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chemistry a level > periodic table and energy > Flashcards

periodic table and energy Flashcards

1
Q

what is the periodic table? (3)

A

the arrangement of elements by:
1) increasing atomic (proton) number
2) periods showing repeating trends in physical and chemical properties (periodicity)
3) in groups having similar chemical properties

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2
Q

what is the definition of first ionisation energy?

A

energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous + ions.

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3
Q

what does ionisation energy measure?

A

measures how easily an atom loses electrons to form positive ions

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4
Q

what factors affect ionisation energy? (3)

A

1) no. of electron shells/ atomic radius
2) nuclear charge
3) electron shielding (inner electron shells blocking charge from the nucleus)

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5
Q

what is the first ionisation energy equation for sodium?

A

Na (g) –> Na+(g) + e-

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6
Q

what is the general rule in successive ionisation energies?

A

no. of electrons = ionisation energies

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7
Q

every time an electron is removed… (3) (successive ionsation energy)

A

1) there is a greater nuclear attraction
2) atomic radius is slightly smaller
3) successive ionisation energies will increase

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8
Q

how can you tell the electronic structure of an element from data which gives ionisation number and energies? (2)

A

1) name where the jump in energy occurs
2) say how many electrons are in its outer shell from that

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9
Q

what do trends in 1st ionisation energy provide evidence for?

A

the existence of shells and sub shells

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10
Q

what is the trend in first ionisation energies down a group? (3)

A

1) first ionisation energy decreases down a group
2) atomic radius increases down a group
3) so shielding increases meaning it is easier for the electron to be lost as attraction is weaker between the nucleus and electron

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11
Q

what is the trend in first ionisation energies across period 2? (3)

A

1) general increasing trend along the period
2) due to decreasing atomic radius and increasing nuclear charge meaning outer electrons are held more strongly
3) which requires more energy to overcome

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12
Q

why is boron an exception to the trend in first ionisation energy across period 2? (2)

A

1) Boron has a lower first ionisation energy than expected due to the energy difference between 2s and 2p sub shells / new sub shell shielding
2) the electron is being removed from a higher energy level that is further from the nucleus so the electron is held less strongly.

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13
Q

why is oxygen an exception to the trend in first ionisation energy across period 2? (3)

A

1) oxygen has a lower first ionisation energy than expected by the general trend due to repulsion within the 2p orbital
2) two electrons with opposite spins are placed in the same orbital
3) allows electron to be removed from oxygen more easily as the repulsion is destabilising

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14
Q

what is the trend in first ionisation energies across period 3? (3)

A

1) general increasing trend
2) due to decreasing atomic radius and increasing nuclear charge
3) meaning outer electrons are held more strongly

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15
Q

why is aluminium an exception to the trend in first ionisation energy across period 3? (3)

A

1) aluminium has a lower first ionisation energy due to the energy difference between the 3s and 3p sub shells
2) the electron is being removed at a higher energy level which is further away from the nucleus
3) held less strongly so requires less energy to remove

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16
Q

why is sulfur an exception to the trend in first ionisation energy across period 3? (3)

A

1) lower first ionisation energy than expected due to the electron repulsion in the 3p sub shell
2) electrons with opposite spins held in the same orbital
3) allows electron to be removed from sulfur more easily as the repulsion is destabilising

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17
Q

what is the general trend for first ionisation energy down a group? (3)

A

1) first ionisation energy decreases down a group
2) more electron shells lead to increasing atomic radius so weaker nuclear attraction
3) shielding also increases down a group which further reduces attraction

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18
Q

what is the general trend for first ionisation energy across a period? (4)

A

1) first ionisation energy increases across a period
2) due to a decreasing atomic radius due to extra nuclear charge
3) greater nuclear attraction to overcome
4) more nuclear charge because of an extra proton

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19
Q

what 3 things affect first ionisation energy? describe the trend in these across a group and down a period

A

1) atomic radius (increases down a group, decreases across a period)
2) nuclear charge (increases across a period, decreases down a group)
3) shielding (constant across a period, increases down a group)

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20
Q

why does shielding remain constant across a period?

A

the number of energy levels is the same

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21
Q

why does atomic radius decrease across a period? (2)

A

1) elements gain an extra proton for the same number of electron shells, which increases nuclear charge
2) attracting the electron closer which decreases atomic radius

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22
Q

how are atomic radius and nuclear charge linked?

A

A higher nuclear charge causes greater attractions to the electrons, pulling the electron shells closer to the nucleus which results in a smaller atomic radius.

