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A-level Chemistry > Module 3: periodic table and energy > Flashcards

Module 3: periodic table and energy Flashcards

1
Q

How could you categorise elements in the early 1800s?

A

There were only two ways, by their physical and chemical properties and by their relative atomic mass

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2
Q

What happened in 450 BC?

A

Democritus believed matter to be composed of small particles with empty space between them; they were referred to as atoms

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3
Q

What did Dobereiner do in 1817?

A

He attempted to group similar elements -> Dobereiner triads. He saw that chlorine, iodine and bromine had similar characteristics and also noticed that other properties of bromine fell halfway between those of chlorine and iodine

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4
Q

What did john newlands notice in 1863?

A

In 1863, newlands noticed that if he arranged elements in order of mass, similar elements appeared at regular intervals- every 8th was similar -> the law of octaves

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5
Q

What did Mendeleev do in 1869?

A

Arranged all the known elements by atomic mass, but left gaps in the table where the next element didn’t seem to fit, so he could keep all elements with similar properties in the same group

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6
Q

What did Mendeleev predict?

A

He was able to predict the properties of undiscovered elements that would go in the gap
- when elements were later discovered with properties that matched mendeleev’s predictions, it showed he got it right

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7
Q

Who introduced the modern periodic table?

A

It was produced by Henry Moseley in 1914, he arranged the elements by increasing atomic number rather than by mass

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8
Q

what is the definition of ionisation energy?

A

the first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous

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9
Q

what types of reaction is ionisation energy?

A

a endothermic reaction as you have to be energy in

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10
Q

how does nuclear charge affect ionisation energy?

A

the more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons

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11
Q

how does atomic radius affect ionisation energy?

A

attraction falls off very rapidly with distance, an electron close to the nucleus will be much more strongly attached than one further away

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12
Q

how does shielding affect ionisation energy?

A

as the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge. this lessening of the pull of the nucleus by inner shells of electrons is shielding

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13
Q

what does it mean if you have a high ionisation energy?

A

it means there is a strong attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron

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14
Q

what happens as you go down a group?

A

the ionisation energies generally fall as its easier to remove an electron
- the extra inner shells shield the outer electrons from the attraction of the nucleus

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15
Q

why does IE fall as you go down a group?

A

elements further down a group have extra electrons shells compared to ones above. the extra shells mean that the atomic radius is larger, so the outer electron are further away from the nucleus, which greatly reduces their attraction to the nucleus.

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16
Q

why does the IE increase as you move across a period?

A

the number of protons is increasing as you go across, as the positive charge of the nucleus increases, the electrons are pulled closer to the nucleus making the atomic radius smaller

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17
Q

what happens to the extra electrons that elements gain across a period?

A

they are added to the outer energy level so they don’t really provide any extra shielding effect

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18
Q

where is the outer electron in group 2 and 3?

A
  • its in the p-orbital rather than a s-orbital
  • a p -orbital has a slightly higher than a s orbital in the same shell, so the electron is on average to be found further from the nucleus
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19
Q

what does the p -orbital have in groups 2 and 3?

A
  • the p-orbital has additional shielding provided by the s electrons
  • these factors override the effect of the increased nuclear charge
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20
Q

what happens to elements in group 5?

A

the electron is being removed from a singly occupied orbital

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21
Q

what happens to elements in group 6?

A

the electron is being removed from an orbital containing two electrons
- the repulsion between two electrons in an orbital means that electrons are easier to remove from shared orbitals

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22
Q

when is a atom or molecule ionised?

A

when an electron is removed

- the energy you need to remove the first electron is called the ionisation energy

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23
Q

Whats successive ionisation energy?

A

You can remove all electrons from an atom, each time you remove a electron theres a successive IE

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24
Q

What happens within each shell with IE?

A

The successive IE increases as the electrons are being removed from an increasingly positive ion and theres also less repulsion amongst the rest of the electrons so they’re hold on stronger

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25
Q

When do the big jumps happen?

