(Green) Bonding II Covalent and Metallic

Question 1

info card (read and recite)

Answer 1

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

Question 2

what is the definition of a covalent bond

Answer 2

A covalent bond is a shared pair of electrons one from each donor atom

Question 3

what is the definition of a Covalent bonding

Answer 3

Covalent bonding is the electrostatic attraction between the nuclei and the bonding pair of

Question 4

what is the definition of a Bond length

Answer 4

The Bond length is the distance between the nuclei of the two atoms that are covalently bonded

Question 5

how does bond length effect strength

Answer 5

the shorter the bond length the stronger the bond

Question 6

what is a double covalent bond

Answer 6

A double covalent bond is one in which two pairs of electrons are shared between two atoms.

Question 7

what is a triple covalent bond

Answer 7

A triple covalent bond is one in which three pairs of electrons are shared between two atoms.

Question 8

what is a dative covalent bond

Answer 8

A dative covalent bond is a shared pair of electrons both from one donor atom

Question 9

is a dative covalent bond shorter and or stronger than a covalent bond

Answer 9

neither

it will have the same length

and also strength

Question 10

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

Answer 10

3

4

electron deficient

Question 11

draw out the diagram for AlCl3

Answer 11

http://kwokthechemteacher.blogspot.com/2009/04/chemical-bonding-dative-covalent-and.html

Question 12

which atoms have to obay the octet rule

and what is the octet rule

Answer 12

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.

Question 13

info card (read and recite)

Answer 13

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

Question 14

info card (read and recite)

Answer 14

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).

Question 15

info card (read and recite)

Answer 15

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.

Question 16

what is the definition of a allotrope

Answer 16

Allotropes – two or more forms of the same element in which the atoms or molecules are arranged in different ways

Question 17

describe the structure of diamond

Answer 17

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.

Question 18

compare the diffrences between diamond and iodine :

appearance

Answer 18

diamond
Extremely hard, transparent, colourless, crystals with a high refractive index

iodine :

grey solid

Question 19

compare the diffrences between diamond and iodine :

bonding

Answer 19

diamond
Each C atom makes 4 covalent single bonds.

iodine :
Covalent bond between iodine atoms and weak attractions between molecules

Question 20

compare the diffrences between diamond and iodine :

Structure

(describe)

Answer 20

diamond
Each C atom is tetrahedrally bonded to four others in a giant 3D lattice.

iodine
Simple molecule I2 packed into a regular lattice

Question 21

compare the diffrences between diamond and iodine :

M.Pt. / B.Pt.

Answer 21

diamond
Very high = 3550oC

iodine :

low = 114oC

Question 22

Why is there such a difference in the M.Pt. of iodine and diamond?

Answer 22

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

Question 23

Silicon is also in group 4 and the element shows the same structure as diamond. Why does silicon have a very high melting point?

Answer 23

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

Question 24

If silicon forms four covalent bonds and oxygen forms two bonds how can a giant structure be formed?

Answer 24

Each Si atom makes single covalent bonds to four O atoms

Each O atom makes a single covalent bond to two Si atoms

Question 25

what are the physical propertys of SiO2

Answer 25

very hard
high melting point
does not conduct
insoluable

Question 26

info card (read and recite)

Answer 26

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)

Question 27

Explain why graphite can conduct heat and electricity

Answer 27

· 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

Question 28

In which direction can graphite conduct and why?

Answer 28

· Graphite only conducts ALONG sheets, electrons are only moving within sheets,

· adjacent sheets/ individual planes are too far apart.

Question 29

· Explain why graphite is a good lubricant or useful in lead pencils?

Answer 29

Weak London forces allow layers to slip leaving a thin layer of graphite on the paper.

Question 30

info card (read and recite)

Answer 30

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.

Question 31

diamond vs graphite :

appearance

Answer 31

Diamond :
Shiny transparent crystal

graphite :
Dull grey / black solid

Question 32

diamond vs graphite :

Ability to conduct electricity
and
Ability to conduct heat

Answer 32

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.

Question 33

diamond vs graphite :

density

Answer 33

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

Question 34

diamond vs graphite :

hardness

Answer 34

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

Question 35

diamond vs graphite :

Solubility

Answer 35

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

Question 36

diamond vs graphite :

uses

Answer 36

diamond :
On surfaces of cutting/drilling tips

graphite :
Good lubricant.

Question 37

info card (read and recite)

Answer 37

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.

Question 38

what is the definotion of metallic bonding

Answer 38

DEFINITION: Metallic bonding is the electrostatic attractions between positive metal cations and the delocalised electrons.

Question 39

what dictates the strngth and melting point of metalic substances

Answer 39

SIZE AND CHARGE

the greater the size and charge the greater the strength and boiling and melting point

Question 40

what is the definition of delocalised electrons

Answer 40

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.

Question 41

State and explain which of the following will have the strongest metallic bonding.

Sodium or potassium?

Answer 41

Na – smaller ionic radius, same no of delocalised electrons therefore greater electrostatic attraction and more energy required

Question 42

State and explain which of the following will have the strongest metallic bonding.

Sodium or magnesium?

Answer 42

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)

Question 43

State and explain which of the following will have the strongest metallic bonding.

Magnesium or aluminium

Answer 43

Al – smaller ion, 3e- delocalised.

Question 44

Explain why the delocalised electrons in metals make them:-

Good electrical conductors:

Answer 44

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.

Question 45

Explain why the delocalised electrons in metals make them:-

Good conductors of heat:

Answer 45

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.

Question 46

Explain why the delocalised electrons in metals make them:-

Explain why the electrical conductivity increases Na —> Mg ——>Al

Answer 46

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.

Question 47

Explain why the delocalised electrons in metals make them:-

Explain why the thermal conductivity increases Na —> Mg ——>Al

Answer 47

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.

Question 48

The stronger the metallic bond the higher the melting point. Most metals are solids.

Which one is liquid at room temperature?

Answer 48

mercury

Question 49

The stronger the metallic bond the higher the melting point. Most metals are solids.

Which other two have the lowest melting points?

Answer 49

Caesium, francium