Chapter 4: Chemical Bonding And Structure Flashcards
(183 cards)
1
Q
What bond is present in inert gases
A
Weak force of attraction
2
Q
What does a bond mean
A
Binding force
3
Q
What is a chemical bond
A
Force of attraction between particles
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
5
Q
How do bonded atoms differ from parent atoms
A
They have very different properties
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
7
Q
How can bonds form foe stable configuration
A
Losing, gaining, sharing, or freeing electrons in the outer shell
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
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
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
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
12
Q
What is an ionic bond
A
Electrostatic force of attraction between positive and negative ions
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
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 ^
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)
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
17
Q
State the equation for NaCl dissolving
A
NaCl(s)+H2O–>NaCl(aq)
NaCl(s)+H2O–>Na+(aq) + Cl-(aq)
18
Q
When is a liquid, other than water, can dissolve an ionic substance the solid is said to be
A
Solvated
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
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”
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
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
23
Q
What is a covalent bond
A
Electrostatic force of attraction between shared electrons and positive nuclei of the two atoms
24
Q
A molecule is
A
Group of atoms held together by covalent bonds
25
What are valence electrons shared in covalent bonds called
Bond pairs
26
What is a pair of electrons not shared but help in giving octet is called
Lone pair
27
Explain the cleaving of an ionic lattice by recalling the diagram
+ - -
- + slide and cleave at planes --> + +
+ -
28
Whic values are covalent bonds characterized by
Bond length
| Bond strength
29
What is bond length
Measure of distance between 2 bonded nuclei
30
What is bond strength
Measure of energy needed to break a bond ie bond enthalpy
31
State the trend down a group relating bond length and strength
- 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
32
State how bond length varies in single and multiple bonds
* >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
33
What is a coordinate covalent bond
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
34
When is a coordinate bond formed
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
35
What are properties of coordinate bonds
It is a typical covalent bond with covalent properties
36
How is a dative bond represented
An arrow pouting from donor to acceptor
37
What is the orbital theory
In electron sharing two atomic orbitals each carrying an unpaired electron overlap to form a single covalent bond
38
What is an orbital overlap
Part of atomic orbitals of two atoms that occupy the same space
39
How is a molecular orbital formed
From overlap of 2 atomic orbitals
40
Shape of molecular orbital depends on
Type of electron orbitals that overlap
41
What are 2 common molecular orbitals
Sigma
| Pi
42
How is a sigma bond formed
Head on overlap of two atomic orbitals along their internuclear line, an imaginary axis, where bonding electrons are likely to be located.
43
In Which cases are single covalent bonds formed
```
s-s
s-p
p-p
Overlaps
Or
Hybridized orbitals
```
44
How is a pi bond formed
Side to side or lateral overlap of two parallel p-orbitals of adjacent atoms.
45
State the difference in amount of overlapping in sigma and pi bonds
Pi is less
46
What makes pi bonds better than sigma
Weaker, hence more reactive because bonding electrons are localised away from the intermolecular axis
47
When can pi bonds form
Alongside sigma bonds
| As in multiple bonds ONLY
48
State electron density/concentration in sigma and pi bonds
Sigma - between nuclei of bonded atoms
| Pi - electron density is concentrated above and below the plane of bond axis
49
What is hybridisation
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
50
Simply explain 2s-2p hybridisation
2 different orbitals in the same shell merge to form hybrid orbitals of same energy and showing characteristics of both s and p orbitals
51
The no of hybrid orbitals is =
Total no of atomic orbitals involved in hybridisation
52
What is sp hybridisation
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
53
How does sp hybridisation come about (eg in Be)
One of the 2s orbitals is promoted to a vacant p orbital. Single occupied s and p orbitals merge to form 2 new orbitals
54
What is sp2 hybridisation
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
55
How does sp2 hybridisation happen eg Bf3
One 2s promoted in excited state to vacant p orbital
| Single occupied s and p(2) merge to form 3 new
56
What is sp3 hybridisation
1 s and 3 p combine to form 4 new hybrid orbitals of the same energy.
