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AS Chemistry > Chapter 4: Chemical Bonding And Structure > Flashcards

Chapter 4: Chemical Bonding And Structure Flashcards

1
Q

What bond is present in inert gases

A

Weak force of attraction

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2
Q

What does a bond mean

A

Binding force

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3
Q

What is a chemical bond

A

Force of attraction between particles

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4
Q

Different ways in which particleS bond and the result of each

A

•small or large numbers •similar or different atoms

Result of combination is a stable association called a chemical bond

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5
Q

How do bonded atoms differ from parent atoms

A

They have very different properties

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6
Q

Why are bonds formed

A

To achieve a more stable outer electronic configuration of 2 OR 8 electrons . This is the Octet rule

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7
Q

How can bonds form foe stable configuration

A

Losing, gaining, sharing, or freeing electrons in the outer shell

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8
Q

State the general number of electrons in metals outer shells and how they attain a stable atom

A

Small valance no. (1,2,3)

Lose electrons to form postive ions called cations

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9
Q

State the general number of electrons in nonmetals outer shells and how they attain a stable atom

A

Higher valance nos. (5,6,7)

Gain electrons to form negative ions called anions

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10
Q

What is the bonding tendency of group 4 elements and why

A

Do not gain or lose electrons so don’t form ions. Because the energy involved to transfer electrons is too large to be favourable. They react to form a different type of bond

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11
Q

What is the ionic bond model

A
  • electrostatic force between ions which are small, spherical, charged particles attracting eachother
  • electrons are transferred completely from on outer shell to another atom’s so do not physically affect electrons of each other
  • total charge for a formula unit of ionic substance is 0
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12
Q

What is an ionic bond

A

Electrostatic force of attraction between positive and negative ions

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13
Q

State the structure of ionic compounds

A
  • giant ionic lattice
  • lattice is an orderly arrangement of particles in a 3d space
  • the lattice is made of postive and negative ions arranged at regular distances beside each other continuously throughout the solid
  • crystals are solids with plane surfaces arranged att regular angles to each other
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14
Q

Describe the structure of NaCl lattice from memory

A

Chlorine- ions are relatively larger than Sodium+ ions

Each chlorine ion is surrounded by 6 sodium ions
Imagine that ^

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15
Q

What are 3 properties of ionic compounds and explain each

A
  • high melting and boiling points due to strong electrostatic force between oppositely charged ions
  • conduct electricity when molten or aqueous but not in solid state, because ions are free to move only when molten or dissolved in water
  • brittle and cleave (spilt as cut) along planes because ions of same charge come closer as crystal planes slide over each other so they repel strongly
  • soluble in polar (water) solvents but not in organic ones (benzene)
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16
Q

Describe how ionic compounds dissolve in water

A

Water is polar. It has a separation of charge though partial. At contact surface molecules are attracted to ions of opp charge in the lattice which may cause ions to dislodge from their positions and become surrounded by molecules are said to hydrated and hence dissolves

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17
Q

State the equation for NaCl dissolving

A

NaCl(s)+H2O–>NaCl(aq)

NaCl(s)+H2O–>Na+(aq) + Cl-(aq)

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18
Q

When is a liquid, other than water, can dissolve an ionic substance the solid is said to be

A

Solvated

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19
Q

Why are non polar substjaces not able to dissolve ionic solids

A

No charge separation. No attraction between liqued and ions so ions remain tightly bound to eachother

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20
Q

Solubility trends are based on

A

Similar chemical nature of solute and solvent to most likely lead to successful interactions

“Like dissolves like”

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21
Q

Why do elements covalently bond

A

When the ionization energies to obtain an inert gas configuration are very high, so much that a single ion of these elements becomes unstable

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22
Q

How is covalent bonding done

A

Non metal atoms share electrons so each atom has a stable configuration in the outer shell

The shared electrons are concentrated between and simultaneously attracted by both nuclei so they can contribute to each atoms structure.

