Physical chemistry

what is the atomic number of an atom?

number of protons (electrons) in the nucleus

what is the mass number of an atom?

number of protons and neutrons in an atom

Definition of an isotope?

• atoms with the same number of protons but different number of neutrons and different masses/

• Same electronic configuration so same chemical reactivity

• Atoms of different isotopes of the same element vary in mass number because of the different number of neutrons

what is the number of electrons in each shell/ main energy level?

2n²

How does the time of flight mass spectrometer work?

• sample of element placed in sample chamber

• Atoms then go through ionisation, where all atoms are converted into positive ions

• Positive ions are attracted to negatively charged plate, causing atoms to accelerate and increase in KE

• Once ions pass through negatively charged plate, they start drifting towards detector down the chamber

• Lighter ions drift at greater velocity (faster)

• As the positive ion passes the detector, it gains an electron

• Transfer of electron causes current to flow

Why do atoms in a TOF mass spectrometer travel at the same speed?

• atoms of the same charge have the same kinetic energy

what can we tell from a TOF mass spectrometer?

• time taken to move down drift chamber = mass of isotope

• Size of current produced when each group of isotope hits detector = abundance of isotope

Condition of TOF mass spectrometer:

• vacuum so ions do not collide with molecules in the air

What is the m/z ratio?

• ratio of mass of each ion to its charge

• Relative atomic mass

What is the order of energy levels for 32 electrons?

1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁶

definition of relative atomic mass:

weighted average mass of an atom of an element, taking into account its naturally occurring isotopes, relative to 1/12 the relative atomic mass of an atom of carbon-12

definition of relative molecular mass:

• Mass of that molecule compared to 1/12 of the relative atomic mass of an atom of carbon-12

what is the ideal gas equation?

pV = nRT

How to convert into SI units:

• 1kPa =1000 Pa

• 100cm³ = 100 × 10^-6m³

• 10 degrees celsius = 10 + 273 K

Percentage atom economy equation:

molecular mass of desired product/sum of molecular masses of all reactants x 100

define ionic bonding:

electrostatic force of attraction between oppositely charged ions in a latice

define a covalent bond:

Electrostatic force of attraction between a shared pair of electron and both nuclei

define dative bond:

shared pair of electrons with both electron supplied from one atom

Properties of metal:

• malleable - can be beaten into shape

• Ductile - they can be pulled into thin wires

what are the different types of crystal structure (with examples)?

• ionic (sodium chloride)

• Metallic (magnesium)

• Macromolecular/ giant covalent (diamond and graphite)

• Molecular (iodine)

Explain simple molecular crystals:

• small molecules held in a regular array by weak intermolecular forces

• Covalent bonds within the molecule holds atom together but doesn’t act between the molecules

• low melting point

structure of diamond:

• carbon covalently bonded to four carbons

• Macromolecular structure, giant covalent

• Each carbon forms tetrahedral shape

• Hard, high melting point

structure of graphite:

• carbon covalently bonded to three other carbons

• Layers of hexagonal carbons, held strong by covalent bonds

• Layers held by van de waals, weak forces

• Soft and has high melting pointstructure

• Delocalised electrons within each layer

Electronegativity definition:

ability for an atom to attract electron density towards itself in an covalent structure

2bp:

linear

180

3bp:

trigonal planar

120

2bp 1lp:

Bent vshape

118

4bp:

tetraheral

109.5

3bp 1lp:

trigonal pyramidal

107

2bp 2lp:

bent v-shape

104.5

5bp:

• trigonal bipyramidal

• 90, 120

4bp 1lp:

see saw

119, 89

3bp 2lp:

Trigonal planar

120

6bp:

octahedral

90

5bp 1lp:

square pyramid

89

4bp 2lp:

square planar

90

Define enthalpy change:

a heat change at constant pressure

definition for the standard molar enthalpy of formation:

enthalpy change associated with the formation of one mole of substance from its constituent elements under standard conditions with all products and reactants in standard states (+ve)

definition for the standard molar enthalpy of combustion:

enthalpy change associated with the complete burning of one mole of a substance in oxygen under standard conditions with all products and reactants in standard state (-ve)

different between heat and temperature:

• temperature is the average KE of the particles

• Heat is the TOTAL energy of all the particles

What is Hess’s law?

the enthalpy change for a chemical reaction is the same whatever route is taken from reactants to products

Lattice enthalpy of formation/ dissociation:

enthalpy change when one mole of a solid ionic compound is formed/ broken up into its constituent ions in the gas phase (f = +ve, d = -ve)

define bond dissociation enthalpy:

enthalpy change when one mole of covalent bonds is broken in the gaseous state (+ve)

define mean bond enthalpy:

energy required to break a given covalent bond averaged over a range of compounds

Chromium electron configuration:

1s2 2s2 2p6 3s2 3p6 3d5 4s1

copper electron configuration:

1s2 2s2 2p6 3s2 3p6 3d10 4s1

Define enthalpy of neutralisation:

enthalpy change associated with the formation of 1 mole of water in a reaction between an acid and an alkali under standard conditions (-ve)

Define ionisation enthalpy:

• first ionisation energy: the enthalpy change when one mole of electrons is removed from one mole of gaseous atoms to form one mole of gaseous +1 ions (+ve)

define electron affinity:

• enthalpy change when one mole of gaseous atoms loses one mole of electrons to form one mole of gaseous -1 atoms (-ve)

define enthalpy of atomisation:

enthalpy change when one mole of gaseous atoms is produced from an element in its standard state (+ve)

define hydration enthalpy:

enthalpy change when one mole of gaseous ions dissolve in water (-ve)

define enthalpy of solution:

enthalpy change when one mole of an ionic solid dissolves in enough water that the dissolved ions are well separated and do not interact with one another

Define enthalpy of vaporisation:

enthalpy change when one mole of liquid turns into gas (+ve)

define enthalpy of fusion:

enthalpy change when one mole of a solid is turned into a liquid (+ve)

How to test for group 1 cations:

• flame test

• Dip nichrome wire in HCl and sample

• Place over fire

• Li+ red

• Na+ yellow

• K+ lilac

How to test for Ca2+:

• flame test

• Red flame

• Add few drops of NaOH to make calcium hydroxide

• White precipitate, DOESN’T redissolve in excess NaOH

how to test for copper ions:

Cu2+

Flame test

Green

Add few drops of NaOH to make copper hydroxide

Blue precipitate

how to test for iron ions:

add few drops of NaOH to make iron hydroxide

Fe2+ = green

Fe3+ = orange/brown

How to test for aluminium ions:

add a few drops of NaOH to form aluminium hydroxide

White precipitate

Add excess NaOH

Re-dissolves

How to test for ammonium ions:

• add a few drops of NaOH to form ammonium hydroxide

• Bubbling and pungent odour (NH3 formed)

• Ammonia turns red litmus paper blue

ionic equation of metal and NaOH general formula:

M+ (aq) + OH- (aq) —> MOH(s)

how to test for carbonate ions:

• add HCl

• If bubbling, collect gas and bubble it through lime water

• Lime water turns foggy

test for sulfate ion:

• add HCl

• If no bubbles, add barium chloride

• White precipitate forms

Ionic equation of barium ion and sulfate ion:

Ba2+ (aq) + SO4 2-(aq) —> BaSO4 (s)

Test for halide ions:

• add nitric acid to remove carbonate ions

• Add silver nitrate (AgNO3)

• White ppt = chloride

• Cream ppt = bromide

• Yellow ppt = iodide