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mod 6 Flashcards

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1
Q

how to calculate the concentration

A
  • find the concentration of standard solution
  • to find moles of it reacted –> standard solution x volume aliquots
  • find moles of analyte reacted with equation
  • find concentration of analyte reacted = moles/volume of titre
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2
Q

everyday life: preventing tooth decay

A
  • acids from foods accelerate tooth decay
  • hence toothpaste is weakly alkaline undergoing neutralisation
  • contains bases (e.g CaCO3, Al2O3, MgCO3)
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3
Q

Arrhenius’ theory is that

A
  • acids in aqueous solution ionise to form hydrogen ions
  • bases in aqueous solution ionise to form hydroxide ions
  • hence neutralisation forms water from hydrogen and hydroxide
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4
Q

what is pH

A

pH scale is based on the concentration of hydronium ions defined as:
pH = -log10[H30+]
[H3O+] = 10^-pH
- the number of sigfigs of the [H3O+] = no. of decimal places for pH

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5
Q

what is a conductivity

A

is the tendency of a material to conduct electricity
- proportional to proportional to [ion]
- larger ions are less conductive

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6
Q

Indicators are

A

substances which change colour based on the pH of the environment
- can determine acid or base and extend based on transition ranges
- are weak acids and bases

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7
Q

how to select an appropriate indicator

A

depends on the salt’s acidity
pH –> analyte + primary standard –> water + salt
- if acidic –> methyl orange
- if neutral –> bromothymol blue
- if basic –> phenolphthalein

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8
Q

equivalence point definition

A

the point when n(H3O) reacted = n(OH) reacted

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9
Q

relationship between pH and pOH

A

pH + pOH = 14

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10
Q

poor primary standards

A
  • HCl, HNO3 are not stable –> can be used as secondary standards titrated twice
  • NaOH, KOH are deliquescent (absorbs water from air to form solution) and reacts with CO2
  • H2SO4 is hygroscopic (absorbs water from air)
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11
Q

what is enthalpy of neutralisation

A

enthalpy change is associated with neutralisation and is the enthalpy change per mole of water formed (exo)
- weak acids have a less negative enthalpy of neutralisation

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12
Q

popular indicators and colour of acidic, transition and basic ranges

A
  1. methyl orange: red, orange, yellow
  2. bromothymol blue: yellow, green, blue
  3. litmus: red, purple, blue
  4. phenolphthalein: colourless, pale pink, pink/magenta
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13
Q

Acids are

A

compounds which form hydrogen ions in solution

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14
Q

titrant def

A

the solution used to determine the concentration of an unknown solution

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15
Q

how to perform a titration

A
  1. transfer 25mL of standard solution into a conical flask using pipette
  2. add few drops of appropriate indicator
  3. fill burette with solution below 0mL mark
  4. mount burette so the tip is just inside the neck of conical flask and read off volume
  5. let the solution into the flask drop wise until colour change lasts longer than 15 seconds and read off volume
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16
Q

what is titration

A

titration is a technique by which the concentration of a solution is determined by measuring the volumes of the solutions in the reactions
- neutralisation reactions are used to determine the concentration of an acid or base

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17
Q

how to calculate enthalpy of neutralisation

A
  • use -q/n(water)
  • q=mc(change in T)
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18
Q

Arrhenius theory advantages

A
  • works for many acids
  • base is defined
  • explains neutralisation
  • can explain differences between strong and weak acids
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19
Q

pharmacy: acetylsalicylic acid (ASA) in aspirin

A
  • ASA is a weak acid hence must be titrated with a strong base (NaOH)
  • NaOH must be used as a secondary standard
  • ASA is a powder so back-titration must be used
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20
Q

what is a conductivity graph

A

measures conductivity vs volume of titrant added

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21
Q

examples of buffers: human blood

A
  1. human blood
    - pH of blood is regulated to 7.35-7.45 by many buffers to maintain homeostasis (body’s eq)
    - CO2 dissolves in blood to form H2CO3 (acid)
    - H2CO3 converts into HCO3 (base)
    - when we exercise [H3O] increases when we hyperventilate [OH] increases
    EQUATION:
  2. H2CO3 + H2O –><– HCO3 + H3O
  3. HCO3 + H2O –><– H2CO3 + OH
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22
Q

bases properties

A
  • bitter taste
  • soapy feel in solution
  • caustic
  • conductive in solution
  • red litmus blue
  • ph>7
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23
Q

pros of pH probe/cons of indicators

A
  • indicators have broad pH range vs probe’s specific pH reading
  • indicators very variable due to colour interpretation vs probe’s highly reproducible measurements
  • indicator’s are destructive and can alter the pH of solution vs probe’s are not destructive
  • indicator’s can’t be used on coloured solutions vs probe’s can be
  • indicators are cheap and portable vs probe’s are expensive and less portable and need high maintenance
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24
Q

Amphoteric meaning

A

substances that act as both base and acid

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25
Q

acid strength Ka and pKa:1
1. strong
2. weak
3. weaker

A
  1. large Ka, negative pKa
  2. small Ka, small positive
  3. smaller Ka, large positive
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26
Q

memorise titration graphs

A

memorise titration curves
- strong acid-strong base
- strong acid-weak base
- weak acid-strong base
- weak acid-weak base