A

A big jump in IE when a new shell os broken into an electron is being removed from a shell closer to the nucleus

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26
Q

what are giant covalent lattices?

A
  • huge networks of covalently bonded atoms (macromolecular structures)
  • carbon atoms form this type of structure because they can each form four, strong covalent bonds
  • different forms of the same element in the same state -> allotropes
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27
Q

describe diamond?

A
  • each carbon atom is covalently bonded to four other carbon atoms and form a tetrahedral shape
    Due to its strong covalent bonds:
  • diamond has a high MP
  • its very hard
  • its a good thermal conductor as vibrations travel easily through the stiff lattice
  • cant conduct electricity as outer e are held in localised bonds
  • wont dissolve in any solvent
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28
Q

describe graphite’s overall structure?

A
  • carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 bonds each; the fourth outer e of each atom is delocalised between the sheets of hexagons therefore current can flow as they’re free to move
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29
Q

describe the sheets of graphite?

A
  • they are bonded together by weak induced dipole-dipole forces, which can be easily broken so the sheets can easily slide over each other. the layers are quite far apart compared to the length of the covalent bonds, so graphite is less dense than carbon
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30
Q

what are the other 2 properties of graphite?

A
  • it has a high MP due to strong covalent bonds in the sheets
  • its insoluble in any solvent as bonds are too strong to break
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31
Q

describe graphene?

A
  • a sheet of carbon atoms joined together in hexagons. the sheet is one atom thick, making it a 2 dimensional compound
  • the delocalised e in graphene are free to move, without layers they can move quickly above and below layers making it the best known electrical conductor
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32
Q

what are the other 2 properties of graphene?

A
  • the delocalised e strengthen the covalent bonds between the carbon atoms which make it v strong
  • a single layer of graphene is transparent and light
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33
Q

what potential does graphene have to do?

A
  • due to its high strength, low mass and good electrical conductivity graphene has potential applications in high speed electronics and aircraft technology.
  • its flexibility and transparency make it a potentially useful material for touchscreens
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34
Q

what is metallic bonding?

A
  • metal elements exist as giant metallic lattice structures. the electrons in the outermost shell of a metal atom are delocalised. the e are free to move which leaves a positively charged metal atom cation
  • the metal cations are electrostatically attracted to the delocalised negative e. they form a lattice of closely packed cations in a sea of delocalised e. -> metallic bonding
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35
Q

how does the number of delocalised electrons per atom affect the melting point of a metal?

A
  • the more there are, the stronger the bonding will be and the higher the melting point. the size if the metal ion and the lattice structure also affects the MP. a smaller ionic radius will hold the delocalised e closer to the nucleus
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36
Q

what type of conductors are metals in metallic bonding?

A
  • the delocalised e pass kinetic energy so make good thermal conductors
  • the delocalised e can move and carry current so make good electrical conductors
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37
Q

why are metals malleable and ductile?

A
  • there are no bonds holding ions together, the metal ions can slide past each other when the structure is pulled, so metals are malleable and ductile
  • they’re also insoluble (except liquid metals) due to strength of metallic bonds
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38
Q

what are simple molecular structures?

A
  • they contain only a few atoms. the covalent bonds between the atoms in the molecule are v strong, but the melting and boiling points depend upon the strength of the induced dipole-dipole forces between the molecules
  • these intermolecular forces are weak and easily overcome, so they have low MP’s and BP’s
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39
Q

how does the number of delocalised electrons per atom affect the melting point of a metal?

A
  • the more there are, the stronger the bonding will be and the higher the melting point. the size if the metal ion and the lattice structure also affects the MP. a smaller ionic radius will hold the delocalised e closer to the nucleus
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40
Q

what type of conductors are metals in metallic bonding?

A
  • the delocalised e pass kinetic energy so make good thermal conductors
  • the delocalised e can move and carry current so make good electrical conductors
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41
Q

why are metals malleable and ductile?

A
  • there are no bonds holding ions together, the metal ions can slide past each other when the structure is pulled, so metals are malleable and ductile
  • they’re also insoluble (except liquid metals) due to strength of metallic bonds
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42
Q

what are simple molecular structures?

