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wace chem > acids + bases > Flashcards

acids + bases Flashcards

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1
Q

Explain 3 characteristics of an acidic solution

A
  • conduct an electric current
  • turn litmus red and Universal Indicator red, orange and/or yellow
  • taste sour
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2
Q

Is HBr strong or weak

A

strong

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3
Q

Is HF strong or weak

A

weak

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4
Q

Is H2SO3 strong or weak

A

weak

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5
Q

is HI strong or weak

A

strong

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6
Q

Is HNO3 strong or weak

A

strong

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7
Q

Are Na2O, K2O and CaO strong or weak bases

A

strong

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8
Q

4 Characteristics of Bases

A
  • also conduct an electric current
  • turn litmus blue and universal indicator blue, indigo and/or violet
  • taste bitter
  • have a slippery feel
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9
Q

Observation of H2S

A

colourless, rotten egg odour

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10
Q

Observation of SO2

A

colourless, pungent

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11
Q

Observation of NH3

A

colourless, pungent

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12
Q

Observation of NO2

A

brown, pungent

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13
Q

Observation of Cl2

A

greenish-yellow, pungent

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14
Q

What does an acid and metal sulfite produce (HCl + Na2SO3)

A

Salt + H₂O + SO₂

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15
Q

What does a base and a non-metal oxide produce (SO2 + KOH)

A

Salt + H₂O (K2SO3 + H2O)

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16
Q

Explain Davy Theory for Acids and Bases

A

Acids: Have replaceable H (hydrogen could be partly or totally replaced by metals)

Hydrochloric acid - 2HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)

Bases: React with acids to form salt + water

Sodium Hydroxide - NaOH(s) + HNO3(aq) →NaNO3(aq) + H2O(l)

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17
Q

Explain Arrhenius Theory for Acids, Bases, Neutralisation

A

Acids: Have H in their formula and produce hydrogen ions (H+) when dissolved in water

Hydrochloric acid: HCl(g) → Cl ion + H ion

Bases: Have OH in their formula and produce hydroxide ions (OH-) when dissolved in water

Sodium Hydroxide: NaOH(s) → Na ion + OH ion

Neutralisation: An acid plus a base produces a neutral solution

H ion + OH ion → H2O

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18
Q

What are 3 problems with Arrhenius Theory

A
  • Some bases produce OH ions in solution and yet do not have OH in their formula (NH3, BaO)
  • Restricted to aqueous solutions
  • Not all salts are neutral
  • Doesn’t allow for the existence of hydronium ions (H3O+)
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19
Q

Explain Bronsted-Lowry Theory for acids, bases, neutralisation

A

Acids: are proton (H ion) donors

  • Strong acid: HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)
    • HCl is donating a proton and acting as an acid
    • H2O is accepting a proton and acting as a base

Bases: are proton (H+) acceptors

  • Ammonia: NH3(g) + H2O(l) ⇌ NH4+(aq) + OH-(aq)
    • NH3 is accepting a proton and acting as a base
    • The H2O is donating a proton and acting as an acid

Neutralisation: reaction between a proton donor and a proton acceptor

  • Carbonate ion plus water: CO3(2-) + H2O ⇌ HCO3- + OH-
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20
Q

How is a conjugate base formed

A

Once an acid has donated a proton it has the potential to act as a base

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21
Q

How is a conjugate acid formed

A

Once a base has accepted a proton it has the potential to act as an acid

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22
Q

What is the conjugate acid of H2O

A

H3O+

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23
Q

What is the conjugate base of HNO2

A

NO2-

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24
Q

What is the trend with acids and their conjugate base

A

Stronger the acid, the weaker its conjugate base

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25
Q

What does the Ka value measure

A

The extent to which an acid ionises in aqueous solutions
It is a measure to which the proton transfer goes to completion

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26
Q

What do large Ka values suggest

A

the greater the tendency of an acid to donate protons to the water
Larger Ka Value → Stronger the acid is → Greater the degree of its ionisation

