module six - revision deck Flashcards
hydrochloric acid
HCl - strong
Hydrofluoric Acid
HF - weak
Sulfuric Acid
H2SO4 - strong
Nitrous Acid
HNO2 - weak
Nitric Acid
HNO3 - strong
phosphoric acid
H3PO4 - weak
carbonic acid
H2CO3 - weak
Sodium Hydroxide
NaOH - strong
Calcium Hydroxide
Ca(OH)2 - strong
Barium Hydroxide
Ba(OH)2 - strong
Sodium Carbonate
Na2CO3 - weak
Sodium Bicarbonate
NaHCO3 - weak
ammonia
NH3 - weak
litmus
red below pH 5
blue above pH 7.6
Methyl Orange
Red below pH 3.1
Yellow above pH 4.4
bromothymol blue
yellow below pH 6
blue above pH 7.6
phenolphthalein
colourless below pH 8.3
pink above pH 10.0
acid/base reaction
acid + base –> salt + water
acid/carbonate reaction
acid + carbonate –> salt + water + carbon dioxide
acid/metal reaction
dilute acid + active metal –> salt + hydrogen gas
applications of neutralisation - household
antacids - contain bases e.g. mg(OH)2 to neutralise excess HCl in the stomach
toothpaste - alkaline to neutralise the acid in the mouth and to remove food particles that produce acid when they decay
formic acid from bee or ant stings is neutralised by using creams containing bases
applications of neutralisation - industry
slaked lime (calcium hydroxide) or limestone (calcium carbonate) is added to neutralise acid in soil for agriculture
gypsum (calcium sulfate) can be used to neutralise alkaline soil
wastewater from industrial processes is acidic or alkaline so needs to be neutralised before entering waterways
lavoisier model of acids
1780 - Antoine Lavoiser
defined acids as substances containing oxygen
e.g. H2SO4
not all acids contain oxygen (HCl)
Davy model of acids and bases
Humphry Davy (1815) concluded that acids contain hydrogen didn't explain why some compounds that contain hydrogen don't behave like an acid
Arrhenius model
Svante Arrhenius (1884) suggested acids were substances that produced hydrogen ions in aqueous solutions, and that bases produced OH ions in aqueous solutions
doesn’t explain why substances that don’t contain hydroxide act like bases (NH3)
limited to reactions in aqueous solution, so doesn’t explain acid-base reactions in other states
Bronsted-Lowry Model
1923
proposed that an acid-base reaction involved protons being transferred between reaction components and states that acids are proton donors and bases are proton acceptors
does not explain the acidity of acid oxides (SO2 and SO3) and their reactions with basic oxides (CaO)
concentration definition
the amount of solute in a specified amount of solution
strength
the degree to which an acid ionises or a base dissociates
- strong acids are acids in which almost all of the acid molecules ionise to produce a hydrogen ion
conjugate acid/base pairs
members of a conjugate acid-base pair differ from each other by the presence or absence of the transferrable H
the conjugate base of an acid has one less proton
the conjugate acid of a base has one more proton
amphiprotic
a substance that can donate or accept a proton
e.g. HSO4, H2O, H2PO4
equivalence point
the point in a titration where the amount of titrant added is enough to completely neutralise the analyte solution
end point
the physical sign that the equivalence point has been reached
i.e. when the indicator changes colour
titrant
substance you want to determine the concentration of
aliquot
volume measured by the pipette
titre
volume delivered by the burette
primary standard
a chemical that can be made up into a solution of accurately known concentration
properties of primary standards
- available in highly pure form
- large molar mass
- stable in air
- does not absorb moisture or CO2 from the air
- readily soluble in distilled water
- reacts readily with solution of known concentration
when back titrations are needed
- it is difficult to determine a definite end point because the reaction occurs too slowly
- the sample is not soluble in water, but will react with another acid
- the sample is toxic
- the sample is volatile
- the sample is gaseous and in a mixture of gases
- the sample is fairly unreactive
Ka
for HA + H2O(l) H3O+ + A-
Ka = [H3O+][A-]/[HA-]
Ka of weak vs strong acids
Ka of weak acids is very small
Ka of strong acids is very large
percentage ionisation
percentage of an acid that has been ionised in water
[A-]/[HA] x 100
pKa
pKa = -log10 Ka
smaller pKa corresponds to greater acid strength
Kb
for B + H2O(l) BH+ + OH-
Kb = [BH+][OH-]/[B]
relationship between Ka and Kb
Ka x Kb = Kw = 1.01x10^-14
acid base techniques - winemaking
wine mainly contains tartaric acid, malic acid and citric acid and has a pH of about 2.5 to 4
back titrations are often used to find the alcohol content in wine
acid base techniques - mining
need to know the composition of ores, metals and alloys
sample are dissolved in acid then titrated to find the composition
acid base techniques - water treatment
both acids and bases are used to treat drinking water and waste water to ensure they have a pH of around 7
HCl is commonly used to lower the pH
Mg(OH)2 or NaOH used to raise the pH
neutralisation process prevents highly acidic or basic water flowing through pipes and corroding them
Acid Base techniques by ATSI people
Soap Tree (Alphitonia Excelsa)
- used by the Kuku Yalanji people of the Daintree
- leaves contain saponins which tend to be acidic
- the wood, bark and leaves contain many organic acids and small amounts of methyl salicylate
- methyl salicylate is thought to be responsible for the analgesic properties of the plant
- one end of the saponin molecule is hydrophilic while the other is hydrophobic, which explains its ability to act as a mild soap
- used on the skin as a cleanser and as an antiseptic for treating rashes and ringworm
buffers
maintain a pH within a certain range despite the addition of an acid or base
occurs in a conjugate acid base pair where the acid is weak
ability of buffers
the ability of buffers to resist changes in pH depends on the conjugate acid/base
- higher initial concentration of conjugate acid and base means a greater capacity for a buffer to resist changes to its pH
- a buffer should have equal ability to resist changes from acids and bases
blood buffers
blood needs to be in a pH range of 7.35 to 7.45
CO2 + H2O H2CO3
H2CO3 + H2O HCO3 - + H3O+
carbon dioxide converts into acid when dissolved in blood
ocean buffer
ocean absorbs CO2 from atmosphere
CO2 + H2O H2CO3
H2CO3 + H2O HCO3 - + H3O+
ocean buffer systems are unable to combat the increase in CO2, increasing the oceans acidity
cells buffers
cells need pH close to neutral
intracellular fluid contains phosphate buffer system
H2PO4 + H2O HPO42- + H3O+