A
  • they contain only a few atoms. the covalent bonds between the atoms in the molecule are v strong, but the melting and boiling points depend upon the strength of the induced dipole-dipole forces between the molecules
  • these intermolecular forces are weak and easily overcome, so they have low MP’s and BP’s
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43
Q

what does it mean the more atoms there are in a molecule?

A
  • the stronger the induced dipole-dipole forces. the noble gases have very low melting and boiling point as they exist as individual atoms, resulting in very weak induced dipole-dipole forces
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44
Q

how is the melting and boiling points of metals such as Na, Mg, Li, Be and Al affected across a period?

A
  • they increases as the metallic bonds get stronger as the ionic radius decreases and the number of delocalised electrons increase
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45
Q

how is the melting and boiling points of giant covalent structures and simple molecular structures affected across a period?

A

GCS - they have strong covalent bonds so need lots of energy to break
SMS - they have weak intermolecular forces to overcome between their molecules. so low M and B points

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46
Q

how is the melting and boiling points of the noble gases affected across a period?

A
  • they have the lowest melting and boiling points in their periods as they are held together by the weakest forces
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47
Q

what is enthalpy change?

A

enthalpy change, delta H, is the heat energy transferred in a reaction at a constant pressure. the units of delta H are Kjmol-1

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48
Q

what is standard conditions?

A
  • 100kpa (about 1 atm) pressure and a temperature of 298K (25v degrees)
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49
Q

what is the activation energy?

A
  • the minimum amount of energy needed to begin breaking reactant bonds and start a chemical reaction
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50
Q

what is the standard enthalpy change of reaction?

A
  • the enthalpy change when the reaction occurs in the molar quantities shown in the equation, under standard conditions
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51
Q

what is the standard enthalpy change of formation?

A
  • the enthalpy change when 1 mole of a compound is formed from its elements in their standard states, under standard conditions
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52
Q

what is the standard enthalpy change of combustion?

A
  • the enthalpy change when 1 mole of a substance is completely burned in oxygen, under standard conditions
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53
Q

what is the standard enthalpy change of neutralisation?

A
  • the enthalpy change when an acid and an alkali react together, under standard conditions, to form 1 mole of water
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54
Q

what is the enthalpy change of reaction?

A
  • you need energy to break bonds, so bond breaking is endothermic and delta H is positive
  • energy is released from bonds that are formed, this is endothermic and delta H is negative
  • the overall effect of these two changes is the enthalpy change of reaction
55
Q

what is attracted to what?

A
  • in ionic bonding positive ions and negative ions are attracted, in covalent molecules the positive nuclei are attracted to the negative charge of the shared electrons in a covalent bonds
56
Q

what is bond dissociation energy?

A
  • you need energy to break this attraction, the stronger the bond the more energy needed to break
  • the amount energy needed per mole is called the bond dissociation energy
  • happens in bond breaking in gaseous compounds
57
Q

what is the equation for enthalpy change of reaction?

A
  • q=mc delta T
  • q is the heat loss or gained (J)
  • m is the mass of water (g)
  • c is the specific heat capacity (4.18)
  • delta T is the change in temperature in (K)
58
Q

what two things do particles in a reaction need for it to take place?

A
  • they need to collide in the right direction, they need to be facing the right way
  • they collide with at least a certain minimum amount of K.E
59
Q

what happens when you increase the temperature/

A
  • the particles will have more kinetic energy and will move faster. so a greater proportion of molecules will have at least the activation energy and be able to react. this pushes the Boltzmann curve to the right
60
Q

what happens when you increase the concentration?

A
  • the particles will be closer together, if they’re closer, they’ll collide more frequently. if there are more collisions they’ll have more chances to react
61
Q

what happens when you increase the pressure?

A
  • if any of the reactants are gases, increasing the pressure will increase the rate of reaction. at higher pressures, the particles will be closer together increasing the chance of successful collisions
62
Q

what is a catalyst?