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27
Q

what is the acid dissociation constant (Ka)

A

The acid dissociation constant is a qualitative scale to differentiate between weak acids (<100% ionisation).
Ka = [H3O][A]/[HA]
HA + H2O –><– H3O + A

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28
Q

Weak acid definition

A

<100% ionisation in water (equilibruim arrow)

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29
Q

naming bases

A
  • element and hydroxide at end
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30
Q

salt acidity of:

A
  • strong acid + strong base = neutral (spectator + almost neutral base)
  • strong acid + weak base = acidic
    (almost neutral base + weak acid)
  • weak acid + weak base = depends
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31
Q

how to explain buffer questions

A
  1. define
  2. write 2 dissociation equations
  3. use LCP to show what happens when a base is added
  4. use LCP to show what happens when an acid is added
  5. therefore pH doesn’t change
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32
Q

Bronsted-Lowry theory of acids

A
  • defines acids as substance which donates protons (H+)
  • defines bases as substances which accepts protons (H+)
  • neutralisation is an exothermic proton transfer of protons from acid to base
33
Q

naming acids

A

(inorganic acids)
- drop the ‘ide’ and add ‘ic acid’
- cations prefix
(oxyanions with oxygen)
- drop ‘ate’ and add ‘ic acid’
- don’t say hydrogen
(oxyanions with reduced state or one less oxygen)
- drop ‘ate’ add ‘urous acid’
- don’t say hydrogen

34
Q

popular indicators and transition ranges

A
  1. methyl orange: TR 3.1-4.4
  2. bromothymol blue: TR 6.0-7.6
  3. litmus: TR 5.5-8.0
  4. phenolphthalein: TR 8.3-10.0
35
Q

water self ionises to form:

A

2H2O –><– H3O+ + OH- (endo)
- equilibrium lies heavily left
- the self-ionisation constant of water (Kw) is [H3O+][OH-] = 1.0x10^-14
- [H3O+] = [OH-] = 1.0x10^-7 in water

36
Q

titration error def

A

the difference between EP and endpoint

37
Q

titre def

A

the minimum volume required to reach the endpoint of neutralisation reaction in titration

38
Q

how to select a primary standard (primary standard prepares a standard solution with known concentration hence must):

A
  • be water soluble
  • have high purity
  • have a definite chemical composition
  • practical
  • stable in air
  • primary standard should be acid is analyte is base and vise versa
  • avoid weak-weak titrations
39
Q

Indigenous methods using neutralisation

A
  1. using clays to neutralise stomach acid
    - many ochres and clays contain hydroxides of transition metals (e.g yellow ochre is FeO(OH).H2O)
    - indigenous people ingest the clay so the hydroxides neutralise stomach acids
    (OH + HCl –> H2O + Cl)
  2. using leaves to neutralise insect stings
    - used alkaloids with basic nitrogenous organic molecules in leaves to neutralise formic acid in sting bites
40
Q

Strong acid definition

A

100% ionisation in water (one directional arrow to completion)

41
Q

Kb properties

A
  • when Kb increases, lies right, high DOI
  • when Ka decreases, lies left, low DOI
  • pKb = -log10(Kb)
  • Kb = 10^-pKb=b
42
Q

weak polyprotic titration graphs

A

have many EP’s to match protic
- except sulfuric acid only has one strong EP at second ionisation

43
Q

analyte def

A

the solution whose concentration is to be determined

44
Q

steps to titration

A
  1. select a primary standard
  2. prepare a standard solution
  3. select an appropriate indicator
  4. rinse the glasses
  5. perform titration
45
Q

Amphiphrotic meaning

A

substances that act as both base and acid due to it’s ability to donate or accept in different environments
(a type of amphoteric)

46
Q

degree of ionisation meaning

A

is how much a substance ionises
= [H30+]/[HA] x 100

47
Q

Ka properties

A
  • when Ka increases, lies right, high DOI
  • when Ka decreases, lies left, low DOI
  • pKa = -log10(Ka)
  • Ka = 10^-pKa
48
Q

what is neutralisation

A

is a special type of reaction that doesn’t produce OH or H

49
Q

given any 2 of pKa, pH or [HA] you can get the other

A

practise 3 types of equations

50
Q

endpoint def

A

the point a sustained colour change is achieved using an appropriate indicator

51
Q

memorise conductivity graphs

A

memorise conductivity graphs
- strong acid - strong base
- strong acid/base (conical) - weak base/acid (burette)
- weak acid/base (conical) - strong base/acid (burette)
- weak acid - weak base

52
Q

strong acids in water

A

HCl
HBr
HI
HNO3
H2SO4

53
Q

standard solution def

A

a solution with known concentration

54
Q

good primary standards

A
  • hydrated oxalic acid is a weak acid for a strong base
  • anhydrous sodium carbonate or sodium hydrogen carbonate are weak bases for strong acids
55
Q

effects of disturbances to weak acids ionisations (DI)