This is only possible if electrons are strongly localised between atoms

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23
Q

What is a covalent bond

A

Electrostatic force of attraction between shared electrons and positive nuclei of the two atoms

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24
Q

A molecule is

A

Group of atoms held together by covalent bonds

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25
Q

What are valence electrons shared in covalent bonds called

A

Bond pairs

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26
Q

What is a pair of electrons not shared but help in giving octet is called

A

Lone pair

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27
Q

Explain the cleaving of an ionic lattice by recalling the diagram

A

+ - -
- + slide and cleave at planes –> + +
+ -

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28
Q

Whic values are covalent bonds characterized by

A

Bond length

Bond strength

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29
Q

What is bond length

A

Measure of distance between 2 bonded nuclei

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30
Q

What is bond strength

A

Measure of energy needed to break a bond ie bond enthalpy

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31
Q

State the trend down a group relating bond length and strength

A
  • atomic radius increases down a group atoms form molecules with longer bonds
  • as a result the shared electron pair is further from pull Of nuclei in larger molecules so bond is weaker and bond enthalpy decreases
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32
Q

State how bond length varies in single and multiple bonds

A
  • > Multiple bonds have greater no of shared electrons so have a stronger force of electrostatic attraction between bonded nuclei
  • > The greater pulling power on nuclei brings them closer resulting in shorter and stronger bonds than single
  • > Hence length in multiple bonds is shorter, but enthalpy is greater
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33
Q

What is a coordinate covalent bond

A

Covalent bond formed between 2 atoms in which one of the 2 atom donate a pair of electrons (from lone pair) to be shared between the 2 ATOKS

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34
Q

When is a coordinate bond formed

A

Between an atom which is short of a stable electronic config by a pair of e and another which has a lone pair to offer

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35
Q

What are properties of coordinate bonds

A

It is a typical covalent bond with covalent properties

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36
Q

How is a dative bond represented

A

An arrow pouting from donor to acceptor

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37
Q

What is the orbital theory

A

In electron sharing two atomic orbitals each carrying an unpaired electron overlap to form a single covalent bond

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38
Q

What is an orbital overlap

A

Part of atomic orbitals of two atoms that occupy the same space

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39
Q

How is a molecular orbital formed

A

From overlap of 2 atomic orbitals

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40
Q

Shape of molecular orbital depends on

A

Type of electron orbitals that overlap

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41
Q

What are 2 common molecular orbitals

A

Sigma

Pi

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42
Q

How is a sigma bond formed

A

Head on overlap of two atomic orbitals along their internuclear line, an imaginary axis, where bonding electrons are likely to be located.

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43
Q

In Which cases are single covalent bonds formed

A 
s-s 
s-p
p-p 
Overlaps 
Or 
Hybridized orbitals
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44
Q

How is a pi bond formed

A

Side to side or lateral overlap of two parallel p-orbitals of adjacent atoms.

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45
Q

State the difference in amount of overlapping in sigma and pi bonds

A

Pi is less

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46
Q

What makes pi bonds better than sigma

A

Weaker, hence more reactive because bonding electrons are localised away from the intermolecular axis

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47
Q

When can pi bonds form

A

Alongside sigma bonds

As in multiple bonds ONLY

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48
Q

State electron density/concentration in sigma and pi bonds

A

Sigma - between nuclei of bonded atoms

Pi - electron density is concentrated above and below the plane of bond axis

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49
Q

What is hybridisation

A

Process by which two or more atomic orbitals merge to form new orbitals of the same energy.

In other words redistribution of energy among orbitals to form new ones

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50
Q

Simply explain 2s-2p hybridisation

A

2 different orbitals in the same shell merge to form hybrid orbitals of same energy and showing characteristics of both s and p orbitals

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51
Q

The no of hybrid orbitals is =

A

Total no of atomic orbitals involved in hybridisation

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52
Q

What is sp hybridisation

A

Process in which one s and one p orbital combine to form two new orbitals of same energy. Angle between the sp orbitals is 180 degrees

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53
Q

How does sp hybridisation come about (eg in Be)

A

One of the 2s orbitals is promoted to a vacant p orbital. Single occupied s and p orbitals merge to form 2 new orbitals