A
  • it increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy. the catalyst is chemically unchanged at the end of the reaction
  • they dont get used up
63
Q

what does the catalyst do?

A
  • it lowers the activation energy meaning there are more particles with enough energy to react when they collide so in a certain amount of time, more particles react
64
Q

what is a heterogeneous catalyst?

A
  • a catalyst that is in a different physical state to its reactants
  • the reaction happens on the surface of the catalyst so increasing the SA of the catalyst increases the number of molecules that can react at the same time -> increasing the rate
65
Q

what is a homogeneous catalyst?

A
  • a catalyst that is in the same physical state as the reactants, usually this catalyst in an aqueous catalyst
  • it works by forming an intermediate species, the reactants combine with the reactants to make an intermediate species, which then reacts to form the products and reform the catalyst
66
Q

why do so many industries rely on catalysts?

A
  • they can dramatically lower production costs
  • give you more product in a shorter time
  • help make better products
67
Q

where is iron used as a catalyst?

A
  • in ammonia production, if there was no catalyst, the temperature would be raised loads to make the reaction happen quick enough. this would be bad for fuel bills and it reduces the amount of ammonia made
68
Q

what are the benefits of using a catalyst environmentally?

A
  • lower temperatures and pressures means energy is saved to less CO2 is released and fossil fuel reserves are preserved
  • they can reduce waste by allowing a different reaction to be used with a better atom economy
69
Q

what are catalytic converters?

A
  • they are on cars and made from alloys of palladium, rhodium and platinum. they reduce the amount of pollution released into the atmosphere by speeding up the reaction
70
Q

what is dynamic equilibrium?

A
  • as the reactants are used up, the forwards reaction slows down and as more product is formed the reverse reaction speeds up. after a while, the forward reaction will be going at exactly the same rate as the backward reaction, so the amounts of reactants and products wont change -> DYNAMIC EQUILIBRIUM
71
Q

where can dynamic equilibrium happen?

A
  • at equilibrium, the concentrations of reactants and products stays constant. it can only happen in a closed system
72
Q

what is Le Chatelier’s principle?

A
  • it tell you how the position of equilibrium will change if a condition changes. if there’s a change, in concentration, pressure or temperature the equilibrium will move to help counteract the change
73
Q

describe the affect of concentration on equilibrium?

A
  • increasing the conc of a reactant, the equilibrium tries to get rid of the extra reactant. it does this by making more product so the equilibrium will shift to the right
  • increasing the conc of a product, the equilibrium tries to get rid of the extra product. it does this by making more reactant so the reverse reaction goes faster. the equilibrium shifts to the left.
  • decreasing the conc has the opposite effect
74
Q

describe the effect of pressure on equilibrium?

A
  • increasing the pressure, shifts the equilibrium to the side with the fewer gas molecules. this reduces the pressure
  • decreasing the pressure, shifts the equilibrium to the side with more gas molecules. this increases the pressure again
75
Q

describe the effect of temperature on equilibrium?

A
  • increasing the temp means adding heat so the equilibrium shifts in the endothermic (positive delta H) direction to absorb the heat
  • decreasing the temp means removing heat so the equilibrium shifts in the exothermic (negative delta H) direction to try to replace the heat
76
Q

describe the effect of a catalyst on equilibrium?

A
  • catalysts have no effect on the position of equilibrium. they speed up the forward and reverse reactions by the same amount. they wont increase yield - but they do mean equilibrium is reached faster
77
Q

what is the problem wit the haber pressure?

A
  • at a high pressure the reaction vessel would have to be very strong and maintaining a high pressure is costly. in addiyion, high pressures carry a risk of explosion, so 200atm is used
  • at a low temperatures, the reaction is very slow, although a high yield of NH3 is made at equilibrium it would take a long time before equilibrium was established. 450 degrees is used which is high enough to give an acceptable rate but not high enough to drive the equilibrium too far left, reducing the yield
78
Q

what is the general equation for Kc?

