Topic 2- bonding and structure

Question 1

What are ions

Answer 1

Charged particles that is formed when an atom loses or gains electrons

Question 2

What are molecular ions

Answer 2

Covalently bonded atoms that lose or gain electrons

Question 3

Which electrons are lost when an atom becomes a positive ion

Answer 3

Electrons in the highest energy levels

Question 4

Do metals usually gain or lose electrons?
Do non metals usually gain or lose electrons?

Answer 4

Metals- lose electrons

Non metals- generally gain electrons

Question 5

Which are the 4 elements that don’t tend to form ions and why?

Answer 5

The elements are beryllium, boron, carbon, and silicon

Requires a lot of energy to transfer outer shell electrons

Question 6

What are the 3 main types of chemical bonds

Answer 6

Ionic
Covalent
Metallic

Question 7

Define ionic bonding

Answer 7

The electrostatic attraction between positive and negative ions

Question 8

Give an example of an ironically bonded substance

Answer 8

NaCl (sodium chloride- salt)

Typically a metal and non metal

Question 9

What determines the strength of an ionic bond

Answer 9

the size of the charge
- the bigger the charge of an ion the stronger the electrostatic attraction between ions
- thus more energy is required to overcome these forces and so they have a high melting and boiling point
- eg KCL is made of K+ and Cl- has a melting point of 770c whereas CaO is made of Ca2+ and O2- has a melting point of 2572c

the size of the ion (ionic radii)
- the smaller the ion the stronger the forces of attraction between ions
- smaller ions can pack together more colsely and more energy is required to overcome the stronger forces. the melting and boiling points are higher.
- NaCl is made from Na+ and Cl- has a melting point of 801C
- KCL is made up of K+ and Cl- has a melting point of 770c (has bigger ionic radius)

smaller ion + big charge —> high charge density

Question 10

Explain the trend in ionic radius down a group

Answer 10

Ionic radius increases going down the group. As the number of electron shells increases and so the ionic radius increases too.

Question 11

what are isoelectric ions
Explain the trend in ionic radius of isoelectric ions

Answer 11

isolectric ions- these are different atoms that have the same number of electrons

as the number of protons increase in isolectric ions the ionic radius decreases. this pulls tin the outer electron shell more. (the number of shells are the same)

Question 12

What are physical properties of ionic compounds

Answer 12

High melting points as there are many, strong electrostatic forces between oppositely charged ions. lots of energy neede to overcome these forces

Non conductor of electricity when solid- ions need to be free to move

Conductor of electricity when in solution or molten

Brittle-layers will slide past eachother.
positives will be next to positives and will repel- and so break apart really easily.

Question 13

do ionic compounds dissolve in water?

Answer 13

Most ionic compounds dissolve in water as water molecules are polar they can attract positive and negative ions and break up the structure.

Question 14

In a solution of CuCrO4 with connected electrodes which electrode will the 2 ions migrate to?

Answer 14

Cu2+- migrate to negative electrode
CrO4 2- migrated to positive electrode

Question 15

Define covalent bonding

Answer 15

Electrostatic attraction between a shared pair of outer electrons and the positve nucleus.

Question 16

what is dative covalent bonding

see notes for example

Answer 16

where one atom donates 2 electrons ro anotherion to form a bond.

represented by an arrow.

Question 17

giant covalent structures include graphite and diamond
describe the structure of graphite?

Answer 17

-layers slide easily as there are weak forces between the layers
-delocalised electrons between the layers allow graphite to conduct electricity as they can carry charge
-layers are far apart in comparision to covalent bond length. this means it has a low density
-lots of strong covalent bonds mean graphite has a very high melting point.
-graphite is insoluble. the covalent bonds are too strong to break.

Question 18

sulphur dioxide has similar structure to diamond describe the structure of diamond

Answer 18

• the tightly packed, rigid arrangement allows heat to conduct well in diamonds
• unlike graphite, diamond can be cut into gemstones
• ver high melting point due to many strong covalent bonds. it is also very hard
• diamond cannot conduct electricicty well as it doesnt have any delocalised electrons
• diamond is insoluble. the covalent bonds are too strong to break

Question 19

graphene is another example of a giant covalent structure.
describe structure and use of graphene

Answer 19

-graphene is 1 layer of graphite. it is one atom thick and made up of hexagonal carbon rings.
-delocalised, free moving electrons make graphene an excellent conductor of electricty as they can carry charge.
- the same delocalised electrons strengthen the covalent bonds, this gives graphite a high strength property
- graphene is only one cell thick, it is transparent and lightweight

uses
aircraft shells
use in super computers and high speed computing
smart phone screens

Question 20

Define metallic bonding

Answer 20

Electrostatic attraction between the positive metal ions and the sea of delocalised electrons

Question 21

metals have a giant mettalic latice structure
explain properties of metallic bonding

see diagram (metallic bonding)

Answer 21

-metals are good thermal conductors as the delocalised electrons can transfer kinetic energy
- metals have high melting point due to the strong electrostatic attractions
- solid metals are insoluble as the metallic bond is too strong to break
- metals are good electrcial conductors as the delocalised electrons are mobile and can carry charge
- the more electrons an atom can donate to the delocalised system the higher the melting point eg magnesium has a higher melting point than sodium as magnesium can donate 2 electrons (group 2) where as sodium only donates 1 (group 1) per atom
- positive metal ions are formed as metals donate electrons to form a ‘sea’ of delocalised electrons.
- there is an electrostatic attraction between postive metal ions and the negative delocalised electrons.