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54
Q

What is sp2 hybridisation

A

Process in which one s orbitals combines with 2 p orbitals to form 3 similar sp2 orbitals. The angle between any two of these is 120 degrees

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55
Q

How does sp2 hybridisation happen eg Bf3

A

One 2s promoted in excited state to vacant p orbital

Single occupied s and p(2) merge to form 3 new

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56
Q

What is sp3 hybridisation

A

1 s and 3 p combine to form 4 new hybrid orbitals of the same energy.
Angle between the hybrids is 109.5

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57
Q

How does sp3 hybridisation happen

A

One of the 2s electrons is printed to vacant 2p. 1s and 3p form 4 new orbitals

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58
Q

How does sp3 hybridisation happen eg methane

A

One of the 2s electrons is printed to vacant 2p. 1s and 3p form 4 new orbitals

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59
Q

How can shapes of molecules and ions be determined

A

Valence shell electron pair repulsion theory

VSEPR theory

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60
Q

What is the VSEPR theory

A

Orbitals around central atom contain electron pairs, and these orbitals experience a force of repulsion. In this theory, electrons arrange themselves around the central atom so that they are as far apart as possible in order to have: minimum repulsion, maximum stability.

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61
Q

Why are multiple bonds counted as pairs of electrons

A

Electrons lie in the same direction

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62
Q

Shape of molecule depends on:

A

The way atoms in it occupy space
Electron pairs or domains(bond and lone)
Ie negative centres around central atom

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63
Q

State the repulsion intensity of each type of electron pairs

A

Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair repulsion

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64
Q

Another name for VSEPR

A

Electron cloud. Used to predict shapes and bond angles of molecule
Can be single or multiple bonds

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65
Q

Structure of diatonic molecule

A

Composed of 2 atoms

Linear in shape as the 2 centres are in straight line

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66
Q

A lone pair is

A

Non bonding pair of electrons

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67
Q

Electron clouds are _ charged and _ eachother

A

Negatively

Repel

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68
Q

What is bond angle

A

Angle which orbitals make with respect ti the central atom

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69
Q

Bond angle for 3 bond pairs is _ ie _

A

120°

Trigonal plana

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70
Q

Bond angle for all 4 bond pairs is _ and shape is _

A

109.5°

Tetrahedral

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71
Q

Bond angle for 5 bond pairs. Shape is

A

120 and 90

Trigonal bypyramidal

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72
Q

6 bonds pairs angle is. Shape is

A

90°

Octahedral

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73
Q

What is HCN and how are atoms in it bonded

A

Hydrogen cyanide

C is central atom. All covalently bonded. Carbon has no lone pairs (1e-H, 3e-N). Bond pairs repel, molecule is linear

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74
Q

HCl is diatomic, _ in shape, angle of _

A

Linear

180

75
Q

AlCl3 is _ in shape, because bond angle is _, and has _ bonds pairs and _ lone pairs

A

Plane triangular
120
3b
0l

76
Q

3 bonds pairs means x atoms, ie 1 central atom and y surrounding atoms

A

4c

3s

77
Q

Methane, _ , is _ in shape due to angle of _ and _ bond pairs

A

CH4
Tetrahedral
109.5
4

78
Q

How is bonding in a molecule of Ammonia

A

NH3
N is central atom. Valence 5, 3 of which bond with 3 Hydrogen atoms. These repel and get as far from each other as possible
2 are unused, ie the lone pair
Lone-bond repulsion is greater than bond-bond so lone pair pushes bond pairs closer than in a tetrahedral arrangement of electron clouds. This distorted tetrahedral is trigonal pyramidal
Bond angle is 107°

79
Q

What is the chemical bonding of water

A

H2O
O is central atom. 2 electrons of O used ti covalently bond with 2 H. These will repel.
4e valence of O will remain as 2 lone pairs, repel eachother equally.
Since lone-lone is greater than lone-bond or bond-bond repulsion, lone pairs repel and push bond pairs closer in distorted tetrahedral arrangement of electron clouds called bent.
Bond angle of bent is 104.5°

