MOR Chemistry Module 2

isotopes

atoms of the same element with different neutron numbers + different mass numbers

relative atomic mass

Ar
weighted mean mass of an atom of an element, relative to 1/12th mass of a carbon-12 atom

why may the Ar of a sample of an element be different from the Ar on the periodic table?

sample may not accurately reflect the relative abundances of all isotopes - the periodic table may use a different sample

example: higher abundance of 1 isotope will influence average mass/Ar

relative isotopic mass

mass of an isotope of an element, relative to 1/12th mass of a carbon-12 atom

relative molecular mass

weighted mean mass of a molecule, relative to 1/12th mass of a carbon-12 atom

relative formula mass
calculation

sum of all RAMs of elements, as given in a formula of non-molecular compounds

calculation: sum up all RAMs of elements

relative mass of an electron

1/1835

relative atomic mass calculation

see image

acid

proton donor (H+)

alkali

A base that is soluble in water

releases OH- ions in solution

base

can neutralise an acid to produce a salt

proton acceptor

disproportionation reaction

same element is simultaneously oxidised and reduced in the same reaction

aufbau principle

lowest energy levels are occupied first

pauli exclusion rule

each orbital cannot contain more than 2 electrons, with opposing spins

hund's rule

single electrons occupy all empty orbitals first, before forming pairs

because electrons repel

4 common acids

nitric acid HNO3
ethanoic acid CH3COOH
sulfuric acid H2SO4
hydrochloric acid HCl

4 common bases

ammonia NH3
potassium hydroxide KOH
sodium hydroxide NaOH
+ all carbonates (e.g. Na2CO3)

strong acid

more H+ ions fully dissociate and completely ionise

more ions left in solution
less molecules left in solution

lower pH

weak acid

less H+ ions fully dissociate and only partly ionise

less ions left in solution
more molecules left in solution

higher pH

neutralisation reaction

H+ ion in the acid is replaced by metal ion or NH4 + ion

acid + base --> salt + water

example: HCl + NaOH --> NaCl H2O

metal carbonate + acid

salt + water + carbon dioxide

metal + acid

salt + hydrogen

acid + base

salt + water

neutralisation ionic equation

H+ (aq) + OH- (aq) >>> H2O (l)

how is a salt formed?

when H+ ion is replaced by a metal ion or NH4 + ion

oxidation

loss of electrons
oxidation number increases

reduction

gain of electrons
oxidation number decreases

describe the currently accepted model of the atom

electrons share fixed energies

electrons move around nucleus in energy levels/shells

what is principal quantum number

each energy shell is assigned a principal quantum number (n)

how does principal quantum number change?

further away from nucleus = larger principal quantum number

s-subshell

1 orbital, 2 electrons

p-subshell

3 orbitals, 6 electrons

d-subshell

5 orbitals, 10 electrons

f-subshell

7 orbitals, 14 electrons

shell 1

1s = 2 electrons

shell 2

2s, 2p = 8 electrons

shell 3

3s, 3p, 3d = 18 electrons

shell 4

4s, 4p, 4d, 4f = 32 electrons

s-orbital shape

spherical

p-orbital shape

dumbell shape

maximum number of electrons in an orbital

2 electrons

what is an orbital?

space an electron moves in
pair of 2e- in an orbital have opposing spins

what are opposing spins

pair of electrons spin in opposite directions (spin-pairing)

2 electrons can only occupy the same orbital if they have opposing spins

subshell notation

example in image

box and arrow model

each box = 1 orbital
each arrow = 1 electron

up/down directions of arrows represent opposing spins of electrons

electron energy level diagrams

shows energies of electrons in different orbitals

expresses:
- number of electrons
- electron arrangement

order of filling

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p

ionic bond

the electrostatic attraction between oppositely charged ions

dot and cross diagram example

image

giant ionic lattice

regular structure
'giant' because the same basic units are repeated over and over

why do giant ionic lattices form?

