Definitions Flashcards
Ionic Bonding
Electrostatic attraction between positively and negatively charged ion.
Covalent Bonding
Electrostatic attraction between shared pair of electrons and bonding nuclei.
Metallic Bonding
Electrostatic attraction between positive metal ions and sea of delocalised electrons
Electronegativity
The tendency of an atom to attract to a pair of electrons in a covalent bond.
Enthalpy
H
A measure of the heat energy in a chemical system (atoms, molecules or ions)
Enthalpy Change
^H = H(products) - H(reactants)
Activation Energy
The minimum energy required for a reaction to take place.
Average Bond Enthalpy
The energy required to break one mole of a specific type of bond in a gaseous molecule.
Ionisation Energy
The energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous +1 ions.
Second Ionisation Energy
Energy required to remove one electron from each atom in one mole of gaseous +1 ions of an element to form one mole of gaseous +2 ions.
Homogeneous Catalyst
Hetrogeneous Catalyst
The same physical state as reactants
Other physical states as reactants
Brownsted Lowry Acid
Proton donor
Browbstead Lowry Bad
Proton Acceptor
Buffer Solution
a system that minimises pH changes when small amounts of an acid or base are added. They contain a weak acid and its conjugate base.
Lattice Enthalpy
The enthalpy change that accompanies the formation of one mole of a ionic compound. Forms gaseous ions under standard conditions.
Enthalpy change of Neutralisation?
The energy change that accompanies the reaction of and acid by a base to form one mole of H20 (l), under standard conditions.
Enthalpy change of Combustion?
The enthalpy change that take place when one mole of a substance reacts completely with oxygen, under standard conditions.
Enthalpy change of Formation?
Enthalpy change that takes place when one mole of a compound is formed from its elements, under standard conditions.
Intermolecular forces:
Permanent dipole- induced dipole
Permanent dipole- permanent dipole
Instantaneous- induces dipole
Hydrogen bonding
Molecule with permanent dipole induce dipole on non molar molecule
Attraction between two polar molecules
Random electron movement creates instantaneous dipole, that induces the neighbouring molecule
Requires a lone pair of electrons as electronegative attraction is created to exposed H nucleus
Halogen Group?
At room temp and pressure, exist as diatomic molecules. Group changes from gas to liquid to solid down the group.
Reactivity decreases down the group. As there is less attraction
Qualities Analysis?
1) Carbonate test - add dilute HNO3, bubbles, can then bubble through lime water
2) sulphate test - add BaNO3 form BaSO4 (s)
3) Halide test - add AgNO3
Chloride = white ppt, soluble in dilute NH3
Bromine = cream ppt, soluble in conc. NH3
Iodine = yellow ppt, insoluble in conc. NH3
Le Chateliers Principle
Conc - if more products are formed, equilibrium shifts to the right
Temp - if endothermic, equilibrium to the right, take in heat energy, minimise increase in temp
Pressure - only for gas, depends on no. Of molecules on each side, e.g. if more molecules on LHS, increase pressure will mean shift to the right with fewer molecules
Kc?
Kc = [C]c[D]d / [A]a[B]b
Kc = 1 equilibrium halfway
Kc > 1 equilibrium towards products
Kc < 1 equilibrium towards reactants
Arrhenius Equation?
K = Ae^-Ea/RT
Log form - LnK = -Ea/RT + LnA
Partial pressure and mole fraction?
Mole fraction = number of moles of A / total number of moles in a gas mixture
Partial pressure = mole fraction of A X total pressure
Kp?
Kp = p[HI]^2 / p[H2] x p[I2]
Exothermic reaction - K will decrease with increasing pressure, raising temp would decrease yield.
This is because the ratio of p[HI]^2 / p[H2] x p[I2] will be greater than Kp, for this to be proportional equilibrium shift to left so [HI]^2 hast to decrease and H2 and I2 has to increase, create equal ratio.
Endothermic - Kp increase with increasing temp, equilibrium shift to the right, increasing yield
pH and H+?
Pka and Ka?
pH = -log(H+)
(H+) = 10^-pH
Pka = -log(Ka)
(Ka) = 10^-Pa
Kw?
Ionic product of water
Kw = [H+][OH]
Kw = 1x10^-14
Entropy?
The greater the entropy, the greater the dispersal of energy and the greater the disorder.
- at 0K no energy to disperse, entropy will be 0
- above 0K energy disperses, substances have positive entropy
• Standard Entropy always POSITIVE
Calculating entropy change?
🔼S* = ES* (products) - ES* (reactants)
Gibbs Free energy?
🔼G, overall change in energy during chemical reactions
🔼G = 🔼H - T🔼S
T = 🔼H/🔼S
🔼G < 0 in order to be feasible