27
Q

Monoprotic meaning

A

can donate only one proton (hydrogen ion) per molecule

28
Q

Diprotic meaning

A

can donate only two protons (hydrogen ion) per molecule

29
Q

Explain successive ionisation

A

process of becoming ions over multiple steps

30
Q

Polyprotic meaning

A

acids which contain two or more acidic protons

31
Q

Are pure metal cations acidic, basic or neutral (Fe3+ , Al3+ , Cr3+ , Sn4+)

A

Acidic

32
Q

Self ionisation of water equation

A

2H2O ⇌ H3O+ + OH-

33
Q

WHat is pure water at any temperature?

A

Neutral

34
Q

What is pH

A

a number indicating how acidic or basic a solution is

35
Q

Definition of a titration

A

Described as an analytical procedure often used to determine the concentration of a solution or the amount of a particular substance present

36
Q

Purpose of a pipette

A

Accurately deliver a known volume (aliquot) of liquid

37
Q

Purpose of a burette

A

Used to accurately deliver a variable volume (titre) of liquid

38
Q

Purpose of a Volumetric Flask

A

Used to prepare an accurately known volume of solution

39
Q

Purpose of a Conical Flask

A

Used to hold the solutions during the titration

40
Q

Purpose of an Analytical Balance

A

Used to weigh out accurately known mass

41
Q

Purpose of an indicator

A

Used to determine when stoichiometrically equivalent amounts of two reactants have reacted - equivalency point

42
Q

What makes an indicator useful

A

end point should match equivalency point

43
Q

Endpoint vs Equivalence point

A

Endpoint

  • Definition: The endpoint is the point in a titration where a visual indicator signals the completion of the reaction. It’s determined by a physical change, often a colour change

Equivalence Point

  • Definition: The equivalence point is the theoretical point in a titration where the equivalents of the acid and base are chemically equal. It’s the stoichiometric point of the reaction. Determined by pH or a back titration
44
Q

What are standard solutions

A

accurately known concentration used to determine concentration of a secondary standard solution

45
Q

5 Characteristics of a standard solution

A
  • Having a known high degree of purity
  • The reactions it is involved in are known
  • Must be stable - not react with air, O2, CO2 and other gases
  • Relatively high molar mass to minimize weighting errors
  • Must not absorb water
    • Not hygroscopic - absorbs water from the atmosphere
    • Not deliquescent - absorbs surrounding water
46
Q

Common primary standards examples

A

anhydrous sodium carbonate, oxalic acid, potassium hydrogenphthalate

47
Q

Methyl Orange pH range, colours in acids and bases, and reactions used for

A
  • 3.1-4.4
  • Red colour in acid
  • Yellow colour in base
    • Strong Acid / Weak Base
48
Q

Phenolphthalein pH range, colours in acids and bases, and reactions used for

A
  • 8.3-10.0
  • Colourless in acid
  • Pink colour in base
    • Strong Acid / Strong Base
    • Weak Acid / Strong Base
49
Q

Explain why we use Phenolphthalein for strong acid / strong base even though the range is different

A

The change in pH is very rapid around the equivalence point so in a strong acid /strong base titration the pH will quickly cross from around 5 to around 9 just in a drop or two, so that means that phenolphthalein changes colour close enough to 7 to work.

50
Q

Write an equation for the reaction between acetic acid and sodium hydroxide in aqueous solution? Explain products and what indicator should be used

A

CH₃COOH + NaOH ⟶ NaCH₃COO + H₂O

CH₃COO- + H₂O ⇌ CH₃COOH + OH-
- Acetate produces hydroxide ions so the equivalence point will be basic [OH-] > [H₃O+]
- Phenolphthalein should be used as it changes colour over a range of basic pH (8.3-10) so the end point will match the equivalence point

51
Q

Write an equation for the reaction between sodium hydrogen carbonate is added to nitric acid in aqueous solution? Explain products and what indicator should be used

