6381062
Description
Flashcards by Kayla Price, updated more than 1 year ago
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Created by Kayla Price
over 7 years ago
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| Question | Answer |
| Define standard lattice enthalpy | The energy transferred when 1 mole of ionic solid if formed from its gaseous ions in standard conditions. |
| How does the strength of ionic bonding change as lattice enthalpy becomes more negative? | The bonding gets stronger |
| How does the lattice enthalpy change when the charge of the ions increases and why? | It becomes more negative (more energy is released) because there are stronger electrostatic forces between ions. |
| How does the lattice energy change as the ionic radii of ions increases and why? | It becomes less negative (less energy is released) because charge density decreases and ions do not pack together as closely, so the electostatic force between ions decreases. |
| Define Hess' law | The total enthalpy change of a reaction is always the same no matter what route is taken, energy is conserved. |
| Define the enthalpy of atomisation | The energy transferred when 1 mole of gaseous atoms is formed from its element in standard states and conditions. |
| Define electron affinity | The energy transferred when 1 electron is added to one mole of gaseous atoms. |
| Define ionisation energy | The energy transferred when 1 mole of singular gaseous atoms are formed from its elements in standard conditions. |
| Define the enthalpy of formation | The energy transferred when one mole of solid is formed from its elements in standard states and conditions. |
| What is the difference between theoretical and experimental lattice energies? | Experimental uses enthalpy values found experimentally and the lattice enthalpy is found using Born-Haber cycles. Theoretical values are found using calculations and the purely ionic model of a lattice. |
| How is comparing the theoretical and experimental values for lattice enthalpy useful? | It shows how ionic a lattice is. The closer the values are, the more purely ionic the lattice is. |
| Why is the experimental lattice enthalpy more negative than the theoretical? | There is some covalent character in the ionic bonds of a lattice, which makes the bonding stronger and so the lattice energy more negative. Theoretical values for lattice energy do not take into account covalent character whereas the experimental does so experimental is usually more negative. |
| What gives an ionic bond more covalent character? | A smaller cation with a higher charge (large charge density) and the larger anion with a large charge (small charge density). This gives the cation more polarising power and distorts the anion by pulling electrons towards it, increasing covalent character. |
| Define electronegativity | The ability for an atom to attract a pair of electrons in a covalent bond. |
| How is the polarity of a bond calculated? | Working out the difference in electronegativity values for the atoms |
| Define the enthalpy of hydration | The enthalpy change when 1 mole of gaseous ions dissolves in sufficient to give and infinitely dilute solution |
| Define the enthalpy of solution | When 1 mole of solid is dissolved in sufficient solvent to ive an infinitely dilute solution |
| What is the enthalpy of solution for a soluble substance? | Negative - lattice enthalpy > enthalpy of hydration |
| How is enthalpy of hydration increased? | Ions have a greater charge and smaller ionic radius. |
| What is entropy? | A measure of disorder of a system, the number of ways particles can be arranged. |
| How does entropy change as disorder increases? | Increases |
| What affects the entropy of a substance? | The state, number of moles and complexity of the substance. |
| How does stability change as entropy increases? | Increases |
| What is the first law of thermodynamics? | Energy can't be created or destroyed |
| State the second law of thermodynamics | The entropy of an isolated system always increases |
| State the third law of thermodynamics | The entropy of a perfect crystalline solid at absolute zero is zero |
| How would you calculate the entropy change of a system? | ∆Ssystem = Sproducts - Sreactants |
| How would you calculate the entropy change of surroundings? | -∆H/T |
| How would you calculate the total entropy change? | ∆Stotal = ∆Ssystem + ∆Ssurroundings |
| What is the entropy change of a spontaneous reaction? | Positive |
| What is the equation used to find the free energy change of a reaction? | ∆G = ∆H - T∆Ssystem |
| How can you tell how feasible a reaction is? | The more negative ∆G is, the more feasible the reaction |
| When ∆H is negative and ∆S is positive, how feasible is the reaction? | The reaction will always occur |
| When ∆H is negative and ∆S is negative, how feasible is the reaction? | If ∆H > T∆S, the reaction will be feasible |
| When ∆H is positive and ∆S is positive, how feasible is the reaction? | If T∆S > ∆H, the reaction will be feasible |
| When ∆H is positive and ∆S is negative, how feasible is the reaction? | The reaction will not occur |
| What equation can be used to find the equilibrium constant of a reaction using Gibb's Free Energy? | K = e^(-∆G/RT) |
| Which direction would a reaction go if the equilibrium constant was greater than 1? | Products favoured |
| Which direction would a reaction go if the equilibrium constant is smaller than 1? | Reactants favoured |
| Why would a reaction with a negative ∆G not occur? | It may require an activation energy |
| Define enthalpy change | The energy transferred in a reaction at a constant pressure in standard states and conditions |
| Define enthalpy of formation | The energy transferred when 1 mole of substance is formed from its elements in standard states and conditions |
| Define enthalpy of combustion | The energy transferred when 1 mole of substance is completely burned in oxygen in standard condition |
| Define enthalpy of neutralisation | The energy transferred when and acid and an alkali react to form 1 mol of water in standard conditions |
| How would you calculate enthalpy change? | q = mcΔt c = 4.18J/g/K |
| State Hess' law | The total enthalpy change of a reaction is always the same regardless of which route is taken |
| Define bond enthalpy | The energy required to break a bond in one mole of gaseous atoms |
| How would you calculate enthalpy change using bond enthalpies? | Sum of bond enthalpies of reactants - sum of bond enthalpies of products |
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