Question 22

Why does giant ionic lattices conduct electricity when liquid but not when solid

Answer 22

In solid state the ions are in fixed position and thus cannot move. When they are in liquid state the ions are mobile and thus can freely carry the charge

Question 23

Giant ionic lattices have high or low melting and boiling point

Answer 23

They have high melting and boiling point becuase a large amount of energy is required to overcome the electrostatic bonds

Question 24

In what type of solvents do ionic lattices dissolve

Answer 24

Polar solvents eg water

Question 25

Explain why both water and carbon dioxide molecules have polar bonds but only
water is a polar molecule.

Answer 25

oxygen is more elctronegative than carbon or hydrogen

which results in a polar bond with oxygen so carbon and hydrogen +

carbn is a symmertrical molecule and so dipole vectors cancel out

lone pairs of electron in oxygen vectors do not cancel out (h20)

Question 26

define periodicity

Answer 26

The study of trends within the periodic table. Often these trends are linked to elements’ electronic configurations.

Question 27

what is Electrongegativity

Answer 27

electronegtivity is the ability for an atom to attract electrons towards itself in a covalent bond.

1. the further right you go in the periodic table the more electronegatib=ve the elemnt is.

fluorine is the most electronegative

Question 28

when can covalent bonds become polar?

Answer 28

covalent bonds can become polar if the atoms attached to it have a difference in electronegativity

Question 29

the bigger the difference in electronegativity…

Answer 29

…the more polar the bond will be

Question 30

when are bonds not polar

Answer 30

atoms bonded with same electronegativity or similar are non-polar. the shared electrons sit in the middle.

Question 31

polar molecules is lead to

Answer 31

uneven distrubution to charge

Question 32

what happens if bonds are arranged symmetrically

Answer 32

then the molecule is non-polar

Question 33

what happens if bonds are arranged asymmetrically

Answer 33

then you have an overall polar molecule

Question 34

name 3 types of intermolecular forces

Answer 34

London
dipole-dipole
hydrogen bonding

Question 35

what are london forces

Answer 35

London forces are the weakest type of intermolecular forces
any electron can form a dipole when they move near to another atom or molecule
london forces are random
this occurs as electrons in a molecule or atom can move from one end to another - hence creating a temporary dipole
this temporary dipole only ecists when 2 molecules or atoms near by.
the slightly + on one atom or molecule will be attracted to the slightly negative on another and a force of attraction is created.

london forces can hold some molecule in crystal structures, iodine is an example

weak london forces - holds the I2 molecule together
strong covalentbonds holds the 2 iodine atoms together

force happens between molecules
bond happens between atoms
bond is much stronger than atom

the bigger the molecule or atom, the morelondon forces as you have a larger electron cloud
we must have enough enrgy to overcome these forces

Question 36

london forces with hydrocarbons

Answer 36

longer straight chain hydrocarbons have more london forces and so more enrgy is needed to overcome these forces. this means the boiling point increases.

hydrocarbons with branches means they cant back together as close. this weakens the london forces between the chains and lowers the boiling points

Question 37

what are dipole dipole forces

Answer 37

permant dipole eg HCL which are stronger than london forces
they also include having london forces
there are weak electrostatic forces that exist between molecules with polarity

if its non-polar only expereience london forces

Question 38

what is hydrogen bonding

see example of ice

Answer 38

hydrogen bonding strongest type of bonding
hydrogen bonds with the lone pair of nitrogen, oxygen or fluorine (the most electronegative elements)

hydrogen bonding also have london forces and permant dipole dipole

Question 39

Explain why hydrogen bonding causes ice to be less dense than liquid water (2 marks)

Answer 39

Ice has more space between molecules
Due to ring structure in ice

Question 40

solubiliy

Answer 40

polar substances can dissolve in polar solvents

the delta positive in H is attracted to the negative ions and the delta negative oxygen is attracted to the positive ion and the structure starts to break down, the waters align them =selves surrounding the ions this is known as hydration

for this to happen the new bonds must be the same strength or greater than those broken. if not then the substance is very unlikely to dissolve.

eg AL2O2 doesnt dissolve as the ionic bonding is too strong even though its ionic

alcohols can dissolve too- they dissolve in polar solvents as they can hydrogen bond with water molecules but the hydrocarbon part is non-polar so cant dissolve in water. the bigger the hydrocarbon part the less soluble the alcohol is.

some polar molecules dont dissolve in water, haloalkanes dont dissolve as their dipoles arent very strong

Question 41

solubility with non-polar solvents

Answer 41

non-polar solvents- these molecules thats dont have a polarity. an example would be alkanes lke butane. they only have london forces

non-polar molecules tend not dissolve in water as water is polar and water forms stronger hydrogen bonds; water interacts with its own molecules as appose to the non-polar one.