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A Level Chemm > Unit 3 > Flashcards

Unit 3 Flashcards

1
Q

What’s meant by first ionisation energy

A

energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of gaseous +1 ions
The lower the ionisation energy the easier it is to form an ion

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2
Q

Explain why the ionisation energy of any element is endothermic

A

because energy is required to overcome the electrostatic attraction between positive nucleus and negative electron

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3
Q

State 3 factors which affect first ionisation energies

A

Nuclear charge, atomic radius, shielding

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4
Q

What’s nuclear charge

A

The more protons there are in the nucleus, the more positively charged the nucleus is and the stringer the nuclear attraction for the electrons

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5
Q

What’s atomic radius

A

Attraction decreases with increasing distance
An electron close to the nucleus will be much more strongly attracted than one further away

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6
Q

What’s shielding

A

As the number of electrons between the outer electrons and the nucleus increases the outer electrons feel less nuclear attraction towards the nuclear charge

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7
Q

Acronym to explain nuclear attraction

A

Nan cars

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8
Q

Why does ionisation energy decrease down the group 1 and 2

A

Number of shells increases so shielding increases
Atomic radius increases down the group
Nuclear charge increases down the group but is outweighed by increase in atomic radius and shielding
Nuclear attraction for outer electron decreases so electrons lost more easily

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9
Q

Explain trend in ionisation energies going across the period

A

nuclear charge increases
Shielding is similar, outer most electron is being removed from the same shell
Atomic radius decreases
Greater nuclear attraction for outer electron

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10
Q

Between group 2 and 3 the ionisation energy decreases explain why using aluminium and magnesium

A

mg has outer electron in 3s subshell and Al has outer electron In 3p subshell
3p subshell has higher energy than 3s subshell
There is a slight increase in shielding which outweighs the increase In nuclear charge

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11
Q

Between groups 5 and 6 the ionisation energy decreases explain why using sulfur and phosphorus and give the electron cofiguration of s and p atoms

A

shielding is identical in phosphurus and sulfur atoms and electron is being removed from 3p subshell.
In sulfur, the electron is easier to remove- lower ionisation energy
orbital has 2 electrons- repulsion between 2 electrons

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12
Q

What’s meant by the second ionisation energy

A

energyneeded to remove 1 electron from each ion in 1 mol of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

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13
Q

Explain the trend in successive ionisation energies

A

ionisationenergy increases as more electrons are removed
electrons are removed from a more positive ion
Attraction between remaining electrons and nucelus increases. Radius decreases
more energy is needed to remove the next electron

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14
Q

What charged ions do group 2 metals from and where do they lose the electrons from

A

2+ ions and lose electrons from s2

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15
Q

What does a group 2 metal form with water and 9bservstions

A

Metal Hydroxide and hydrogen
Fizzing and solid dissolves

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16
Q

State the reaction between a group 2 metal and acid (hcl)

A

Mg + 2HCL -> MCl2 +H2
Forms salt and hydrogen solid dissolves and fizzing

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17
Q

Describe and explain the trend in group 2 metal reactivity going down the group

A

Increases down the group

number of shells increases so shielding increases
atomic radius increases down the group
nuclear charge increases down the group but is outweighed by the increase In atomic radius and shielding
nuclear attraction for outer electron decreases so less energy is required to remove electrons

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18
Q

Group 2 metals react with oxygen to form group 2 metal oxides, give the appearance of them

A

White powder

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19
Q

Metal oxides increase In reactivity with water down the group what’s the affect on solubility and hence pH going down the group

A

Metal oxides increase in reactivity with water down the group
The resulting metal hydroxide is more soluble.
Stronger alkaline solution and higher pH

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20
Q

State and explain 3 uses of group 2 compounds

A

Ca(OH)2 - used in agriculture no neutralise acidic soils
Mg(OH)2 and CaCO3 - used in indigestion tablets antacids, neutralise excess stomach acid
BaSO4- barium meals as BaSO4 is insoluble

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21
Q

Why do successive ionisation energies increase with ionisation number

A

the radius decreases and the attraction between the electrons and the nucleus increases

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22
Q

how do you calc uncertainty of pipette and burette

A

pipette you dodnt times the top mum by 2
burette you do and you divide by volume of mean titre

