Chemistry: Module 6 - Acid Base Reactions

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94 Terms

1
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what is an inorganic Acids and give examples

simple and have a hydrogen and an ion: HCl, HNO3, H3PO4

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define Organic Acids

compounds that are organic but have acidic properties (e.g. alkanoic acids: have -COOH)

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what is the nature of Alkalis

Bases soluble in water → include soluble hydroxides and ammonia (all Group 1 metals)

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Monoprotic

Not all hydrogens in an acid molecule are ionizable: monoprotic = can donate one H+

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define Diprotic

Can donate two H+

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hydronium ion

Is the more stable version of an aqueous H+(aq): H+(aq)+ H2O(l) → H3O+(aq)

H+(aq) = H3O+(aq)

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properties of acids (6)

Produce H+ ions in water

Have a pH < 7

Sour taste

Conducts electricity in solution

Turns blue litmus red

Corrosive

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acid dissociation equation

HA → H+ + A-

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base properties (6)

Produce hydroxide ions (OH-) in solution

pH > 7

Soapy feel, bitter taste

Conducts electricity in solution (if soluble)

Turns red litmus blue

Corrosive (caustic)

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types of bases and an example (4)

Metal Hydroxides

Metal Oxides

(non-metal oxides like CO2 and SO2 are acidic)

Metal Carbonates/Hydrogen Carbonates

Organic Bases (include amines: -NH2)

NaOH, Ba(OH)2

Na2O, CuO

K2CO3, NaHCO3

caffeine, nicotine, morphine

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what is the nature of ammonia?

a weak base

12
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metal oxides are either:

bases or amphoteric

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what is the acetate ion and its nature?

a weak base, CH3COO-

14
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what is amphoteric

  • can act as either an acid or a base (Al2O3) → will be the opposite of whatever it is reacting with

base: Al2O3 (s) + 6HCl(aq) → 2AlCl3 (aq) + 3H2O(l)

acid: Al2O3 (s) + 2NaOH(aq) → 2AlCl3 (aq) + 2Na(Al(OH)4)(aq)


<ul><li><p><span>can act as either an acid or a base (Al<sub>2</sub>O<sub>3</sub>) → will be the opposite of whatever it is reacting with</span></p></li></ul><p style="text-align: center"><span>base: Al<sub>2</sub>O<sub>3 (s)</sub> + 6HCl<sub>(aq)</sub> → 2AlCl<sub>3 (aq)</sub> + 3H<sub>2</sub>O<sub>(l)</sub></span></p><p style="text-align: center"><span>acid: Al<sub>2</sub>O<sub>3 (s)</sub> + 2NaOH<sub>(aq)</sub> → 2AlCl<sub>3 (aq)</sub> + 2Na(Al(OH)<sub>4</sub>)<sub>(aq)</sub></span></p><p><br></p>
15
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what is amphiprotic:

  • Amphiprotic: can donate or accept a proton 

H2O(l) + HCl → H3O+(aq) + Cl-(aq) (water accepted the H+, HCl donated)

H2O(l) + NH3 (aq) → NH4+(aq) + OH-(aq) (water donated H+, NH3 accepted)

16
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what is the purpose of an indicator and what is an indicator’s nature?

used to measure how acidic or basic something is

  • An indicator in any substance that will be different colours in different pH conditions (qualitative change)

  • Indicators are weak acids or bases = exist in two forms

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what is a universal indicator?

  • most commonly used → useful because it is a whole range of different colours depending on the pH - combination of indicators

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purpose of the pH

  • Describes the acidity of a solution [H+]

    • Acids have a high [H+] = low pH

    • Bases have a low [H+] = high pH

    • pH =7: neutral solution

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pH formula

  • pH = -log10[H+]

20
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how does an indicator work?

  • Indicators are weak acids or bases = exist in two forms that are different colours due to its reversibility

HInd ⇌ H+ + Ind-

  • H is the hydrogen atom and Ind is the indicator - the rest of the organic molecule

  • A chemical equilibrium exists in solution between HInd and Ind-: and they are different colours

    • Qualitatively shows where the equilibrium lies

    • If a base is added, [H+] ↓, so eqm shifts right

    • If acid is added, [H+] ↑, so eqm shifts left

21
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methyl orange’s pH range and colour at each pH

Methyl orange

red

3.1- 4.4

yellow

22
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litmus’s pH range and colour at each pH

Litmus (vegetable extract)

red

4.5-8.3

blue

23
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bromothymol blue’s pH range and colour at each pH

Bromothymol blue

yellow

6.0-7.6

blue

24
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phenolphthalein’s pH range and colour at each pH

Phenolphthalein

colourless

8.2-10

pink

25
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how to you prepare and test a natural indicator?

