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What is enthalpy change, ΔH?
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All Flashcards in Topic 4.1
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4.1.112 cards
What is enthalpy change, ΔH?
The **heat energy** released or absorbed by a reaction at **constant pressure** (units: kJ mol⁻¹).
What is an exothermic reaction?
A reaction that **releases** energy to the surroundings, so they get **hotter**; **ΔH is negative**.
What is an endothermic reaction?
A reaction that **absorbs** energy from the surroundings, so they get **colder**; **ΔH is positive**.
Sign of ΔH for exothermic vs endothermic?
Exothermic → **ΔH < 0** (negative); endothermic → **ΔH > 0** (positive).
Is breaking bonds endo- or exothermic?
**Endothermic** — energy must be **put in** to break a bond.
Is making bonds endo- or exothermic?
**Exothermic** — energy is **released** when a new bond forms.
When is a reaction overall exothermic?
When **making** the new bonds releases **more** energy than **breaking** the old bonds absorbed (net energy out).
What is activation energy, Eₐ?
The **minimum** energy reactants need to react — the reactant level up to the **peak** of the profile.
How do you read ΔH off an energy profile?
It is the energy gap between the **reactant** and **product** levels (down for exothermic, up for endothermic).
Which products are more stable, exo or endo?
**Exothermic** products are **lower** in energy and so **more stable** than the reactants.
Surroundings cool down — what type of reaction?
**Endothermic** — energy is absorbed from the surroundings, so **ΔH is positive**.
Two examples of exothermic reactions?
**Combustion** and **neutralisation** (also respiration) — they release energy.
4.1.212 cards
What is calorimetry?
Measuring the **temperature change** of a known mass of water (or solution) to find the heat transferred by a reaction.
What is specific heat capacity, c?
The energy needed to raise **1 g** of a substance by **1 K** (1 °C). For water, **c = 4.18 J g⁻¹ K⁻¹**.
Equation for heat transferred?
$Q = mc\Delta T$ — heat (J) = mass (g) × specific heat capacity × temperature change.
How do you find ΔT?
ΔT = **T_{final} − T_{initial}**. A change of 1 °C equals a change of 1 K, so the number is the same.
How do you get ΔH per mole from Q?
$\Delta H = -\dfrac{Q}{n}$ — divide Q (in kJ) by the amount that reacted, and add the sign.
Temperature rises — exo or endo, and the sign?
**Exothermic** — heat released to the water — so **ΔH is negative**.
Temperature falls — exo or endo, and the sign?
**Endothermic** — heat absorbed from the water — so **ΔH is positive**.
Which mass goes into Q = mcΔT?
The mass of **water** (the substance heated), **not** the mass of fuel or reactant.
Why convert J to kJ in calorimetry?
Q from $mc\Delta T$ is in **joules**; enthalpy changes are quoted in **kJ mol⁻¹**, so divide by 1000.
Main source of error in combustion calorimetry?
**Heat loss** to the surroundings and apparatus — so the measured ΔH is **less exothermic** than the true value.
Two assumptions in the Q = mcΔT calculation?
All the heat goes to the **water**, and the **specific heat capacity** (and density) of the solution equals that of water.
Order of steps in a calorimetry calculation?
ΔT → **Q = mcΔT** → ÷1000 for kJ → **÷ n** for per mole → add the **sign**.
Topic 4.1 study notes
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