bonding Flashcards
1
Q
define metallic bonding
A
- The electrostatic force of attrcation between the nuclei metal cations and a sea of delocalised electrons
- Electrons are free to move throughout the structure
2
Q
2bp 0lp
A
linear
180
3
Q
3bp 0lp
A
trigonal planar
120
4
Q
2bp 1lp
A
bent
118
5
Q
4bp 0lp
A
tetrahedral
109.5
6
Q
3bp 1lp
A
trigonal pyramidal
107
7
Q
2bp 2lp
A
bent
104.5
8
Q
5bp 0lp
A
trigonal bipyramidal
120,90
9
Q
4 bp, 1 lp
A
see-saw
119, 89
10
Q
3bp 2lp
A
t-shape
120, 89
11
Q
6 bp 0lp
A
octahedral
90
12
Q
5bp 1lp
A
square pyramid
89
13
Q
4bp 2lp
A
90
14
Q
phydical properties of metalic bonding
A
- high melting temperatures
- electrical conductivity
- thermal conductivity
- malleability and ductility
15
Q
describe how high melting temperatured are a property of metalic bonding
A
- Metals have giant lattice structure with many forces to be overcome
- Bonds between the metal cations and delocalised electrons take a lot of energy to break therefore have a high melting point
- Number of delocalised electrons per cation determine the melting temperature of a metal eg. Group 1 metals have low melting points because each cation donates only 1 electron, group.gp 0 2 have higher and metals in the d-block have high melting temps because they have more delocalised electrons
- Another factor that affects the melting temp is size of cation, the smaller the cation the closer the delocalised electrons to the cation so and increase in forces of attraction between nuceli and delocalised electron and so an increase in melting temperature
16
Q
describe how electrical conductivity is a property of metalic bonding
A
- Conductive because the delocalised are free to move throughout the metal carrying charge
- Will be attracted to the positive end of the metal
- Movement of electrons constitutes and electric current
17
Q
describe how thermal conductivity is a property of metalic bonding
A
- Free moving delocalised electrons pass kinetic energy along the metal
- Cations are closely packed and pass kinetic energy from one cation to another
18
Q
describe how maleability and ductility is a property of metalic bonding
A
- Metals can be hammered or pressed into different shapes (malleability), and drawn into a wire (ductility)
- When a stress is applied to a metal the layers of cations may slide over one another
- Because delocalised electrons are free moving they move with the cation an prevent strong forces of repulsion forming between the cations in one layer and the cations in another layer
19
Q
factors affecting strength fo metalic bonding
A
- More valance e- give stringer bonding as we have a higher charge cation and so more e- in the cloud
- Smaller metal ions giver stronger bonding because their nuceli are closer to the e- close and they can pack more tightly
20
Q
definition of ionic bonding
A
- The strong electrostatic force of attrcation between positive cations and negative anions
- Arranged in a regular lattice structure
- Electrostatic force of attraction isn’t directional so all ions feel some electrostatic force of attraction between each othe
21
Q
physical properties of ionic bonding
A
- high melting point
- brittleness
- electrical conductivity
- solubility
22
Q
describe how high melting points are a property of ionic bonding
A
- Ion solids consist of a giant lattice network of oppositely charged ions
- Many ions and the combined elctrostatic force of attraction among al of the ions is large so a large amount of enegry is required to overcome the forces of attraction between ions and break out of the giant lattice structure and slide past each other
23
Q
describe how brittless is a property of ionic bonding
A
- If stress applied to an ionic solid the layers of ions may slide over one another, ions of the same charge are now side by side and repel one another so the ionic lattice sturcture breaks
24
Q
