chem: Add strong/weak acids/bases, properties of acids/bases, Louis theory
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@ -577,6 +577,30 @@ $\Delta G = -RT\ln K_c$
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## Acids and bases
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!!! definition
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- An **amphiprotic** acid/base is one that can **either accept or donate** $\ce{H+}$ depending on the situation.
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- A **monoprotic** acid/base is one that can only accept/ionise one $\ce{H+}$ ion.
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- An **amphoteric** chemical may act as an acid or base depending on the situation.
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- An **alkali/alkaline** solution is an aqueous solution of a base, which may **not** necessarily be a **basic solution**.
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An **acid** and **base** are any two corrosive chemicals that react to form water and a salt. They also dissociate/ionise (depending on theory) in water to form electrolytes that conduct electricity.
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Acids:
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- taste sour
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- have a pH less than 7 in aqueous solutions at 25°C
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- stain litmus paper **red**
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- react with active metals to produce $\ce{H2_{(g)}}$ based on the activity series
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- react with carbonates to form $\ce{CO2 + H2O}$
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Bases:
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- taste bitter
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- have a pH greater than 7 in aqueous solutions at 25°C
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- feel slippery as they react with fats/oils to form soap
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- stain litmus paper **blue**
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- react with ammonium salts to product $\ce{NH3 + H2O}$
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### Arrhenius theory
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An acid **dissociates** in water to produce $\ce{H+}$ ions (protons).
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@ -585,20 +609,29 @@ A base **dissociates** in water to produce $\ce{OH-}$ ions.
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### Bronsted-Lowry theory
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The Bronsted-Lowry theory focuses on reactions with water and less the acid and base ions themselves.
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The Bronsted-Lowry theory focuses on reactions with water and less the acid and base ions themselves, so they **ionise** instead of **dissociate**.
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An acid is any compound that can **donate a proton ($\ce{H+}$) to water** to form a hydronium ion.
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An acid is any compound that can **donate/ionise a proton ($\ce{H+}$) to water** to form a hydronium ion.
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$$\ce{acid + H2O -> acid- + H3O+}$$
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!!! info
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In practice, the acid must contain a hydrogen atom attached by an easy-to-break bond (usually $\ce{H-O}$), but any high electronegativity difference polar bond would work as well.
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A base is any compound capable of **removing a proton ($\ce{H+}$) from an acid**.
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A base is any compound capable of **accepting/removing a proton ($\ce{H+}$) from an acid**.
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$$\ce{acid + base -> acid- + base+}$$
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!!! info
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The proton usually comes from water. The base must be able to accept an $\ce{H+}$ ion to form a **dative covalent bond**, so they must contain **lone pairs**.
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Polyprotic acids ionise their $\ce{H+}$s one by one **sequentially**.
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!!! example
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$\ce{
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H3PO4 + H2O <=> H2PO4- + H3O+ \\
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H2PO4- + H2O <=> HPO4^2- + H3O+ \\
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HPO4^2- + H2O <=> PO4^3- + H3O+
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}$
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#### Conjugate acids/bases
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The result of a base obtaining a proton is a **conjugate acid**.
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@ -613,8 +646,42 @@ The result of an acid losing a proton is a **conjugate base**.
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### Louis theory
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A Lewis **acid** is any species that **accepts** an electron pair to form a dative covalent bond.
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A Lewis **base** is any species that **donates** an electron pair to form a dative covalent bond.
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### Strong/weak acids/bases
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**Strong** acids/bases will **completely** dissociate/ionise in an aqueous solution. This means that the initial concentration of acid will be equal to the end concentration of $\ce{H+ or H3O+}$.
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All strong polyprotic acids initially have a one-way reaction then follow with equilibrium reactions.
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!!! warning
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Strength is a property of an acid and has nothing to do with its concentration.
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**Weak** acids/bases will only **partially** dissociate/ionise in an aqueous solution, leaving behind most of the initial acid ($\ce{[acid] > [H+]}$ at equilibrium).
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!!! warning
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Measuring pH only returns $\ce{[H+] or [H3O+]}$, so it cannot be used to determine the concentration, identity, or strength of an acid.
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All weak polyprotic acid reactions are equilibrium reactions.
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!!! example
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The following is a list of strong and weak acids:
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| Strong acid | Weak acid | Strong base | Weak base |
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| --- | --- | --- | --- |
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| $\ce{HClO4}$ | any $\ce{COOH}$ | $\ce{LiOH}$ | $\ce{NH3}$ |
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| $\ce{HCl}$ | $\ce{CO2}$ | any $\ce{group\ 1 + OH}$ | $\ce{Al(OH)3}$ |
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| $\ce{HBr}$ | $\ce{SO2}$ | any $\ce{group\ 2 + (OH)2}$ | |
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| $\ce{HI}$ | $\ce{HF}$ | | |
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| $\ce{H2SO4}$ | $\ce{HCN}$ | | |
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| $\ce{HNO3}$ | $\ce{H3PO4}$ | | |
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To experimentally distinguish between a strong or weak acid/base, if their concentrations are equal, total **ion** concentration or $\ce{H3O+}$ concentration can be compared since the stronger acid ionises more.
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Practically, this means comparing the rate of reaction with a metal or water or measuring conductivity as they reflect total ion count.
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### pH and pOH
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## Organic chemistry
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