A
  • aA + bB -> cC + dD
79
Q

what does it mean the larger or smaller the Kc value

A
  • the larger the value of Kc, the further to the right equilibrium lies and the more products there are relative to the reactants
  • the smaller the value of Kc, the further to the left equilibrium lies and the more reactants there are relative to the products
80
Q

What is the oxidation number?

A

Tells you how many electrons an atom has donated to accepted form an ion, or to form part of a compound

81
Q

What is the oxidation number like of elements?

A
  • they all have an Oxidation Number (ON) of 0. Elements that are bonded to identical atoms eg O2 will also have an ON of 0.
82
Q

What is the oxidation number like of an ion of a single atom?

A

The oxidation number is the same as its charge

83
Q

What is the oxidation number like of a molecular ion?

A
  • its the sum of the oxidation numbers is the same as the overall charge of the ion. So each of the atoms will have an oxidation number of its own, which will add up to overall charge
84
Q

What is the oxidation number for a neutral compound?

A
  • the overall oxidation number is 0. If the compound is made up of more than one element, each element will have its own ON
85
Q

What is oxygens oxidation number?

A
  • its nearly always -2. Except in peroxides (O2 2-) where its -1 and molecular oxygen (O2) where its O
86
Q

What is hydrogens oxidation number?

A
  • it always has an oxidation number of +1, except in metal hydrides (MH= where M = metal) where its -1 and in molecular hydrogen where its 0.
87
Q

When are roman numerals?

A
  • if an element can have multiple oxidation numbers or isnt in its normal oxidation state, its oxidation number can be shown with Roman numerals.
    Eg in iron (III) sulfate, irok has an oxidation number of +2. Formula = FeSO4
88
Q

What is significant about compounds ending in -ate?

A
  • ions with names ending in ate contains oxygen
  • sometime the other element in the ion can exist with different oxidation numbers. In these cases, the ON is attached as a roman numeral after the name of the -ate compound.
  • the roman numerals corresponds to the non-oxygen element in the -ate compound
89
Q

What is oxidation?

A
  • a loss of electrons is called oxidation. A oxidising agent accepts electrons and gets reduced
90
Q

What is reduction?

A
  • a gain of electrons is called reduction. A reducing agent donates electrons and gets oxidised
91
Q

When does the oxidation number for an atom increase or decrease?

A
  • the oxidation number for an atom will increase by 1 for each electron lost
  • the oxidation number for an atom will decrease by 1 for each electron gained
92
Q

What happens when metals and non metals form compounds?

A

Metals - they generally donate electrons to form positive ions they have positive oxidation numbers
Non metals - generally gain electrons meaning they usually have negative oxidation numbers

93
Q

What happens to metals and hydrogen ions in MASH reactions?

A
  • metal atom are oxidised losing electrons to form positive metal ions
  • hydrogen ions in solution are reduced, gaining electrons and forming hydrogen molecules
94
Q

What are halogens?

A
  • they exist as diatomic molecules, their boiling and melting points increase down the group due to increasing strength of the induced dipole dipole forces as the soze and the relative atomic mass of the atoms increase
95
Q

How do halogen atoms react?

A

By gaining an electron in their outer shell to form 1- ions, they’re reduced. As they’re reduced, they oxidise another substance so they are oxidising agents

96
Q

Describe the reactivity of halogens?

A
  • going down the group the atomic radii increase as the outer electrons are further from the nucleus. The outer e are also shielded more from the attraction of the positive nucleus, because there are more inner electrons.
  • this makes it harder for larger atoms to attract the electron needed to form an ion, so larger atoms are less reactive and become less oxidising
97
Q

What do halogens displace?

A

They displace less reactive halide ions from solution
- when these reactions happen, there are colour changes. You can make the changes easier to see by shaking the reaction mixture with an organic solvent eg hexane. The halogen present will dissolve readily in the organic solvent, which settles out as a distinct layer

98
Q

What will chlorine, bromine and iodine displace?

A
  • chlorine displaces Br- and I-
  • bromine displaces I-
  • iodine will displace nothing
99
Q

What is the test for halides?

