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A-level Chem > Bonding > Flashcards

Bonding Flashcards

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1
Q

Ionic bonding definition

A

electrostatic force of attraction between oppositely charged ions formed by electron transfer

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2
Q

metals and non metals form which type of ions each?

A

metals form +ve ions
non metals form -ve ions

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3
Q

What is ionic crystal structure?

A

Giant lattices of ions

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4
Q

When ions are smaller or have higher charges, what is affected and how?

A

ionic bonding is affected and becomes stronger

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5
Q

positive ions are larger/smaller than their atoms and the reason why

A

smaller
one less shell of electrons
ratio of protons to electrons has increased
greater net force on remaining electrons to hold them more closely

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6
Q

negative ions are larger/smaller than their atoms and why (6)

A

larger
more electrons than in the atom
but same number of protons
so pull of the nucleus is shared over more electrons
the attraction per electron is less
making the ion bigger

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7
Q

What do N3- , O2-, F-, Na+, Mg2+ and Al3+ have in common? Why do their ionic radii decrease?

A

the same electronic structures (neon)

increasing no of protons from N to F and Na to Al but same number of electrons
Nuclear attraction per electron increases and ions get smaller

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8
Q

Covalent bond definition

A

A shared pair of electrons between 2 atoms

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9
Q

define dative covalent bond and what is an alternative name?

A

one forming when a shared pair of electrons in the covalent bond icome from only one of the bonding atoms, aka a coordinate bond

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10
Q

Give 3 examples of molecules with a dative bond and state which bond is dative in that molecule.

A

NH4+ (N-> H)
H3O+ (O-> H)
NH3BF3 (N-> B)

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11
Q

Metallic bonding definition

A

Electrostatic force of attraction between the positive metal ions and the delocalised electrons

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12
Q

3 factors affecting strength of metallic bonding

A
  1. number of protons/strength of nuclear attraction - the more protons the stronger the bond
  2. number of delocalised e- per atom - the more delocalised e- the stronger the bond
  3. size of the ion - the smaller the ion, the stronger the bond
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13
Q

Why does Mg have stronger metallic bonding than Na and a higher boiling point? (5)

A

In Mg there are more electrons in the outer shell of each atom released to the sea of electrons
The Mg ion is also smaller
It also has one more proton
There is a stronger electrostatic attraction between the positive metal ions and the delocalised electrons
More energy is required to break bonds

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14
Q

4 different molecular structures, their bonding and examples for each

A

Giant ionic lattice
ionic
NaCl
MgO

Simple molecular (with intermolecular forces)
Covalent
Iodine
Ice
CO2
Water
Methane

Macromolecular
Covalent
Diamond
Graphite
Silicon dioxide
Silicon

Giant metallic lattice
Metallic
Mg
(all metals)

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15
Q

When to use the words molecules/intermolecular forces

A

When dealing with simple molecular substances

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16
Q

Ionic properties
Melting/boiling points
Solubility in water
Conductivity when solid
Conductivity when molten
General description

A

High because of giant lattice of ions with strong electrostatic forces between oppositely charged ions

Generally good

Poor, ions can’t move when fixed in lattice

Good, ions can move

Crystalline solids

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17
Q

Simple molecular properties
Melting/boiling points
Solubility in water
Conductivity when solid
Conductivity when molten
General description

A

Low, because of weak intermolecular forces between molecules

Poor

Poor, no ions to conduct and electrons are localised (fixed in place)

Poor, no ions

Mostly gases and liquids

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18
Q

Macromolecular properties
Melting/boiling points
Solubility in water
Conductivity when solid
Conductivity when molten
General description

A

high, because of many strong covalent bonds in the structure which take a lot of energy to break

Insoluble

Diamond and sand poor as electrons are localised
Graphite is good as free electrons move between layers

Poor

Solids

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19
Q

Metallic properties
Melting/boiling points
Solubility in water
Conductivity when solid
Conductivity when molten
General description

A

High due to strong electrostatic forces between cations and delocalised electron sea

Insoluble

Good, delocalised e- move through structure

Good - “ “

Shiny metal, malleable due to planes of identical ions sliding over each other as attractive forces in the structure are the same whichever ions are adjacent to each other