Angle between the hybrids is 109.5
57
How does sp3 hybridisation happen
One of the 2s electrons is printed to vacant 2p. 1s and 3p form 4 new orbitals
58
How does sp3 hybridisation happen eg methane
One of the 2s electrons is printed to vacant 2p. 1s and 3p form 4 new orbitals
59
How can shapes of molecules and ions be determined
Valence shell electron pair repulsion theory
VSEPR theory
60
What is the VSEPR theory
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.
61
Why are multiple bonds counted as pairs of electrons
Electrons lie in the same direction
62
Shape of molecule depends on:
The way atoms in it occupy space
Electron pairs or domains(bond and lone)
Ie negative centres around central atom
63
State the repulsion intensity of each type of electron pairs
Lone pair-lone pair > lone pair-bond pair > bond pair-bond pair repulsion
64
Another name for VSEPR
Electron cloud. Used to predict shapes and bond angles of molecule
Can be single or multiple bonds
65
Structure of diatonic molecule
Composed of 2 atoms
| Linear in shape as the 2 centres are in straight line
66
A lone pair is
Non bonding pair of electrons
67
Electron clouds are _ charged and _ eachother
Negatively
| Repel
68
What is bond angle
Angle which orbitals make with respect ti the central atom
69
Bond angle for 3 bond pairs is _ ie _
120°
| Trigonal plana
70
Bond angle for all 4 bond pairs is _ and shape is _
109.5°
| Tetrahedral
71
Bond angle for 5 bond pairs. Shape is
120 and 90
| Trigonal bypyramidal
72
6 bonds pairs angle is. Shape is
90°
| Octahedral
73
What is HCN and how are atoms in it bonded
Hydrogen cyanide
| C is central atom. All covalently bonded. Carbon has no lone pairs (1e-H, 3e-N). Bond pairs repel, molecule is linear
74
HCl is diatomic, _ in shape, angle of _
Linear
| 180
75
AlCl3 is _ in shape, because bond angle is _, and has _ bonds pairs and _ lone pairs
Plane triangular
120
3b
0l
76
3 bonds pairs means x atoms, ie 1 central atom and y surrounding atoms
4c
| 3s
77
Methane, _ , is _ in shape due to angle of _ and _ bond pairs
CH4
Tetrahedral
109.5
4
78
How is bonding in a molecule of Ammonia
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
What is the chemical bonding of water
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
Sulfur dioxide, SO2, has 4 bond pairs around S. What is the bond angle
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
Structure of graphite
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
Electrical conductivity of graphite
Good. Because it has one non bonded delocalised electron per atom that gives electron mobility
83
Graphite thermal conductivity
Bad. Unless heat can be forced to conduct in direction parallel to layers
84
Graphite thermal conductivity
Bad. Unless heat can be forced to conduct in direction parallel to layers
85
Appearance of graphite
Not shiny, Grey, crystalline solid
86
Special properties of graphite
Soft and slippery due to layers
Brittle
Very high melting point
Most stable carbon allotrope
87
Uses of graphite
Lubricant
Pencils
Electrode rods in electrolysis
88
Structure of diamond
Each C is sp3 hybridized and covalent bonded to 4 others. Tetrahedral in regular repetitive pattern. Bond angles 109.5
89
Diamond electrical conductivity
Non conductor. All electrons bonded so non mobile
90
Diamond thermal conductivity
Very efficient conductor, better than metals
91
Diamond appearance
Highly transparent, shiny crystal
92
Special properties of diamond
Hardest known natural substance
Cannot be scratched
Brittle
High melting point
93
Uses of diamond
Polished for ornamentation
| Tools and machinery for grinding and cutting glass
94
Structure of fullerene C60
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
Fullerene electrical conductivity
Semiconductor at rtp due yo some electron mobility, easily accepts electrons to form negative ions
96
Thermal conductivity of fullerene
Very low
97
Appearance of fullerene
Yellow crystalline solid, soluble in benzene
98
Special properties of fullerene
Very light and strong, reacts with potassium yo make superconducting crystalline material, low melting point
99
Uses of fullerene
Lubricant, medical and industrial deviced for binding specific target molecules, related forms used to make nanotubes used as capacitors in electronics, and catalysts
100
Structure of graphene
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
Electrical conductivity of graphene
Very good, one delocalised electrons gives electron mobility across layers
102
Graphene thermal conductivity
Best thermal conductor, better than diamond
103
Appearance of Graphene
Almost completely transparent
104
Special properties of Graphene
Thickness of just one atom thinnest of any material (light), but also strongest-100× steel, very flexible, very high melting point, inert,
105
Uses of graphene
TEM, grids, high performance electronics, touch screens, many still being developed
106
Electron domain geometry 2,3,4
2: 180, linear
3: 120, triangular planar,
4: 109.5, tetrahedral
107
How is a nanotubes made
Graphene, single separated layer of graphite. Rolled up and closed with half a fullerene.