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23
Q

what is the periodic trend in electron configurations across Periods 2 and 3? (2)

A

1) across period 2, the 2s sub shell fills with 2 electrons and then the 2p subs shells fill with 6 electrons
2) in period 3 the pattern repeats itself with the 3s and 3p sub-shells.

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24
Q

How are elements classified into SPDF block elements?

A

classification is based on the name of orbitals which receive the last electron.

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25
Q

what can successive ionisation data be used for? (3)

A

1) predict or confirm the simple electronic configuration of elements (the number of ionisation energies shows the number of electrons)
2) confirm the number of electrons in the outer shell of an element
3) deduce the group an element belongs to in the periodic table

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26
Q

what is electron shielding?

A

when the inner shells create a barrier that blocks the attractive forces

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27
Q

what is metallic bonding?

A

strong electrostatic attraction between cations and delocalised electrons

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28
Q

what is ionisation energy measured in?

A

kJmol-1

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29
Q

when do successive ionisation energies occur?

A

when further electrons are removed after first ionisation energy

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30
Q

why does successive ionisation energy require more energy? (2)

A

1) as electrons are removed, nuclear attraction between the positive nucleus and the negative outer shell electron increases
2) more energy is needed to overcome this attraction, so ionisation energy increases

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31
Q

what is the structure of metals? (3)

A

1) metal loses all of its outer shell electrons causing a sea of delocalised electrons
2) lattice of cations
3) strong electrostatic attraction between the oppositely charged particles

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32
Q

what should you remember when drawing a diagram of metallic structure?

A

when asked to draw a specific metal, add the charge to the cation
e.g
Ca would be 2+

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33
Q

what are the properties of metals? (3)

A

1) conductor of electricity as delocalised electrons are able to move and carry charge through their structure
2) high melting and boiling points due to strong electrostatic attraction between cations and electrons which require a lot of energy to break
3) insoluble as metallic bonds are too strong to be overcome

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34
Q

which 3 elements form giant covalent structures?

A

Boron, Carbon and Silicon

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35
Q

what are the properties of giant covalent structures? (3)

A

1) high melting and boiling points due to a high proportion of strong covalent bonds which require a lot of energy to break
2) insoluble as covalent bonds are too strong to be overcome
3) non conductors of electricity as there are no delocalised electrons which can move and carry charge.

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36
Q

what is the periodic trend in melting points across period 2? (4)

A

1) increases between Li and Be as Be forms 2+ ions so there is a stronger metallic bonds as there are more electrons
2) increases between Be to B to C as bonding switches to giant covalent structures
3) decreases from C to N, O and F as they have simple molecular structure
4) decreases most at Ne as it is a group 0 element

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37
Q

what are the simple molecular structures in period 2?

A

N2, O2, F2, Ne

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38
Q

what are the simple molecular structures in period 3?

A

P4, S8, Cl2, Ar

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39
Q

what is a giant covalent lattice?

A

a solid giant covalent lattice is a network of atoms bonded by many strong covalent bonds

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40
Q

what are the properties of giant covalent structures? (3)

A

1) high melting and boiling points due to each atom having multiple strong covalent bonds
2) insoluble as covalent bonds are too strong to be overcome
3) non conductors of electricity as there are no delocalised electrons which can move and carry charge

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41
Q

what are the allotropes of carbon?

A

diamond, graphite and graphene

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42
Q

what are the features of diamond? (3)

A

1) each carbon makes 4 covalent bonds
2) causes a rigid/ hard structure
3) non conductor of electricity as there are no delocalised electrons which can move and carry charge

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43
Q

what are the features of graphite? (3)

A

1) each carbon makes 3 covalent bonds in flat sheets
2) one delocalised electron per atom of carbon which can move between layers so it can conduct electricity
3) weak London forces between molecules allow layers to slide over each other easily

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44
Q

what are the features of graphene? (4)

A

1) 2D sheets of graphite
2) form hexagonal carbon rings
3) strong, rigid, light structure
4) conductor of electricity due to only 3 covalent bonds

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45
Q

what causes the variation in melting point in period 2?

A

due to different bond strength and structure

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46
Q

what is the melting point like for Li and Be in period 2? (3)

A

1) melting point increases from Li to Be because Be has a greater positive charge of 2+ when forming ions
2) more free electrons in Be lattice
3) electrostatic forces are greater than that in lithium

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47
Q

what is the melting point like for B and C in period 2? (2)

A

1) B and C form giant covalent lattices with very strong covalent bonds
2) covalent bonds require a lot of energy to break, causing high melting points

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48
Q

what is the melting point like for N2 O2, F2 and Ne in period 2? (2)

A

1) simple covalent molecules with weak London forces
2) relatively similar low melting and boiling point as not much energy is required to overcome these forces

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49
Q

what causes the variation in melting point in period 3?