A
  • you add dilute nitric acid to remove ions that might interfere with the test. Then you add silver nitrate solution (AgNO3)
    Ag+ + X- —> AgX
100
Q

What are the results from adding dilute nitric acid to halide ions?

A
  • Chloride ions form a white precipitate,
  • bromine ions form a cream precipitate
  • ## iodine ions form a yellow precipitate
101
Q

What can you add futher to the dilute nitric acid?

A

Ammonia solution, the larger the ion is the more difficult it is to dissolve

  • with chlorine ions it dissolves in dilute NH3
  • With bromine ions it dissolved in conc NH3
  • With iodine ions its insoluble in conc NH3
102
Q

What’s disproportion?

A

When you react halogens with cold dilute alkali solutions, in these reactions, the halogen is both oxidised and reduced—> disproportion
- X2 + 2NaOH —> NaXO + NaX + H2O
X2 + 2OH- —> XO- + X + H2O

103
Q

How do you form bleach?

A

If you mix chlorine gas with cold, dilute aqueous sodiuk hydroxide the above reaction takes place and you get sodium chlorate (I) solution, NaCLO

104
Q

What happens when you mix chlorine with water?

A

It undergoes disproportion, it forms a mixture of hydrochloric acid and chloric (I) acid (hypochlorus acid)
- Cl2 + H2O —> HCl + HClO

105
Q

What happens when aqueous chloric acid ionises?

A

It makes chlorate ions (I), hypochlorite ions

HClO + H2O —> ClO- + H3O+

106
Q

Why is chlorine an important part of water treatment?

A
  • it kills disease-causing microorganisms
  • some chloring remains in water, and prevents reinfectiom futher down the supply
  • prevents the growth of algae, eliminating tastes and smells and removes discolouration bu organic compounds
107
Q

What are rhe risks to using chlorine in water?

A
  • Cl2 gas is v harmful if its breathed in to the respiratory system
  • liquid chlorine on skin or eyes causes severe chemical burns
  • organic compounds found in water can react to form chlorinated hydrocarbons many which are carcinogenic but their risk is smaller than untreated water
  • ethical issues too, people dont get a choice about having water chlorinated -> mass medication
108
Q

What are the alternatives to chlorine?

A
  • ozone (O3) a strong oxidising agent, good at killing microorganisms however its expensive to produce and has a short half life in water
  • UV light kills microorganisms by damaging their DNA, but it is ineffective in cloudy water and won’t stop water being contaminated futher down the line
109
Q

describe group 2 elements reactivity?

A
  • they lose their outer electrons to form 2+ ions, as you go down the group, the ionisation energies decrease.
  • this is due to increasing atomic radius and shielding effect. the more reactive the element, the easier it is ti lose electrons and so reactivity increases down the group
110
Q

how do group 2 metals react with water?

A
  • they form metal hydroxides and hydrogen
    eg M + 2H20 -> M(OH)2 + H2
    ON: 0 +2
111
Q

how do group 2 metals react with oxygen?

A
  • they burn in oxygen to form white solid oxides
    eg 2M + O2 -> 2MO
    ON of metal: 0 +2
    ON of O2: 0 -2
112
Q

how do group 2 metals react with dilute acid?

A
  • they react to produce a salt (metal hydroxide) and hydrogen
    eg M + 2HCL -> MCL + H2
    ON: 0 +2
113
Q

what do the metal hydroxides formed from group 2 metals do?

A
  • they dissolve. the hydroxide ions, OH-, make these solutions strongly alkaline.
  • magnesium oxide is an exception, it only reacts slowly and the hydroxide isnt very soluble. the oxides form more strongly alkaline solutions as you go down the group as the hydroxides get more soluble
114
Q

what are group 2 metals known as?

A
  • alkaline earth metals as many of their common compounds are used for neutralising acids
  • calcium hydroxide (slaked lime, Ca(OH)2) is used in agriculture to neutralise acidic soils
  • magnesium hydroxide (Mg(OH)2) and calcium carbonate CaCO3 are used in indigestion tablets as antacids
115
Q

what is the ionic equation for neutralisation?