108
Silicon structure
Group 4 element, valancy 4, each silicon atom bonded to four others tetrahedrally. Giant lattice like diamond
109
SiO2 aka _ structure is _ ie...
Silica/quartz
Tetrahedral giant lattice
1 silicon atom + 4 oxygen atoms, and each oxygen to 2 silicon atoms
110
SiO2 refers to
Ratio of atoms
| 1Si : 2O
111
Properties of SiO2 and reason
*strong
*insoluble
*high melting point
*non conductor of electricity
Because atoms are strongly held tetrahedral positions that involve all four silicon valance electrons
112
What is polarization
Distortion of charge cloud around anion by cation as cations pull electrons away from anion which are not strongly held by nuclei
113
Ability of cation to polarize anion is called it's _
| Ability of anion to get polarized is called
Polarizing power
| Polarizability
114
What does polarization introduce?
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
Ability of cation to polarize anion depends on
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
Polarization in ionic bond is increased by which factors that introduce covakent character
* small cation radius with high ionic charge
| * large anion radius with high ionic charge
117
What is electronegativity
Ability of atom to pull/attract shared electrons in a covakent bond towards itself when it's in a molecule.
118
Main factor influencing electronegativity is
Atomic radius
119
Electronegativity trend across period
Effective nuclear charge increases, atomic radius decreases. So attraction for bonded electrons and hence electronegativity increases
120
Electronegativity trend across period
Effective nuclear charge increases, atomic radius decreases. So attraction for bonded electrons and hence electronegativity increases
121
Electronegativity trend down a group
Atomic radius increases. So attraction for bonded electrons reduces and hence electronegativity decreases
122
Most electronegative atom :
| Least electronegative atom :
Fluorine
| Caesium
123
A pure covalent bond between 2 atoms of same element, where both atoms have _ and _
Same electronegativity and bonded electrons are equally shared
124
Covalent bond between 2 different atoms
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
Polar covalent bond introduces
Ionic character in covalent bond
126
What is a polar covalent bond
Where there is unequal sharing of electrons
127
What makes a compound behave more ionicly
More the polarity, due yo more electronegativity difference
128
What estimated electronegativity difference makes a bond more ionic than covalent
1.7
129
What is a dipole (moment)
Measure of the polarity of a bond. Product of charge on the atoms of bond divided by distance between charges
130
Polar molecules are also called ... Due to ...
Dipoles
| Separated opposite electric charges
131
Apart from representing dipoles with p+ or p- what is other way
Arrow/vector showing pull direction to more electronegative from less
132
What are instantaneous dipoles
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
Strength of temporary dipole increases with (2)
* 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
Straight chain isomers have _ boiling points than _ isomers because...
Higher
Branched
Are more closely points ie have more contact points hence stronger stronger induced dipole forces
135
Like dissolves like. Elaborate.
Polar solutes are soluble in polar solvents. Because dipole-dipole interactions occur between the 2 polar molecules
136
van der Waal's forces refers to _
```
Forces between molecules that do not involve electrostatic interactions between ions.
These are:
-london dispersion forces
-dipole-dipole
(-dipole-induced)
```
137
Hydrogen bonds are strongest or weakest of intermolecular forces?