80
Q

Sulfur dioxide, SO2, has 4 bond pairs around S. What is the bond angle

A

4 bond pairs really count as 2. 1 in each direction. S has 1 lone pair still which repels bond pairs, and since it is greater than bond-bond they are pushed closer reducing bond angle from 120 (planar) to close to 118

81
Q

Structure of graphite

A

Each C is sp2 hybridized and covalently bonded to 3 others forming hexagons in layers with bond angles 120. Layers held by weak London dispersion forces so they can slide over each other

82
Q

Electrical conductivity of graphite

A

Good. Because it has one non bonded delocalised electron per atom that gives electron mobility

83
Q

Graphite thermal conductivity

A

Bad. Unless heat can be forced to conduct in direction parallel to layers

84
Q

Graphite thermal conductivity

A

Bad. Unless heat can be forced to conduct in direction parallel to layers

85
Q

Appearance of graphite

A

Not shiny, Grey, crystalline solid

86
Q

Special properties of graphite

A

Soft and slippery due to layers
Brittle
Very high melting point
Most stable carbon allotrope

87
Q

Uses of graphite

A

Lubricant
Pencils
Electrode rods in electrolysis

88
Q

Structure of diamond

A

Each C is sp3 hybridized and covalent bonded to 4 others. Tetrahedral in regular repetitive pattern. Bond angles 109.5

89
Q

Diamond electrical conductivity

A

Non conductor. All electrons bonded so non mobile

90
Q

Diamond thermal conductivity

A

Very efficient conductor, better than metals

91
Q

Diamond appearance

A

Highly transparent, shiny crystal

92
Q

Special properties of diamond

A

Hardest known natural substance
Cannot be scratched
Brittle
High melting point

93
Q

Uses of diamond

A

Polished for ornamentation

Tools and machinery for grinding and cutting glass

94
Q

Structure of fullerene C60

A

Each C is sp2 hybridized and bonded in a sphere of 60 C atoms, consisting of 12 Pentagon’s and 20 hexagons. Spherical cage in which every C is bonded to 3 others. Not a giant molecule, has fixed formula

95
Q

Fullerene electrical conductivity

A

Semiconductor at rtp due yo some electron mobility, easily accepts electrons to form negative ions

96
Q

Thermal conductivity of fullerene

A

Very low

97
Q

Appearance of fullerene

A

Yellow crystalline solid, soluble in benzene

98
Q

Special properties of fullerene

A

Very light and strong, reacts with potassium yo make superconducting crystalline material, low melting point

99
Q

Uses of fullerene

A

Lubricant, medical and industrial deviced for binding specific target molecules, related forms used to make nanotubes used as capacitors in electronics, and catalysts

100
Q

Structure of graphene

A

Each C is covalently bonded to 3 others forming hexagons with bond angles 120. But it is single layer and exists as 2d material. It is like a chicken wire

101
Q

Electrical conductivity of graphene

A

Very good, one delocalised electrons gives electron mobility across layers

102
Q

Graphene thermal conductivity

A

Best thermal conductor, better than diamond

103
Q

Appearance of Graphene

A

Almost completely transparent

104
Q

Special properties of Graphene

A

Thickness of just one atom thinnest of any material (light), but also strongest-100× steel, very flexible, very high melting point, inert,

105
Q

Uses of graphene

A

TEM, grids, high performance electronics, touch screens, many still being developed

106
Q

Electron domain geometry 2,3,4

A

2: 180, linear
3: 120, triangular planar,
4: 109.5, tetrahedral

107
Q

How is a nanotubes made

A

Graphene, single separated layer of graphite. Rolled up and closed with half a fullerene.