Preparing the Natural Indicator

  1. Tear or slice one red cabbage leaf/flower petal/tea/fruit leaf into small pieces and place them in a beaker

  2. Cover it in 250mL boiling water for 15 minutes

  3. Pour the mixture through a strainer and collect the indicator solution in a clean beaker

Testing the Natural Indicator(observe colour of indicator in solutions with a KNOWN pH)

  1. Pour indicator into 3 small test tubes until they are half-full

  2. Add 1mL of 0.5 HCl into the first tube and agitate to mix (acid)

  3. Add 1mL of 0.5 NaOH into the second tube and agitate to mix (base)

  4. Add 1mL of water into the third tube and agitate to mix (neutral)

  5. Record the colour of each solution

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Acid + base

salt + water

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acid + metal hydroxide

salt + water 


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Acid + Metal Oxides →

salt + water

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Acid + Metal Carbonate →

salt + water + carbon dioxide

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Acid + hydrogen carbonate →

salt + water + carbon dioxide

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Dilute Acid + reactive metal →

salt + hydrogen gas

  • reactive metal is higher than hydrogen on the standard potential table

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Acid + ammonia

ammonium salt

33
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what is the nature of the ammonium ion

a weak acid

34
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what is the nature of ammonia

weak base

35
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what does enthalpy mean

  • how much energy is found within a substance or groups of substances

    • the change of enthalpy (∆H) is measured in kJ/mol

    • ∆H = H(products) - H(reactants)

36
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what is the enthalpy of neutralisation?

  • change in enthalpy per one mole of H2O formed

37
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what is the net ionic equations of strong acids and bases?

  • H+(aq) + OH-(aq) → H2O(l)   ∆Hn = -57kJ/mol

38
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what is the overall enthalpy of neutralisation of weak acids/bases?

  • less exothermic, possibly endothermic

    • They don't completely dissociate in water

    • Heat produced is used to dissociate the weak acid/base into ions

    • ∆H (observed) = ∆H (neutralisation) + ∆H (dissociation of acid/base): negative + positive

39
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how do you find the heat produced?


q=mc∆T

q= heat produced (J)

m = mass of reaction (kg)

c= specific heat capacity of water (4.18 x 103)

∆T = change in temp of water (units don’t matter)

40
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how do you find the change in enthalpy (formula)


∆H = -q/n

∆H=enthalpy change (kJ/mol)

q= heat produced (kJ)

n = moles of water produced

41
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what is acid strength?

the extent of ionisation/dissociation in water (NOT based on the concentration - its the nature of the particle)

  • Acids are covalent molecular compounds that split into ions (ionises) when it interacts with water

42
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what do strong acids do?

Dissociation goes to completion/ionises further (→ )

HA(aq) + H2O(l) → H3O+(aq) + A-(aq) 

HA(aq) →  H+(aq) + A-(aq)

Strong electrolytes (conductors of electricity in solution)

43
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what do weak acids do?

Dissociation is reversible: reaches a dynamic equilibrium (⇌)

% of products/reactants formed depends on the concentration 

HA(aq) + H2O(l) ⇌ H3O+(aq) + A-(aq) 

HA(aq) ⇌ H+(aq) + A-(aq)

Weak electrolytes

44
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what are the 5 main strong acids?

HCl, HBr, HI, H2SO4, HNO3

45
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what does a strong base do?

Strong Base ~ ionic

Dissociates completely in water when it dissolves, producing free OH- ions (→)

Strong electrolytes

NaOH(s) → Na(aq)+ + OH(aq)-

46
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what does a weak base do?

Weak Base

Reacts partially with water to form hydroxide ions in solution (⇌) (include water in equation)

  • If ionic: will dissociate into OH-


Some weak bases are anions inside an ionic compound

  • First dissolution to produce ions (→)

  • Show anion producing OH- from water (⇌)

Weak electrolytes

Ionic

NH3 (aq) + H2O(l) ⇌ NH4+(aq) + OH-(aq) 

Anions in an ionic compound

Na2CO3 (s) → 2Na+(aq) + CO32-(aq)

CO32-(aq) + H2O(l) ⇌ HCO3-(aq) + OH-(aq)

47
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examples of strong bases

Group 1, 2 metal hydroxides: NaOH, KOH, Ca(OH)2

48
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what is Ka?

