PHYSICAL CHEMISTRY - AS LEVEL MOCKS

Chapter 1 + 8

atomic structure and periodicity

Rutherford Model - nuclear model vs current day model

current model has both neutrons and protons whereas nuclear model doesn’t

nuclear model has electrons surrounding nucleus whereas current model has electrons in different shells orbiting nucleus

nuclear model vs plum pudding model

• nuclear model has a nucleus whereas plum pudding model has no nucleus

• nuclear model has electrons orbiting nucleus whereas plum pudding model has electrons scattered in positive cloud

• nuclear model has dense positive mass at the centre (nucleus) plum pudding model has evenly spread mass

• nuclear model has empty space whereas plum pudding model doesn’t

relative mass of electron

1/1840

arrangement of atom

• protons and neutrons held together in nucleus via strong nuclear force

• electrostatic force of attraction between nucleus and orbiting electrons

isotopes definition

atoms with same number of protons different number of neutrons

mass number of atom definition

number of protons and neutrons in atom

proton number of atom definition

number of electrons/protons in atom

Ar definition

average mass of an atom of an element /(1/12) mass of one atom of C12

Mr defintion

average mass of a molecule / (1/12) mass of one atom of C12

function of mass spectrometer

to determine the masses of separate atoms/molecules

what are the steps for mass spectrometry:

vacuum, ionisation, acceleration, ion drift, detection, data analysis

vacuum

prevent molecules from air colliding with ions formed during mass spectrometry

what are the two parts to ionisation

• electrospray ionisation

• electron impact

electrospray ionisation

sample is dissolved in volatile solvent and injected in fine hypodermic needle forming mist

tip of needle attached to positive terminal at high-voltage power supply

each particle gains a proton and becomes ionised

equation for electrospray ionisation

X(g) + H^+ → XH^+

electron impact

• sample vaporised

• high energy electrons fired from electron gun

• knocks of electron of each particle forming uni-positive ion

why is it necessary to ionise particles during mass spectrometry

• ions will interact with and be accelerated by an electric field

• ions can create a current when hitting the detector

Acceleration

• positive ions accelerate towards negatively charged plate

• lighter ions/highly charged ions travel faster

Ion Drift

• all ions have same kinetic energy

• lighter ions travel faster so reach detector sooner

Detection

• positive ions accelerate towards negatively charged plate

• positive ion gains electron from plate and is discharged

• generates movement of electrons causing TOF to be measured

• ion abundance proportional to size of current

Data Analysis

• signal from detector sent to computer to be analysed

why might the Ar calculated from mass spectrometry be different to the periodic table

the Ar in the periodic table takes into account different isotopes

what are the 4 sub-shells

S,P,D,F

S orbital

holds up to 2 electrons - 1 orbital

P orbital

holds up to 6 electrons - 3 orbitals

D orbital

holds up to 10 electrons - 5 orbitals

F orbtial

holds up to 14 electrons - 7 orbitals

state the full electron arrangement

1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2

why is the 4s shell before the 3d shell

as it is of a lower energy level

when forming ions what shell is emptied first

4s as it is off a lower energy level than 3d

name a property and its function for electrons

spin —> used to overcome repulsion between electrons in same orbital

what are the two exceptions in electron notation and what is the exception

Cu and Cr → have a full 3d shell instead of 4s as it more stable

Ionisation energy definition

the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous unipositive ions

equation for nth ionisation energy:

X^(n-1)+(g) → Xⁿ⁺ (g) + e^-

explain the general trend across period 2/3 for IE

• as you go across period generally IE increases

• as nuclear charge increases and all atoms have similar shielding

• so stronger electrostatic force of attraction between nucleus and outermost electron

• so higher IE

explain deviation (1) in periods 2/3 for IE

1. for period 2 - B/Be

2. for period 3 - Al/Mg

• (B/Al) lower than (Be/Mg)

• as outermost electron in new (2/3)p sub-shell

• this is more shielded and of a higher energy level than 2s so easier to remove an electron

