Module 6 Flashcards
Important acids to know:
6.1.1 Investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases
Hydrochloric acid: HCl
Nitric acid: HNO3
Ethanoic (acetic) acid: CH3COOH
Sulfuric acid: H2SO4
Carbonic acid: H2CO3
Phosphoric acid: H3PO4
Citric Acid: C6H8O7
Important bases to know:
6.1.1 Investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases
Sodium hydroxide: NaOH
Ammonia: NH3
Potassium hydroxide: KOH
Calcium hydroxide: Ca(OH)2
Magnesium hydroxide: Mg(OH)2
Sodium carbonate: Na2CO3
How to name organic acids
Acids with oxyanions
Oxyanions with the central atom being in a reduced state
6.1.1 Investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases
If in aqueous form, drop the ide at the end for the anion and change it to ic and then add acid to the end. The word hydrogen changes to hydro as a prefix.
E.g. Hydrochloric acid or HF (hydrofluoric acid)
If it ends with ate, it is dropped and changed to ic. The hydrogen isn’t included in the name.
H2SO4: sulfuric acid
The suffix ite is dropped with ous.
H2SO3: sulfurous acid
HNO2: nitrous acid
Properties of acids and bases
6.1.1 Investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases
Acid:
Sour taste
Can burn skin (corrosive)
Conductive in solution
Turn blue litmus red
pH<7
Base:
Bitter taste
Soapy feel in aqueous solution
Can burn skin (caustic)
Conductive in solution
Turn red litmus blue
pH.7
Acid + base reaction
Acid + carbonate reaction
Acid + metal reaction
6.1.1 Investigate the correct IUPAC nomenclature and properties of common inorganic acids and bases
Acid + base -> salt + water
Acid + carbonate ->salt + water + carbon dioxide
Acid + metal -> salt + hydrogen
Note: name of salt is cation base and anion acid
Define indicator
Examples of indicators and their characteristics range
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
Substances which change colour based on the pH of the environment
Indicator. pH Acidic range Transition range Basic range
transition
range
Methyl 3-4 Red (M<3.1) Orange (3.1<M<4.4). Yellow (M>4.4)
orange
Bromothymol. 6-7.5 Yellow (B<6). Green (6<B<7.6) Blue (B>7.6)
blue
Litmus. 5.5-8 Red. Purple Blue
Phenolphthalein. 8-10. Colourless. Pale pink Pink/magenta
Example of natural indicator and how it functions
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
Red cabbage contains weak acids called anthocyanins. Acids change colour based on the no of removable protons that remain attached to the molecule. For a diprotic anthocyanin molecule:
H2Antho(aq)⇌ H+(aq) + HAntho-(aq)
red. blue
HAntho-(aq)⇌ H+(aq) + Antho2-(aq)
blue yellow
According to LCP, as [H+] changes, the eq will shift. This will elicit a colour change hence acting as a pH indicator.
When an acid is added, the [H+] increases and the eq shifts left towards H2Antho. Thus, as the environment becomes more acidic, the colour tends to approach the rend end of the spectrum.
Conversely, when a base is added, the [OH-] increases and consumes free H+ causing [H+] to decrease. By LCP, the eq shifts right producing more Antho2-. This causes the colour to approach the yellow end of the spectrum as the environment becomes more alkaline.
Note: many acid-base indicators are also weak acids and function like this
Compare a natural indicator to a universal indicator and justify its use
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
The red cabbage indicator is successfully able to classify substances into pH ranges based on colour changes, and hence fulfils the essential function of an indicator. However, it is less precise in identifying a specific pH for an approximately neutral substance in comparison to universal indicator (UI). For instance, the pH 4-7, the red cabbage indicator shows various shades of purple which are difficult to distinguish whilst UI moves through a number of more distinctive colour changes. However, red cabbage indicator can be better for substances with pH <3 or pH>11 but these substances are less common. Thus, red cabbage indicator is most likely to be useful in broadly categorising substances as acidic, basic or approximately neutral but less useful in comparing substances with marginally different pH values.