  • Equilibrium constant for the dissociation (ionisation) of an acid into a hydrogen or hydronium ion and anion

  • Indicates the extent of dissociation/strength of an acid

  • The stronger the acid, the ↑ Ka (eqm lies on the right, more ionisation)

  • Polyprotic acids have more than one Ka value (for each H+ they donate)

49
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what is Kb?

  • Indicates the extent of dissociation/strength of base

  • The stronger the base, the ↑Kb

50
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water is both

amphoteric and amphiprotic (acts as both base and acid and can donate/accept a proton)

  • Water can acts as either a B/L acid or a B/L base

51
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give an example of water acting as an acid

  • Molecular acids dissolve in water to produce ions -- according to the Bronsted/Lowry theory this occurs because a proton is donated from the acid to the water molecule to produce the hydronium ion

NH3 (g) + H2O(l) ⇌ NH4+(aq) + OH-(aq)



52
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give an example of water acting as a base

HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)

  • The HCl = B/L acid as it it a proton donor, water is the B/L base as it is a proton acceptor

53
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pH =

-log10[H+]


54
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pOH =

-log10[OH-]

55
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pH + pOH =


14 (for aqueous solutions at 25℃)

56
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what does the pOH range look like?

57
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what did Lavoisier propose for the definition of acids?

  • Acids contain oxygen - the presence of oxygen gives acids their properties

  • Based on compounds like HNO3 and H2SO4

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what is the limitations of Lavoisier’s definition

Couldn't explain acids that didn’t contain oxygen (HCl)

59
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what is Davy’s definition of an acid?

  • Stated acidity is due to the presence of hydrogen

  • Showed that HCl is strongly acidic

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what is Davy’s limitation in his definition?

  • Doesn’t define a base

  • Lacks an explanation of the chemical behaviour of bases

  • Didn’t explain why other compounds that contained H weren’t acidic (CH4)

61
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what is arrhenius’ definition of acids and bases?

Acids as substances that increases the concentration of H+ ions in aqueous solutions

HCl → H+ + Cl-

  • Strongest acids are the most dissociated


Bases are substances that  increase the concentration of OH- ions in aqueous solutions

NaOH → Na+ + OH-

  • Allowed for an explanation of neutralisation

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what is the limitation of arrhenius’ definition?

  • Only for aqueous solutions

  • Doesn’t explain reactions in non-aqueous solvents or gas phase reactions

  • Cannot explain bases that don’t contain OH- like NH3 

  • Couldn’t explain why some neutralisation reactions produced unneutral solutions

63
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what is bronsted and Lowry’s definition of acids and bases?

Acids are proton (H+) donors

Bases are proton acceptors (H+

  • Explains a wide range of acid-base behaviour

  • Explains how ammonia works as a base because it accepts a proton

  • Explains that ionic compounds are basic because they dissolve first and then ionise

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what is the benefit of bronsted lowry defintions?

Works in both aqueous and non-aqueous environments (any solvent)

  • Introduces the idea of conjugate acid-base pairs

NH3 (g) + HCl(g) → NH4Cl(s)

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NH3 + H2O ⇌ NH4+ + OH-

what is the B/L base and the B/L acid

  • NH3 acts as a B/L base as it accepts a proton to become NH4+

  • H2O acts as a B/L acid as it donates a proton to become OH-

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what is the limitation of B/L definition

  • Doesn’t explain acid-base reactions that don’t involve proton transfer (Lewis acid-base theory) e.g. AlCl3 and BF3

  • Can’t explain reactions between acidic oxides and basic oxides that occur without a solvent

  • HOWEVER, it is still used because it is sufficient in covering most situations

67
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what is the lewis acid-base theory?

Lewis acid: an electron pair acceptor

Lewis base: an electron pair donor


68
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what is the order of acid/base definitions from old to new?

lavoisier, davy, arrhenius, bronsted-lowry, lewis

69
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explain the self-ionisation of water at 25 degrees?

H2O(l) + H2O(l) ⇌ H3O+(aq) + OH-(aq)

  • its pH = 7, so the [H+] = 10-7 mol/L

  • Kw = [H3O+][OH-] = 10-14 (from data sheet)

  • ∴ [OH-]  = 10-7

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Kw =

[H3O+][OH-] = 10-14 (at 25℃)

71
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what is conjugate acid-base pairs?

  • If something is an acid, something else is a base 

  • When a B/L acid donates its proton to the base, the anion of the acid is proton deficient - this species can act as a base as it could accept a proton and reform the original acid

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when writing a conjugate acid base pair what do you write first?

When writing the pairs, write the acid first: so here: H2O/ OH- and NH3/ NH4+

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what is the base dissociation formula?