A

difference in bond strength and structure

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50
Q

what is the melting point like for Na, Mg and Al in period 3? (3)

A

1) melting point increases from Na to Al due to greater positively charged metal ions
2) more electrons are released as free electrons across the metals
3) electrostatic attraction between electrons and metal cations increases from Na to Al

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51
Q

what is the melting point like for Si in period 3? (2)

A

1) giant covalent structure
2) strong covalent bonds which require a lot of energy to break causes high melting point

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52
Q

what is the melting point like for P4, S8 and Cl2 in period 3? (3)

A

1) simple covalent molecules
2) weak London forces between them
3) low similar boiling points as it does not require a lot of energy overcome intermolecular forces

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53
Q

what is the melting point like for Ar in period 3? (2)

A

1) noble gas with a full outer shell of electrons
2) stable molecule so London forces are weak

54
Q

what valence shell configuration do group 2 metals have and how does this link to how they react? (2)

A

1) s2
2) group 2 metals lose these 2 electrons in redox reactions to form 2+ ions

55
Q

what happens when group 2 elements are exposed to air? (2)

A

1) tarnish in air
2) forming a coat of metal oxide

56
Q

what is the trend in reactivity down group 2? (3)

A

1) increases down the group
2) because ionisation energies decrease down the group
3) due to increased electron shielding and atomic radius which makes outer electrons easier to lose in order to react

57
Q

explain the trend in reactivity in terms of the first and second ionisation energies of group 2 elements down the group? (2)

A

1) both first and second ionisation energy of group 2 elements decrease down the group
2) due to greater atomic radius and increased shielding which makes outer shell electron easier to lose

58
Q

describe the relative reactivities of group 2 elements Mg–> Ba shown by their redox reactions with water? (2)

A

1) group 2 metals react with water in a redox reaction to produce a metal hydroxide and hydrogen gas
2) metal hydroxide forms as an alkaline solution
e.g Mg + 2H2O –> Mg(OH)2 + H2
Mg is oxidised from 0 to +2
H is reduced from +2 to 0

59
Q

describe the trend in reactivity of group 2 elements down the group with water (3)

A

1) increases down the group
2) Ca, Sr and Ba all react with water with increasingly vigorous reactions
3) reactions all form metal hydroxide and hydrogen gas

60
Q

describe the relative reactivities of group 2 elements Mg–> Ba shown by their redox reactions with oxygen

A

group 2 metals react with oxygen to form metal oxides
e.g
2Mg + O2 –> 2MgO

61
Q

how do Sr and Ba react with oxygen? give an example (2)

A

1) Sr and Ba can react with excess oxygen and heat energy to form metal peroxides
e.g
Sr + O2 –> SrO2

62
Q

describe the relative reactivities of group 2 elements Mg–> Ba shown by their redox reactions with dilute acids?
give examples with all 3 acids

A

group 2 metals react with dilute acids to produce hydrogen gas and a salt
e.g
X + 2HCl –> XCl2 + H2
X + H2SO4 –> XSO4 + H2
X + 2HNO3 –> X(NO3)2 + H2

63
Q

what is the solubility of group 2 oxides like? (3)

A

1) group 2 oxides increase in solubility down the group
2) more hydroxide ions released in solution
3) more alkaline solution down the group

64
Q

describe the action of water on group 2 oxides and the approximate pH of any resulting solutions, including the trend in increasing alkalinity (3)

A

1) group 2 oxides react with water forming an alkaline solution of a metal hydroxide
e.g
CaO + H2O –> Ca(OH)2
2) hydroxides are only slightly soluble in water- once the solution is saturated any further, hydroxide ions form solid precipitate
3) more alkaline solution down the group as more OH- release as solubility increases

65
Q

what is the outer shell configuration for group 2?
how does this influence the most common type of reaction of group 2 elements? (2)

A

1) s2 electron configuration
2) the loss of these electrons occurs in redox reactions to form 2+ ions

66
Q

explain the trend in reactivity in terms of the first and second ionisation energies of Group 2 elements down the group (3)

A

1) Group 2 elements react losing 2 electrons, forming 2+ cations
2) The easier it is to lose electrons (i.e. the lower the first and second ionisation energies), the more reactive the element.
3) So because ionisation energy decreases down the group (due to increase in atomic radius+shielding) reactivity increases down the group.

