(Green) Bonding II Covalent and Metallic Flashcards
(Green) Bonding II Covalent and Metallic
info card (read and recite)
COVALENT BONDS
In the Electron structure pack, you learnt that it becomes more and more difficult to remove electrons from an atom to form positive ions, a process that requires increasing amounts of energy. Loss of one or two electrons is not a problem, loss of three represents the limit of what can be realistically achieved, but losing four electrons in a chemical reaction simply does not happen as there is no compensating, energy releasing process. Atoms also require increasing amounts of energy in order to gain additional electrons and form negative ions.
In order to overcome these problems and to attain a stable, noble gas electron configuration, atoms can share a pair of electrons, so forming a covalent bond
what is the definition of a covalent bond
A covalent bond is a shared pair of electrons one from each donor atom
what is the definition of a Covalent bonding
Covalent bonding is the electrostatic attraction between the nuclei and the bonding pair of
what is the definition of a Bond length
The Bond length is the distance between the nuclei of the two atoms that are covalently bonded
how does bond length effect strength
the shorter the bond length the stronger the bond
what is a double covalent bond
A double covalent bond is one in which two pairs of electrons are shared between two atoms.
what is a triple covalent bond
A triple covalent bond is one in which three pairs of electrons are shared between two atoms.
what is a dative covalent bond
A dative covalent bond is a shared pair of electrons both from one donor atom
is a dative covalent bond shorter and or stronger than a covalent bond
neither
it will have the same length
and also strength
AlCl3
How many pairs of electrons does the Al have in its
outer shell?
How many electron pairs represent a full outer shell?
what is it described as
3
4
electron deficient
draw out the diagram for AlCl3
http://kwokthechemteacher.blogspot.com/2009/04/chemical-bonding-dative-covalent-and.html
which atoms have to obay the octet rule
and what is the octet rule
carbon, nitrogen, oxygen, and the halogens
The octet rule is a chemical rule of thumb that reflects the theory that main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas.
info card (read and recite)
Most covalent structure are molecules which are held together with strong bonds within the molecule but weak forces between the molecules examples include CO2, CH4, H2O, at room temperature most of these are liquids or gases
However compounds can also form molecular solids, especially if the temperature is low enough . The two examples are ice and iodine
info card (read and recite)
Iodine is a diatomic molecule held together by weak
Intermolecular forces. In the solid form the molecules
are held in a regular lattice. What happens to the structure
as it is heated?
As the temperature increases the molecules gain
kinetic energy and vibrate more, enough to overcome the weak intermolecular forces between the molecules (the strong covalent bonds remain unchanged).
info card (read and recite)
Another example of a molecular solid is ice, again held together by weak intermolecular forces. (you will study these in a later topic). Other molecular solids are Sulphur (S8), phosphorus (P4) C60 etc. Note that they have a specified number of atoms in the molecule. This differentiates them from Giant covalent lattices.
When melting a molecular lattice it is the weak intermolecular forces that are overcome NOT the strong covalent bonds within the molecule.
what is the definition of a allotrope
Allotropes – two or more forms of the same element in which the atoms or molecules are arranged in different ways
describe the structure of diamond
DIAMOND - a giant, covalently bonded structure.
In a diamond, each carbon atom is bonded to four other carbon atoms oriented around it in a tetrahedral manner to form a giant molecule. The bonds are covalent and because the bonding electrons are localised close to the nucleus, they are exceptionally strong. Diamond is the hardest natural substance known.
compare the diffrences between diamond and iodine :
appearance
diamond
Extremely hard, transparent, colourless, crystals with a high refractive index
iodine :
grey solid
compare the diffrences between diamond and iodine :
bonding
diamond
Each C atom makes 4 covalent single bonds.
iodine :
Covalent bond between iodine atoms and weak attractions between molecules
compare the diffrences between diamond and iodine :
Structure
(describe)
diamond
Each C atom is tetrahedrally bonded to four others in a giant 3D lattice.
iodine
Simple molecule I2 packed into a regular lattice
compare the diffrences between diamond and iodine :
M.Pt. / B.Pt.
diamond
Very high = 3550oC
iodine :
low = 114oC
Why is there such a difference in the M.Pt. of iodine and diamond?
Diamond : giant covalent structure need to break many covalent bonds in 3D
Covalent bonds need to be broken
I2 molecule: only has weak attractions (London forces) between molecules
Little energy is needed to overcome these weak forces of attraction
Silicon is also in group 4 and the element shows the same structure as diamond. Why does silicon have a very high melting point?
Si has a giant covalent structure and a lot of heat energy is needed to break the 4 covalent bonds between each Si atom
It has a lower M.Pt. than diamond as the atoms have a larger radius and so the covalent bonds will be longer and weaker
If silicon forms four covalent bonds and oxygen forms two bonds how can a giant structure be formed?
Each Si atom makes single covalent bonds to four O atoms
Each O atom makes a single covalent bond to two Si atoms
what are the physical propertys of SiO2
very hard
high melting point
does not conduct
insoluable
info card (read and recite)
GRAPHITE Chemguide description of graphite (and other giant structures)
Graphite is one of three crystalline forms (allotropes) of carbon. It is an unusual non-metal as it is able to conduct both heat and electricity. The reason lies in its structure. It is composed of carbon atoms in rings of six, bonded together and formed into sheets. (The sheets are then stacked in a layer structure).
A carbon atom has four electrons in its outer shell.
