elements of life
the atom
atoms are made up of protons, neutrons and electrons
electrons
have a -1 charge
they are found in shells
shells make up the majority of the atom
nucleus
most of the mass of the atom is concentrated at the nucleus
the diameter of the nucleus is very small compared to the rest of the atom
the nucleus is where protons and neutrons are found
mass= 1/2000
neutron
mass=1
has a charge of 0
proton
mass=1
has a charge of +1
nuclear symbols show numbers of subatomic particles
mass number (top number)= the total of protons and neutrons in the nucleus
atomic number (bottom number)= number of protons in the nucleus - all atoms of the same element have the same number of protons
ions have different numbers of protons and electrons
negative ions
negative charge means theres more electrons than protons
the number next to the - indicates number of electrons added
positive ions
positive charge means theres fewer electrons than protons
the number next to the + indicates number of electrons lost
isotopes are atoms of the same element with different numbers of neutrons
isotopes= atoms of an elements with the same atomic number but different mass numbers
caused by the numbers of neutrons in an atom
can cause a variation in physical properties which do depend on the mass of the atom (density, boiling points)
atomic models
the accepted model of the atom has changed throughout hustory
some ancient Greeks thought matter was made from invisible particles
at the beginning of the 9th century Dalton described atoms as solid spheres and said that different spheres make up different elements
new experiments constantly changed ideas
experimental evidence showed that atoms weren't solid spheres
in 1897 JJ THOMPSON did a series of experiments which disproved Dalton's theory
his measurements of CHARGE and MASS showed that an atom must contain smaller negatively charged particles- electrons
created PLUM PUDDING MODEL- negative electrons embedded into a sphere of positive charge
Rutherford proved plum pudding wrong
In 1909 Rutherford with his students Geiger and Marsden conducted the alpha scattering (gold leaf) experiment
they fired alpha particles at a thin sheet of gold
they expected (due to plum pudding) for the alpha particles to be very slightly deflected by the positive sphere
in fact most of the alpha particles passed straight through and a small number were deflected backwards
so Rutherford created the nuclear model
nuclear model
tiny positively charged nucleus at the centre of atom- where most of mass is concentrated
nucleus is surrounded by a cloud of freely orbiting electrons
most of atom is empty space
the nuclear model was modified
Mosely discovered the charge of the nucleus increased from one element to another by one
Rutherford also discovered that the nucleus contained positively charged protons
Chadwick discovered the neutron
the Bohr model
electrons existed in fixed orbitals (shells)
each shell had fixed energy
when an electron moves between shells electromagnetic radiation is absorbed or emitted
because the energy of the shells are fixed, the radiation will have a fixed frequency
said that a shell can only hold a certain number of electrons and the electrons determined reactivity
relative mass
relative masses are masses of atoms compared to carbon 12
relative atomic mass= the average mass of an atom of an element (on a scale where an atom of carbon 12 is 12
relative isotopic mass= the mass of an atom of an isotope of an element (on a scale where carbon 12 is 12)
relative molecular mass= the average mass of a molecule or formula unit (on a scale where carbon 12 is 12)
relative masses can be measured by a MASS SPECTROMETER
1) VAPORISATION= the sample is turned into gas using an electrical heater
2) IONISATION= the gas particles collide with high energy electrons to ionise them. electrons are knocked off the particles leaving them as positive ions
3) ACCELERATION= the positive ions are accelerated by an electric field
4)DETECTION= time taken for positive ions to reach the detector is measured. (depends on charge and mass o ion)- light, highly charged ions will reach detector first
mass spectrum
the y axis= abundance of ions or the relative isotopic abundance for elements
the x axis= relative mass
the mole and equations
a mole is a very large number of particles
amount of substance is measured in moles
1 mole= 6.02x 1023 (avogadros constant