Indicator vs pH probe
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
Indicators pH probe
Accuracy: Relatively broad pH range Specific pH usually to 2dp
Reliability: Inter-trial variability due to colour. Highly reproducible measurements
range interpretation
Inter-investigator variability (changing person)
Validity: Destructive (1) Non-destructive
Not ideal for coloured solutions (2) Valid for all aqueous settings
Adding too much indicator changes pH (3)
Other: Cheap and portable Expensive and less portable
No maintenance Very susceptible to environment, requires calibration with
buffer solutions (4) and storage in electrolyte (5)
Compare indicators and pH probe through these questions
How are indicators destructive
Why aren’t indicators not ideal for coloured solutions
How do indicators affect pH
Importance of buffer solutions
How is it stored
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
1: Indicators are destructive as they contaminate the sample solution whilst pH probes don’t
2. Acid-base indicators aren’t ideal for coloured solutions as their colour can interfere with interpretation. Excludes if coloured solutions are diluted such that the colour is sufficiently faint.
3. Indicators are themselves weak acids or bases so only few drops should be added
4. Buffer solutions are solutions with stable pH. Usually, three buffer solutions are used with stable pHs of 4, 7 and 10 to calibrate the pH probe before use.
5. pH probes should be stored in a solution of KCl to maintain the internal electrodes. They must never be stored dry as this will cause the internal electrode to degrade.
Why do oxides of non-metal elements (C, N, S) tend to be acidic. GIve 4 examples
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
Since they form acids when reacted with water
Co2(g) + H2O(l) ⇌H2CO3(aq) carbonic acid
2NO2(g) + H2O(l) ⇌HNO2(aq) + HNO3(aq) nitrous acid and nitric acid
SO2(g) + H2O(l) ⇌H2SO3(aq) sulfurous acid
SO3(g) + H2O(l) ⇌H2SO4(aq) sulfuric acid
Why do oxides of metal elements (Na, K, Ca) tend to be basic
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
Since they form hydroxide bases when reacted with water
Na2O(s) + H2O(l)⇌2NaOH(aq) sodium hydroxide
K2O(s) + H2O(l) ⇌wKOH(aq) potassium hydroxide
CaO(s) + H2O(l) ⇌Ca(OH)2 (aq) calcium hydroxide
`Equations of
CO2(g) + NaOH (aq)
H3PO4(aq) + CaO(s)
SO2(g) + K2O
6.1.1 Conduct an investigation to demonstrate the preparation and use of indicators as illustrators of the characteristics and properties of acids and bases and their reversible reactions
Acid-base reactions can be written using acidic or basic oxides as one or both reagents. The products are salt and water or simply salt. When an acidic oxide reacts with a base, the anion of the product salt is derived from the aqueous acid that the acidic oxide forms if it were reacted with water
CO2(g) + NaOH (aq) ⇌2Na2CO3(aq) + H2O(l)
H3PO4(aq) + CaO(s) ⇌2Ca3(PO4)2 (s) + 3H2O(l)
SO2(g) + K2O (s)⇌2 K2SO3(s) as reactant states are gas and solid meaning water cannot exist
State Lavosier’s theory. Give an example that works and doesn’t. Assess limitations and strengths
State Davy’s theory. Give an example that works and doesn’t. Assess limitations and strengths
6.16 Explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model, including but not limited to:
Arrhenius theory
Brønsted-Lowry
Acids were substances which contained oxygen such as HNO3.
Pros: Raised need to define acids and bases
Cons: Was wrong CaO is basic.
Acids are substances which contained replaceable hydrogen meaning hydrogen in the compound could be partially or totally removed. Ca(s) + 2HCl (aq) –> CaCl2(aq) + H2(g)
Pros: Worked for many common acids
Cons: Couldn’t explain acidic or basic oxides such as SO2
State Arrhenius theory. Assess limitations and strengths
6.16 Explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model, including but not limited to:
Arrhenius theory
Brønsted-Lowry
Acids are substances which in aqueous solution ionised to form H+ ions.
HA(aq) –> H+(aq) + A-(aq)
Bases are substances which in aqueous solution disassociate to form OH- ions.
XOH(aq) –> X+(aq) + OH-(aq)
Pros:
Works for most bases and acids.
Has a common fundamental mechanism of acid-base neutralisation reactions.
Able to explain potency differences between strong and weak acids by relating it to the degree of ionisation
Limitations:
Doesn’t recognise role of solvent in the relative weakness or strength of an acid where this arises from its nature and the solvent. HCl is a strong acid in water and weak in other solvents such as diethyl ether.
Many metal oxides and carbonates are basic but don’t have OH nor do they liberate OH- in aqueous solution.
Only accounted for acid-base neutralisation reactions in ionised form. Some don’t occur in this form e.g. HCl(g) + NH3(g) –> NH4Cl(s) can occur dissolved in benzene or in gaseous phase with molecular rather than ionised reactants
