chem revise flashcards

Question 1

isotope

Answer 1

Isotopes are atoms with the same number of protons, but different numbers of neutrons.

Isotopes have similar chemical properties because they have the same electronic structure. They may have slightly varying physical properties because they have different masses.

Question 2

what is mass spectrometer is used from

Answer 2

The mass spectrometer can be used to determine all the isotopes present in a sample of an element and to therefore identify elements.

Question 3

what is electron impact

Answer 3

*A vaporised sample is injected at low pressure
*An electron gun fires high energy electrons at the sample
*This knocks out an outer electron
*Forming positive ions with different charges e.g. Ti (g) Ti+ (g)+ e–

Question 4

Electro spray Ionisation

Answer 4

• The sample is dissolved in a volatile, polar solvent
• injected through a fine needle giving a fine mist or aerosol
• the tip of needle has high voltage
• at the tip of the needle the sample molecule, M, gains a proton, H+, from the
  solvent forming MH+
• M(g) + H+  MH+(g)
• The solvent evaporates away while the MH+ ions move towards a negative plate

Question 5

Acceleration

Answer 5

Given that all the particles have the same kinetic energy, the velocity of each particle depends on its mass. Lighter particles have a faster velocity, and heavier particles have a slower velocity

Question 6

first ionisstion energy

Answer 6

The first ionisation energy is the enthalpy change when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge

Question 7

second ionisaion enegry

Answer 7

The second ionisation energy is the enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge

Question 8

Factors that affect ionisation energy

Answer 8

1.The attraction of the nucleus
2. The distance of the electrons from the nucleus
(The bigger the atom the further the outer electrons are from the nucleus and the
weaker the attraction to the nucleus) 3. Shielding of the attraction of the nucleus

Question 9

Why are successive ionisation energies always larger?

Answer 9

The second ionisation energy of an element is always bigger than the first ionisation energy. When the first electron is removed a positive ion is formed.
The ion increases the attraction on the remaining electrons and so the energy required to remove the next electron is larger.

Question 10

Why has helium the largest first ionisation energy?

Answer 10

Its first electron is in the first shell closest to the nucleus and has no shielding effects from inner shells. He has a bigger first ionisation energy than H as it has one more proton

Question 11

Why do first ionisation energies decrease down a group?

Answer 11

As one goes down a group, the outer electrons are found in shells further from the nucleus and are more shielded so the attraction of the nucleus becomes smaller

Question 12

Why is there a general increase in first ionisation energy across a period?

Answer 12

As one goes across a period the electrons are being added to the same shell which has the same distance from the nucleus and same shielding effect. The number of protons increases, however, making the effective attraction of the nucleus greater.

Question 13

Why has Na a much lower first ionisation energy than neon?

Answer 13

This is because Na will have its outer electron in a 3s shell further from the nucleus and is more shielded. So Na’s outer electron is easier to remove and has a lower ionisation energy.

Question 14

Why is there a small drop from Mg to Al?

Answer 14

Al is starting to fill a 3p sub shell, whereas Mg has its outer electrons in the 3s sub shell. The electrons in the 3p subshell are slightly easier to remove because the 3p electrons are higher in energy and are also slightly shielded by the 3s electrons

Question 15

Why is there a small drop from P to S?

Answer 15

With sulfur there are 4 electrons in the 3p sub shell and the 4th is starting to doubly fill the first 3p orbital.
When the second electron is added to a 3p orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove.

Question 16

poisitve ion radiii

Answer 16

Positive ions are smaller compared to their atoms because it has one less shell of electrons and the ratio of protons to electrons has increased so there is greater net force on remaining electrons holding them more closely.

Question 17

dative covalent bond

Answer 17

A dative covalent bond forms when the O
shared pair of electrons in the covalent bond come from only one of the bonding atoms. A dative covalent bond is also called co-ordinate bonding

Question 18

factos affectsing metallic bonding

Answer 18

1. Number of protons/ Strength of nuclear attraction.
   The more protons the stronger the bond
2. Number of delocalised electrons per atom (the outer shell electrons are delocalised)
   The more delocalised electrons the stronger the bond
3. Size of ion.

