Chemistry AS

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ATOMIC STRUCTURE AND THE PERIODIC TABLE

Structure of the atom and isotopes

atoms have protons, neutrons and electrons. protons and neutrons are found in the nucleus and have a relative mass of 1 whilst electrons are found in energy levels (quantum shells) around the nucleus and have a relative mass of 1/1840. protons have a relative charge of +1, electrons -1 and neutrons 0.

the atomic number (Z) of an atom is the number of protons in the nucleus of an atom of a particular element.

the mass number of an atom is the sum of the number of protons and neutrons in the nucleus of an atom

isotopes are atoms of the same element with different masses/numbers of neutrons in the nucleus

Mass spectrometry and relative masses

relative atomic mass (Ar) is the weighted mean (average) mass of an atom of an element compared to 1/12 the mass of an atom of carbon-12

relative isotopic mass is the mass of an individual atom of a particular isotope relative to 1/12 the mass of an atom of carbon-12.

relative isotopic mass, as well as percentage abundance is used to calculate the relative atomic mass of a sample of an element. this occurs in a mass spectrometer. which produces positive ions that are deflected by a magnetic field according to their m/z ratio, it also calculates the relative abundance of each positive ion and displays it as a percentage. the positive ions could be atoms, molecules or fragments of molecules.

you can obtain the relative molecular mass of an element or compound (covalently bonded together) by observing the peaks with the highest m/z ratios e.g Chlorine (Cl2)

two peaks corresponding to isotopic masses of 35 and 37, height ratio being 3:1 so realitve atomic mass is ((3x35)+(1x37)) / 4 = 35.5

the peaks at 70, 72 and 74 are due to the molecules- 35+35=70, 35+37=72, 37+37=74.

there is a 3/4 chance of selecting a Cl-35, so the chance of selecting two together is 3/4 x 3/4 = 9/16

the chance of combining Cl-37 and Cl-35 is 1/4 x 1/3 = 3/16 X2 = 6/16

there is a 1/4 chance of selecting a Cl-37, so the chance of selecting two together is 1/4 x 1/4 = 1/16

the whole number ratio for this is 9:6:1

using this you can calculate the relative molecular mass (Mr) by:

calculating the weighted mean of the various molecules present - Mr (Cl2) = (9x70) + (6x72) + (1x74) / 16 = 71.

OR- by adding together the relative atomic masses of the the elements in the molecule so 2 x 35.5 = 71

OR- by considering the molecular ion peak (peak with the highest m/z ratio in the mass spectrum, the M peak) but ignoring the M+1 peak as a result of the presence of the carbon-13 isotope, seen in molecules with large masses

Atomic orbitals and electronic configurations

a quantum shell defines the energy level of an electron

electrons in the first quantum shell (closest to the nucleus) have the least energy, going out from that they gain more energy

each quantum shell, bar the first is divided further into subshells of slightly different energy levels:

1s

2s 2p

3s 3p 3d

4s 4p 4d 4f (electron energies increasing in order 4s<4p<4d<4f)

each subshell contains orbitals which is a region within an atom that can hold up to 2 electrons with opposite spins

number of electrons in:

s subshell is 1x2= 2

p subshell is 3x2= 6

d subshell is 5x2= 10

f subshell is 7x2= 14

number of electrons in:

1st quantum shell = 2

2nd quantum shell = 8

3rd quantum shell = 18

4th quantum shell = 32

the electron configuration of an atom shows the number of electrons in each sublevel in each energy level of the atom, generally filling in the lowest energy orbitals first.

there are exceptions to this rule such as potassium (K) is 1s2 2s2 2p6 3s2 3p6 4s1 and not 1s2 2s2 2p6 3s2 3p6 3d1 this is because for all elements after Ca on the periodic table, the energy of the 3d orbital is less than the energy of the 4s orbital

Hund's rule states that electrons will occupy orbitals singly before pairing takes place

Pauli Exclusion Principle states that two electrons can't occupy the same orbital unless they have opposite spins, usually shown by arrows

Ionisation Energies

Spectroscopes is used to analyse the EM radiation by forming an electromagnetic spectrum, with specific frequencies unique to each element.

the First Ionisation Energy of an element is the energy required to remove an electron from each atom in one mole of atoms in the gaseous state

A(g) --> A+(g) + e-

the Second Ionisation Energy of an element is the energy required to remove an electron from each singly charged positive ion in one mole of positive ions in the gaseous state

A+(g) --> A2+(g) + e-

successive ionisation energies provide evidence for quantum shells as there is a large increase in ionisation energy from one shell to the next

Ionisation energy (IE) = energy of e- when removed - energy of e- when in orbital

H and He

H= 1s1

He= 1s2 - increased electron shielding of nuclear charge effect (increases energy) < has double nuclear charge (decreases energy)

He and Li

He= 1s2

Li= 1s2 2s1 - larger nuclear charge (decreases energy) < 2s electron in a higher energy level (increases energy) and shielding from inner e- (increases)

factors:

orbital electron exists in

nuclear charge of atom (proton no.)

