physical chemistry

atomic number

the number of protons in the nucleus

mass number

the total number of protons and neutrons in the nucleus

isotope

atoms of the same element with the same number of protons and electrons but a different number of neutrons

relative atomic mass

the average mass of an atom compared to 1/12 of a carbon-12 atom

relative formula mass

the sun of the relative atomic masses of atoms in a molecule

relative molecular mass

the average mass of a molecule of an element or compound relative to 1/12 of a carbon-12 atom

avogadro’s number

6.022×10²³

stages of mass spectrometry

vaporisation, ionisation, acceleration, ion drift, detection

electron impact

bombarded with high energy electrons, forms M^+ ions

electrospray ionisation

sprayed through a hypodermic needle with large voltage, forms MH^+ ions

acceleration stage of tof

all ions are accelerated such that they have the same kinetic energy. ions are accelerated by a negatively charged acceleration plate. all have same kinetic energy

ion drift

lighter ions travel faster and meet the detector first. heavier ions travel slower

detection stage of tof

ions hit the detector and gain an electron. the more abundant a particular ion, the greater the current generated

avogadros number meaning

the number of ions present in a sample with a mass equal to its relative formula mass

mass of one ion equation

mass = relative isotopic mass x 10^-3 / 6.022×10²³

kinetic energy equation

KE = ½mv²

velocity equation simple

v= d/t

velocity equation combing kinetic energy

v = √m/2KE

tof comparing different isotopes equations

mv = mv

m/t² = m/t²

calculation for max number of electrons in energy level

2n² where n is the energy level

number of orbitals in the s subshell

1

shape of s orbital

spherical

number of orbitals in p subshell

3

shape of p orbital

dumbbell shaped

number of orbitals in d subshell

5

number of orbitals in f subshell

7

the Aufbau principle

electrons always fill up the lowest energy subshell first

hund’s rule

electrons will always go into empty orbitals before pairing because they repel each other

the Pauli exclusion principle

there is a maximum of two electrons with opposite spin in one orbital

electronic configuration of copper

[Ar] 4s^1 3d^10

electronic configuration of chromium

[Ar] 4s^1 3d^5

isoelectronic

atoms or ions of different elements that have the same electronic configuration

are 4s or 3d electrons lost first?

4s because it is a higher energy level

enthalpy change of 1st ionisation energy

when 1 mole of gaseous atoms lose 1 mole of electrons to form 1 mole of gaseous 1+ ions

factors that affect ionisation energy

1. atomic radius - greater distance - less attraction

2. nuclear charge - greater charge - greater attraction

3. electron shielding - greater shielding - less attraction

trend in ionisation energy down the group

1. greater radius - loses electrons more easily

2. greater charge - more attraction

3. more shielding - less attraction

trend in ionisation energy across a period

1. smaller radius - greater attraction

2. greater charge - more attraction

3. same shielding - neutral effect

number of moles equation

n = mass / Mr

empirical formula

simplest whole number ratio of atoms in a compound

percentage yield

how much product was produced compared to the expected amount

reasons for percentage yield of less than 100%

• impure reactants

• reversible reaction

• side reactions

• loss of mass during transfer

percentage yield calculation

(actual yield / theoretical yield) x 100

atom economy

the measure of the amount of useful products in relation to the reactants

atom economy calcualtation

(total mass of useful products / total mass of reactants) x 100

number of moles calculation with liquids

n = cv

ideal gas equation

PV = nRT

assumptions made for ideal gas equation to work

• all the particles are perfect spheres

• there are no attractive forces between the particles

• all collisions between particles are elastic so kinetic energy stays the same

ionic bonding

the electrostatic force of attraction between oppositely charged ions

factors that determine strength of an ionic bond

1. ionic radius - smaller radius - stronger

2. ionic charge - greater charge - stronger

covalent bond

the electrostatic force of attraction between a pair of shared electrons and a positive nucleus

factor that determines the strength of a covalent bond

bond length - larger atoms - longer bonds- weaker

dative covalent bond

one atom supplies both electrons to form a covalent bond

conditions for a dative covalent bond to form

1. one atom must have a vacant orbital

2. another atom must have a lone pair of electrons

dimer

when two of the same small molecule join together to form a larger molecule via dative covalent bonding