Question 19

ionic properties

Answer 19

high mp and bp
soluble in water
cant conduct when solid

Question 20

simple comleuclar properties

Answer 20

low mp and bp
poor solubility
poor conductivity when solid
poor conductivity when molten

Question 21

macromolecular properties

Answer 21

high mo and bp
soluble
poor conductivty

Question 22

metallic properties

Answer 22

high mp and bp
insolble in water
good conductvity

Question 23

linear

Answer 23

2
0
180

Question 24

trigonal planar

Answer 24

3
0
120

Question 25

tetrahderal

Answer 25

4
0
109.5

Question 26

trigonal planar

Answer 26

3
1
107

Question 27

bent

Answer 27

2
2
104.5

Question 28

trigonal bipyramidal

Answer 28

5
0
120 and 90

Question 29

octahedral

Answer 29

6 bp
0lp
90

Question 30

factors affecting electroneg

Answer 30

Electronegativity increases across a period as the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more.
It decreases down a group because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases

Question 31

Formation of a permanent dipole

Answer 31

A polar covalent bond forms when the elements in the bond have different
electronegativities

When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends.

Question 32

Symmetric molecules

Answer 32

A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar.

Question 33

anomalously high boiling points of H2O, NH3 and HF are caused by

Answer 33

the hydrogen bonding between the molecules
The general increase in boiling point from H2S to H2Te is caused by increasing Van der Waals forces between molecules due to an increasing number of electrons.
Alcohols, carboxylic acids, proteins, amides all can form hydrogen bonds

Question 34

Main factor affecting size of Van der Waals

Answer 34

The more electrons there are in the molecule the higher the chance that temporary dipoles will form. This makes the Van der Waals stronger between the molecules and so boiling points will be greater
The increasing boiling points of the halogens down the group 7 series can be explained by the increasing number of electrons in the bigger molecules causing an increase in the size of the Van der Waalsbetweenthemolecules

Question 35

Permanent dipole-dipole forces

Answer 35

Permanent dipole-dipole forces occurs between polar molecules
*It is stronger than Van der Waals and so the compounds have higher boiling points
*Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds) *Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.

Question 36

what is emthalpy change

Answer 36

is the amount of heat energy taken in or given out during any change in a system provided the pressure is constant.

Question 37

stsandard enthalpy of formation

Answer 37

The standard enthalpy change of formation of a compound is the enthalpy change when 1 mole of the compound is formed from its elements under standard conditions (298K and 100kpa), all reactants and products being in their standard states.

Question 38

standard enthalpy of combustion

Answer 38

The standard enthalpy of combustion of a substance is defined as the enthalpy change that occurs when one mole of a substance is combusted completely in oxygen under standard conditions. (298K and 100kPa), all reactants and products being in their standard states.

Question 39

hess law

Answer 39

Hess’s law states that total enthalpy change for a reaction is independent of the route by which the chemical change takes place

Question 40

H reaction =

Answer 40

Σ fH products - Σ fH reactantsi

Question 41

Mean Bond energies

Answer 41

Definition: The mean bond energy is the enthalpy needed to break the covalent bond into gaseous atoms, averaged over different molecules.

Question 42

what is coliision theory

Answer 42

Reactions can only occur when collisions take place between particles having sufficient energy. The energy is usually needed to break the relevant bonds in one or either of the reactant molecules.
This minimum energy is called the activation energy.

Question 43

what is rate of reaction

Answer 43

as the change in concentration of a substance in unit time. The usual unit is mol dm-3s-1

Question 44

effects of inc conc inc pressure

Answer 44

At higher concentrations(and pressures) there are more particles per unit volume and so the particles collide with a greater frequency and there will be a higher frequency of effective collisions.

Question 45

effect of inc temp

Answer 45

At higher temperatures the energy of the particles increases. The pariticles collide more frequently and more often with energy greater than the activation energy. More collisions result in a reaction.
As the temperature increases, the graph shows that a significantly bigger proportion of particles have energy greater than the activation energy, so the frequency of successful collisions increases.

Question 46

effect of catalyst

Answer 46

They do this by providing an alternative route or mechanism with a lower activation energy.