AB —> A+ + B-

B- + H2O → HB + OH-

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what is the strength of an acid?

the degree of ionisation of an acid or dissociation of a base → can be determined by calculating the [H+] of equal concentration of acids

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how does the dissociation of polyprotic acids work?

in subsequent reactions (NOT all at once)

  • the reactions strength can be different

76
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what is the advantages and disadvantages of using a pH probe?

advantages

disadvantages

  • More precise

  • Doesn’t change composition of sample (unlike indicators)

  • Has to be calibrated to ensure accuracy

  • Is initially expensive

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what is the experiment to indicate the difference in strengths between acids and bases

  1. Add 100mL of 0.1mol/L of each solution in three separate, dry, 250mL beakers

  2. Label the beakers with the solutions it contains

  3. Place the calibrated pH probe into the first beaker, and record the pH

  4. Remove the probe from the solution and rinse it with distilled water and blot dry

  5. Repeat steps 2 and 3 for the remaining solutions

  6. Repeat the procedure for 0.01M and 0.001M of the solutions

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what is the relative strength of conjugate pairs?

  • The pairs have inverse strength: strong acid has extremely weak (almost neutral) conjugate base

    • Weak acid will have relatively strong conjugate base


Weak acid

Relatively strong conjugate base

Weak base

Relatively strong conjugate acid

Strong acid

Relatively weak/neutral conjugate base

Strong base

Relatively weak/neutral conjugate acid

79
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why do conjugate pairs have inverse strength?

  • A strong acid has a strong tendency to give up H+ and completely ionises in water

    • Equilibrium lies far to the right

    • The reverse reaction will barely proceed as the conjugate base has no tendency to accept H+ as the reaction is not reversible

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what is concentration?

Concentration = the amount of solute in a specified amount of solution (oft. Expressed as moles/L)

  • Concentrated = total concentration is high

  • A dilute solution is one which the total concentration is low (e.g 0.1M)

81
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what is the strength of a solution?

The degree of ionisation of an acid or dissociation of a base → can be determined by calculating the [H+] of equal concentration of acids


82
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draw a model of a concentrated weak acid, concentrated strong acid, dilute weak acid, dilute strong acid?

83
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is the pH of the final solution in a neutralisation always neutral?

Acid + Base → salt + water

pH of final solution is not always neutral

  • The salt could also be acidic or basic.

  • The excess acid/base would react with water to produce ions

84
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what type of salt is formed if: strong acid + strong base

neutral salt

85
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what type of salt is formed if: strong acid + weak base

acidic salt

86
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what type of salt is formed if: weak acid + strong base

basic salt

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what type of salt is formed if: weak acid + weak base

Depends on which is the stronger of the acid and base

  • Compare the dissociation value

  • If Ka > Kb = acidic

If Kb > Ka = basic

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how to calculate the pH of a weak acid/base?

  1. Write equation and Ka expression

  2. Use ICE table to determine eqm conc

  3. Let x be the change in concentration

  4. Assume Ka is small, so it = 1 (prove it later)

  5. Solve for x

  6. Check assumption is < 5%

  7. Calculate pH

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if an acid is multiprotic how do you calculate the pH?

If the acid is multiprotic, do calculations of the H+ concentrations of each dissociation and then add the concentrations together


90
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what is the formula for degree of ionisation?

% ionised = ([H+]eqm / [HA] initial) x 100

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how to calculate the pH of diluted strong acids/bases?

  1. Find the concentration of the solution diluted (use C1V1 = C2V2)

  2. Find the concentration of diluted H+

  3. Calculate the pH

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how to find the pH of diluted weak acids/bases?

  1. Write the equation (⇌) and the Ka formula

  2. Determine the new diluted concentration (use C1V1 = C2V2)

  3. Use a Ka or Kb calculation (using an ICE table) to find the concentration of diluted H+

  4. Calculate the pH

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how to calculate the pH of the mixing of a strong acid and base (no reaction)

  1. Calculate the moles of H+ or OH- from each solution

  2. Calculate the total moles of H+

  3. Calculate the final volume

  4. Calculate the final [H+] and find the pH

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how to calculate the pH of a neutralisation reaction b/w strong acid and strong base?

  1. Write a balanced equation

  2. Find the moles of the acids and bases 

  3. Find the limiting reagent/ excess reagent

  4. Calculate the leftover moles of the excess reagent

nexcess = ninitial - nexcess used with limiting reagent - use mole ratio

  1. Calculate the concentration of excess reagent using the total volume

  2. Calculate the pH