number of moles= number of particles you have / number of particles in a mole (avogadros constant)
molar mass is the mass of one mole
molar mass= relative molecular/formula mass
number of moles= mass of substance / molar mass
in a solution concentration is measured in mol/dm-3
number of moles= concentration x volume(cm3) / 1000
ionic equations
only reacting particles are included
empirical and molecular formulas
empirical and molecular formulas are ratios
e formula gives the smallest whole number ratio of a compound
m formula gives the actual numbers of atoms in a molecule
percentage yield in never 100%
percentage yield = actual yield/ theoretical yield x100
titrations
a standard solution has a known concentration
1) work out how many moles of solute needed by using moles= concentration x volume/ 1000
2) work out how many grams of solute is needed using mass= moles x molar mass
3) weigh the empty beaker and then the beaker with the correct mass of solute
4) add a small amount of distilled water and stir until dissolved
5) tip solution into a volumetric flask
6) rinse the beaker with distilled water then add to flask
7) top up flask to 250cm3 with distilled water
8) stopper the bottle and turn upside down 5 times
making a standard solution from a more concentrated solution
vol to use= final conc / initial conc x vol required
use this equation to get desired concentration
then transfer this into volumetric flask
titrations must be accurate
allow you to find out exactly how much acid is needed to neutralize and alkali
you measure out some alkali using a pipette and put in a flask with indicator (phenolphthalein)
first do a rough titration to get an idea where point of neutralization is
take a reading of initial volume of acid and slowly add acid to the alkali swirling solution and stop when there is a colour change
work out amount of acid needed to neutralize the alkali (the titre)
repeat and work out mean
indicators
methyl orange- turns yellow to red
phenolphthalein - turns red to colourless
titration calculations
number of moles = concentration x volume / 1000
number of moles= mass / molar mass(mr)
electronic structure
electron shells are made up of sub shells and orbitals
electrons move around nucleus in shells which are given numbers known as PRINCIPAL QUANTUM NUMBERS
Shells further away from nucleus have a greater energy level than shells closer to the nucleus
s sub shell
1 orbital
maximum number of electrons= 2
p sub shell
3 orbitals
maximum number of electrons = 6
d sub shell
5 orbitals
maximum number of electrons= 10
f sub shell
7 orbitals
maximum number of electrons = 14
orbitals have shapes
electrons in each orbital have to spin in opposite electrons (spin pairing)
s orbitals = spherical
p orbitals = dumbbell shapes
using the periodic table to workout electron configurations
s block= groups 1 & 2
p block= boron to neon and below
d block= transition metals
for an ion
write electronic structure of atom
add or remove electrons from the highest occupied sub shell
4S SUB SHELL FILLS BEFORE 3D SUB SHELLS BUT EMPTIES BEFORE IT
ionic bonding
ionic bonding= when ions are joined by electrostatic attractions
ions are formed when electrons are transferred from one atom to another
generally the charge on the ion is equal to its group number
electrostatic attractions= forces that hold positive and negative ions together (very strong)
giant ionic lattices
a lattice is just a regular structure
the structure is giant because it is the same basic unit repeated over and over again
in the lattice ions with different charges attract eachother and the ions with the same charge repel
ions arrange themselves in order to maximize attractions and minimise repulsions
ionic structure influences properties
ionic compounds conduct electricity when molten or dissolved (not solid)
the ions in a liquid/ solution are free to move and carry charge
ionic compounds have high melting points
giant ionic lattices are held together by strong electrostatic forces
lots of energy is needed to overcome these forces
ionic compounds are often soluble
water molecules are POLAR
water molecules pull ions away from lattice and cause it to dissolve
covalent bonding
covalent bonds hold atoms in molecules together
two atoms share electrons o they both have full outer shells
both positive nuclei are attracted electrostatically to the shared electrons
repulsion between positive nuclei - maintains a balance in forces
molecular substances have some typical properties
