chem revise flashcards
(125 cards)
1
Q
isotope
A
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.
2
Q
what is mass spectrometer is used from
A
The mass spectrometer can be used to determine all the isotopes present in a sample of an element and to therefore identify elements.
3
Q
what is electron impact
A
*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–
4
Q
Electro spray Ionisation
A
- 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
5
Q
Acceleration
A
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
6
Q
first ionisstion energy
A
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
7
Q
second ionisaion enegry
A
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
8
Q
Factors that affect ionisation energy
A
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
9
Q
Why are successive ionisation energies always larger?
A
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.
10
Q
Why has helium the largest first ionisation energy?
A
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
11
Q
Why do first ionisation energies decrease down a group?
A
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
12
Q
Why is there a general increase in first ionisation energy across a period?
A
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.
13
Q
Why has Na a much lower first ionisation energy than neon?
A
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.
14
Q
Why is there a small drop from Mg to Al?
A
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
15
Q
Why is there a small drop from P to S?
A
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.
16
Q
poisitve ion radiii
A
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.
17
Q
dative covalent bond
A
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
18
Q
factos affectsing metallic bonding
A
- Number of protons/ Strength of nuclear attraction.
The more protons the stronger the bond - Number of delocalised electrons per atom (the outer shell electrons are delocalised)
The more delocalised electrons the stronger the bond - Size of ion.
19
Q
ionic properties
A
high mp and bp
soluble in water
cant conduct when solid
20
Q
simple comleuclar properties
A
low mp and bp
poor solubility
poor conductivity when solid
poor conductivity when molten
21
Q
macromolecular properties
A
high mo and bp
soluble
poor conductivty
22
Q
metallic properties
A
high mp and bp
insolble in water
good conductvity
23
Q
linear
A
2
0
180
24
Q
trigonal planar
A
3
0
120
25
tetrahderal
4
0
109.5
26
trigonal planar
3
1
107
27
bent
2
2
104.5
28
trigonal bipyramidal
5
0
120 and 90
29
octahedral
6 bp
0lp
90
30
factors affecting electroneg
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
31
Formation of a permanent dipole
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.
32
Symmetric molecules
A symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar.
33
anomalously high boiling points of H2O, NH3 and HF are caused by
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
34
Main factor affecting size of Van der Waals
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
35
Permanent dipole-dipole forces
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.
36
what is emthalpy change
is the amount of heat energy taken in or given out during any change in a system provided the pressure is constant.
37
stsandard enthalpy of formation
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.
38
standard enthalpy of combustion
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.
39
hess law
Hess’s law states that total enthalpy change for a reaction is independent of the route by which the chemical change takes place
40
H reaction =
Σ fH products - Σ fH reactantsi
41
Mean Bond energies
Definition: The mean bond energy is the enthalpy needed to break the covalent bond into gaseous atoms, averaged over different molecules.
42
what is coliision theory
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.
43
what is rate of reaction
as the change in concentration of a substance in unit time. The usual unit is mol dm-3s-1
44
effects of inc conc inc pressure
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.
45
effect of inc temp
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.
46
effect of catalyst
They do this by providing an alternative route or mechanism with a lower activation energy.
47
electron acceptor
element thats reduced in the equation
48
electron donor
reducing agent
49
enthalpy of fromation e
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
50
enthalpy pf atomisation
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
51
bond enthalpy of dissociaion
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)
52
first ionitsation enthalpy
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
53
first elevton afffintity
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
54
enthalpy of lattice formation
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
55
second iomnisatioj enthalpt
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.
56
second electron affintiy
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
57
enthalpy of hydration
Enthalpy change when one mole of gaseous ions become aqueous ions
58
enthalpy of solution
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.
59
trends in lattice enthalpy
The sizes of the ions
The charges on the ion
60
perfect ionic model
Theoretical lattice enthalpies assumes a perfect ionic model where the ions are 100% ionic and spherical and the attractions are purely electrostatic.
61
sspontaneous process
(e.g. diffusion) will proceed on its own without any external influence.