A
  • H+ (aq) + OH- (aq) –> H2O (l)
116
Q

what do giant metallic structures include?

A

lithium, berillium, sodium, magnesium, aluminium

117
Q

what do giant covalent structures include?

A

boron, carbon, and silicon

118
Q

what do simple molecular structures include?

A

nitrogen, oxygen, fluorine, neon, phosphorus, sulfur, chlorine and argon
- they have weak london forces between molecules

119
Q

what are the trend of group 2?

A
  • their reactivity increases going down the group

- going down the group their solubility, pH and alkalinity increases

120
Q

what are the trends of group 7?

A
  • their reactivity decreases going down the group as atomic radius decreases
  • ## their boiling points increase going down
121
Q

what is the test for carbonates?

A
  • add dilute nitric acid to the solution to be tested
  • if you see bubbles it could be a carbonate
  • to prove that the gas bubbles are CO2, you use the limewater test. bubble the gas through limewater, CO2 forms a fine white ppt of CaCO3 which turns limewater cloudy
122
Q

what is the test for sulfates?

A
  • Barium sulfate is insoluble so forms the basis for the test, in which barium ions are added to a solution. a dense, white ppt forms.
123
Q

what is the test for halides?

A
  • add aq silver nitrate to a solution of a halide
  • silver chloride makes a white ppt, silver bromide makes a cream ppt and silver iodide makes a yellow ppt
  • add aq ammonia to test the solubility of the ppts
124
Q

what order should the test be done in?

A
  1. carbonate
  2. sulfate
  3. haildes
125
Q

why should tests be done in a certain order?

A
  • neither sulfate or halide tests make bubbles, so the carbonate test can be done without the possibility of an incorrect conclusion
  • in sulfate test if you carry it out on a carbonate you will get a white ppt too which is the same result for carbonates. so carrying out the carbonate test before this one is important
  • silver carbonate and silver sulfate are both insoluble in water and will form ppts in the halide test which is why it is carried out last to see whether it is a carbonate or sulfate
126
Q

how do you test for ammonium NH4+ ions?

A
  • aq sodium hydroxide is added so a solution
  • ammonia gas is made, the mixture is warmed and ammonia gas is released
  • test with pH indicator paper, ammonia is alkaline and will turn the paper blue
127
Q

what is an endothermic reaction?

A

a reaction where energy is transferred to the system from the surroundings, so the chemical system gains energy. delta H is positive

128
Q

what is an exothermic reaction?

A

a reaction where energy is transferred from the system to the surroundings, so the chemical system loses energy. delta H is negative

129
Q

why may an enthalpy change of combustion value not be accurate? ie less exothermic than expected

A
  1. heat loss to the surroundings other than water
  2. incomplete combustion of the compound
  3. happened non-standard conditions
130
Q

what is bond breaking and bond making?

A
  • bond breaking is endothermic as energy is needed to break bonds and so delta H is negative
  • bond making is exothermic as energy is released when bonds are formed and so delta H is positive
131
Q

what are the reactant molecules like in a heterogenous catalyst?

A

the catalysts are usually solids in contact with gaseous reactants or reactants in solution. the reactant molecules are adsorbed (weakly bonded) onto the surface of the catalyst where the reaction takes place

132
Q

what is desorption?

A

in a heterogenous catalyst, after a reaction has finished, the product molecules leave the surface of the catalyst by desorption

133
Q

what are the molecules like in a boltzmann distribution curve at higher temps?

A
  • more molecules have an energy higher than the Ea
  • a greater proportion of collisions will lead to a reaction, increasing the rate
  • collisions will be more frequent as molecules are moving faster due to increased energy
134
Q

what are the molecules like in a boltzmann distribution curve with a catalyst?

A
  • a catalyst provides an alternative reaction route with a lower Ea
  • a greater proportion of molecules now have an energy equal to or greater than the lower Ea
  • more molecules will then react

A-level Chemistry (5 decks)

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