AQA CHEMISTRY A LEVEL

Plum pudding model

It was initially thought that atoms consisted of a sphere of positive charge with small negative charges distributed evenly within it.

What do atoms consist of?

-          Small, dense central nucleus

-          Surrounded by electrons

-          Surrounded by electron shells

what does the nucleus consist of

protons and neutrons this gives the overall charge and contains the entire mass of the atom

proton

relative charge - +1

relative mass - 1

neutron

relative charge - 0

relative mass - 1

electron

relative charge - -1

relative mass - 1/1840

mass number

sum of protons and neutrons in an atom

atomic number

equal to the number of protons in an atom

relative atomic mass (Ar)

The mean mass of an atom of an element divided by one twelfth of the mean mass of an atom of the carbon-12 isotope – it takes the relative abundances of the different isotopes of an element into account

Isotopes

atoms of the same element with the same atomic number but with a different number of neutrons resulting in a different mass number

isotopes characteristics

  Neutral atoms of isotopes will react chemically in the same way as their proton number and electron configuration is the same. The sharing and transfer of electrons is unaffected. However the different mass number means they will have different physical properties

Ions

formed when an atom loses or gains electrons meaning it is no longer neutral and will have an overall charge

Mass spectrometry

this is an analytical technique used to identify different isotopes and find the overall relative atomic mass of an element

Time of flight mass spec (TOF)

form of mas spec that records the time it takes for ions of each isotope to reach a detector. Using this, spectra can be produced showing each isotope present.

steps of TOF

1. ionisation

2. acceleration

3. ion drift

4. detection

5. analysis

TOF ionisation

sample of an element is vapourised and injected into the mass spectrometer, high voltage is passed over the chamber. This causes electrons to be removed from the atoms (it is ionized) leaving 1+ charged ions.  During the ionization process a 2+ charged ion may be produced. This means it will be affected more by the magnetic field producing a curved path of smaller radius. As a result, its mass to charge ratio (m/z) is halved and this can be seen on spectra as a trace at half the expected m/z value.

TOF acceleration

the positively charged ions are then accelerated towards a negatively charged detection plate

TOF ion drift

the ions are deflected by a magnetic field into a curved path. The radius of their path is dependent on the charge and mass of the ion

TOF detection

when the positive ions hit the negatively charged detection plate, they gain an electron producing a flow of charge. The greater the abundance the greater the current produced

TOF analysis

These current values are then used in combination with flight times top produce a spectra print out with the relative abundance of each isotope displayed

RAF calculation

m/z x abundance / Total abundance

chlorine spectra

-          Characteristic pattern

-          Appear in a 3:1 ratio for Cl+ ions

-          Appear in a 3:6:9 ratio for Cl2 + ions

-          This is because one isotope is more common than the other and the chlorine molecule can form different combinations

electronic configuration

Electrons are held in a cloud of negative charge called orbitals

S orbital

-          spherical shaped

-          groups 1 and 2

can hold up to 2 electrons

P-orbital

-          dumbbell shaped

-          groups 3,4,5,6,7,0

-          can hold up to 6 electrons

D-orbital

-          Van Cleef shaped

-          Transition metals

-          Can hold up to 10 electrons

Spin

-          In an orbital atoms pair up with opposite spins – ensures atom is stable

-          Electrons in the same orbital must have opposite spin

-          Spin is represented by arrows

3 rules for electronic configuration

1. lowest energy level is filled first

2. electrons with the same spin fill up an orbital first before pairing begins

3. no single orbital holds more than 2 electrons

electronic configuration for unstable elements

If  electron spins are unpaired and therefore unbalanced, it produces a natural repulsion between the electrons making the atom very unstable. If this is the case, electrons may take on a different arrangement to improve stability. The 3p4 orbital contains a single pair of electrons with opposite spin making it unstable. Therefore the electron configuration changes to become 3p3 4s1 which is more stable