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23
Q

whats periodicity

A

Repeating pattern in chemical and physical properties

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24
Q

Describe how atoms are arranged in the periodic table

A

The periodic table is arranged in elements in increasing atomic number

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25
state and explain the trend in atomic radius across the period
Decreases across a period Number of protons and nuclear charge increases Outer electrons are in the same shell and have similar shielding Stronger nuclear attraction and electrons are drawn in more
26
state the trend in atomic radius down the group
Increases down the group Number of shells increases
27
Q is about noble gases. explain the mp and bp down group 8
Melting and boiling points - Increases down group 8 Number of electrons and atomic radius increases Strength of induced dipole-dipole interactions between atoms increases which require more energy to overcome
28
Describe mp and bp across period 3
Na to Al increases as number of delocalised electrons increases. charge on positive ion increases. metallic bonding gets stronger from al to si, increases to si, boiling point very high for Si as it has giant covalent lattice structure with lots of strong covalent bonds from p4 to ar decreases to p4 as it has molecular weak induced dipole dipole interactions increases to s8 nore atoms more electrons so more stronger induced dipole dipole interactions which require more energy to overcome decreases to ar as this has simple atomic structure and has weaker induced dipole dipole interactions which require less energy to overcome
29
are halogens oxidising or reducing agents
halogens are oxidising agents as theyre reduced and part of redox reaction
30
whats the colour of iodine and state
Grey - gas
31
whats the colour of fluorine and state
pale yellow- gas
32
whats the colour and state of bromine
brown liquid
33
Describe the Trend in reactivity for halogens going down the group
Halogens are less reactive and less oxidising down the group. Atomic radius increases. Outer electrons are further away from the nucleus and are more shielded from the nucleus. Less nuclear attraction, so more difficult for larger atoms to attract electrons needed to form an ion
34
describe a reaction that shows the oxidising strengths of halogens and explain how this works
oxidising strengths seen in displacement reactions with halide ions more reactive halogens will oxidise and displace halide ions of less reactive halogens
35
displacement reaction of chlorine with bromide (bromide doesnt exist on own, but other ion spectator ion not involved) and displacement reactions chlorine with iodine
Cl more reactive so displacement takes place aq- pale green to orange organic solvent- pale green to orange With iodine chlorine more reactive displacing aq- col change pale green to brown organic solvent - pale green to violet no visible change for both. if added to chlorine
36
describe reaction taking place when bromine reactd with bromine and bromine with iodide
no visible change for first for second, bromine more reactive displacing col change- aq orange to brown organic solvent - orange to violet
37
describe what happens in displacement reactions iodide reacting with bromine and iodine
no displacement reaction and no visible change
38
what are the observations from displacement reactions
we can shake a reaction mixture with an organic solvent, ie cyclohexane halogen dissolves in the organic solvent- solvent layer- separated using funnel
39
describe how the organic layer containing the halogen can be separated from the solution
using separating funnel
40
Describe how to test for THE PRESENCE OF halide ions
1. Dissolve unknown halide in water with nitric acid 2. Aqueous solution of silver nitrate added 3. Silver ions from the silver nitrate reacts with the halide ions forming a silver halide precipitate 4. Silver halide precipitates are coloured with the colour intensity increasing from Cl to I Colour tells you which halide ion is present. V 5. If it is too difficult to decide on a colour, add ammonia solution, NHa) as the halide ions have different solubilities in ammonia.
41
what do we form, observe and solubility of silver reacted with chloride bromide and iodide
chloride - forms silver chloride. white precipitate. soluble in dilute NH3 bromide - silver bromide, cream precipitate, soluble in concentrated NH3 Iodide - silver iodide, yellow precipitate, insoluble in concentrated NH3
42
whats an example of precipitation reactions
aqueous solution where aqueous ions react together to form a solid precipitate
43
what to we add to reactions to remove any carbonates and whats produced
dilute HNO3 can be added to remove any carbonates which could produce white solid
44
what can we use to produce white solid
nitric acid as h2so4 and hcl produce white solid (false positive ag2so4 and agcl)
45
Whats disproportionation