low melting points
intermolecular forces between molecules are weak
don't conduct electricty (no electrons free to move)
insoluble in water
the polar water molecules are more attracted to eachother than the molecular substance
dative covalent bonding= where both electrons come from one atom
giant covalent and metallic structures
some covalently bonded substances have giant structures
have a network of covalently bonded atoms- electrostatic attractions holding the atoms together are stronger than in simple covalent molecules
e.g. diamond, graphite, silicon dioxide
properties of giant covalent structures
very high melting points
needs a lot of energy to break bonds before substance melts
very hard
strong bonds throughout lattice
good thermal conductors
vibrations travel easily through stiff lattices
insoluble
don't contain ions
cant conduct electricty
no delocalised electrons to carry charge throughout the structure (all are bonded)
metals have giant structures
metal elements exist as giant metallic lattice structures
the electrons in the outer shell are delocalised (leaves a positive metal ion)
positive metal ions are attracted to the delocalised negative electrons- form a lattice of packed positive ions in a sea of electrons
properties
high melting points
strong metallic bonding (the more electrons there are the stronger the bonding)
the size of the metal ion and the structure of the lattice effect it
ductile
no bonds holding specific ions together
the metal ions slide over eachother when the structure is pulled
good thermal conductors
delocalised electrons can pass kinetic energy to eachother
good electrical conductors
delocalised electrons are free to move and carry charge
impurities can reduce this (electrons transfer to impurities are form anions)
insoluble
strength of metallic bonds
shapes of molecules
molecular shape depends on electron pairs around central atom
electron pairs repel eachother
electrons are negatively charged so repel as much as possible
lone pairs repel more than bonding pairs
the greatest angles are between lone pairs of electrons
bonding pair angles are often reduced because they are pushed together by lone pair repulsion
the shape of the molecule depends on the type of electron pairs as well as the number
electron pair repulsion theory is used to predict shapes
you can use electron pairs to predict shapes of molecules
draw dot and cross diagram first
then work out shape of molecule
types of molecules
2 electron pairs on central atom
LINEAR
treat double bond same as single
3 electron pairs on central atom
TRIGONAL PLANAR
no lone pairs
BENT/ non linear
one lone pair
4 pairs of electrons on central atom
180 degrees angle
117.5 degrees
120 degrees
TETRAHEDRAL
no lone pairs
109.5 degrees
TRIGONAL PYRAMIDAL
one lone pair
107 degrees
BENT/ non linear
2 lone pairs
104.5 degrees
5 electron pairs on central atom
TRIGONAL BIPYRAMIDAL
90 and 120 degrees
6 electron pairs on central atom
OCTAHEDRAL
90 degrees
periodic trends
the periodic table arranges elements by atomic number
all elements within a period have the same number of electron shells
all the elements within a group have the same number of electrons in their outer shells
periodic trends are patterns (periods 2 &3)
metals
melting points increase across period (metallic bonding gets stronger)
bonds get stronger due to increasing number of delocalised electrons and a decreasing ionic radius
creates a higher charge density- attracts ions more strongly
elements w/ giant covalent structures
a lot of energy is needed to break these strong bonds
carbon and silicon have the highest melting points in their periods
simple molecular structures
their melting points depend on intermolecular forces
intermolecular forces are weak and easily broken
have low melting points
noble gases have the lowest melting points as they exist as individual atoms
more atoms in a molecule means stronger intermolecular forces- higher melting point
ionisation is the removal of one or more electrons
first ionisation enthalpy= the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions
X(g) ----> X+ + e-
the lower the ionisation enthalpy the easier it is to remove the outer electron
main things that affect size of ie's
atomic radius
the further the outer shell electrons from the positive nucleus the less they will be attracted
it will be lower
nuclear charge
this the positive charge on the nucleus caused by protons
the more protons the more they will attract electrons