Question 47

electron acceptor

Answer 47

element thats reduced in the equation

Question 48

electron donor

Answer 48

reducing agent

Question 49

enthalpy of fromation e

Answer 49

The standard enthalpy change of formation of a compound is the energy transferred when 1 mole of the compound is formed from its elements under standard conditions (298 K and 100 kpa), all reactants and products being in their standard states

Question 50

enthalpy pf atomisation

Answer 50

The enthalpy of atomisation of an element is the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state

Question 51

bond enthalpy of dissociaion

Answer 51

The bond dissociation enthalpy is the standard molar enthalpy change when one mole of a covalent bond is broken into two gaseous atoms (or free radicals)

Question 52

first ionitsation enthalpy

Answer 52

The first ionisation enthalpy is the enthalpy change required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a +1 charge

Question 53

first elevton afffintity

Answer 53

The first electron affinity is the enthalpy change that occurs when 1 mole of gaseous atoms gain 1 mole of electrons to form 1 mole of gaseous ions with a –1 charge

Question 54

enthalpy of lattice formation

Answer 54

The enthalpy of lattice formation is the standard enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form

Question 55

second iomnisatioj enthalpt

Answer 55

The second ionisation enthalpy is the enthalpy change to remove 1 mole of electrons from one mole of gaseous 1+ ions to produces one mole of gaseous 2+ ions.

Question 56

second electron affintiy

Answer 56

The second electron affinity is the enthalpy change when one mole of gaseous 1- ions gains one electron per ion to produce gaseous 2- ion

Question 57

enthalpy of hydration

Answer 57

Enthalpy change when one mole of gaseous ions become aqueous ions

Question 58

enthalpy of solution

Answer 58

The enthalpy of solution is the standard enthalpy change when one mole of an ionic solid dissolves in a large enough amount of water to ensure that the dissolved ions are well separated and do not interact with one another.

Question 59

trends in lattice enthalpy

Answer 59

The sizes of the ions
The charges on the ion

Question 60

perfect ionic model

Answer 60

Theoretical lattice enthalpies assumes a perfect ionic model where the ions are 100% ionic and spherical and the attractions are purely electrostatic.

Question 61

sspontaneous process

Answer 61

(e.g. diffusion) will proceed on its own without any external influence.

Question 62

Entropy, S ̊

Answer 62

Substances with more ways of arranging their atoms and energy (more disordered) have a higher entropy

Question 63

An increase in disorder and entropy will lead to

Answer 63

o a positive entropy change ∆S ̊ = +ve

Question 64

increase in entropy and temp on feasbility

Answer 64

If the reaction involves an increase in entropy (∆S is +ve) then increasing temperature will make it more likely that ∆G is negative and more likely that the reaction

Question 65

if the reaction has a ∆S close to zero then

Answer 65

temperature will not have a large effect on the feasibility of the reaction as - T∆S will be small and ∆G will not change much

Question 66

electrochemicl cell

Answer 66

*A cell has two half–cells.
*The two half cells have to be connected with a salt bridge.
*Simple half cells will consist of a metal (acts an electrode) and a solution of a compound containing that metal (eg Cu and CuSO4).
*These two half cells will produce a small voltage if connected into a circuit. (i.e. become a Battery or cell).

Question 67

why does a voltage form in an electrochemical cell

Answer 67

When connected together the zinc half-cell has more of a tendency to oxidise to the Zn2+ ion and release electrons than the copper half-cell. (ZnZn2+ + 2e-)
More electrons will therefore build up on the zinc electrode than the copper electrode.
A potential difference is created between the two electrodes.
The zinc strip is the negative terminal and the copper strip is the positive terminal

Question 68

why use a high resistance voltmeter

Answer 68

he voltmeter needs to be of very high resistance to stop the current from flowing in the circuit. In this state it is possible to measure the maximum possible potential difference (E).
The reactions will not be occurring because the very high resistance voltmeter stops the current from flowing.

Question 69

what a salt bridge is made from and used for

Answer 69

The salt bridge is used to connect up the circuit. The free moving ions conduct the charge.
A salt bridge is usually made from a piece of filter paper (or material) soaked in a salt solution, usually potassium nitrate.
The salt should be unreactive with the electrodes and electrode solutions. E.g. potassium chloride would not be suitable for copper systems because chloride ions can form complexes with copper ions.
A wire is not used because the metal wire would set up its own electrode system with the solutions.

Question 70

What happens if current is allowed to flow?

Answer 70

If the voltmeter is removed and replaced with a bulb or if the circuit is short circuited, a current flows. The reactions will then occur separately at each electrode. The voltage will fall to zero as the reactants are used up.

Question 71

which electrode wil undergo reduciton

Answer 71

The most positive electrode will always undergo reduction.

Question 72

which electrode will undergo oxidation

Answer 72

The most negative electrode will always undergo oxidation.

