electrons
negatively charged particles which move around the nucleus in shells
nucleus
central part of the atom with most the mass made up of protons and neutrons
mass number
larger number which is the total of protons and neutrons
atomic number
smaller number which is the number of protons or electrons
negative ions
have more electrons than protons
postive ions
have fewer electrons than protons
isotope
elements with the same atomic number but different mass number due to different amounts a neutrons, with the same chemical properties but different physical properties (density, diffusion)
ancient greeks
though that all mater was made from indivisible particles
John dalton
at the start of the 19th century it was suggested atoms were solid spheres
JJ thomson
in 1897 experiments concluded that atoms must have smaller negatively charged particles, electrons, he created the plum pudding model
Ernest Rutherford
Geiger-Marsden experiment where alpha particles were fired at gold sheet and where deflected and some went through, nuclear model with a positive centre and cloud of electrons
Henry Moseley
charge of nucleus increased by one form element to element, lead to Rutherford discovering the proton
James Chadwick
conclude there are neutrons with a mass but no charge
Bohr model (Niels Bohr)
proposed that elections are in fixed orbits with a fixed energy, when electrons move shells they emit or absorb electromagnetic radiation
quantum model
model which uses quantum mechanics to predict where electrons will be
relative mass
the molecular mass is relative to carbon 12
mass spectrometer
sample is vaporised
ionisation to positive ions by bombardment of high energy electrons
the ions are accelerated by an electric shield
time for the ions to reach detector is measured with lighter ones been quicker
calculate relative atomic mass
(abundance x mass) ÷ 100
mole
amount of a substance
avogadros constant
6.02 x10 ^23
number of moles
= number of particles you have ÷ avogadros constant
number of moles
= mass ÷ Mr
concentration
= moles ÷ volume
empirical formula
smallest whole number ratio of a compound, worked out from moles and then dividing by smallest
water of crystallisation
compounds with incorporated water
hydrated
compounds containing water of crystallization
anhydrous
compounds without incorporated water
calculate water of crystallization
find moles of water lost from mass
find moles of compound from mass
divide by smallest
percentage yield
actual yield ÷ theoretical yield x 100
standard solution
a solution with a known concentration
6, 3, 4, 9, 8, 2, 7, 1, 5
order statements making standard solution
place stopper in bottle and invert
rinse the beaker and stirring rod with distilled water and add it to flask
work out how many grams of solute needed using mass = moles x mr
measure out the mass of the solute first measuring beaker
check meniscus again and add water if needed
work out moles of solute using using moles = conc x volume ÷ 1000
top up flask with distilled water till at meniscus using pipette
top up solution into volumetric flask using funnel
add distilled water to beaker to dissolve solute
diluting solution
making a standard solution from more conc solution
volume to use
final conc ÷ initial conc x volume required
1, 4, 7, 2, 3, 5, 6,
add acid to burette to the 0cm3 line
carry out rough titration to find the endpoint
work out amount of acid used
add alkali to flask using 25cm3 pipette
carry out accurate titration till you get concordant results
calculate mean
add indicator to flask
methyl orange
turns yellow to red/peach when adding acid to alkali
phenolphthalein
turns red to colourless when adding acid to alkali
s
sub shell with 2 electrons
p
sub shell with 8 electrons
d
sub shell with 10 electrons
f
sub shell with 14 electrons
spin pairing
electrons spin in opposite directions
s orbital
spherical orbitals
p orbital
dumbbell shaped orbital, made up of 3
s block
group 1 and 2
d block
transition metals
p block
group 3-8
ionic bonding
electrons are transferred making 2 ions and then hold each other together by electrostatic attractions, non metal and metal
dot and cross
diagram showing electron distribution in shells, use square brackets and charge for ionic compounds, shells overlap for covalent
giant ionic lattice
compounds which form crystals with regular repeating structure
ionic properties
conduct when molten or dissolved
high melting points