Strongest
138
Hydrogen bonds can be formed where? Ie in molecules, and compounds
Intermolecular or intramolecular
| Compounds of hydrogen with highly electronegative element eg F, O, N
139
Why can H bonds be formed in compounds of hydrogen with highly electronegative element eg F, O, N
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
H bonds is in essence a particular case of
Dipole-dipole attraction
141
Strength of hydrogen bonds determined by :
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
Define a hydrogen bond
Intermolecular force of attraction between a very polar hydrogen atom of one molecule and highly electronegative atom of another molecule
143
H bond can be formed between molecules when
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
Variation in bp of hydrides of groups 4,5,6,7 as a consequence of hydrogen bonding
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
Water is liquid at room temp due to
Hydrogen bonds
146
Compare bp of organic molecules of same molar mass but yet have diff bp due to
Hydrogen bonding between molecules ie isomers
147
Anomalous behavior of water is...
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
Explain the anomalous behavior of water
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
Unexpected high mp and bp of water are due to
Hydrogen bonds
150
Water exists as the only...
Only liquid around hydrides of groups 4,5,6,7
151
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 _
Surface tension
Spherical
Meniscus
Capillary action
152
Waters acts as solvent by
| Ie hydration of ionic substances
Making and breaking of hydrogen bonds
153
Hydrogen bonds have unique role in _
| Eg
Structure determination
| Carbs, proteins, nucleic acids eg DNA held together by them
154
Recall water structure
K
155
Properties of simple molecular structures
```
#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
Properties of giant covalent/molecular structures
```
#hard
#high melting and boiling points
#bad conductors of electricity
```
157
Physical properties of molecular substances are governed by
Intermolecular forces-force attracting a molecule to neighbours - van der Waal's or hydrogen
Not covalent. These are irrelevant
158
Molecular substances tend to exist as _ bcos of _
Gas, liquid, or low mp solid
Weak intermolecular forces.
No covalent bonds need to be broken
159
Size of mp/bp depends on _. These are _ and hence _ to break them
Strength of intermolecular forces
- hydrogen bonds present
- larger molecule, more vanderwaals attractions possible
More energy needed
160
Molecular substances that do dissolve in water...
React with it or form H bonds
161
Why doesn't methane dissolve?
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
Why does ammonia dissolve in water
Can form H bonds.
Those between water broken and replaced by equivalent bonds
Also reacts to form NH4+ and OH-
163
Molecular substances can't conduct electricity. Bcos... Even if there are
No free electrons
Even if there are delocalised electrons, not enough contact between molecules to allow electrons to move through whole solid/liquid
164
Volatility of ionic, polar c, non polar c, giant c
I: low
P: higher
NP: highest
G: L
165
Solubility of ionic, polar c, non polar c, giant c in polar solvent
I: sol
P: sol increases as polarity increases
NP: non sol
G: non sol
166
Solubility of ionic, polar c, non polar c, giant c in nonpolar solvent
I: nonsol
P: sol increases as polarity decreases
NP: sol
G: non sol
167
Electrical conductivity of ionic, polar c, non polar c, giant c
I: when molten or dissolved l/aq
P: not
NP: not
G: non except graphite, graphene. Semi Si and fullerene
168
Metallic bond formation
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
Strength of metallic bond depends on
- 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
Strength of metallic bonds confirmed by
Melting points
171
Trend of mp of metals
Decrease down group as strength of metallic bonds decrease due to cation size increases, reducing attraction between +charges and e-
172
Why do transition elements have strong metallic bonds
Large no of electrons that become delocalised
173
Why are metals good conductors of electricity
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
Metals are good heat conductors because
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
Metals are malleable because
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
Metals are ductile because
Slipping of layers of atoms over each other when metal is pulled
177
High mp of metals due to
Lot of energy needed to break metallic bonds and separate atoms
178
Shiny appearance of metals due to
Delocalised electrons in crystal structure reflect light
179
Application of metal electrical conductivity
Electrical circuits use copper
180
Application of metal thermal conductivity
Cooking utensils
181
Application of metal malleability and ductility
M-moulded to form machinery and structural components of building cars etc
D-electric wires n cables
182
Application of metal high bp
High speed tools and turbine engines
| Tungsten has highest bp
183
Application of metal being shiny
Ornaments