Question 73

solid vertical line in cell diagrams

Answer 73

represents the boundary between phases e.g. solid (electrode) and solution (electrolyte)

Question 74

double line in cell diagram

Answer 74

*The double line represents the salt bridge between the two half cells
*the voltage produced is indicated
*the more positive half cell is written on the right if possible

Question 75

systems that dont include metals

Answer 75

If a system does not include a metal that can act as an electrode, then a platinum electrode must be used and included in the cell diagram. It provides a conducting surface for electron tra

Question 76

how potential of electrodes are measured

Answer 76

The potential of all electrodes are measured by comparing their potential to that of the standard hydrogen electrode.
The standard hydrogen electrode (SHE) is assigned the potential of 0 volts.

Question 77

Components of a standard hydrogen electrode.

Answer 77

1. Hydrogen gas at pressure of 100kPa
2. Solution containing the hydrogen ion at 1.0 mol dm-3 (solution is usually 1 mol dm-3 HCl)
3. Temperature at 298K 4. Platinum electrode

Question 78

Effect of concentration on Cell voltage Ecell

Answer 78

Looking at cell reactions is a straight forward application of le Chatelier. So increasing concentration of ‘reactants’ would increase Ecell and decreasing them would cause Ecell to decrease

Question 79

effect of temp on e cell

Answer 79

Most cells are exothermic in the spontaneous direction so applying Le Chatelier to a temperature rise to these would result in a decrease in Ecell because the equilibrium reactions would shift backwards.

Question 80

what is a fuel cell

Answer 80

uses the enrgy from the reaction of a fuel with oxygen to create a voltage

Question 81

conditions of fule cells

Answer 81

high temps and constant voltage continuously fed with fresh O2 and H2 so contant conc
Higher temperatures are therefore used to increase rate but the reaction is exothermic

Question 82

advantage of fuel cells

Answer 82

over conventional petrol or diesel-powered vehicles
(i) less pollution and less CO2. (Pure hydrogen emits only water whilst hydrogen-rich fuels produce only small amounts of air pollutants and CO2).
(ii) greater efficiency

Question 83

kimitations of hyfrogen fuel cells

Answer 83

(i) expensive
(ii) storing and transporting hydrogen, in terms of safety, feasibility of a pressurised liquid and a limited life cycle of a solid ‘adsorber’ or ‘absorber’
(iii) limited lifetime (requiring regular replacement and disposal) and high production costs,
(iv) use of toxic chemicals in their production

Question 84

ethanol fuel cells

Answer 84

Compared to hydrogen fuel cells they have certain advantages including. Ethanol can be made from renewable sources in a carbon neutral way.
Raw materials to produce ethanol by fermentation are abundant.
Ethanol is less explosive and easier to store than hydrogen. New petrol stations would not be required as ethanol is a liquid fuel.

Question 85

atomic radius trend

Answer 85

Atomic radii decrease from left to right across a period, because the increased number of protons create more positive charge attraction for electrons which are in the same shell with similar shielding.

Question 86

1st ionisatio energy trend

Answer 86

There is a general trend across to increase. This is due to increasing number of protons as the electrons are being added to the same shell.
There is a small drop between Mg + Al. Mg has its outer electrons in the 3s sub shell, whereas Al is starting to fill the 3p subshell. Al’s electron is slightly easier to remove because the 3p electrons are higher in energy.
There is a small drop between phosphorous and sulfur. Sulfur’s outer electron is being paired up with an another electron in the same 3p orbital.
When the second electron is added to an orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove

Question 87

melting and boiling point trend

Answer 87

For Na, Mg, Al- Metallic bonding : strong bonding – gets stronger the more electrons there are in the outer shell that are released to the sea of electrons. A smaller sized ion with a greater positive charge also makes the bonding stronger. Higher energy is needed to break bonds.
Si is Macromolecular: many strong covalent bonds between atoms, high energy needed to break covalent bonds– very high mp +bp

Question 88

period 2 trend

Answer 88

Similar trend in period 2
Li,Be metallic bonding (high mp) B,C macromolecular (very high mp) N2,O2 molecular (gases! Low mp as small v der w)
Ne monoatomic gas (very low mp)

Question 89

atomic radius

Answer 89

Atomic radius increases down the group.
As one goes down the group, the atoms have more shells of electrons making the atom bigger

Question 90

melting points down the group

Answer 90

Melting points decrease down the group. The metallic bonding weakens as the atomic size increases. The distance between the positive ions and delocalized electrons increases. Therefore the electrostatic attractive forces between the positive ions and the delocalized electrons weaken

Question 91

first ionisation energy

Answer 91

The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells.
In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons.

