Chemistery

Enthalpy votes for releasing heat. Entropy votes for spreading out. When the two disagree — and they constantly do — chemistry needs a referee. That referee has a name, an equation, and one beautifully simple rule. The short answer: Gibbs free energy change (ΔG) combines enthalpy, entropy, and temperature into one number: ΔG = ΔH − TΔS . If ΔG is negative, the process is spontaneous (it can happen on its own); if positive, it isn't; if zero, the system sits at equilibrium. What "free energy" actually means The "free" in free energy doesn't mean it costs nothing — it means available . ΔG measures how much of a reaction's energy change is available to do useful work, after entropy has taken its share. A reaction with ΔG = −100 kJ/mol could, in principle, deliver up to 100 kJ of work per mole — driving a battery, contracting a muscle, powering a cell. And spontaneous is a technical word, not a speed claim. It means "thermodynamically allowed to ...

Two-thirds of the way into every thermodynamics unit, the same question hits: "Wait — is enthalpy the disorder one, or is that entropy?" They're the two quantities that together decide whether a reaction happens, and mixing them up scrambles everything downstream. The short answer: enthalpy change (ΔH) measures the heat a process releases or absorbs at constant pressure, in kJ/mol. Entropy change (ΔS) measures how much more spread out energy and matter become, in J/(mol·K). Nature favours lower enthalpy and higher entropy — and when the two disagree, temperature settles the argument. Quick comparison at a glance Feature Enthalpy (H) Entropy (S) What it tracks Heat content; energy released or absorbed Dispersal of energy and matter; number of ways to arrange Symbol for change ΔH ΔS Units kJ/mol J/(mol·K) — note the joules, not kilojoules Favourable direction Negative (heat released) Positive (more spread out) Everyday signal Flask feel...

You've felt enthalpy changes your whole life: the campfire that warms your hands, the cold pack that soothes a sprain. Chemistry just puts a number and a sign on what your skin already knows — and that number is where thermochemistry calculations begin. The short answer: enthalpy change (ΔH) is the heat a system absorbs or releases during a process at constant pressure, usually reported in kJ/mol. A negative ΔH means heat is released (exothermic); a positive ΔH means heat is absorbed (endothermic). What enthalpy actually is Every substance carries a store of chemical energy — in its bonds, its particle motion, its interactions. Chemists call the heat-related bookkeeping of that store at constant pressure the substance's enthalpy (H) . You can never measure H itself, and you never need to: what reactions reveal is the change , ΔH = H(products) − H(reactants). If the products end up lower in enthalpy than the reactants, the difference left the system as heat: ΔH < 0,...

Every thermochemistry question starts with an unspoken decision: which part of the world are we talking about? Get that boundary wrong and every sign in the problem flips. Students lose more marks to this quiet setup step than to any equation. The short answer: the system is the part of the universe you're studying — usually the reacting chemicals themselves. The surroundings are everything else that can exchange energy or matter with it: the solvent, the flask, the air, the thermometer, you. Together they make up the universe: universe = system + surroundings. Quick comparison at a glance Feature System Surroundings What it is The chemicals or process being studied Everything outside the system Example in a beaker reaction The reacting substances Water, beaker, bench, air, thermometer Where ΔH refers to Energy change of the system — Exothermic reaction Releases energy Absorbs that energy — warms up Endothermic reaction Absorbs energy Supplies...

Leave a metal spoon and a mug of water in the sun and the spoon gets hot long before the water does. Same sunshine, very different responses — because every substance has its own price tag for warming up. The short answer: specific heat capacity (c) is the amount of energy needed to raise the temperature of 1 g of a substance by 1 °C, measured in J/(g·°C). The heat involved in any temperature change is calculated with q = mcΔT — mass × specific heat capacity × temperature change. What specific heat capacity actually means Think of specific heat capacity as a substance's thermal stubbornness . A high value means the substance soaks up a lot of energy for each degree it warms — and hands a lot back as it cools. A low value means its temperature swings easily. Some benchmark values worth recognizing: Substance c in J/(g·°C) Water (liquid) 4.18 Ethanol 2.44 Ice 2.09 Aluminum 0.897 Iron 0.449 Copper 0.385 Lead 0.128 Water's 4.18 is unusu...

A sparkler burns at over 1000 °C, yet a stray spark on your hand barely stings. A bathtub of water at 40 °C could warm you up for an hour. If temperature told the whole story, the sparkler should be far more dangerous — so clearly heat and temperature are not the same thing. The short answer: heat is energy being transferred from a hotter object to a colder one, measured in joules (J). Temperature is a measure of the average kinetic energy of the particles in a substance, measured in °C or kelvin — it tells you which way heat will flow, not how much energy there is. Quick comparison at a glance Feature Heat Temperature What it is Energy in transit between objects Measure of average particle kinetic energy Symbol q T Units joules (J), kilojoules (kJ), calories °C, K (kelvin), °F Depends on amount of substance? Yes — more matter can transfer more energy No — a drop and a bucket of boiling water are both 100 °C Measured with Calorimeter (calculated, ...