Looking at the diagram, to how many other carbon atoms is each carbon atom bonded? 3 others
Each of these bonds is a strong covalent bond.
This means that each carbon atom has a valence. electron free which is not localised. These electrons are mobile and move freely across the plane of the sheet. They are said to be delocalised as they form a delocalised p molecular orbital.
Between each layer are weak intermolecular forces.
(called London forces)
Explain why graphite can conduct heat and electricity
· Delocalised electrons are free to move ALONG a layer (not between layers)
· When a p.d.is applied they move to the +ve terminal conducting electicity.
· In a hot region they move faster colliding with other electrons so transferring KE
In which direction can graphite conduct and why?
· Graphite only conducts ALONG sheets, electrons are only moving within sheets,
· adjacent sheets/ individual planes are too far apart.
· Explain why graphite is a good lubricant or useful in lead pencils?
Weak London forces allow layers to slip leaving a thin layer of graphite on the paper.
info card (read and recite)
In a diamond, each carbon atom is bonded to four other carbon atoms oriented around it in a tetrahedral manner to form a giant structure. The bonds are covalent and because the bonding electrons are localised close to the nucleus, they are exceptionally strong. Diamond is the hardest natural substance known.
diamond vs graphite :
appearance
Diamond :
Shiny transparent crystal
graphite :
Dull grey / black solid
diamond vs graphite :
Ability to conduct electricity
and
Ability to conduct heat
diamond :
electrisity : No – all electrons are localised within covalent bonds.
heat : Yes but no delocalised e- , it is a quantum effect – you do not need to know about this
graphite :
heat conductivity - C atoms covalently bonded to three others forming sheets of hexagons, fourth electron is delocalised within the sheet.
Electrical conductivity – good along layers, delocalised e- free to move with applied p.d., current = flow of electrons. Heat energy transferred by e- gaining KE and transferring it along layers.
diamond vs graphite :
density
diamond :
3.53 g cm-1 High, tetrahedrally bonded C atoms held close together so a large number of atoms per unit volume
graphite :
2.25 g cm-1 Low, due to large spaces between hexagonally bonded sheets
diamond vs graphite :
hardness
diamond :
Hardest natural substance known. Tetrahedrally bonded C atoms form a strong rigid structure
graphite :
Soft, layers rub off. Sheets weakly held to adjacent ones by weak dispersion forces
diamond vs graphite :
Solubility
diamond :
Insoluble in polar and non-polar solvents as it is a giant structure
graphite :
Insoluble in polar and non-polar solvents as it is a giant structure
diamond vs graphite :
uses
diamond :
On surfaces of cutting/drilling tips
graphite :
Good lubricant.
info card (read and recite)
Because metals have large atoms the electrons are far from the nucleus. Metals have low ionisation energies. This means that metal atoms have a high tendency to lose electrons and form positive ions.
The properties of metals depend on the way the atoms are packed (structure) and the bonding of the atoms making up the metal structure.
Most metals have fairly high melting points suggesting that the attractive forces between cations and delocalised electrons are strong. Metallic atoms readily lose their outer shell electrons which move randomly throughout the crystal structure, are shared by all the neighbouring positive ions and are said to be delocalised.
what is the definotion of metallic bonding
DEFINITION: Metallic bonding is the electrostatic attractions between positive metal cations and the delocalised electrons.
what dictates the strngth and melting point of metalic substances
SIZE AND CHARGE
the greater the size and charge the greater the strength and boiling and melting point
what is the definition of delocalised electrons
In chemistry, delocalized electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or a covalent bond.
State and explain which of the following will have the strongest metallic bonding.
Sodium or potassium?
Na – smaller ionic radius, same no of delocalised electrons therefore greater electrostatic attraction and more energy required
State and explain which of the following will have the strongest metallic bonding.
Sodium or magnesium?
Magnesium ion/ Mg2+ is smaller (than sodium ions) (1)
The magnesium ion contributes two electrons (to the sea of electrons) compared to sodium’s one. (1)
More energy/heat required to overcome (attractive) forces/bonds (between cations and “sea” of electrons) in magnesium (compared to sodium)
State and explain which of the following will have the strongest metallic bonding.
Magnesium or aluminium
Al – smaller ion, 3e- delocalised.
Explain why the delocalised electrons in metals make them:-
Good electrical conductors:
delocalised electrons are free to move in any direction
when a potential difference is applied the electrons move to the positive terminal
a flow of electrons is an electric current.
Explain why the delocalised electrons in metals make them:-
Good conductors of heat:
delocalised electrons are free to move in any direction
in hot areas electrons have more K.E., move faster
electrons collide more, passing on K E to other electrons.
Explain why the delocalised electrons in metals make them:-
Explain why the electrical conductivity increases Na —> Mg ——>Al
More electrons are delocalised Na(1) à Mg(2) à Al(3) are free to move in any direction.
when a p.d. is applied the electrons move to the positive terminal
larger flow of electrons results in a larger current and so greater conductivity.
Explain why the delocalised electrons in metals make them:-
Explain why the thermal conductivity increases Na —> Mg ——>Al
More electrons are delocalised Na(1) à Mg(2) à Al(3)
in hot areas electrons have more K.E.
more electrons collide more passing on K E to other electrons more efficiently.
The stronger the metallic bond the higher the melting point. Most metals are solids.
Which one is liquid at room temperature?
mercury
The stronger the metallic bond the higher the melting point. Most metals are solids.
Which other two have the lowest melting points?
Caesium, francium
