From: Andrew Lorimer Date: Sun, 31 Mar 2019 09:43:50 +0000 (+1100) Subject: [chem] acid/base equilibria and tidy section on K_c value X-Git-Tag: yr12~176 X-Git-Url: https://git.lorimer.id.au/notes.git/diff_plain/d1a774977e5a53fd74716e2aae7d31059b14dabd [chem] acid/base equilibria and tidy section on K_c value --- diff --git a/chem/graphics/rxn-complete.png b/chem/graphics/rxn-complete.png new file mode 100644 index 0000000..89130f0 Binary files /dev/null and b/chem/graphics/rxn-complete.png differ diff --git a/chem/graphics/rxn-incomplete.png b/chem/graphics/rxn-incomplete.png new file mode 100644 index 0000000..63c008b Binary files /dev/null and b/chem/graphics/rxn-incomplete.png differ diff --git a/chem/reactions.md b/chem/reactions.md index 7e8d080..67f0a19 100644 --- a/chem/reactions.md +++ b/chem/reactions.md @@ -1,3 +1,10 @@ +--- +header-includes: + - \usepackage{mhchem} +columns: 2 +geometry: margin=2cm +--- + # Rates and Equilibria ## Energy profile diagrams @@ -7,6 +14,10 @@ $$E_A = E_{\text{max}} - E_{\text{initial}}$$ - Energy always needed to initiate reaction (break bonds of reactants) - Reactant particles must collide at correct angle, energy etc - Most collisions are not fruitful +- Energy must be greater than or equal to $E_A$ + +**Endothermic** (products > reactants, $\Delta H > 0$) +**Exothermic** (reactants > products, $\Delta H < 0$) ![](graphics/endothermic-profile.png) ![](graphics/exothermic-profile.png) @@ -19,7 +30,7 @@ $$E_A = E_{\text{max}} - E_{\text{initial}}$$ ## Kinetic energy -Temperature - measure of _avg_ kinetic energy of particles. Over time each particle will eventually have enough energy to overcome $E_A$. +**Temperature** - measure of _avg_ kinetic energy of particles. Over time each particle will eventually have enough energy to overcome $E_A$. Note same distribution indicates same temperature. ![](graphics/ke-temperature.png) @@ -28,6 +39,9 @@ Note same distribution indicates same temperature. - alternate reaction pathway, with lower $E_A$ - increased rate of reaction - involved in reaction but regenerated at end +- does not alter $K_c$ or extent of reaction +- attracts reaction products +- removal/addition of catalyst does not push system out of equilibrium **Homogenous** catalyst: same state as reactants and products, e.g. Cl* radicals. **Hetrogenous** catalyst: different state, easily separated. Preferred for manufacturing. @@ -38,33 +52,55 @@ Haber process (ammonia producition) - enzymes are catalysts for one reaction eac ## Equilibrium systems -**Equilibrium** - the stage at which quantities of reactants and products remain unchanged +*Equilibrium* - the stage at which quantities of reactants and products remain unchanged + Reaction graphs - exponential/logarithmic curves for reaction rates with time (simultaneous curves forward/back) -## Equilirbium constant $K_C$ +![](graphics/rxn-complete.png){#id .class width=20%} +**Complete reaction** - all reactant becomes product + +![](graphics/rxn-incomplete.png){#id .class width=20%} +**Incomplete reaction** - goes both ways and reaches equilibrium + +- All reactions are equilibrium reactions, but extent of backwards reaction may be negligible +- Double arrow indicates equilibrium reaction +- At equilibrium, rate of forward reaction = rate of back reaction. + +### States (not in course) + +- **Homogeneous** - all states are the same +- **Heterogeneous** - states are different -For reaction $aA + bB + cC + dD + \dots \leftrightarrow zZ + yY + xX + \dots$: +## Equilibrium constant $K_c$ -$$K_c = {{[Z]^z [Y]^y [X]^x \dots} \over {[A]^a [B]^b [C]^c [D]^d \dots}}$$ +For \ce{$\alpha$A + $\beta$B + $\dots$ <=> $\chi$X + $\psi$Y + $\dots$}: -Indicates extent of reaction. If value is high ($> 10^4$), then [products] > [reactants]. If value is low ($< 10^4$), then [reactants] > [products]. +$$K_c = {{[\ce{X}]^\chi \cdot [\ce{Y}]^\psi \cdot \dots} \over {[\ce{A}]^\alpha \cdot [\ce{B}]^\beta \cdot \dots}}$$ -If $K_c$ is small, equilibrium lies *to the left*. +More generally, for reactants $n_i \ce{R}_i$ and products $m_i \ce{P}_i$: -**$K_c$ depends on direction that equation is written (L->R)** +$$K_c = {{\prod\limits^{|P|}_{i=1} [P_i]^{m_i}} \over {\prod\limits^{|R|}_{i=1} [R_i]^{n_i}}} \> | \> i \in \mathbb{N}^*$$ -## Reaction constant $Q$ +Indicates extent of reaction +If value is high ($> 10^4$), then [products] > [reactants] +If value is low ($< 10^4$), then [reactants] > [products] -Same for as $K_C$. If $Q=K_c$, then reaction is at equilibrium. +- **$K_c$ depends on direction that equation is written (L to R)** +- If $K_c$ is small, equilibrium lies *to the left* +- aka *equilibrium expression* + +## Reaction constant (quotient) $Q$ + +Proportion of products/reactants at a give time (specific $K_C$). If $Q=K_c$, then reaction is at equilibrium. ## Le Châtelier’s principle > Any change that affects the position of an equilibrium causes that equilibrium to shift, if possible, in such a way as to partially oppose the effect of that change. -### Changing volume +### Changing volume / pressure -1. $\Delta V \implies [\Sigma \text{particles}] \uparrow$, therefore system reacts in direction that produces less particles -2. $\Delta V \implies [\Sigma \text{particles}] \uparrow$, therefore system reacts in direction that produces more particles +1. $\Delta V < 0 \implies [\Sigma \text{particles}] \uparrow$, therefore system reacts in direction that produces less particles +2. $\Delta V > 0 \implies [\Sigma \text{particles}] \downarrow$, therefore system reacts in direction that produces more particles 2. $n(\text{left}) = n(\text{right})$ (volume change does not disturb equilibrium) ### Changing temperature @@ -73,14 +109,33 @@ Only method that alters $K_c$. Changing temperature changes kinetic energy. System's response depends on whether reaction is exothermic or endothermic. -- Exothermic - increase in temperature decreases $K_c$ -- Endothermic - increase in temperature increases $K_c$ +- Exothermic - increase temp decreases $K_c$ +- Endothermic - increase temp increases $K_c$ Time-concentration graph: smooth change +### Changing concentration + +- Decreasing "total" concentration of system causes a shift towards reaction which produces more particles + ## Yield -$$\text{yield %} = {{text{actual mass obtained} \over {theoretical maximum mass}} \times 100$$ +$$\text{yield \%} = {{\text{actual mass obtained} \over \text{theoretical maximum mass}} \times 100}$$ + +## Acid/base equilibria + +Strong acid: $\ce{HA -> H+ + A-}$ +Weak acid: $\ce{HA <=> H+ + A-}$ + +For weak acids, dilution causes increase in % ionisation. +$\therefore [\ce{HA}] \propto 1 \div \text{\% ionisation}$ +(see 2013 exam, m.c. q20) +$$\text{\% ionisation} = {{[\ce{H+}] \over [\ce{HA}]} \times 100}$$ +When a weak acid is diluted: +- amount of $\ce{H3O+}$ increases +- equilibrium shifts right +- overall $[\ce{H3O+}]$ decreases +- therefore pH increases diff --git a/chem/reactions.pdf b/chem/reactions.pdf new file mode 100644 index 0000000..8419a6e Binary files /dev/null and b/chem/reactions.pdf differ