the simultaneous oxidation and reduction of the same element. all halogens except fluorine can exist in oxidation states other than 0 and -1 oxidation states
46
conditions for halogens to undergo disproportionation/ synthesise bleach
when they react with cold dilute alkali solutions eg NaOH or KOH
47
How do we make bleach
bubble chlorine gas with cold, dilute aqueous NaOH makes sodium chlorate (I) solution equation: Cl2(g) +2NaOH(aq) -> NaCL(aq)+ NaCLO(aq)+ H20(l) ionic eq- Cl2(g) +2OH-(aq) -> Cl-+CLO-(aq)+H20
48
some uses of bleach
water treatment, bleaching paper and textiles and cleaning toilets
49
describe reaction of chlorine and water
chlorine can used to to treat water. disproportionate reaction are hydrochloric acid and chloric (l) acid Cl2(g) + H20(l) -> HCl(aq) + HCLO(aq) chloric acid ionises to make chlorate ions HCLO(aq) +H20 -> ClO-(aq) + H30+
50
2 advantages of chlorination of water
Chlorine added to drinking water to kill bacteria prevents growth of algae and prevents discolouration caused by organic compounds
51
state 3 disadvantages of chlorination of water
chlorine is toxic- irritates the respiratory system chlorine can react with organic matter to form carcinogenic chlorinated hydrocarbons liquid chlorine causes chemical burns on skins/eyes
52
ethical concerns about chlorination of water
risk from untreated water outweigh the ridk to cancer form of mass medication - we have no choice but to
53
alternatives to chlorination of water adv and diasdb
ozone(O3) adv - strong oxidising agent, which can kill microorganisms. diasdv - expensive to produce and its short half life means the treatment isnt permanent ultraviolet light- adv- kills microorganisms disadv- ineffective in cloudy water and is only effect short term
54
test for carbonates
Ad md dilute HNO3, observation is fizzing due to production of CO2. equation: CO3 2-(aq)+ 2H(aq) -> CO2(g) + H20(l)
55
test for sulfate ions
add Ba(NO3)2(aq) observation- insoluble white precipitate BaSO4(s) SO4 2- (aq)+ Ba2+ (aq)-> BaSO4(s)
56
test for NH4 ions
heating with NaOH(aq) BUT NOT AMMONIA observation- NH3, litmus paper turns blue equation: NH4+(aq) + OH-(aq) -> NH3(g) + H20(l)
57
whats the test order to avoid false positive results
carbonates, sulfates then halides
58
which element has induced dipole dipole interactions in its solid lattice?
sulfur
59
why do successive ionisation energies increase with ionisation number 1 mark
radius decreases between between electrons and nucleus increases
60
whats the shape around the carbon atoms in graphene
trigonal planar
61
what type of reaction makes calcium carbonate go to calcium oxide
thermal decomposition
62
how can you determine in which group of the periodic table an element would be found when given an ionisation energy table of the elements
Large increase between the _nd and _rd ionisation energy values
63
what happens to ionisation energy down a group and across the period
ionisation energy decreases down the group and increases across the period.
64
explain why magnesium has a higher melting point than sodium.
Magnesium has more outer electrons. Magnesium ions have a greater charge Magnesium has a greater attraction between ions and delocalised electrons.
65
why are silicon chlorine and oxygen all classed as p block elements?
Highest energy electron in a P subshell
66
Explain how the successive ionisation energies in a table provide evidence for the electron shells in sodium atoms
large increase shows a new shell large increase between - ionisation energy and -- ionisation energy and -----
67
what would we observe if we reacted less reactive metal with water
Less fizzing solid dissolves slowly
68
What's Enthalpy change
The heat energy transferred in a reaction at constant pressure unit -kJ mol-1
69
What are standard states? and the standard conditions?
Physical states they are in under standard conditions Standard conditions ΔH⦵ : 100kPa and 298K
70
Standard enthalpy change of reaction ΔHr ⦵
The enthalpy change when a reaction occurs in the molar quantities shown in the chemical equation with all reactants and products in their standard states, under standard conditions. eg: CO (g) + NO2 (g) → CO2 (g) + NO (g
71
Standard enthalpy change of neutralisation ΔHneut⦵
The enthalpy change when solutions of an acid and an alkali react together to form 1 mole of water, under standard conditions. ✓ Eg; 0.5 H2SO4 (aq) + NaOH (aq) → 0.5 Na2SO4 (aq) + H2O (l) – negative ΔH
72
Standard enthalpy change of combustion ΔHc⦵
Is the enthalpy change when 1 mole of a substance is completely burned in oxygen, under standard conditions. ✓ Eg C2H5OH (l) + 3O2 (g) → 2CO2 (g) + 3H2O (l) – negative ΔH
73
Standard enthalpy change of formation ΔHf ⦵
The enthalpy change when 1 mole of a compound is formed from its elements in their standard states, under standard conditions. Eg: 2 C(s) + 3 H2 (g) + 0.5 O2 (g) → C2H5OH (l) – usually negative ΔH
74
Activation Energy
The minimum amount of energy needed for a reaction to start.
75
Exothermic reactions