108
Q

Silicon structure

A

Group 4 element, valancy 4, each silicon atom bonded to four others tetrahedrally. Giant lattice like diamond

109
Q

SiO2 aka _ structure is _ ie…

A

Silica/quartz
Tetrahedral giant lattice
1 silicon atom + 4 oxygen atoms, and each oxygen to 2 silicon atoms

110
Q

SiO2 refers to

A

Ratio of atoms

1Si : 2O

111
Q

Properties of SiO2 and reason

A

*strong
*insoluble
*high melting point
*non conductor of electricity
Because atoms are strongly held tetrahedral positions that involve all four silicon valance electrons

112
Q

What is polarization

A

Distortion of charge cloud around anion by cation as cations pull electrons away from anion which are not strongly held by nuclei

113
Q

Ability of cation to polarize anion is called it’s _

Ability of anion to get polarized is called

A

Polarizing power

Polarizability

114
Q

What does polarization introduce?

A

Covalent character yo ionic bonds ie intermediate character as electrons appear yo be shares between cation and anion as electrons are pulled towards cation

115
Q

Ability of cation to polarize anion depends on

A

Strength of electric field used to pull electrons away from anion.
Electric field is based on charge density of cation.
Charge density is proportional to the charge to radius ratio, q/r, of cation.

So ion with small radius and high charge has high charge density and polarizing power

116
Q

Polarization in ionic bond is increased by which factors that introduce covakent character

A
  • small cation radius with high ionic charge

* large anion radius with high ionic charge

117
Q

What is electronegativity

A

Ability of atom to pull/attract shared electrons in a covakent bond towards itself when it’s in a molecule.

118
Q

Main factor influencing electronegativity is

A

Atomic radius

119
Q

Electronegativity trend across period

A

Effective nuclear charge increases, atomic radius decreases. So attraction for bonded electrons and hence electronegativity increases

120
Q

Electronegativity trend across period

A

Effective nuclear charge increases, atomic radius decreases. So attraction for bonded electrons and hence electronegativity increases

121
Q

Electronegativity trend down a group

A

Atomic radius increases. So attraction for bonded electrons reduces and hence electronegativity decreases

122
Q

Most electronegative atom :

Least electronegative atom :

A

Fluorine

Caesium

123
Q

A pure covalent bond between 2 atoms of same element, where both atoms have _ and _

A

Same electronegativity and bonded electrons are equally shared

124
Q

Covalent bond between 2 different atoms

A

Bonding electrons not shared equally due yo electronegativity difference, so more electronegative one pulls shared electrons more. This more electronegative atoms acquires slight - charge, and the less electronegative one gains partial + charge. It is a polar covalent bond

125
Q

Polar covalent bond introduces

A

Ionic character in covalent bond

126
Q

What is a polar covalent bond

A

Where there is unequal sharing of electrons

127
Q

What makes a compound behave more ionicly

A

More the polarity, due yo more electronegativity difference

128
Q

What estimated electronegativity difference makes a bond more ionic than covalent

A

1.7

129
Q

What is a dipole (moment)

A

Measure of the polarity of a bond. Product of charge on the atoms of bond divided by distance between charges

130
Q

Polar molecules are also called … Due to …

A

Dipoles

Separated opposite electric charges

131
Q

Apart from representing dipoles with p+ or p- what is other way

A

Arrow/vector showing pull direction to more electronegative from less

132
Q

What are instantaneous dipoles

A

Weak intermolecular forces.
Aka Induced dipole forces. Aka London dispersion forces.

Electron cloud movements result in temporary dipole, as charge is more on one side of molecule for a moment.
This induces dipole on neighbouring molecule. As a result there are forces of attraction between p+ and p- between dipoles.

133
Q

Strength of temporary dipole increases with (2)

A
  • increase in molecular size, ie more no of electrons per molecule increases hence the probability of developing tempo dipoles
  • increas in no of contact points-where molecules come close together
134
Q

Straight chain isomers have _ boiling points than _ isomers because…

A

Higher
Branched
Are more closely points ie have more contact points hence stronger stronger induced dipole forces

135
Q

Like dissolves like. Elaborate.

A

Polar solutes are soluble in polar solvents. Because dipole-dipole interactions occur between the 2 polar molecules

136
Q

van der Waal’s forces refers to _

A 
Forces between molecules that do not involve electrostatic interactions between ions.
These are: 
-london dispersion forces 
-dipole-dipole 
(-dipole-induced)
137
Q

Hydrogen bonds are strongest or weakest of intermolecular forces?