67
Q

describe the uses of some group 2 compounds as bases, including equations (2)

A

1) Ca(OH)2 in agriculture to neutralise acid soils- has to be used in moderation
2) Mg(OH)2 and CaCO3 as ‘antacids’ to neutralise stomach acids/ treating indigestion, CaCO3 is an insoluble base so it can be swallowed in indigestion tablets.
equation:
- Mg(OH)2(s) + 2 HCl(aq) → MgCl2(aq) + 2 H2O(l)
- CaCO3(s) + 2 HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)
- H+(aq) + OH-(aq) -> H2O(l)

68
Q

what are some features of the halogens? (4)

A

1) most reactive non-metallic group
2) do not occur in elemental form
3) occur as stable halide ions dissolved in sea water/ combined with Na or K as solid deposits
4) all halogens exist as diatomic molecules

69
Q

what is the trend in boiling point for the halogens? (3)

A

1) increases down the group
2) London forces get stronger due to more electrons
3) so bigger dipoles form bewitch require more energy to overcome

70
Q

describe the outer shell electron figuration of the halogens and how this affects how they react

A

1) s2p5
2) gain one electron in many redox reactions to form 1- ions

71
Q

what do halogens form in their solid state?

A

lattices with simple molecular structures

72
Q

what is the appearance of the halogens at room temperature? (4)

A

1) F2–> pale yellow gas
2) Cl2 –> pale green has
3) B2–> red-brown liquid
4) I2–> shiny black-grey solid

73
Q

what is the most common type of reaction for the halogens? (2)

A

1) redox
2) each halogen atom is reduced and acts as an oxidising agent (gains an electron from another element which loses it)

74
Q

what is the general half equation for the reduction of halogens in halogen displacement reactions?

A

X2(aq) + 2e- –> 2X- (aq)
(more reactive one)

75
Q

what is the general half equation for the oxidation of halides in halogen displacement reactions?

A

2X- (aq) –> X2 (aq) + 2e-
(less reactive one)

76
Q

in aqueous solutions what do Cl2, Br2 and I2 appear as? (3)

A

1) Cl2 (aq) –> pale green
2) Br2 (aq) –> orange
3) I2 (aq) –> brown

77
Q

why is cyclohexane added to halogens in aqueous solutions?

A

to make the colour more clear

78
Q

in the cyclohexane layer, what do Cl2, Br2 and I2 appear as? (3)

A

1) Cl2 (aq) –> pale green
2) Br2 (aq) –> orange
3) I2 (aq) –> violet

79
Q

what do the results of halogen displacement reactions show?

A

that the reactivity of halogens deceases down the group

80
Q

what happens during a halogen displacement reaction?

A

a solution of each halogen is added to aqueous solutions of the other halides
(halogen and halide will be different elements)

81
Q

give an example of a halogen displacement reaction and give the ionic and half equations

A

Cl2 (aq) + 2NaBr (aq) –> 2NaCl (aq) + Br
chlorine is more reactive than bromine
ionic equation:
Cl2 (aq) + 2Br- (aq) –> 2Cl- (aq) + Br2 (aq)
half equations:
Cl2(aq) + 2e- –> 2Cl-
2Br- –> Br2 (aq) + 2e-

82
Q

what is the general rule for writing ionic equations in halogen displacement reactions?

A

eliminate metal ion- only thing that remains unchanged/ spectator ion

83
Q

explain the trend in reactivity of the halogens from the decreasing ease of forming 1- ions (4)

A

1) decreases down the group as it is harder to gain an electron/ their oxidising ability decreases
2) due to extra electron shells which leads to a larger atomic radius
3) more shielding
4) weaker attraction to the nucleus so it is harder to gain the electron they need to react

84
Q

explain the trend in reactivity of the halogens illustrated by reactions with other halide ions (2)

A

1) The reactivity of halogens can be shown by their displacement reactions with other halide ions in solutions
2) A more reactive halogen can displace a less reactive halogen from a halide solution of the less reactive halogen

85
Q

what is a disproportionation reaction?

A

the same element is both oxidised and reduced in a reaction, can be illustrated by analogous reactions.

86
Q

what is one example of a disproportionation reaction?