• so lower IE

explain deviation (2) in periods 2/3 for IE

1. for period 2 - N/O

2. for period 3 - P/S

• (O/S) lower than (N/P)

• as 2 electrons in (2/3)p need to pair

• pairing causes repulsion between electrons so easier to remove an electron

• so lower IE

Trends in IE down a group

• generally IE decreases as you go down a group

• more shells as you go down

• so more shielding as you go down

• so weaker electrostatic force of attraction between nucleus and outermost electron

• so easier to remove an electron

how do chemical properties of isotopes differ -if they do?

they don’t differ as they have the same electron configuration

why is the ionisation energy of every element endothermic

energy needed to overcome the electrostatic force of attraction between positive nucleus and negative electrons

why does atomic radius decrease across a period?

• nuclear charge increases across a period

• similar shielding

• so strength of electrostatic force of attraction between nucleus and outermost electrons increase

• atomic radius decreases

why does melting point increase from sodium to aluminium

• as ionic charge increases

• so smaller ions and more delocalised electrons

• so more energy required to overcome stronger electrostatic force of attraction between nucleus and sea of delocalised electrons

CHAPTER 3

bonding and structure

ionic bonding definition

the electrostatic force of attraction between oppositely charged ions that extends in every direction throughout compound

ionic bonding occurs between…

non-metals and metals

how does ionic bonding occur:

• electrons are transferred from metal atoms to non-metal atoms forming ions

• electrostatic force of attraction between oppositely charged ions

covalent bond definition

electrostatic force of attraction between nucleus and shared electrons in covalent bond

covalent bonding occurs between…

non-metal atoms only

molecule definition

small group of covalently bonded atoms

dative covalent bonding definition

when one atom donates its lone pair of electrons with another atom to form a covalent bond

metallic bonding definition

the electrostatic force of attraction between positive metal ions and sea of delocalised electrons

why is the melting point of copper chloride lower than copper iodide

• chloride ion is smaller

• so weaker efoa to copper ion

electronegativity definition

the power of an atom to attract the shared pair of electrons in a covalent bond

electron density definition

the way electrons are distributed within a cloud

factors that affect electronegativity:

1. atomic radius

2. nuclear charge

3. shielding

how does atomic radius affect electronegativity:

1. as atomic radius decreases electrostatic force of attraction between nucleus and outermost electrons increase

how does nuclear charge affect electronegativity:

1. increase in nuclear charge

2. so similar shielding

3. smaller atomic radius

4. so increase in electrostatic force of attraction between nucleus and shared pair fo electrons

5. so increased electronegativity

how does shielding affect electronegativity:

1. more shells means more shielding

2. so electrostatic force of attraction decreases between nucleus and outer electron

3. so decreased electronegativity

TREND - why does electronegativity increase across a period

• nuclear charge increases as you go across

• similar shielding

• stronger electrostatic force of attraction between nucleus and outermost electron

• smaller atomic radius

TREND - why does electronegativity decrease as you go down a column

• each element has an extra shells so more shielding

• increased atomic radius

• weaker electrostatic force of attraction between nucleus and outermost electron

4 most electronegative elements:

F,Cl,O,N

polarity definition

the unequal sharing of electrons between atoms that are covalently bonded together

how does a bond become a permanent dipole:

1. one atom is more electronegative than the other

2. so one atom delta+ the other is delta-

3. so permanent dipole formed

name the 3 types of intermolecular forces from strongest to weakest:

1. Hydrogen bonds

2. Permanent-Dipole-Permanent Dipole

3. Van Der Waals Forces

how do permanent dipole-dipole forces arise

• difference in electronegativity leads to bond polarity

• attraction between delta+ side on one molecule and delta- side on another

how does hydrogen bonding occur?

1. high difference in electronegativity between Hydrogen and other element

2. so produces permanent dipole with delta+ and delta- side

3. bond forms between lone pair on electronegative element and delta+ hydrogen

state one condition for hydrogen bonding to occur

Hydrogen bonded to highly electronegative element

how does strength of dipole-dipole forces increase?