Ionisation energy

the minimum energy required to remove one mole of electrons from one mole of atoms in a gaseous state

successive ionisation energies

-          occur when further electrons are removed

-          Requires more energy because electrons are removed closer to the nucleus meaning there is stronger attraction between electrons and the nucleus

trends in ionisation energy along a period

-          Ionization energy increases

-          Similar shielding

-          Smaller atomic radius

-          Larger proton number

trends in ionisation energy down a group

-          Ionization energy decreases

-          Increased shielding

-          Larger atomic radius

exceptions to increasing ionisation energies across a period

Aluminium - first ionization energy is lower as its outer electron is in the p orbital which is further away from the nucleus and therefore less attracted to the nucleus

Relative molecular mass

The mean mass of a molecule of a compound divided by one twelth of the mean mass of an atom of carbon 12 isotope – for ionic compounds it is known as the relative formula mass

mole

a unit measurement for substances – it always contains the same number of particles

mole calculation 1

mass/mr

mole calculation 2

volume x concentration

avagradros constant

6.022 × 10²3

number of particles calculation

6.022 × 10²3 x moles

Ideal gas equation

PV=nRT

Ideal gas equation measurements

• P = pressure in pascals

• V= volume in m³

• n=moles

• R=ideal gas constant

• T= temperature in kelvin

empirical formula

The simplest whole number ratio of atoms of each element in a compound

molecular formula

The true number of each atom in the molecule

Percentage yield

actual yield/theoretical yield x 100

atom economy formula

Mr of desired products/mr of reactants x 100

why is a high atom economy desirable

When reactions have higher atom economies it means that they are more economically viable for industrial scale manufacture and have little waste products also uses less energy

Ionic bonding

• occurs between metal and nonmetal

• happens when metal loses electron to non metal to get an outershell full

• strong electrostatic forces between oppositely charged ions

• example - NaCl

Covalent bonding

• Forms between two non-metals

• electrons are shared between the two outer shells in order to achieve a full outer electron

• Multiple electron pairs can be shared to form multiple covalent bonds

dative covalent bond / coordinate bond

occurs when both of the electrons in the shared pair are supplied form a single atom in a covalent bond this is indicated as an arrow from the lone pair.

Metallic bonding

• giant lattice of positive ions in a sea of delocalised electrons

• strong electrostatic forces between particle

• the greater the charge on the positive ion the stronger the attraction between the sea

• ions larger in size have weaker attraction due to a greater atomic radius

physical properties

• boiling/melting point

• solubility

• conductivity

types of crystal structures

ionic, metallic, macromolecular, simple molecular

ionic structure properties

• high mp and bp

• can conduct electricity when molten or in solution

• brittle

metallic structure properties

• good conductors - due to sea of delocalised electrons

• malleable - layers able to slide over

• high mp and bp

• nearly always solid at room temp apart from mercury which is liquid

simple molecular structure properties

• consists of covalently bonded molecules held together by van der waal forces

• low mp and bp

• very poor conductors

macromolecular structure properties

• covalently bonded into a giant lattice

• each atom has multiple covalent bonds

• high mp and bp

• rigid

• examples - diamond and graphite

diamond

• made up of carbon atoms each bonded further to 4 carbon atoms

• one of the strongest materials known

graphite

• Made up of carbon atoms however each carbon atom is bonded to three other flat sheets.

• The electrons not used in bonding are released as free electrons which move between layers

• slide over eachother so soft and slippery

• conducts electricity

sulfate ion formula

SO4²-

hydroxide ion formula

OH-

carbonate ion formula

CO3²-

Nitrate ion formula

NO3-

Ammonium ions formula

NH4+

shapes of molecules

• The shape of a molecule is determined by the number of electron pairs around the central atom

• Each electron pair naturally repels eachother so that the largest bond angle possible exists between the covalent bonds

• Any lone pairs present around the central atom provide additional repulsive forces, which changes the bond angle