A

Strongest

138
Q

Hydrogen bonds can be formed where? Ie in molecules, and compounds

A

Intermolecular or intramolecular

Compounds of hydrogen with highly electronegative element eg F, O, N

139
Q

Why can H bonds be formed in compounds of hydrogen with highly electronegative element eg F, O, N

A

It difficult to polarise H atom.
It is small, has one electron closest to nucleus and no shielding electrons to lessen effect of proton. Hence a super electronegative atom to attract it.

And hydrogen exerts strong attractive force on lone pair of electronegative atom of neighbouring molecule

140
Q

H bonds is in essence a particular case of

A

Dipole-dipole attraction

141
Q

Strength of hydrogen bonds determined by :

A

1) extent of polarization within molecule ie electronegativity
2) electronic structure of atom ie having lone pair to point to H
3) angle between 3 atoms involved. Strongest at 180. When less than 140 H bonding not possible

142
Q

Define a hydrogen bond

A

Intermolecular force of attraction between a very polar hydrogen atom of one molecule and highly electronegative atom of another molecule

143
Q

H bond can be formed between molecules when

A

One molecule has H atom covalently bonded to highly electronegative atom

Other molecule has an atom which is also electronegative, small in size, carries lone pair

144
Q

Variation in bp of hydrides of groups 4,5,6,7 as a consequence of hydrogen bonding

A

In general, all four groups show increase in bp down the group (increasing periods) as molar mass increases.Anomalies being NH3, HF, H2O, which are in P1 but have highest bp which does not correspond with their molar mass. This can only be explained by hydrogen bonding in these molecules

145
Q

Water is liquid at room temp due to

A

Hydrogen bonds

146
Q

Compare bp of organic molecules of same molar mass but yet have diff bp due to

A

Hydrogen bonding between molecules ie isomers

147
Q

Anomalous behavior of water is…

A

Ice has lower density than water and floats.
Water freezes to ice and increases volume, it contracts when temp is 0-4°C instead of expand. Water has max density at 4. This is unexpected as liquids increase in volume when temp increased/vv

148
Q

Explain the anomalous behavior of water

A

Due to crystal structure of ice. In solid ice molecules are near to form H bonds. Each water molecule surrounded by 4 others, tetrahedrally. Results in open spaces in ice lattice. Ice occupies more volume than the sane mass of water does. Hence density lower than liquid water.
When water heated from 0-4 H bonds break and free molecules escape into spaces causing shrink in volume and increase in density. This change dominates normal volume expansion up to 4°C. Father temp inc change is usual ie volume is increased

149
Q

Unexpected high mp and bp of water are due to

A

Hydrogen bonds

150
Q

Water exists as the only…

A

Only liquid around hydrides of groups 4,5,6,7

151
Q

Strong hydrogen bonds in water produce high _, as a result droplets are _ and water bends yo form _ in small tubes. Water rises up tunes as pressure is exerted, called _

A

Surface tension
Spherical
Meniscus
Capillary action

152
Q

Waters acts as solvent by

Ie hydration of ionic substances

A

Making and breaking of hydrogen bonds

153
Q

Hydrogen bonds have unique role in _

Eg

A

Structure determination

Carbs, proteins, nucleic acids eg DNA held together by them

154
Q

Recall water structure

A

K

155
Q

Properties of simple molecular structures

A 
#exist as solid, liquid or gas with weak intermolecular forces 
#volatile, with low melting and boiling points due to weak intermolecular force
#bad conductors of electricity
#soluble in non polar solvents and insoluble in polar ones eg water
156
Q

Properties of giant covalent/molecular structures

A 
#hard
#high melting and boiling points
#bad conductors of electricity
157
Q

Physical properties of molecular substances are governed by

A

Intermolecular forces-force attracting a molecule to neighbours - van der Waal’s or hydrogen
Not covalent. These are irrelevant

158
Q

Molecular substances tend to exist as _ bcos of _

A

Gas, liquid, or low mp solid
Weak intermolecular forces.
No covalent bonds need to be broken

159
Q

Size of mp/bp depends on _. These are _ and hence _ to break them

A

Strength of intermolecular forces

  • hydrogen bonds present
  • larger molecule, more vanderwaals attractions possible

More energy needed

160
Q

Molecular substances that do dissolve in water…

A

React with it or form H bonds

161
Q

Why doesn’t methane dissolve?

