Proton Transfer Reaction

**What is a Brønsted-Lowry acid

** species that donates a proton (H⁺)

**What is a Brønsted-Lowry base

** species that accepts a proton (H⁺) using a lone pair of electrons

**What is a conjugate acid-base pair

** two species that differ by exactly one H⁺

**What does amphiprotic mean

** a species that can both donate and accept a proton (H⁺)

**What does amphoteric mean

** a species that can act as both an acid and a base

**Are all amphoteric substances amphiprotic

** no, but all amphiprotic substances are amphoteric

**Example of amphiprotic substance

** H₂O (water)

**Example of amphoteric but not amphiprotic substance

** Al₂O₃ (aluminium oxide)

**What is the pH formula

** pH = –log[H⁺]

**What is the [H⁺] of a solution with pH 3

** 1 × 10⁻³ mol dm⁻³

**How does a pH change when acid is diluted by a factor of 10

** increases by 1 unit

**How is pH affected when water is heated

** pH decreases (water becomes more acidic because it naturally undergoes auto-ionization aka split into H+ and OH- because of its polarity). More acidic = lower pH

**What is Kw at 298 K

** 1.0 × 10⁻¹⁴ mol² dm⁻⁶

**How are [H⁺] and [OH⁻] related in water

** [H⁺][OH⁻] = Kw

**What is a strong acid

** an acid that completely dissociates in aqueous solution

**What is a weak acid

** an acid that partially dissociates in aqueous solution

**Examples of strong acids

** HCl, HNO₃, H₂SO₄

How to spot weak acids with chemical formula

Exception: HF
Organic acids
Oxoacids with fewer oxygens (HNO2, HClO)
Not a binary acid HX (hydrochloric). Sterngth increases going up a group (HCl > HF)

**What is a strong base

** a base that completely dissociates in aqueous solution

**What is a weak base

** a base that partially dissociates in aqueous solution

How to spot strong bases?

Group 1 + 2 metal hydroxides (OH)
Ex: NaOH or KOH or Ba(OH)₂

How to spot weak bases?

NH₃ (ammonia)
Organic amines
Metal oxides/hydroxides of transition metals

**What does pH depend on

** concentration of H⁺ ions

How do strong and weak acids differ in conductivity

strong acids conduct better due to more ions

How do strong and weak acids differ in reactivity

Strong acids reacts more vigorously than weak acids because it dissacoiates in aqueous solution completely. This means that there are weaker bonds holding the H+ protons, making them easier to lose causing a reactive chemical reaction. However, weak acids partially dissolve in aqueous solutions resulting in a weaker hold onto their protins = less reactive chemical reaction.

**Neutralisation reaction general formula

** acid + base → salt + water

**Ionic equation for neutralisation (equation that shows ions that react during acid-base neutrailisation)

** H⁺(aq) + OH⁻(aq) → H₂O(l)

**How to make sulfates

** react sulfuric acid with metal oxides/hydroxides/carbonates

**How to make nitrates

** react nitric acid with a base

**How to make chlorides

** react hydrochloric acid with a base

**What are the products of acid + metal

** salt + hydrogen gas

**What are the products of acid + metal oxide

** salt + water

**What are the products of acid + metal hydroxide

** salt + water

**What are the products of acid + metal carbonate

** salt + water + carbon dioxide

**What are the products of acid + metal hydrogencarbonate

** salt + water + carbon dioxide

**What is a pH curve

** a graph showing pH change as a base is added to an acid (or vice versa)

How are conjugate strength and acid/base strength related:

Strength of acid and strengrh of its conjugate base are inversely related because if an acid easily loses a proton (strong acid) means that base is very stable (doesn’t accept the proto) or vice versa.

Define the enthalpy of neutralisation:

enthalpy change when 1 mole of water is formed from neutralisation

What does the equivalence point on a pH curve represent:

the point at which acid and base have completely reacted in stoichiometric amounts

How does the pH change after the equivalence point in a strong acid-strong base titration:

pH rises rapidly and then slowly levels off. This is because after equivalence point, all acid is complety neutralized by the base. Extra base stays in solution which increases concentration of OH which increases pH (solution becomes more basic)

How do you calculate pH before equivalence point:

find excess H⁺, then pH = –log\[H⁺]

How do you calculate pH after equivalence point:

find excess OH⁻, then pOH = –log\[OH⁻], and pH = 14 – pOH