• the greater the different in electronegativity the more polar the molecule

• and the stronger the forces

how do van der waals forces work?

• constant movement of electrons forms instantaneous dipole

• delta+ side of dipole attract delta- side of nearby molecules

• instantaneous molecule induces a dipole

• molecules are symmetrical

• the force between them is VDW

how does strength of van der waals forces increase?

• the larger the number of electrons the larger the instantaneous dipole

• so greater partial charges

• stronger VDW

electron pair repulsion theory definition

• electron pairs arrange themselves as far apart as possible to minimise repulsion

how do lone pairs affect bond angles:

• lone pairs cause extra repulsion

• reduce bond angles by 2.5*

giant ionic lattice structure

• repeating pattern between oppositely charged ions

• uniform arrangement

• cubic structure

Giant Ionic Lattice melting points:

• high amount of energy required to overcome strong electrostatic forces of attraction between oppositely charge ions

Giant Ionic Lattice conductivity:

• able to conduct electricity when molten or aqueous only

• as charged ions able to move throughout structure and carry charge

Simple molecular structure melting points

• low melting points

• little energy required to overcome weak intermolecular forces between molecules

Simple molecular structure conductivity

• unable to conduct electricity

• no freely charged particles able to carry charge throughout structure

Metals melting point

• high melting point

• lots of energy required to overcome strong electrostatic forces of attraction between positive metal ions and sea of delocalised electrons

Metals conductivity (electricity)

• able to conduct electricity

• sea of delocalised electrons carry charge throughout structure

metals conductivity (heat)

• high heat conductivity

what does the strength of metals depend on:

• charge of ion → greater the charge of ion, the greater the delocalised electrons so stronger efoa

• size of ion → smaller the ion the stronger the efoa

TREND → metal boiling points across a period

• increases

• as ion charge increases

• so more delocalised electrons

• so stronger electrostatic force of attraction between positive ions and sea of delocalised electrons

• which requires more energy to overcome

TREND → metal boiling points down a column

• decreases

• increase in size of ion

• so weaker electrostatic force of attraction between positive nucleus and sea of delocalised electrons

• which requires less energy to overcome

metals ductile property

layers of metal ions able to slide over eachother

Macromolecular structure melting points

requires a lot of energy to overcome strong covalent bonds

so high melting point

Macromolecular structure conductivity

• unable to conduct electricity

• no freely charged particles able to carry charge throughout structure

Diamond Properties:

very hard → as carbon covalently bonded to 4 other carbon atoms, so layers unable to slide over each other

high melting points as lots of energy required to overcome strong covalent bonds

Doesn’t conduct electricity as no freely charged particles to carry it

Diamond Structure:

• each carbon forms 4 covalent bonding

• 4 bonding electron pairs repel each other equally

• 109.5

• tetrahedral

• giant macromolecular structure

diamond structure:

• each carbon forms 4 covalent bonding

• 4 bonding electron pairs repel each other equally

• 109.5

• tetrahedral

• giant macromolecular structure

Graphite Properties:

• soft and flaky → weak van der waals between planes allowing them to slide over each other

• high melting points → due to strong covalent bonding and

• conducts electricity due to sea of delocalised electrons

TREND → boiling point from HCl to HI

• increases

• strength of VDW increases as number of electrons increases

• so requires more energy to overcome

why doesn’t HF follow this trend

• HF is able to hydrogen bond

• so has stronger IMF which requires more energy to overcome

• so higher boiling point than HCl

why does AlCl3 not ionic bond

to little difference in electronegativity

SiO2 structure

• tetrahedral

• 109.5

• giant macromolecular structure

which ion is smaller Na^+ or Mg²+:

• Mg2^+

• as greater nuclear charge/more protons

• same shielding between Na and Mg

• so stronger electrostatic force of attraction

Dynamic equilibrium definition

the rate of the forward reaction and the reverse reaction are equal, concentrations of the products and reactants remain constant

macroscopic properties definition

properties that don’t depend on the total quantity of matter

le chateliers principle definition

if a change occurs to the equilibrium system, position of equilibrium shifts to counteract the change