it will be higher
electron shielding
the inner electron shells shield the outer shell electrons from the more attractive force of the nucleus
it will be lower
decreases down the group
increase across a period
group 2
group 2 elements react with water and oxygen
they react with water to produce hydroxides
gives a metal hydroxide + hydrogen
e.g. M + 2H20 -----> M(OH)2 + H2
they burn in oxygen to form oxides
produces white solids
e.g. 2M + O2 ----> 2MO
group 2 oxides and hydroxides are bases
they form alkaline solutions in water
the oxides of group 2 react with water to form metal hydroxides which dissolve
the hydroxide ions (OH-) make the solution strongly alkaline
More alkaline down the group
they neutralize acids
form solutions of corresponding salts
DEACREASE IN SOLUBILITY DOWN THE GROUP
group 2 carbonates decompose to form CO2 and metal oxides
thermal decomposition= when a substance breaks down when heated
e.g. MCO3 -----> MO + CO2
Thermal stability of carbonates changes down group
the more thermally stable, the more heat needed to break it down
carbonate ions are large anions and can be made unstable by the presence of a cation (polarizes it) and distorts it (the greater the distortion the less stable)
large cations cause less distortion than small ones (have a lower charge density)
so the further down the group the larger the cations, the less distortion caused and the more stable the carbonate
making salts
salts are ionic compounds
acid + base ---> salt + water (neutralization)
salts are formed from cations and anions so product is neutral
cations:
-NH4+
-Cu2+
-Zn2+
-Pb2+
-Fe2+
-Fe3+
anions:
-NO3-
-OH-
-HCO3-
-SO42-
-CO32-
Salts can soluble or insoluble in water
soluble= nitrates, most chlorides and bromides, most sulfate, lithium sodium potassium ammonium salts
insoluble= carbonates, hydroxides
tests for ions
flame tests
1) dip nichrome loop into HCl
2) dip wire loop into sample and place in flame
3) observe colour
colours:
-Li+= crimson
-Na+= orange
-K+= lilac
-Ca2+= brick red
-Ba2+= green
-Cu2+=blue/ green
add sodium hydroxide and look for coloured precip
colours:
-Ag+= brown
-Ca2+= white
-Cu2+= blue
-Pb2+= white
-Fe2+=green
-Fe3+= red brown
-Zn2+= begins white then colourless
-Al3+=begins white then colourless
hcl can detect carbonates
test for limewater
test for sulfates
add dilute HCl followed by BARIUM CHLORIDE SOLUTION
if white precip forms it contains sulfate
litmus paper going blue= ammonia or hydroxide is present
test for halides with silver nitrate
Cl= white precip
Br= cream precip
I= yellow precip
nitrates
heat solution with sodium hydroxide and aluminium foil
if present ammonia is produced
used litmus paper to test for ammonia
atomic spectra and nuclear radiation
electromagnetic spectrum
radiation increases in frequency and decreases in wave length from left to right
radio waves, micro waves, infra-red, visible light, ultra violet, x rays, gamma rays
electrons absorb or release energy
atoms in their ground state have electrons in the lowest possible shells
if an atom takes in energy they can move to higher energy levels further from nucleus
electrons can also release energy by dropping from a higher energy level to a lower one
the energy levels all have certain fixed values- electrons can jump by absorbing or releasing a fixed amount of energy
transition between energy levels produces line spectra
absorption spectra
energy is related to frequency
when em radiation is passed through a gaseous element the electrons only absorb certain frequencies corresponding to different energy levels
the missing frequencies show up as dark lines
emission spectra
when electrons drop to lower energy levels they give out certain amounts of energy
each element has a different electron arrangement so the frequencies of radiation absorbed and emitted are different
has coloured lines
atomic spectra and nuclear radiation
both types of spectra are made up of sets of lines
absorption and emission spectra have lines in the same position for an element
lines represent electrons moving from or to a different energy level
the lines on both get closer as frequency increases
energy is related to frequency
E= h x v
speed and frequency related to wavelength
c= v x y
nuclear fusion releases lots of energy
nuclear fusion=when two small nuclei combine under high temp and pressure to make one larger nucleus
in stars hydrogen nuclei combine to make helium nuclei
can only be done in stars