A

reaction of chlorine with water used to purify water
Cl2 (g) + H2O(l) –> HCl (aq) + HClO (aq)
(hypochlorous acid)
Cl is both reduced and oxidised

87
Q

what is another example of a disproportionation reaction?

A

reaction of chlorine with aqueous sodium hydroxide
Cl2 (g) + 2NaOH (aq) –> NaClO (aq) + NaCl (aq) + H2O (l)
(sodium hypochlorite)
Cl is both oxidised and reduced

88
Q

what is the trend in oxidising power down the halogens? (4)

A

decreases
1) ability to attract electrons decreases
2) more electron shells so greater atomic radius
3) more shielding
4) weaker electrostatic attraction to attract electron and be reduced

89
Q

what are the pros and cons of chlorine use in water treatment? (3)

A

pros: kills bacteria
cons: risks with hazards of toxic chlorine gas
possible risks from formation of chlorinated hydrocarbons

90
Q

what is the problem with the test for halide ions? how can you resolve this? (2)

A

1) colours are hard to tell apart
2) add dilute ammonia solution

91
Q

describe the test for halide ions and the ionic equation

A

1) add nitric acid
2) add silver nitrate
3) results:
AgCl- –> white precipitate
AgBr- –> cream precipitate
AgI–> yellow precipitate
ionic equation:
Ag+ (aq) + X- –> AgX (s)

92
Q

describe the results of adding dilute ammonia solution after the halide test

A

Cl –> precipitate dissolves
Br –> no change/ cream precipitate
I –> no change/ yellow precipitate

93
Q

describe the results of adding concentrated ammonia solution after the halide test

A

Cl –> precipitate dissolves
Br –> precipitate dissolves
I –> no change/ yellow precipitate

94
Q

describe the test for carbonates (4)

A

1) add a dilute acid (preferably nitric),
2) CO32- will react with H+ (aq) forming CO2 (g)
3) bubbles of CO2 will be produced
4) bubble through limewater which will turn cloudy

95
Q

what is the ionic equation for the carbonate test

A

2H + (aq) + CO3 2-(aq) –> CO2 (g) + H2O (l)

96
Q

describe the test for sulfates (3)

A

1) add nitric acid to eliminate any carbonate ions
2) add barium chloride or barium nitrate
3) will result in a white precipitate with Ba2+ (aq) forming CO2 (g)

97
Q

what is the ionic equation for the sulfate test?

A

Ba 2+ (aq) + SO4 2- (aq) –> BaSO4 (s)

98
Q

what is the general ionic equation for the halide test?

A

Ag+ (aq) + X- (aq) –> AgX (s)

99
Q

how would you identify an unknown solution? (6)

A

1) test for carbonates
2) test for sulfates
3) test for halides
4) use nitric acid instead of sulfuric acid/ hydrochloric in the tests to avoid adding Cl- and SO4 2- ions to sample
5) use barium nitrate instead of barium chloride in the test for sulfates to avoid adding Cl- ions
6) This will prevent false results as both BaCO3 and Ag2SO4 are insoluble.

100
Q

describe the test for ammonium ions (NH4+)
`(4)

A

1) add dilute sodium hydroxide and heat
2) will produce ammonia and water
3) hold damp red litmus paper in the mouth of test tube
4) paper will turn blue (alkali)

101
Q

what is the ionic equation for the test for ammonium ions?

A

NH4 + (aq) + OH- (aq) –> NH3 (g) + H2O (l)

102
Q

why should you do the test for carbonates before the test for sulfates?

A

to eliminate any CO32- ions as they will give a positive result for the sulphate test- remove false positives

103
Q

what is enthalpy and its symbol?

A

measure of heat energy in a chemical system
symbol: H

104
Q

what are some chemical reactions accompanied by?
why? (2)

A

1) enthalpy changes that are either exothermic (negative energy change) or endothermic (positive energy change)
2) When chemical reactions take place, changes in chemical energy take place and therefore the enthalpy changes

105
Q

describe the features of an exothermic reaction (6)

A

1) the enthalpy of the reactants is higher than the products
2) Heat energy is given off by the reaction to the surroundings
3) The temperature of the environment increases
4) The energy of the system decreases
5) There is an enthalpy decrease during the reaction so ΔH is negative
6) more energy is released making bonds than breaking bonds

106
Q

what is a chemical system?

A

a system of atoms, molecules or ions which make up the chemicals

107
Q

what is an enthalpy change and what is its symbol?

A

the difference between the enthalpies of the reactants and products
symbol: ΔH

108
Q

what is the general equation for enthalpy change?

