Unit 9 Flashcards
Bronsted Lowry Theory
defines acids and bases in terms of proton transfer between chemical compounds
- A Brønsted-Lowry acid is a species that gives away a proton (H+)
- A Brønsted-Lowry base is a species that accepts a proton (H+) using its lone pair of electrons
Conjugate acid-base pairs
is two species that are different from each other by a H+ ion. The acid and base are related to each other by one proton difference
Amphiprotic species
Species that can act both as proton donors and acceptors
Amphiprotic vs amphoteric
- amphoteric means it has both basic and acidic character, when compound reacts with acid, it shows basic character and vice versa
- amphiprotic means it can act as a proton donor and as a proton acceptor
pH formula
pH = – log10[H+]
- [H+] is the concentration of H+ in mol dm–3
- The pH scale is a logarithmic scale with base 10. This means that each value is 10 times the value below it
pH scale
numerical scale that shows how acidic or alkaline a solution is
- Acidic solutions always have more H+ than OH- ions, H+ greater than 10^-7 mol / dm^3. pH always below 7, the higher the H+ of acid, lower the pH
- Basic solutions always have more OH- than H+ ions, H+ lower than 10^-7, pH always above 7 and higher the (OH+) means higher pH
- Water at 298K has equal amounts of OH- and H+ ions with concentrations of 10-7 mol dm-3, meaning pH always 7
How to measure pH using pH meter
- The most accurate way to determine the pH is by reading it off a pH meter
- The pH meter is connected to the pH electrode which shows the pH value of the solution
How to measure pH using universal indicator
- The universal indicator paper is dipped into a solution of acid upon which the paper changes colour
- The colour is then compared to those on a chart which shows the colours corresponding to different pH values
How to calculate Ion product of water (Kw)
Kw (ion product of water) = Kc x [H2O] = 1.00 10-14 at 298K
- The product of the two ion concentrations is always 1.00 x 10–14
How does temperature affect ion product of water (Kw)
As ionisation of water is an endothermic process. Increase in temperature will result in the forward reaction being favoured
- This causes an increase in the concentration of the hydrogen and hydroxide ions
- This leads to the magnitude of Kw increasing
- Therefore, the pH will decrease
Increasing the temperature decreases the pH of water
Decreasing the temperature increases the pH of water
Strong acids
- An acid that dissociates almost completely in aqueous solutions. The position of the equilibrium is so far over to the right that you can represent the reaction as an irreversible reaction
- solution formed is highly acidic due to the high concentration of the H+/H3O+ ions
Weak acids
- an acid that partially dissociates in aqueous solutions. The position of the equilibrium is more towards the left and an equilibrium is established
- The solution is less acidic due to the lower concentration of H+ / H3O+ ions
What are examples of strong acids
HCl: Hydrochloric acid
HBr: Hydrobromic acid
HI: Hydroiodic acid
HNO3: Nitric acid
H2SO4: Sulfuric acid
HClO4: Perchloric acid
What are examples of weak acids
HF: Hydrofluoric acid
H2CO3: Carbonic acid
H3PO4: Phosphoric acid
HNO2: Nitrous acid
H2SO3: Sulfurous acid
HCN: Hydrogen Cyanide
CH3COOH: Ethanoic acid
Organic acids are considered as weak acid.
Strong base
- A strong base is a base that dissociates almost completely in aqueous solutions. The position of the equilibrium is so far over to the right that you can represent the reaction as an irreversible reaction
- The solution formed is highly basic due to the high concentration of the OH– ions
Weak base
- A weak base is a base that partially dissociates in aqueous solutions. The position of the equilibrium is more to the left and an equilibrium is established
- The solution is less basic due to the lower concentration of OH– ions
Examples of strong bases
LiOH: Lithium hydroxide
NaOH: Sodium hydroxide
KOH: Potassium hydroxide
RbOH: Rubidium hydroxide
CsOH: Cesium hydroxide
Ca(OH)2: Calcium hydroxide
Sr(OH)2: Strontium hydroxide
Ba(OH)2: Barium hydroxide
Examples of weak bases
Be(OH)2: Beryllium hydroxide
Mg(OH)2: Magnesium hydroxide
NH3: Potassium hydroxide
CH3NH2: Methyl amine
Organic bases are considered as weak base.
How to distinguish between strong and weak acids
pH value: The stronger the acid, the greater the concentration of H+ and therefore the lower the pH
Electrical conductivity: stronger acid has a higher concentration of H+ it conducts electricity better. Thus greater electrical conductivity. Measured using conductivity meter
Reactivity: Strong and weak acids of the same concentrations react differently with reactive metals. Greater concentration of H+ is greater in strong acids compared to weak acids as more H2 gas produced in shorter time.
Neutralization reactions
A neutralization reaction is one in which an acid (pH <7) and a base/alkali (pH >7) react together to form water (pH = 7) and a salt
Formula
- acid + base (alkali) → salt + water
- H+ (aq) + OH– → H2O (l)
- spectator ions which are not involved in the formation of water, form the salt
Metal and acid reactions
acid + metal → salt + hydrogen
Extent of metal and acid reactions
depends on the reactivity of the metal and the strength of the acid
- Very reactive metals would react dangerously with acids
- Metals low in reactivity do not react at all
- Stronger acids will react more vigorously with metals than weak acid, shown through more fizzing, metal dissolving faster, more exothermic
Metal and oxide reactions
acid + metal oxide → salt + water
Metal and hydroxide reactions
acid + metal hydroxide → salt + water
Metal and carbonate reactions
acid + metal carbonate → salt + water + carbon dioxide
Metal and hydrogen carbonate reactions
acid + metal hydrogencarbonate → salt + water + carbon dioxide
pH curve
a graph showing how the pH of a solution changes as the acid or base is added in a strong acid - strong base titration.
- All pH curves show an s-shape curve
- midpoint of the inflection is called the equivalence point
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