 types of intermolecular forces?
● Hydrogen bonding ● Permanent dipoles ● London forces
87
Does boiling point increase or decrease down the noble gas group? Why?
Boiling point increases because the number of electrons increases and hence the strength of London forces also increases
88
What conditions are needed for hydrogen bonding to occur?
O-H, N-H or F-H bond, lone pair of electrons on O, F, N Because O, N and F are highly electronegative, H nucleus is left exposed Strong force of attraction between H nucleus and lone pair of electrons on O, N, F
89
Are alkanes soluble in water? Explain your answer.
Insoluble because hydrogen bonds in water are stronger than alkanes’ London forces of attraction
90
What kind of intermolecular forces do alcohols have? Why?
Hydrogen bonding, due to the electronegativity difference in the OH bond
91
How do alcohols’ melting point and boiling point compare to other hydrocarbons’ of similar C chain lengths? Why?
Higher, because they have hydrogen bonding (strongest type of intermolecular force) → stronger than London forces
92
Are alcohols soluble in water? Why does solubility depend on chain length?
Soluble when short chain - OH hydrogen bonds to hydrogen bond in water Insoluble when long chain - non-polarity of C-H bond takes precedence
93
What is a coordinate bond?
A covalent bond that is formed when both electrons are donated by the same atom Once formed, acts in the same way as a regular covalent bond
94
The shorter a covalent bond....
… the stronger the bond | shorter length, higher enthalpy
95
What is bond length measured in?
bond length is measured in picometers (pm)

chemistry (13 decks)

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