describe how electrical conductivity is a property of ionic bonding
A
- As a solid don’t conduct electricity since no delocalised electrons and ions are not free to move since in a giant ionic lattice structure
- When molten or aqueous will conduct since the ions are now mobile
Exception is solid lithium nitride will conduct because of electro
25
describe how solubility is a property of ionic bonding
- Many ionic compounds are soluble in water
- Ionic boning broken by ion-dipole interactions forming between H2O molecules and seperate ions, hydration - water moelcules surround the ion
26
Evidence for existence of ions
- Ability for an ionic compound to conduct when molten or aqueous
- Eg when a direct current is passed through molten sodium chloride sodium is formed at the negative electrode and chlorine is formed at the posiitve electrode
- This is because positive sodium ions migrate towards the negatuve electrode where they gain electrons and become the sodium atoms
- Negative chloride ions migrate towards the positive electrode where they lose electrons and become chloride molecules
27
definition of ionic radii
a measure of the space occupied by an ion in a crystal lattice. Ionic radii are also useful in explaining and understanding chemical properties
28
describe the change in ionic radii accross a period
decrease in radius as number of protons increases
number of shells stays the same
increase in nuclear charge so stronger forces of attraction between nucleus and outer electron
29
what does isolectronic mean
ions that have the same electronic configuration
30
describe the change in ionic radii down a group
increases because numbe ro fshells increases
31
why are the ionic radii of positive ions are smaller than their corresponding atomic radii
Because they have lost their outer shell electrons so have fewer electron shells.
32
why are the ionic radii of negative ions are greater than their corresponding atomic radii
Although they have the same numbers of shells as the corresponding atoms, they have more electrons so there are increased repulsions and more electrons being attracted to the same nuclear charge.
33
lattic energy definition
the energy change when one mole of an ionic solid is formed from its gaseous ions.
More energy released means stronger ionic bonding.
34
defintion of covalent bonding
the ellectrostatic force of attraction between two nuclei and the shared bonded pairs
35
waht is dative bonding
when one atom in a covalent bond shares a pair of electrons
36
37
whats a pi bond
sideways overlap of two p-orbitals
only exist after a sigma bond is formed
weaker than sigma bond
38
reltionship etween bond length and bond strength
the shorter the bond the greater the bond strength
result in increase in electtrostatic attaction between 2 nuclei and electros in overlapping atomic orbitals
39
what is electrongetativity
th ebality of an atom to attrcat a onding pair of electrons
decreases down a goroup
increase accross a period
40
continum of bonding
0-0.5 - pure covalent
0.5-1.7 - polar covalent
1.1.7-2 - ioinic with covalent character
2 - 4
41
how to find out if a bond is polar
electronegativity difference has ot be equal or greater than to 0.5
42
how to figure out if a molecule is polar or not
whether the molecule is symmetrical or assymetrical
do the charges cancel out
43
how to describe the shape of a molecule
draw the shape of molecules/ 3D diagram
*(central atom)* has *_x _* bondign pairs and *_x _* lone pairs. These are arrabged ot maxime seperation/minimsie repulsion. Shape *_x _*, bond angle *_x _*
44
strength of repulsion in order
lone pair + lone pair > bond pair + lone pair > bond pair + bond pair
45
waht si the valance shell electron repulsion theory
* The shapes of molecules are caused by the repulsion between the pairs of electrons, both bonding and lone pairs, that surround the central atom.