Question 92

group 2 w o2

Answer 92

The group 2 metals will burn in oxygen. Mg burns with a bright white flame. The MgO appears as a white powder.
2 Mg + O2  2 MgO
MgO is a white solid with a high melting point due to its ionic bonding.

Question 93

reactivity of group 2 metals

Answer 93

increases down the group.

Question 94

group 2 and steam

Answer 94

Magnesium reacts in steam to produce magnesium oxide and hydrogen. The Mg would burn with a bright white flame. The MgO appears as a white powder.

Question 95

``group 2 and warm water

Answer 95

giving a different magnesium hydroxide product.
Mg + 2 H2OMg(OH)2 + H2
This is a much slower reaction than the reaction with steam and there is no flame.

Question 96

generwl observatios withbgroup 2 and water

Answer 96

fizzing, (more vigorous down group)
the metal dissolving, (faster down group) the solution heating up (more down group)
with calcium a white precipitate appearing (less precipitate forms down group with other metals)

Question 97

extracting titanium

Answer 97

TiO2 (solid) is converted to TiCl4 (liquid) at 900C:
The TiCl4 is purified by fractional distillation in an argon atmosphere.
The Ti is extracted by Mg in an argon atmosphere at 500C

Question 98

why titanium is expenice

Answer 98

1. The expensive cost of the magnesium
2. This is a batch process which makes it expensive because the process is slower (having to fill up and empty reactors takes time) and requires more labour and the energy is lost when the reactor is cooled down after stopping
3. The process is also expensive due to the argon, and the need to remove moisture (because TiCl4 is susceptible to hydrolysis).

Question 99

solubility of groip 2 hydorxides

Answer 99

become more soluble down the group.
All Group II hydroxides when not soluble appear as white precipitates

Question 100

magnesiym hydroxide use

Answer 100

Magnesium hydroxide is used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation.

Question 101

calcium hydroxide

Answer 101

is classed as partially soluble in water and will appear as a white precipitate It is used in agriculture to neutralise acidic soils.
A suspension of calcium hydroxide in water will appear more alkaline (pH 11) than magnesium hydroxide as it is more soluble so there will be more hydroxide ions present in solution.
An aqueous solution of calcium hydroxide is called lime water and can be used a test for carbon dioxide. The limewater turns cloudy as white calcium carbonate is produced.

Question 102

barium hydroixde

Answer 102

would easily dissolve in water. The hydroxide ions present would make the solution strongly alkaline.

Question 103

solubility of sulfates

Answer 103

Group II sulfates become less soluble down the group. BaSO4 is the least soluble.

Question 104

testing for sulfate ions

Answer 104

BaCl2 solution acidified with hydrochloric acid is used as a reagent to
test for sulfate ions.
If acidified barium chloride is added to a solution that contains sulfate ions a
white precipitate of barium sulfate forms.

Question 105

flourine chlorine bromine and iodine colours

Answer 105

Fluorine (F2): very pale yellow gas. It is highly reactive
Chlorine : (Cl2) greenish, reactive gas, poisonous in high concentrations Bromine (Br2) : red liquid, that gives off dense brown/orange poisonous fumes Iodine (I2) : shiny grey solid sublimes to purple gas.

Question 106

halogens melting and boiling point

Answer 106

Increase down the group
As the molecules become larger they have more electrons and so have larger van der waals forces between the molecules. As the intermolecular forces get larger more energy has to be put into break the forces. This increases the melting and boiling points.

Question 107

halogens trend in electronegativty

Answer 107

Electronegativity is the relative tendency of an atom in a molecule to attract electrons in a covalent bond to itself.
As one goes down the group the electronegativity of the elements decreases.
As one goes down the group the atomic radii increases due to the increasing number of shells. The nucleus is therefore less able to attract the bonding pair of electrons.

Question 108

The reactions of halide ions with silver nitrate.