A

HA + H2O –><– A- + H3O+ (endo)
- adding A- would shift left and decrease DI = [H3O+]/[HA]
- adding H3O+ would shift left and decrease DI = [A]/[HA]
- dilution would shift right increase DI
- increase temperature will shift right increase DI and increase Ka

56
Q

how to calculate pH of dilute solutions

A
  1. write auto-ionisation equation of water
  2. use an ICE table add the conc of H3O and OH to waters (10^-7) for initial
  3. solve equation by making change x
57
Q

Base Dissociation constant

A

Base Dissociation constant is similar to the Ka is a qualitative scale to differentiate between weak bases (<100% ionisation)
Kb = [OH][BH]/[B]
B + H2O –><– BH + OH

58
Q

how buffer’s work

A

a buffer solution is formed when similar amounts of a weak acid and it’s conjugate base are mixed
1. HA + H2O –><– BOH + H3O
2. BOH + H2O –><– HA + OH

59
Q

calculating pOH

A

pOH = -log10[OH-]
[OH-] = 10^-pOH

60
Q

acids properties

A
  • sour taste
  • corrosive
  • conductive in solution
  • blue litmus red
  • pH<7
61
Q

how to prepare the standard solution

A
  • dry solid in desiccator and weigh
  • dissolve with demin water
  • transfer to volumetric flask and dilute until meniscus touches gradation line
  • put stopper on and invert 10 times
62
Q

proticity is

A
  • classified by the number of protons a substance can donate
  • monoprotic = 1 H+ to donate (e.g HCl)
  • acids with more than one proton to donate ionise in steps (e.g H2SO4)
63
Q

acid + base –>
acid + carbonate –>
acid + metal –>

A

salt + water
salt + water + CO2 (limewater test)
salt + H2 (pop test)

64
Q

factors that affect pH

A
  1. concentration
    - higher [H3O+] = lower pH
  2. strength
    - strong acids pH < weak aicds pH
    - due to different DI
  3. proticity (SA)
    - strong acids with higher proticity pH < strong acids with lower proticity pH
  4. degree of ionisation (WA)
    - weak acids with higher DOI pH < weak acids with lower DOI pH
    - citric acid DOI > acetic acid DOI
65
Q

what is titration curve

A

a titration curve plots the pH of the reaction mixture in the conical flask against the volume of titrant added
- has to have an identifiable EP POI
- shape depends on if the base is in burette or flask and flip graph sideways for vise versa

66
Q

relationship between Ka and Kb of conjugate pairs

A

Ka x Kb = Kw = 1.0 x 10^-14
pKa + pKb = 14

67
Q

how to rinse glassware correctly

A
  • rinse volumetric flask and conical flask with demin
  • rinse pipette and burette with solution it will be containing
68
Q

relationship between pH and pOH

A

pH + pOH = 14

69
Q

weak acids in water

A

HF
CH3COOH
H2CO3
H3PO4
C6H8O7

70
Q

Bronsted-Lowry theory advantages

A
  • explains the behaviour of acids, bases in non-aqueous solutions
  • considers the role of solvent in determining strength
  • explains how some species acts as both acids and bases (amphiphrotic)
  • an acid and base can only be defined in relation to each other
71
Q

what is back-titration

A

back-titration is used to find the concentration of substances when direct titration can’t be used i.e
- titrating a solid
- insoluble acid/base
- gaseous acid/base

72
Q

what happens when you add an acid or base to the buffer solution

A
  1. adding an acid
    - shift equilibrium left to minimise changes (show in eq 1.)
    - hence pH is stable
  2. adding a base
    - shift equilibrium right to minimise change (show in eq 2.)
    - hence pH is stable
73
Q

industrial: acid and base spills

A

neutralisation reaction are used to neutralise spills that could harm the environment
factors to consider:
- acid or base
- extent of spill
- strength or concentration
characterisation of of neutralising agents
- cheap, practical
- weak
- solid powers
- amphiprotic if spilt substance is unknown (e.g sodium hydrogen carbonate)

74
Q

how to back-titrate

A
  1. add excess known strong acid to react with base
  2. titrate excess acid with standard solution of base
  3. find moles of excess
  4. find moles of strong acid originally reacted = n(strong acid) total - n(strong acid) excess
  5. write chemical equation
  6. using molar ratio to find unknown base concentration
75
Q

how to show an amphiphrotic substance?

A

write 2 equations
- one with it reacting with a H3O
- one with it reacting with a OH

76
Q

Bases are

A

compounds which form hydroxide ions in solution
- alkali’s are water soluble bases

77
Q

what is a conjugate acid/base pair

A

when an acid donates a proton to a base, it is it’s conjugate base and vise-versa
- weak acid = weak base
- weak base = weak acid
- strong acid = almost neutral base
- strong base = almost neutral acid

78
Q

Arrhenius theory disadvantages

A
  • doesn’t recognise solvent’s role in acids strength
  • doesn’t account for ALL acids or bases
  • can’t explain salt acidity
  • cant’ explain neutralisation in gaseous or solid forms
79
Q

what is a buffer

A

a buffer is composed of a weak acid/base and it’s conjugate with similar concentrations which resists pH changes

chemistry yr 12 (8 decks)

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