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20
Q

Linear
No. bonding pairs
No. lone pairs
Bond angle
Examples

A

2
0
180°
CO2, CS2, HCN, BeF2

21
Q

Trigonal planar
No. bonding pairs
No. lone pairs
Bond angle
Examples

A

3
0
120
BF3, AlCl3, SO3, NO3-, CO32-

22
Q

Tetrahedral
No. bonding pairs
No. lone pairs
Bond angle
Examples

A

4
0
109.5°
SiCl4, SO42-, ClO4-, NH4+

23
Q

Trigonal bipyramidal
No. bonding pairs
No. lone pairs
Bond angle
Example

A

5
0
120° and 50°
PCl5

24
Q

Bent
No. bonding pairs
No. lone pairs
Bond angle
Examples

A

2
2
104.5°
OCl2, H2S, OF2, SCl2

25
Q

Trigonal pyramidal
No. bonding pairs
No. lone pairs
Bond angle
Examples

A

3
1
107°
NCl3, PF3, ClO3, H3O+

26
Q

Octahedral
No. bonding pairs
No. lone pairs
Bond angle
Example

A

6
0
90°
SF6

27
Q

How to explain shape in 6 steps

A

State number of bonding pairs and lone pairs
State that electron pairs repel to get as far apart as possible
If there are no lone pairs, state that electron lairs repel equally
else, state the lone pairs repel more than bonding pairs
State actual shape and bond angle(s)

28
Q

If lone pairs are in a structure, how much do they reduce bond angles by?

A

around 2.5°

29
Q

State where Van der Waals form and where they do not

A

all molecular substances and noble gases
not in ionic substances

30
Q

What are VDWs also known as?

A

transient induced dipole dipole interactions

31
Q

How do induced dipole dipole interactions occur?

A

e- move constantly and randomly in any molecule
e- density fluctuates
small temporary dipoles can form
these can lead to induced dipole dipoles that have the opposite sign to the original dipole

32
Q

2 main factors affecting VDWs?

A

more e-, more chance temporary dipoles will form, and more chance induced dipoles form, leading to stronger VDWs between molecules and greater boiling points

Shape of molecule - Long chain alkanes have a larger surface area of contact between molecules for VDW to form compared to branched ones, so have stronger VDW

33
Q

Explain the increase in boiling points down group 7 and in the alkane homologous series

A

Increasing no of e- in bigger molecules down groups
Larger VDW between them in bigger molecules
This is why I2 is a solid but Cl2 is a gas

34
Q

What is a permanent dipole dipole force and when does it occur? How is its strength relative to VDWs?
When do polar bonds form?

A

A force occuring between polar molecules with a permanent dipole (eg C-Cl, C=O)
Stronger than VDW
They form in asymmetrical polar molecules where there is a significant difference in electronegativity between the atoms

35
Q

When does hydrogen bonding occur? What always needs to be shown in H bond diagrams?

A

In compounds with a hydrogen attached to N,O or F
There is a large electronegativity difference

The Lone pair must be shown on the NOF, and all the partial charges must be shown

36
Q

In a diagram containing a graph of boiling point (y) vs molecular mass (x) of a range of compounds of elements bonded with hydrogen, why are H2O, NH3, and HF so anomolously high?
Why is there an increase in boiling point from H2S to H2Te?

A

Hydrogen bonding in those 3 molecules

Increasing VDWs between them due to increasing numbers of electrons

37
Q

Which 4 organic molecules can form hydrogen bonds?

A

alcohols
carboxylic acids
amides
proteins

38
Q

What is the arrangement of carbon atoms in diamond?

A

tetrahedral

39
Q

what is the arrangement of simple molecular molecule iodine?

A

regular arrangement of I-I held together by weak VDWs

40
Q

How to draw simple molecular ice diagram?
How is ice of a lower density then liquid water?

A

show central water molecule with 2 covalent bonds, and 2 hydrogen bonds in a tetrahedral arrangement

Molecules are held further apart in ice than liquid water explaining the lower density of ice

41
Q

electronegativity definition

A

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

42
Q

Electronegativity scale name, most electronegative element and it’s value on the scale

A

Pauling scale
Fluorine
4.0

43
Q

Trends of electronegativity across a period and down a group

A

Increases across a period BECAUSE
the number of protons increases
atomic radius decreases because electrons in the same shell are pulled in more

Decreases down a group BECAUSE
distance between (+vely charged) nucleus and electrons in pair increases and the shielding increases

44
Q

What determines a covalent bond?

A

The compound contains elements of a similar electronegativity, and hence a small electronegativity difference

45
Q

What determines an ionic bond?

A

The compound contains elements of a very different electronegativity, and hence a very large electronegativity difference

46
Q

How is a permanent dipole formed?

A

it forms due to an unequal distribution of electrons in the bond, and produces a charge separation due to the elements in the bond, having different electronegativities

47
Q

Why is CCl4 non polar?

A

It’s a symmetrical molecule (all identical bonds and no lone pairs) and all the dipoles cancel out
There is no net dipole
The molecule is non polar

48
Q

Explain why the oxidation state of chromium in Cr(PF3)6 is 0

A

PF3 is neutral
The complex is neutral overall

A-level Chem (35 decks)

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