* The electrons pairs arrange themselves around the central atom to minimise repulsions and maximise seperation
46
types of intermollecular forces
- london force (induced dipole induced dipole interactions)
- permanant dipole interactions
- hydrogen bonding
47
describe london forces
- The electron charge cloud in non-polar molecules or atoms are constantly moving
- During this movement, the electron charge cloud can be more on one side of the atom or molecule than the other
- This causes a temporary dipole to arise
- This temporary dipole can induce a dipole on neighbouring molecules
- When this happens, the δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other
- Because the electron clouds are moving constantly, the dipoles are only temporary
- all molecules have london forces
48
weakest to strobgets bond type
london forces, induced dipole, hydrogen bonding, covalent bonding
49
relative strength of london forces
For small molecules with the same number of electrons, permanent dipoles are stronger than induced dipoles
so induced dipoles are mostly weakest
however become more the stronger force the longer the chain becomes
50
describe permanant dipole interactions
- al polar molecule shave them
- the molecule will always have a negatively and positively charged end
- Forces between two molecules that have permanent dipoles are called permanent dipole - dipole forces
- The δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other
51
describe hydrogen bonding
- Hydrogen bonding is the strongest form of intermolecular bonding
- For hydrogen bonding to take place a species which has an O, N or F (very electronegative) atom bonded to a hydrogen
- When hydrogen is covalently bonded to an O, N or F, the bond becomes highly polarised
- The H becomes so δ+ charged that it can form a bond with the lone pair of an O, N or F atom in another molecule
52
giant covalent lattics
- diamond
- graphite
- graphine
- silcon(IV) oxide
53
describe diamond
in diamond, each carbon atom bonds with four other carbons, forming a tetrahedron
All the covalent bonds are identical, very strong and there are no intermolecular forces
high melting point because agreat numbe rof C-C bond shave ot be broken in order to melt it
54
describe graphite
Each carbon atom in graphite is bonded to three others forming layers of hexagons, leaving one free electron per carbon atom
These free electrons migrate along the layers and are free to move and carry charge, hence graphite can conduct electricity
The covalent bonds within the layers are very strong, but the layers are attracted to each other by weak intermolecular forces, so the layers can slide over each other making graphite soft and slippery
failry good conductor of electricity because the delocalsied electrons betwent he layers arre free to move in its layer
high mleting point because a great number of C-C bons have ot be broken in order to melt it
55
describe graphene
Graphene consists of a single layer of graphite which is a sheet of carbon atoms covalently bonded forming a continuous hexagonal layer
It is essentially a 2D molecule since it is only one atom thick
It has very unusual properties make it useful in fabricating composite materials and in electronics
56
describe simple mollecular structures
Covalent substances tend to have small molecular structures, such as Cl2, H2O or CO2
These small molecules are known as simple molecules
Hydrogen (H2), chlorine (Cl2), oxygen (O2), nitrogen (N2), hydrogen chloride (HCl), water (H2O), ammonia (NH3) and methane (CH4) are also examples of simple molecules
57
propoerties of giant covalent lattices
Giant covalent lattices have very high melting and boiling points
These compounds have a large number of covalent bonds linking the whole structure
A lot of energy is required to break the lattice
The compounds can be hard or soft
Graphite is soft as the forces between the carbon layers are weak
Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds
Most compounds are insoluble with water
Most compounds do not conduct electricity however some do
Graphite has delocalised electrons between the carbon layers which can move along the layers when a voltage is applied
Diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom are involved in a covalent bond so there are no freely moving electrons available
58
describe ionic lattices
The ions form a lattice structure which is an evenly distributed crystalline structure
Ions in a lattice are arranged in a regular repeating pattern so that positive charges cancel out negative charges
The attraction between the cations and anions is occurring in all directions
Each ion is attracted to all of the oppositely charged ions around it
Therefore the final lattice is overall electrically neutral
59
describe metallic lattices
Metals form giant metallic lattices in which the metal ions are surrounded by a ‘sea’ of delocalised electrons
The metal ions are often packed in hexagonal layers or in a cubic arrangement
This layered structure with the delocalised electrons gives a metal its key properties
60
how does branching affect intemrolecular forces
The larger the surface area of a molecule, the more contact it will have with adjacent molecules
The surface area of a molecule is reduced by branching
The greater its ability to induce a dipole in an adjacent molecule, the greater the London (dispersion) forces and the higher the melting and boiling points
61
how does the number of electrons affect intemrolecular forces
The greater the number of electrons (or the greater the molecular mass) in a molecule, the greater the likelihood of a distortion and thus the greater the frequency and magnitude of the temporary dipoles
The dispersion forces between the molecules are stronger and the enthalpy of vaporisation, melting and boiling points are larger
