Definitions Flashcards

(31 cards)

1
Q

Ionic Bonding

A

Electrostatic attraction between positively and negatively charged ion.

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2
Q

Covalent Bonding

A

Electrostatic attraction between shared pair of electrons and bonding nuclei.

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3
Q

Metallic Bonding

A

Electrostatic attraction between positive metal ions and sea of delocalised electrons

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4
Q

Electronegativity

A

The tendency of an atom to attract to a pair of electrons in a covalent bond.

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5
Q

Enthalpy

A

H

A measure of the heat energy in a chemical system (atoms, molecules or ions)

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6
Q

Enthalpy Change

A

^H = H(products) - H(reactants)

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7
Q

Activation Energy

A

The minimum energy required for a reaction to take place.

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8
Q

Average Bond Enthalpy

A

The energy required to break one mole of a specific type of bond in a gaseous molecule.

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9
Q

Ionisation Energy

A

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.

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10
Q

Second Ionisation Energy

A

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.

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11
Q

Homogeneous Catalyst

Hetrogeneous Catalyst

A

The same physical state as reactants

Other physical states as reactants

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12
Q

Brownsted Lowry Acid

A

Proton donor

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13
Q

Browbstead Lowry Bad

A

Proton Acceptor

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14
Q

Buffer Solution

A

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.

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15
Q

Lattice Enthalpy

A

The enthalpy change that accompanies the formation of one mole of a ionic compound. Forms gaseous ions under standard conditions.

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16
Q

Enthalpy change of Neutralisation?

A

The energy change that accompanies the reaction of and acid by a base to form one mole of H20 (l), under standard conditions.

17
Q

Enthalpy change of Combustion?

A

The enthalpy change that take place when one mole of a substance reacts completely with oxygen, under standard conditions.

18
Q

Enthalpy change of Formation?

A

Enthalpy change that takes place when one mole of a compound is formed from its elements, under standard conditions.

19
Q

Intermolecular forces:
Permanent dipole- induced dipole

Permanent dipole- permanent dipole

Instantaneous- induces dipole

Hydrogen bonding

A

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

20
Q

Halogen Group?

A

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

21
Q

Qualities Analysis?

A

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

22
Q

Le Chateliers Principle

A

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

23
Q

Kc?

A

Kc = [C]c[D]d / [A]a[B]b

Kc = 1 equilibrium halfway
Kc > 1 equilibrium towards products
Kc < 1 equilibrium towards reactants

24
Q

Arrhenius Equation?

A

K = Ae^-Ea/RT

Log form - LnK = -Ea/RT + LnA

25
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
26
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
27
pH and H+? Pka and Ka?
pH = -log(H+) (H+) = 10^-pH Pka = -log(Ka) (Ka) = 10^-Pa
28
Kw?
Ionic product of water Kw = [H+][OH] Kw = 1x10^-14
29
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
30
Calculating entropy change?
🔼S* = ES* (products) - ES* (reactants)
31
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