metallic bonding

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

factors that affect the strength of metallic bonds

1. ionic radius

2. ionic charge

valence shell electron pair repulsion theory

• bonding pairs repel other bonding pairs equally and as much as possible

• lone pairs repel other lone pairs equally and as much as possible

• lone pairs repel more than bonding pairs

way to find base shape of molecule

valence electrons of central atom + number of bonds formed all divided by 2

bond angles with lone pairs

take 2.5 degrees away for every lone pair

polar bonds

ionic character of covalent bonds

electronegativity

the ability of an atom to withdraw electron density towards itself in a covalent bond

factors affecting electronegativity

1. nuclear charge

2. atomic radius

trend in electronegativity across a period

• nuclear charge increases

• atomic radius decreases

  —> greater electronegativity

trend in electronegativity down a group

• atomic radius increases

• nuclear charge increases

  —> lower electronegativity

most electronegative atom

fluorine

requirement for ionic bonds with covalent character

1. a polarising cation - small highly charged cations which distort the shape of the anion

2. a polarisable anion - large anions

van der waal’s forces

electron movement in first molecule induces a dipole in another, resulting to a temporary attraction between d+ on one molecule and d- on another

factors that determine the strength of van der Waal’s forces

1. size of the molecule - greater size - greater imf

2. shape of the molecule - branched chains don’t pack as closely together - lower imf

factors that determine the presence of dipole-dipole interactions

1. presence of electronegative atoms resulting in polar bonds

2. shape of the molecule - the distribution of charge must be asymmetric

   the molecule must have an overall dipole moment

order of increasing strength of intermolecular forces

1. van der Waal’s forces

2. dipole-dipole interactions

3. hydrogen bonds

factors for hydrogen bonds to occur

1. a hydrogen atom must be directly bonded to one of the most electronegative elements eg. F, N, or O

2. the electronegative element must have a lone pair of electrons

way to draw hydrogen bonds

1. show all lone pairs of electrons

2. show all partial charges

3. linear arrangement of covalent bond and hydrogen bond

periodicity

the study of trends shown by the elements which repeat themselves in each period of the periodic table

oxidation

loss of electrons

reduction

gain of electrons

reducing agent

(itself is oxidised) loses electrons

oxidising agent

(itself is reduced) gains electrons / oxidises another atom or ion

neutral compound oxidation state

equals 0

name and bond angle of this shape

name: linear

bond angle: 180

name and bond angle of the shape

name: trigonal planar

bond angle: 120

name and bond angle of this shape

name: bent

bond angle: 117.5

name and bond angle of this shape

name: tetrahedral

bond angles: 109.5

name and bond angle of this shape

name: trigonal pyramidal

bond angle: 107

name and bond angle of this shape

name: bent

bond angle: 104.5

name and bond angles of this shape

name: trigonal bipyramidal

bond angle: 90,120

name and bond angles of this shape

name: sawhorse or seesaw

bond angles: 87.5, 117.5

name and bond angles of this shape

name: T shaped

bond angles: 90

name and bond angles of this shape

name: linear

bond angle: 180

name and bond angles of this shape

name: octahedral

bond angles: 90

name and bond angles of this shape

name: square pyramidal

bond angles: 87.5

name and bond angles of this shape

name: square planar

bond angles: 90

name and bond angles of this shape

name: T shape

bond angle: 87.5

name and bond angles of this shape

name: linear

bond angle: 180

Exothermic reaction

Releases energy to the surroundings

• energy of the products is less than the reactants

• Characterised by an increase in temperature

• Bond formation is exothermic

Endothermic reaction

Absorbs energy from the surroundings

• energy of the products is greater than the reactants

• Characterised by a decrease in temperature

• Bond breaking is endothermic

Enthalpy change (delta H)

Heat change at constant pressure

Standard conditions (°)

298K, 100kPa, 1.00 moldm-3

Standard enthalpy change of formation

Enthalpy change when 1 mole of a substance is formed from its constituent elements with all reactants and products in standard states under standard conditions

Standard enthalpy change of combustion

Enthalpy change when 1 mole of a substance is completely burned in oxygen with all reactants and products in standard states under standard conditions

Mean bond enthalpy

The enthalpy change required to break 1 mole of a covalent bond in gas phase, averaged over a range of compounds