A

It is a gas, with separated moleculesm no need for water to pull apart
But hydrogen bonds between water, so methane can’t force its way in between

162
Q

Why does ammonia dissolve in water

A

Can form H bonds.
Those between water broken and replaced by equivalent bonds
Also reacts to form NH4+ and OH-

163
Q

Molecular substances can’t conduct electricity. Bcos… Even if there are

A

No free electrons

Even if there are delocalised electrons, not enough contact between molecules to allow electrons to move through whole solid/liquid

164
Q

Volatility of ionic, polar c, non polar c, giant c

A

I: low
P: higher
NP: highest
G: L

165
Q

Solubility of ionic, polar c, non polar c, giant c in polar solvent

A

I: sol
P: sol increases as polarity increases
NP: non sol
G: non sol

166
Q

Solubility of ionic, polar c, non polar c, giant c in nonpolar solvent

A

I: nonsol
P: sol increases as polarity decreases
NP: sol
G: non sol

167
Q

Electrical conductivity of ionic, polar c, non polar c, giant c

A

I: when molten or dissolved l/aq
P: not
NP: not
G: non except graphite, graphene. Semi Si and fullerene

168
Q

Metallic bond formation

A

Metals are elements with low electronegativity
Electrons in vacancy shells loosely held to nucleus and kernel(inner shells). These move away into vacant orbitals
Giant network of regularly spaced +ve ions surrounded by sea of free delocalised electrons

Formed by electrostatic attraction between + ions and free electrons

169
Q

Strength of metallic bond depends on

A
  • atomix radius
  • no of valance electrons delocalised
  • charge on cation

Greater no of free electrons, smaller the cation, greater the binding force between them

170
Q

Strength of metallic bonds confirmed by

A

Melting points

171
Q

Trend of mp of metals

A

Decrease down group as strength of metallic bonds decrease due to cation size increases, reducing attraction between +charges and e-

172
Q

Why do transition elements have strong metallic bonds

A

Large no of electrons that become delocalised

173
Q

Why are metals good conductors of electricity

A

Presense of mobile delocalised electrons in metallic lattice. Electrons move along wires and make up electric current. Flow of e into metal pushes the e near to it, pushing from one end of metal yo other. Electrons are free to move hence

174
Q

Metals are good heat conductors because

A

Heating one end increases kinetic energy of delocalised electrons. They vibrate to hit other nearby e. This continues from hotter end to colder until electrons transfer heat ti colder end. Metal is heat carrier

175
Q

Metals are malleable because

A

Diff metals have diff atom arrangement in structure. In each there’s strong metallic bonds between fixed +ve ions which in layers. When strong force applied, layers slip over to new positions keeping metallic bond intact. Or delocalised electrons move nondirectionally and randomly through lattice, bond intact. Thickness is reduced. Can similarly be bent/moulded into shapes

176
Q

Metals are ductile because

A

Slipping of layers of atoms over each other when metal is pulled

177
Q

High mp of metals due to

A

Lot of energy needed to break metallic bonds and separate atoms

178
Q

Shiny appearance of metals due to

A

Delocalised electrons in crystal structure reflect light

179
Q

Application of metal electrical conductivity

A

Electrical circuits use copper

180
Q

Application of metal thermal conductivity

A

Cooking utensils

181
Q

Application of metal malleability and ductility

A

M-moulded to form machinery and structural components of building cars etc

D-electric wires n cables

182
Q

Application of metal high bp

A

High speed tools and turbine engines

Tungsten has highest bp

183
Q

Application of metal being shiny

A

Ornaments

AS Chemistry (11 decks)

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