soluble in water
covalent bonds
share electrons been 2 non metals, help by electrostatic attraction between nuclei and electrons as well as repulsion between nuclei
dative covalent bonding
covalent bond where both bonding electrons are from one element
giant covalent structures
structures like carbon and silicon dioxide
properties of giant covalent structures
very high melting point
very hard
good thermal conductors
don’t dissolve
can’t conduct
metallic bonding
bonding between metals where they lose electrons making them delocalised leaving a positive ion, attraction between ion and electrons hold them together
metallic bonding properties
high melting points due to sea of delocalised electrons
ions able to slide over each other
good thermal and electrical conductors due to sea of delocalised electrons
insoluble due to high bond strength
shape
depends on the number pf pairs of electrons in outer shell
electron effect
pairs of electrons are negative and will repel each other to get as fair apart from each other as possible
lone pairs
pairs of electrons that repel more than bonding pairs
wedges
shows bonds pointing out page
dashes
shows bonds pointed into page
linear molecules
2 bonding pairs bond angle of 180 degrees
trigonal planar
3 bonding pairs, bond angle of 120˚
non linear
2 bonding pairs with double bonds cancel out effect of 1 lone pair, bond angle of 120˚
tetrahedral
4 bonding pairs, bond angle of 109.5˚
trigonal pyramidal
3 bonding pairs, 1 lone pair, bond angle of 106˚
bent
2 bonding pairs, 2 lone pairs, bonding angle of 104.5˚
trigonal bipyramidal
5 bonding pairs, bonding angle of 120˚ and 90˚
octahedral
6 bonding pairs, bonding angle of 90˚d
determine shapes
draw a dot and cross diagram, allows you to see numbers of bonding and lone pairs
calculate shape
find group number of central atom (only one of)
add number of atoms around atom to its group number
divide by 2 to get electron pairs
if there is the same number of electron pairs as surrounding atoms then no lone pairs.
if ion then subtract or add its charge from its group
group
column where all elements have the same number of electrons in outer shell
period
row where all elements have the same number of shells
melting points
a trend which increases across a period peaking at group 4, related to number of delocalised electrons, giant covalent structures and intermolecular forces
first ionisation enthalpy
energy needed to remove 1 electron from each atom in one mole of gaseous atoms to become 1 mole of gaseous 1+ ions
lower enthalpy
easier it is to remove on outer electron and form an ion
factors effecting first ionisation enthalpies
atomic radius- dist electrons from nucleus
nuclear charge- more positive nucleus attracts electrons more
electron shielding- inner electrons shield outer electrons.
group 1 and 2 first ionisation enthalpy
trend which decreases down a group due to increasing shells causing shielding
across a period
direction where first ionisation enthalpy increases due to increasing protons
s block metals
block with low first ionisation enthalpies due to few protons a limited outer shells
hydroxide
product of water and group 2 metal, also produces H2
oxides
product of when oxygen combines with a group 2 metal
strong alkaline solutions
when group 2 oxides react with water forming metal hydroxides which dissolve, expect magnesium oxide, increases down group
neutralise acids
both are bases so react with acids to form salts and water
group 2 hydroxide
solubility increases down group (single negative charge)
group 2 oxide
solubility decreases down group (double negative charge)
thermal decomposition
when heat is added to group 2 carbonates they form the oxide and CO2
thermal stability increases
change in thermal stability down group due to smaller cations (group 2) distorting the large carbonate more due to charge density
salts
neutral ionic compounds with a cations and anions
acids
substances with a pH less than 7
bases
substances with a pH more than 7
HN4+
ammonium
NO3-
nitrate
HCO3-
hydrogencarbonate
(SO4)2-
sulfate
(CO3)2-
carbonate
cation first
order when naming salts
most sulfates
soluble except barium, calcium, lead which form white ppt
lithium, sodium, potassium, ammonium and nitrates
all soluble salts
most chlorides, bromides, iodides
soluble salts except silver halides, copper iodide (white ppt), lead chloride and bromide (white ppt) and lead iodide (yellow ppt)