Reaction that releases energy with a negative ∆H Example: CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔcH⦵ = -890 kJ mol-1
76
Endothermic reactions
Reaction that absorbs energy with a positive ∆H Eg: CaCO3(s) → CaO(s) + CO2(g) ΔrH⦵ = +179 kJ mol-1 to help think end on a positive.
77
look at the graphs of an endothermic and exothermic reaction
78
when looking at endo and exothermic graphs in which direction does the activation energy point and why, and enthalpy change too.
(activation energy) Ea arrows always point up, as it is an endothermic process, whereas ΔH can point up or down as enthalpy change can be either endothermic or exother
79
Calculating enthalpy changes
can be calculated directly using experiments. ✓ However, it is not always possible to measure the enthalpy change directly from experiment for the following reasons: - The activation energy of the reaction is too high. ✓ - There is a slow reaction rate. ✓ - More than one reaction may be taking place (and so more than one product is formed)
80
How can We calculate enthalpy change indirectly
1. Average bond enthalpy data ✓ 2. Standard enthalpy changes of combustion or formation data ✓
81
Experiments to calculate enthalpy change
To measure an enthalpy change, we need to find: 1. Number of moles of what's reacting, n. ✓ 2. Change in temperature, ΔT. ✓ 3. Measure the mass of the solution/ water to be heated. ✓ Method: 1. Measure the mass of the solution/ water to be heated. ✓ 2. Record the temperature change of the solution/ water. ✓ Note: calculations are done on the limiting reactant.
82
Calculating enthalpy change directly through experiments equation
1. Calculate the energy released or absorbed Q= mcΔT Q= energy released, m- mass, c- specific heat capacity ΔT temp change 2. Calculate the number of moles of the substance reacting (limiting reagent if more than one 3. Calculate the enthalpy change (you need to convert Q into kJ) ∆H = Q/n
83
Reasons why standard enthalpy changes of combustion determined experimentally are less exothermic than the calculated theoretical values. ✓
1. Heat released to the surroundings. ✓ 2. Incomplete combustion OR incomplete reaction (depending on experiment). ✓ 3. Non-standard conditions ✓
84
Experiments to calculate enthalpy change for combustion reactions
Calculating enthalpy changes of solids reacting with solutions m to calculate energy absorbed/ released = total mass of solution. i.e. 100 cm3 of NaOH with the assumption that the total mass does not change, m = 100 g To calculate amount of substance in moles: take m of limiting reagent to calculate mol of limiting reagent. i.e. 100 cm3 of 0.1 mol dm3 NaOH, nNaOH = 0.1 x 0.1 = 0.01 mol. Calculating enthalpy changes of neutralisation m to calculate energy absorbed/ released = total mass of solution. i.e. 100 cm3 of NaOH and 100 cm-3 of HCl, m = 200 g To calculate amount of substance in moles: take c and v to calculate n of limiting reagent, then find amount of water produced, in mol, using stoichiometry. i.e. 100 cm3 of 0.1 mol dm3 NaOH, nNaOH = 0.1 x 0.1 = 0.01 mol. nWater = 0.01 x 1 = 0.01 mol. Calculating enthalpy changes of combustion m to calculate energy absorbed/ released = mass of water heated. i.e. 100 cm3 of H2O , m = 100 g as water has a density of 1 g cm-3 . To calculate amount of substance in moles: take change in mass by minitial – mfinal of fuel. i.e. 1.2g of butan-1-ol, nbutan-1-ol = 1.2/ 74 = 0.0189 mol.
85
Assumptions made when measuring the temperature of solutions rather than water
density of solution is 1 g cm-3 / same as water. ✓ c of solution is same as water. ✓ no water lost/evaporated.
86
Average bond enthalpies
Average enthalpy change when one mole of gaseous covalent bonds is broken Bond enthalpies can have positive values – Bond breaking is endothermic. ✓ Bond enthalpies can have negative values – Bond forming is exothermic.
87
Determining whether a reaction is endothermic or exothermic
Some reactions can be endothermic ✓ More energy is required for breaking bonds. ✓ than is released by forming bonds. Some reactions can be exothermic ✓ More energy is released by forming bonds. ✓ than energy required when breaking bonds. ✓
88
Limitations with using bond enthalpies
Standard enthalpy changes will be different from that calculated using average bond enthalpies 1. Actual bond enthalpies may be different from average values. ✓ 2. Conditions are not standard.
89
Calculating enthalpy change using average bond enthalpies BR–P
ΔH = Total energy absorbed (Total bonds broken) – Total energy released (Total bonds formed) Example: Calculate the enthalpy change for the following reaction: N2 + 3 H2 → 2 NH3 ∆H = [(1 x 945) + (3 x 436)] – (6 x 391) ∆H = 2253 – 2346 = -93 kJ mol-1
90
Hess’s Law
The total enthalpy change of a reaction is always the same, no matter which route is taken. note; when going along the arrow we add values when going against the arrow we subtract values
91

A Level Chemm (3 decks)

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