62
Entropy, S ̊
Substances with more ways of arranging their atoms and energy (more disordered) have a higher entropy
63
An increase in disorder and entropy will lead to
o a positive entropy change ∆S ̊ = +ve
64
increase in entropy and temp on feasbility
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
65
if the reaction has a ∆S close to zero then
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
66
electrochemicl cell
*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).
67
why does a voltage form in an electrochemical cell
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. (ZnZn2+ + 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
68
why use a high resistance voltmeter
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.
69
what a salt bridge is made from and used for
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.
70
What happens if current is allowed to flow?
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.
71
which electrode wil undergo reduciton
The most positive electrode will always undergo reduction.
72
which electrode will undergo oxidation
The most negative electrode will always undergo oxidation.
73
solid vertical line in cell diagrams
represents the boundary between phases e.g. solid (electrode) and solution (electrolyte)
74
double line in cell diagram
*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
75
systems that dont include metals
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
76
how potential of electrodes are measured
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.
77
Components of a standard hydrogen electrode.
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
78
Effect of concentration on Cell voltage Ecell
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
79
effect of temp on e cell
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.
80
what is a fuel cell
uses the enrgy from the reaction of a fuel with oxygen to create a voltage
81
conditions of fule cells
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
82
advantage of fuel cells
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
83
kimitations of hyfrogen fuel cells
(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
84
ethanol fuel cells
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.
85
atomic radius trend
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.
86
1st ionisatio energy trend
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
87
melting and boiling point trend
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
88
period 2 trend
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)
89
atomic radius
Atomic radius increases down the group.
As one goes down the group, the atoms have more shells of electrons making the atom bigger
90
melting points down the group
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
91
first ionisation energy
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.
92
group 2 w o2
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.
93
reactivity of group 2 metals
increases down the group.
94
group 2 and steam
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.
95
``group 2 and warm water
giving a different magnesium hydroxide product.
Mg + 2 H2OMg(OH)2 + H2
This is a much slower reaction than the reaction with steam and there is no flame.
96
generwl observatios withbgroup 2 and water
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)
97
extracting titanium
TiO2 (solid) is converted to TiCl4 (liquid) at 900C:
The TiCl4 is purified by fractional distillation in an argon atmosphere.
The Ti is extracted by Mg in an argon atmosphere at 500C
98
why titanium is expenice
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).
99
solubility of groip 2 hydorxides
become more soluble down the group.
All Group II hydroxides when not soluble appear as white precipitates
100
magnesiym hydroxide use
Magnesium hydroxide is used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation.
101
calcium hydroxide
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.
102
barium hydroixde
would easily dissolve in water. The hydroxide ions present would make the solution strongly alkaline.
103
solubility of sulfates
Group II sulfates become less soluble down the group. BaSO4 is the least soluble.
104
testing for sulfate ions
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.
105
flourine chlorine bromine and iodine colours
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.
106
halogens melting and boiling point
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.
107
halogens trend in electronegativty
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.
108
The reactions of halide ions with silver nitrate.
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
109
flouride and chloride with h2so4
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
110
bromide w h2so4
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
111
iodine and h2so4
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,
112
chlroine and water
Reaction with water in sunlight
If the chlorine is bubbled through water in the presence of bright sunlight a different reaction occurs.
2Cl2 + 2H2O4H+ + 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
113
soidum and water
Sodium reacts with cold water. It fizzes around on surface etc.
114
magensium and water
Magnesium reacts very slowly with cold water to form the hydroxide but reacts more readily with steam to form the oxide
115
soidium and oxygen
Sodium burns with a yellow flame to produce a white solid.
116
mg al si p and oxygen
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
117
ionic oxide properties
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.
118
macromoleuckar proeprites
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
119
simple molculear properties
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
120
Ionic oxides
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.
121
how aluminium metal is poretdcted from corrosion
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.
122
ionic oxides and water
tend to react with water to form hydroxides which are alkaline
pH 13 (This is a vigorous exothermic reaction)
123
why zinc is not a transition metal
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.
124
what is a complex
is a central metal ion surrounded by ligands.
125
what is a ligand
an atom, ion or molecule which can donate a lone electron pair