A

HCO3- + H+⟶ CO₂ + H₂O

CO₂ + H₂O ⇌ H₂CO₃

H₂CO₃ + H₂O ⇌ HCO3- + H₃O+

  • The reaction results in a slightly acidic solution due to presence of H3O+ which results in concentrations of [H₃O+] > [OH-]
  • Indicator to use is methyl orange as it changes colour over the range of 3.0-4.5, so end point will match equivalence point
52
Q

Write an equation for the reaction between ethanoic acid and ammonia in aqueous solution? Explain products and what indicator should be used

A

CH₃COOH + NH₃ ⟶ CH₃COO- + NH₄+

NH₄ + H₂O ⇌ NH₃ + H₃O+

CH₃COO- + H₂O ⇌ OH- + CH₃COOH

  • The solution will be roughly neutral depending on which of the above equations occurs more so the [H₃O+] ≈ [OH-]
53
Q

What are buffers

A

solution that resists changes in its pH when small amounts of an acid or base or added

54
Q

What are buffers made of

A

mixture of a weak acid and its conjugate base or a weak base and its conjugate acid

55
Q

Explain the 2 factors of buffer capacity

A
  • Relative Concentration
    • Occurs between weak acid (base) and its conjugate base (acid)
    • Equal concentrations of these two offers the best buffering capacity
  • Absolute Concentration
    • Occurs between weak acid (base) and its conjugate base (acid)
    • The greater the concentrations (in general) the greater the buffering capacity
56
Q

Explain blood buffering when it becomes too acidic

A

H₂CO₃ + H₂O ⇌ HCO3- + H₃O+

  • If blood becomes too acidic, this increases the [H₃O+]
  • This increases the frequency of successful collisions between H₃O+ and HCO3- by decreasing the distance between these particles and so increasing the rate of the reverse reaction according to the equation

HCO3- + H₃O+ ⟶ H₂CO₃ + H₂O

  • This will consume H₃O+ and decrease its concentration
  • Equilibrium is then re-established where [H₃O+] and [OH-] is very close to what it was originally, thus maintaining pH
57
Q

Explain blood buffering when it becomes too basic

A

H₂CO₃ + H₂O ⇌ HCO3- + H₃O+

  • If blood becomes too basic, this increases the [OH-]
  • This decreases the distance between OH- and H₂CO₃ particles, increasing rate of successful collisions according to the reaction:

H₂CO₃ + OH- ⟶ HCO3- + H₂O

  • This will consume OH- and thus, decrease [OH-]
  • Equilibrium is then re-established where [H₃O+] and [OH-] is very close to what it was originally, thus maintaining pH
  • Although the pH is close to its original value, its now slightly higher as the excess HCO3- produced now shifts the equilibrium left, reducing [H₃O+] slightly

H₂CO₃ + H₂O ⇌ HCO3- + H₃O+

58
Q

What is carbonic hydrase and what is its purpose

A
  • An enzyme found in blood
  • catalyses the formation of carbonic acid from carbon dioxide, which ensures the equilibrium adjusts quickly to changes in H₂CO₃ or CO₂
59
Q

Explain buffering in an ethanoic acid solution when a small amount of acid is added

A

CH₃COO- + H₃O+ ⟶ H₂O + CH₃COOH

  • Adding acid increases [H₃O+]
  • The added H₃O+ will react with the CH₃COO- so that [H₃O+] will return very close to where it started, just slightly higher
  • Thus, the pH will return very close to its original value, just slightly lower
60
Q

2 Sources of Systematic Error in Titration

A

Calibration error in burette/pipette

61
Q

2 Sources of Random Error in Titration

A

Difficulty determining end point
Parallax error reading volume of burette

62
Q

Why is NaOH not a good primary standard

A

It cannot be accurately measured
NaOH is deliquiscent and hydroscopic and reacts with CO2 in air

63
Q
A

wace chem (5 decks)

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