A

H(products)-H(reactants)

109
Q

can enthalpy be measured? (2)

A

1) enthalpy can not be measured directly
2) however enthalpy change can be measured by an experiment

110
Q

what is the conservation of energy? (2)

A

1) energy cannot be created or destroyed
2) energy can only be transferred between the system and the surrounding

111
Q

describe the features of an endothermic reaction (6)

A

1) the enthalpy of the reactants is lower than the products
2) heat energy is taken in from the surroundings
3) the temperature of the environment decreases
4) the energy of the system increases
5) There is an enthalpy increase during the reaction so ΔH is positive
6) more energy is taken in breaking bonds than is released when making new bonds

112
Q

what is activation energy?

A

the minimum amount of energy required to start a reaction

113
Q

what is one reason that some reactions may happen slowly?

A

only a small proportion of reactant particles have the activation energy

114
Q

what are the standard conditions? (4)

A

1) standard pressure: 100kPa
2) standard temperature: 298k (25 degrees)
3) standard concentration: 1 mol dm-3
4) standard state: the state of the substance at standard conditions.

115
Q

what is the enthalpy change of reaction?

A

enthalpy change that accompanies a reaction in the molar quantities shown in a chemical equation, under standard conditions

116
Q

what is the enthalpy change of formation?

A

enthalpy change that takes place when one mole of a compound is formed from its elements under standard conditions

117
Q

give an example of the enthalpy change of formation and explain

A

1/2 N2 (g) + 3/2 H2 (g) –> NH3 (g)
the compound needs to remain as one mole so in order to balance it is necessary to balance the reactants in fractions

118
Q

what is the enthalpy change of combustion?

A

the enthalpy change that takes place when one mole of a substance reacts completely with oxygen under standard conditions

119
Q

give an example of the enthalpy change of combustion and explain

A

CH4 (g) + 2O2 (g) –> CO2 (g) + 2H2O (l)
substance needs to remain as one mole but everything else can be balanced.

120
Q

what is the enthalpy change of neutralisation?

A

the enthalpy change when the reaction of an acid and a base form one mole of water under standard conditions

121
Q

give an example of the enthalpy change of neutralisation and explain

A

HNO3 (aq) + 1/2Ba(OH)2 (aq) –> 1/2 Ba(NO3)2 (aq) + H2O (l)
water needs to remain as one mole so everything else needs to be balanced with fractions

122
Q

what equation do you use to calculate heat energy which you need to calculate enthalpy changes?

A

q = mcΔT
q- heat energy (J)
m- mass of water being heated (g or cm3)
c- specific heat capacity of water (4.18)
ΔT- temperature change (K/C)

123
Q

what equation do you use to calculate enthalpy change after calculating heat energy?

A

q/mol. of limiting reactant = Jmol-1
divide by 1000 to get kJmol-1

124
Q

what 4 things do you need to measure when carrying out an enthalpy change of combustion experiment? (2)

A

1) weigh spirit burner before and after for mol of limiting reactant
2) measure temperature of water before and after

125
Q

how accurate is the experimental value when doing the enthalpy change of combustion experiment? (4)

A

1) heat is lost to the surroundings (not all heat goes toward heating the water)
2) evaporation of alcohol from the wick (mass is bigger than it should be on the balance)
3) assuming complete combustion has occured
4) may have occurred under non-standard conditions

126
Q

what experiment do you carry out to determine the enthalpy change of reaction? (4)

A

calorimetry
1) reaction is carried out in a polystyrene cup with a lid
2) thermometer is placed in polysterene cup
3) measure temperature
4) plot a graph time on x axis and extrapolate for temp. change

127
Q

how do you calculate the enthalpy change of reaction? (2)

A

1) q= mcΔT
2) ΔrH = q/mol of limiting reactant
- if question gives 2 liquids then add the two volumes together
- if question doesn’t specify limiting reactant, work it out by calculating the moles of each reactant then writing out the symbol equation

128
Q

how can you improve experimental values in determining the enthalpy change of reaction? (3)

A

1) take temperature every 30 seconds/ frequent time intervals
2) plot a graph of temperature v time (time on the x axis)
3) extrapolate to correct for heat loss

129
Q

bond enthalpies are always…

A

endothermic (+ve enthalpy value)

130
Q

what are the limitations of average bond enthalpies? (2)

A

1) the actual bond enthalpy can vary depending on the chemical environment of the bond
2) an average bond enthalpy is calculated from the actual bond enthalpies in different chemical environments

chemistry a level (6 decks)

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