every ion is electrostatically attracted to oppositely charged ions in all directions

electrical conductivity of ionic compounds

conduct electricity when molten/in solution
cannot conduct as solids, because ions are in fixed positions (due to strong electrostatic forces), so cannot carry a charge

melting/boiling points of ionic compounds

high melting/boiling points, because giant ionic lattices are held together by strong electrostatic forces

a lot of energy is required to overcome these forces

solubility of ionic compounds

usually dissolve in water
water is polar (partly charged)
water molecules PULL ions away from the lattice = dissolves

define a compound

2+ atoms bonded together

covalent bond

the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms

single covalent bond

only 1 pair of electrons is shared
1e- is donated to bonding pair

double or triple covalent bonds

2 or 3 pairs of electrons are shared between atoms

describe a 'stable arrangement'

outer shell USUALLY filled to 8 electrons (octet)

+ exceptions exist

dative (coordinate) covalent bond

1 atom provides both electrons in the shared pair

expressing dative covalent bonds in diagrams

image

special exceptions to octet in outer shell

some compounds use d-orbitals to 'expand their octet' (can contain more than 8 electrons in outer shell) (e.g. SF6)

some compounds have less than 8 electrons in their outer shell (e.g. BF3)

radical

contains a single unpaired electron
very reactive

determining covalent bond strength

average bond enthalpy
stronger bond = more energy needed to break bond = higher bond enthalpy value

linear

2 bonding regions
no lone pairs
180 degrees

trigonal planar

3 bonding regions
no lone pairs
120 degrees

trigonal pyramidal

3 bonding regions
1 lone pair
107 degrees

tetrahedral

4 bonding regions
no lone pairs
109.5 degrees

trigonal bipyramidal

5 bonding regions
no lone pairs
120 + 90 degrees

non-linear

2 bonding regions
2 lone pairs
104.5 degrees

octahedral

6 bonding regions
no lone pairs
90 degrees

repulsion of lone pairs

repel more

repulsion of bonded regions

repel equally

drawing shapes of molecules

image

repulsion in shapes of molecules

pairs of electrons repel
repel to become as far from each other
assumed shape minimises repulsion

electronegativity

the ability of an atom to attract electrons in a covalent bond

what does a dipole contain?

both δ+ and δ- ends

if both ends were the same delta charge, then they would cancel out = no dipole

trend in electronegativity

increases as we move towards F on Periodic Table

why is CO2 non-polar?

symmetrical
both ends have the same dipole (δ-) = cancels out
polar bonds point away from each other

why is H2O polar?

asymmetrical
contains both δ+ and δ- ends = do not cancel out
polar bonds point in the same direction as each other

what do van der Waals' forces consist of? (2)

induced dipole-dipole interactions

permanent dipole-dipole interactions

induced dipole-dipole interactions

1) electron cloud is mobile = uneven distribution of electrons

2) temporary dipole = induces dipoles in neighbouring molecules

2) dipole δ+ and δ- ends ATTRACT each other = form dipole-dipole interactions

induced dipole-dipole interactions TREND

larger molecule → more electrons → stronger dipole-dipole interactions → higher boiling point

permanent dipole-dipole interactions

stronger than induced

1) difference in electronegativity leads to a PERMANENT dipole + polar bonds

2) attraction between δ+ and δ- ends → WEAK IMF formed

define redox reaction

reduction and oxidation occur simultaneously
no net gain/loss of electrons

hydrogen bonds

strongest intermolecular force

3 molecules which form hydrogen bonds

HF
H2O
NH3

why do HF/H2O/NH3 form hydrogen bonds?

difference in electronegativity between H atoms and very electronegative N/F/O

properties of water

ice is less dense than liquid water
H2O molecules are arranged further apart in a lattice structure = floats

unusually high boiling point
due to hydrogen bonds

surface tension
due to strong H bonds on surface

half equations

example: oxidation of magnesium

Mg --> Mg 2+ + 2e -

reducing agent

is oxidised (donates electrons)

reduces other

oxidising agent

is reduced (accepts electrons)

oxidises other

name the 3 molecules which form hydrogen bonds

H2O
HF
NH3

H is not very electronegative
O, F and N are very electronegative
= large difference in electronegativity

questions on boiling point guide e.g. why does water have a higher boiling point than ethane?

1) identify IMFs in molecule

2) compare strengths of IMFs

3) amount of energy required to overcome (more/less)

questions on shapes of molecules + repulsion guide e.g. explain why ammonia has a trigonal pyramidal shape

1) name of molecule's shape

2) describe structure - lone pairs, bonding pairs (+ angle if needed)

3) describe repulsion