Answer 108

This reaction is used as a test to identify which halide ion is present. The test solution is made acidic with nitric acid, and then silver nitrate solution is added dropwise

Question 109

flouride and chloride with h2so4

Answer 109

NaF(s) + H2SO4(l)NaHSO4(s) + HF(g) Observations: White steamy fumes of HF are evolved. NaCl(s) + H2SO4(l)NaHSO4(s) + HCl(g) Observations: White steamy fumes of HCl are evolved

Question 110

bromide w h2so4

Answer 110

Br- ions are stronger reducing agents than Cl- and F- and after the initial acid-base reaction, the bromide ions reduce the sulfur in H2SO4 from +6 to + 4 in SO2

Question 111

iodine and h2so4

Answer 111

Observations:
White steamy fumes of HI are evolved. Black solid and purple fumes of Iodine are also evolved
A colourless, acidic gas SO2
A yellow solid of sulfur
H2S (Hydrogen sulfide), a gas with a bad egg smell,

Question 112

chlroine and water

Answer 112

Reaction with water in sunlight
If the chlorine is bubbled through water in the presence of bright sunlight a different reaction occurs.
2Cl2 + 2H2O4H+ + 4Cl- + O2
The same reaction occurs to an equilibrium mixture of chlorine water when standing in sunlight. The greenish colour of chlorine water fades as the Cl2 reacts and a colourless gas (O2) is produced

Question 113

soidum and water

Answer 113

Sodium reacts with cold water. It fizzes around on surface etc.

Question 114

magensium and water

Answer 114

Magnesium reacts very slowly with cold water to form the hydroxide but reacts more readily with steam to form the oxide

Question 115

soidium and oxygen

Answer 115

Sodium burns with a yellow flame to produce a white solid.

Question 116

mg al si p and oxygen

Answer 116

Mg, Al, Si and P burn with a white flame to give white solid smoke.
S burns with a blue flame to form an acidic choking gas

Question 117

ionic oxide properties

Answer 117

The metal oxides (Na2O, MgO, Al2O3) are ionic. They have high melting points. They have ionic giant lattice structures: strong forces of attraction between oppositely charged ions : higher mp. They are ionic because of the large electronegativity difference between metal and O
The increased charge on the cation makes the ionic forces stronger (bigger lattice enthalpies of dissociation) going from Na to Al so leading to increasing melting points.
Al2O3 is ionic but does show some covalent character. This can be explained by the electronegativity difference being less big or alternatively by the small aluminium ion with a high charge being able to get close to the oxide ion and distorting the oxide charge cloud.

Question 118

macromoleuckar proeprites

Answer 118

SiO2 is macromolecular: It has many very strong covalent bonds between atoms. High energy needed to break the many strong covalent bonds – very high mp +bp

Question 119

simple molculear properties

Answer 119

P4O10 (s), SO2 (g) are simple molecular with
weak intermolecular forces between molecules (van der waals + permanent dipoles) so have lower mp’s. They are covalent because of the small electronegativity difference between the non-metal and O atoms. P4O10 is a molecule containing 4P’s and 10 O’s. As it is a bigger molecule and has more electrons than SO2 it will have larger van der waals forces between molecules and a higher melting poin

Question 120

Ionic oxides

Answer 120

The metal oxides (Na2O, MgO, Al2O3) are ionic. They have high melting points. They have ionic giant lattice structures: strong forces of attraction between oppositely charged ions : higher mp. They are ionic because of the large electronegativity difference between metal and O
The increased charge on the cation makes the ionic forces stronger (bigger lattice enthalpies of dissociation) going from Na to Al so leading to increasing melting points.
Al2O3 is ionic but does show some covalent character. This can be explained by the electronegativity difference being less big or alternatively by the small aluminium ion with a high charge being able to get close to the oxide ion and distorting the oxide charge cloud.

Question 121

how aluminium metal is poretdcted from corrosion

Answer 121

n moist air by a thin layer of aluminium oxide. The high lattice strength of aluminium oxide and its insolubility in water make this layer impermeable to air and water.

Question 122

ionic oxides and water

Answer 122

tend to react with water to form hydroxides which are alkaline
pH 13 (This is a vigorous exothermic reaction)

Question 123

why zinc is not a transition metal

Answer 123

Zinc can only form a +2 ion. In this ion the Zn2+ has a complete d orbital and so does not meet the criteria of having an incomplete d orbital in one of its compounds.

Question 124

what is a complex

Answer 124

is a central metal ion surrounded by ligands.

Question 125

what is a ligand

Answer 125

an atom, ion or molecule which can donate a lone electron pair