PHYSICAL CHEMISTRY YEAR 1

Chapter 1 + 8

Atomic structure and trends in period 3

nuclear model vs plum pudding model

1. nuclear model has a nucleus whereas plum pudding model has no nucleus

2. nuclear model has electron orbiting nucleus whereas plum pudding model has electrons scattered inside positive cloud

3. nuclear model has dense positive charge at centre of atom whereas plum pudding mass evenly spread throughout

4. nuclear model has empty space whereas plum pudding model doesn’t

relative mass of electron

1/1840

describe the current atomic structure

• nucleus made up of protons and neutrons

• protons and neutrons held together by strong nuclear force

• electrons orbit nucleus

• strong electrostatic force of attraction between nucleus and outermost electrons

mass number definition

the number of protons and neutrons in an atom

Relative atomic mass definition

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

relative molecular mass definition

the average mass of a molecule / (1/12) mass of one atom of carbon-12

isotopes definition

atoms of an element that have the same number of protons different number of neutrons

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

function of mass spectrometer

to determine the masses of separate atoms/molecules

6 steps of mass spectrometer

1. vacuum

2. ionisation

3. 3acceleration

4. ion drift

5. detection

6. data analysis

vacuum function

prevent molecule of the air colliding with ions formed during mass spectrometry

2 parts to ionisation

1. electrospray ionisation

2. electron impact

electrospray ionisation process (3 marks)

1. sample is dissolved in volatile solvent

2. sample injected through needle at high voltages

3. each particle / molecule gains a proton forming an ion

electron impact ionisation process (3 marks)

1. sample vaporised

2. high energy electrons fire from electron gun

3. an electron is knocked off each particle

acceleration process

• positive ions accelerate towards negatively charged plate

• lighter ions/highly charged ions travel faster

ion drift process

• all ions have same kinetic energy

• lighter ions travel faster so reach detector sooner

detection process

• 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

why is it necessary to ionise particles

1. so ions may be accelerated by an electric field

2. so ions may produce a current when hitting the detector allowing them to be detected

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

periodic table is the average of all 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 - has 3 orbitals

D orbital

holds up to electrons - has 5 orbitals

F orbital

holds up to 14 electrons - has 7 orbitals

why is the 4s shell before the 3d shell

as its of a lower energy level

when forming ions what shell is emptied first

4s as its of 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

have a full 3d shell instead of 4s as it more stable

explain the general trend across period 2/3 for IE

• as you go across period IE increases

• bigger nuclear charge

• similar shielding

• stronger electrostatic force of attraction between nucleus and outermost electron

Explain deviation 1 - period 3

• Al has lower IE than Mg

• outermost electron in 3p sub shell for Aluminium

• this is of a higher energy level

• so weaker e.f.o.a. and easier to remove

Explain deviation 2 - period 3

• S has a lower IE than P

• 2 electrons need to pair in an orbital of 3p sub shell

• repulsion between paired electrons so makes it easier to remove one

Trends in IE down a group

• IE decreases as you go down a group

• more shells so larger atomic radius 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

which element has highest melting point in period 3

silicon

• Giant covalent structure

• lots of energy required to overcome strong covalent bonds between silicon molecules

• so high melting points

state and explain the period 3 element with the highest first ionisation energy

• Argon

• biggest nuclear charge

• similar shielding

predict the element in period 3 which has the highest second ionisation energy

• sodium

• electron removed from 2p orbital

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

metals and non-metals

covalent bond definition

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

covalent bonding occurs between…

non-metal atoms only

molecule definition

a small group of atoms which are covalently bonded together

define a dative covalent bond

when one atom donates its lone pair of electrons to 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 smaller 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

factors that affect electronegativity

1. atomic radius

2. nuclear charge

3. shielding

how does atomic radius affect electronegativity

1. as atomic radius decreases electronegativity increases

2. this is because stronger electrostatic force of attraction between nucleus and shared electrons

how does nuclear charge affect electronegativity

1. nuclear charge increases

2. similar shielding

3. smaller atomic radius

4. stronger electrostatic force of attraction between nucleus and shared electron

5. increased electronegativity

how does shielding affect electronegativity

1. increase in shells means more shielding

2. larger atomic radius

3. so weaker electrostatic force of attraction between nucleus and shared electrons

4. so decreased electronegativity

TREND - electronegativity across a period

1. as you go across a period electronegativity increases

2. nuclear charge increases

3. similar shielding

4. smaller atomic radius

5. stronger electrostatic force of attraction between nucleus and shared electron

TREND - electronegativity down a group

1. electronegativity decreases as you go down a group

2. each element has an extra shell

3. so more shielding

4. so larger atomic radius

5. so weaker electrostatic force of attraction between nucleus and shared electron

polarity definition

the uneven distribution 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 delta-

3. so permanent dipole formed

name 3 types of intermolecular forces strongest to weakest

1. hydrogen bonds

2. permanent dipole-dipole forces

3. van der waals forces

how do permanent dipole-dipole forces work

• difference in electronegativity leads to bond polarity

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

how does the strength of dipole-dipole forces increase

1. the greater the difference in electronegativity the more polar the molecule

2. the stronger the partial charges

how does hydrogen bonding occur

1. high difference in electronegativity between hydrogen and other element

2. so produces a permanent dipole with delta + hydrogen and delta- element

3. lone pair on element forms hydrogen bond with delta+ hydrogen

how do van der waals forces occur?

• 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 the strength of VDW forces increase?

1. the larger the number of electrons the larger the instantaneous dipole

2. so greater partial charges

3. stronger VDW

how do lone pairs affect bonding angles

• lone pairs cause extra repulsion

• reduce bond angles by 2.5*

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

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

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

Chapter 4 + 17

Energetics and Thermodynamics

Exothermic reaction definition

heat energy has been released to the surroundings

Endothermic reaction definition

heat energy has been absorbed from the surroundings

Enthalpy change definition

the measure of heat content of a substance under constant pressure

what are the standard conditions for enthalpy change:

1. pressure → 100kPa or 1 atm

2. temperature → 298K or 25 degrees celsius

3. elements are in standard states

4. any solution will have a concentration of 1 mol dm^-3

in exothermic products have __________ energy than reactants

less - as heat energy is released to the surroundings

in endothermic reactions products have _______ energy than the reactants

more as heat energy is absorbed from the surroundings

enthalpy change in exothermic reactions

negative enthalpy change

• the energy required to break bonds is greater than the energy required to make bonds

enthalpy change in endothermic reactions

positive enthalpy change

• the enthalpy change required to break bonds is lesser than the energy required to make bonds

name 4 factors of enthalpy change:

1. temperature

2. pressure

3. concentration of solution

4. physical states

bond energy changes

the energy required to break a bond between 2 atoms

enthalpy change formula

reactants enthalpy - products enthalpy

bond breaking is an __________ reaction

endothermic - energy is needed to break bonds s absorbed from the surroundings

bond making is an _______________ reaction

exothermic - energy is released to the surroundings when bonds are made

standard molar enthalpy change of formation definition

the enthalpy change when one mole of compound is formed under standard conditions from elements in their standard states