• For every lone pair present, the bond angle between covalent bonds reduce by 2.5 degrees

how to find a shape of a molecule

1.      Find the number of electron pairs

2.      Determine how many of the pairs are bonding pair and how many are lone pairs

3.      Bonding pairs indicate the basic shape and lone pairs indicate the additional repulsion

linear shape (straight line)

bp-2

lp-0

bond angle -180

example - BeCl2

Bent shape (upside down v shape)

bp -2

lp -2

bond angle - 104.5

example - water

trigonal planar shape (triangle)

bp -3

lp- 0

bond angle-120

example - Boron fluoride

Trigonal pyramidal shape (upside down diamond shape)

bp -3

lp-1

bond angle - 107

example - ammonia

tetrahedral shape

bp -4

lp-0

bond angle - 109.5

example - methane

trigonal bipyramidal

bp-5

lp-0

bond angle- 90 and 120

example - phosphorus pentachloride

octahedral shape

bp -6

lp -0

bond angle -90

example - sF6

bond polarity

The negative charge around a covalent bond is not evenly spread around the orbitals of the bonded atoms

electronegativity

The power of an atom to attract negative charge towards itself within a covalent bond. This power is different for every atom depending on its size and nuclear charge

trends

electronegativity increases along a period as atomic radius decreases and it decreases down a group as shielding increases

polar bond

when two atoms that are bonded have different electronegativities- example - HF

permenant dipole

The more electronegative atom draws more of the negative charge towards itself and away from the other atom producing a delta negative and delta positive region

induced dipole

Forms when the electron orbitals around a molecule are influenced by another charged particle

3 types of intermolecular forces

1. van der waal

2. dipole-dipole

3. hydrogen

van der waal forces

• weakest intermolecular force

• acts as an induced dipole between molecules

• varies depending on mr larger mr = stronger van der waals

• Straight chain molecules have stronger van der waal forces than branched molecules as they are more closely packed together reducing the distance in which the force acts over

dipole-dipole

• Acts between molecules with polar bonds

• The delta negative and delta positive regions attract eachother and hold the molecule together in a lattice like structure

hydrogen bonding

• strongest intermolecular force

• occurs between hydrogen and the three most electronegative elements (nitrogen, fluorine and oxygen)

• high mp and bp

• show as a dotted line

Enthalpy change (delta H)

-          In a reaction, bonds are broken and the made

-          For bonds to be broken energy is taken in and for bonds to be made energy is given out

-          The overall energy change of the reaction depends on how much energy is transferred during these processes

endothermic

When energy is taken in from the surroundings, the enthalpy change is positiveex

exothermic

When energy is given out to the surroundings the enthalpy change Is negative

enthalpy change formula

energy to break bonds +energy to make bonds

Enthalpy of formation

the enthalpy change when one mole of a substance is formed from its constituent elements under standard states and standard conditions

Enthalpy of combustion

The enthalpy change when one mole of a substance is completely burnt in oxygen under standard conditions

mean bond enthalpies

The bond enthalpy values calculated this way often differ from data book values due to being averaged values

calorimetry

-          An experimental method for finding enthalpy change by measuring temperature change over time

-          When observed and plotted on a graph it can be extrapolated to give an accurate value for the change in temperature

-          The measured change in temperature is proportional to the energy change

calorimetry experiment

• polystyrene cup to minimise heat loss

• lid to minimise heat loss

• thermometer

• place acid along with powder and immediately shut lid

• mix with thermometer and measure highest and lowest temperature

Calorimetry equation

Q=MCdeltaT

Calorimetry equation components

Q- heat loss

M- mass

C- specific heat capacity

deltaT - change in temperature

specific heat capacity

The energy required to raise 1g of a substance by 1k without a change in state

ethalpy change equation

q/moles

convection

occurs when particles with a lot of heat energy in a liquid or gas move nd take the place of particles with less heat energy

Hesses law

states that the overall enthalpy change of a reaction is the same regardless of the route taken

Enthalpy of formation hess cycle

arrows point out from the central product C as both A and B are formed from elements at C

Enthalpy of combustion hesses law

Arrows point toward the central product which is always H2O and CO2 as both A and B burn to form products at C