shielding

Trends across a PERIOD =overall increase in IE

- nuclear charge increases (decreases energy) more significant

- shielding increases (increases energy)

Trends down a GROUP =overall decrease in IE

-nuclear charge increases (decreases energy)

-one new quantum shell - higher energy value (increases energy)

- increased shielding (increases energy) more significant

Periodicity

Periodicity is a regularly repeating pattern of atomic, chemical and physical properties with increasing atomic number

Atomic Radii

measure of the size of atoms. distance from centre of nucleus to boundary of electron cloud

calculated by determining the distance between the two nuclei and divide by two.

van der waals has to be used for neon and argon as they don't bond with other elements

metallic radius is used for metals

radius decreases across a period as does the nuclear charge

Melting and Boiling temperatures

elements with giant lattice structures have larger melting/boiling points than molecular structures, more so if covalent

First Ionisation Energies

H to He

increase in nuclear charge > shielding

ANOMALIES

Be 1s2 2s2

B 1s2 2s2 2p1 - higher nuclear charge < 2p orbital has more energy than 2s & more shielding

similar with Mg and Al

N 1s2 2s2 2px1 2py1 2pz1

O 1s2 2s2 2px2 2py1 2pz1 less energy required to remove 2px electron due to shielding within orbital > nuclear charge

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CHEMICAL BONDING AND STRUCTURE

Metallic Bonding

typically have 1-3 outer electrons and have low ionisation energies, the electrical conductivity generally increases as the no. of outer (delocalised) electrons increases.

Delocalised Electrons are electrons not associated with any single atom or any single covalent bond.

Metallic Bonding is the electrostatic force of attraction between the nuclei of metal cations and delocalised electrons

MELTING TEMPERATURE

high as the energy needed to overcome the attraction between the nuclei of the cations and delocalised electrons so that the cations are free to move around the structure

group 1 metals have low melting temperatures

group 2 metals have higher melting temperatures

d block metals have high melting temperatures (most delocalised electrons per ion)

the smaller the cation, the closer the electrons, increasing forces of attraction therefore increasing melting temperature

ELECTRICAL CONDUCTIVITY

when a p.d is applied to a metal, the delocalised electrons will move towards the positive terminal of the cell, causing a movement in electrical charge and therefore an electrical current

THERMAL CONDUCTIVITY

free-moving delocalised electrons pass kinetic energy along the metal

closely packed cations pass kinetic energy amongst themselves

MALLEABILITY & DUCTILITY

hammered/pressed & drawn into a wire

stress causes layers of cations to slide and electrons move with them (free-moving) preventing strong forces of repulsion between cations

Ionic Bonding

solid materials consisting of a regular array of oppositely charged ions extending through a giant lattice network

strong electrostatic forces between the the ions. arranged so that attraction > repulsion between charges.

STRENGTH

the smaller the ions, the stronger the electrostatic attractions

the higher the charges, the stronger the electrostatic attractions

IONIC RADII

going down the group, there are more electron shells so they ions get larger

going across a period, there are more protons so ionic radius gets smaller

as the positive charge of the nucleus increases, the electrons are pulled closer to the nucleus

Properties

HIGH MELTING TEMPERATURE

large amount of energy is required to overcome the forces of attraction sufficiently for the ions to break free from the lattice and be able to slide past one another

BRITTLENESS

when a stress is applied to a crystal of ionic solid, the layers of ions may slide over one another, now with ions of the same charge side by side, they will repel and break apart the crystal

ELECTRICAL CONDUCTIVITY

generally don't as there are no delocalised electrons/ free ions

BUT molten/aqueous will conduct as ions will migrate to opposite signs when a p.d is applied. if a d.c is applied, the compound will undergo electrolysis and ions will discharge at the electrodes.

Li3N is an exception and is used in batteries

SOLUBILITY energy required to break apart the lattice structure can be supplied by the hydration of the separated ions produced. Both positive and negative ions are attracted to water molecules due to the polarity of water molecules

Formation of Cations and Anions

Sodium chloride can be formed by burning sodium in chlorine:

2Na(s) + Cl2(g) --> 2NaCl(s)

2Na --> 2Na+ + 2e- and Cl2 + 2e- --> 2Cl-

Ion existence evidence:

ability of an ionic compound to conduct electricity and undergo electrolysis when molten/in aqueous solution

e.g when d.c is applied through molten sodium chloride, sodium is formed at negative electrode and chlorine is formed at the positive electrode

2NaCl --> 2Na + Cl2

passing direct current through copper(II) chromate(VI) solution

Cu2+(aq) are blue and CrO42-(aq) are yellow

Covalent Bonding

formed by the overlap of 2 atomic orbitals each containing a single electron

sigma (σ) is an end-on overlap

single covalent bond between two atoms

can occur with s and p orbitals

highest electron density between two nuclei

pi (π) is a sideways overlap of two p-orbitals.

it can't form until a σ bond has been formed

atoms with double/triple bonds

high electron density above and below the molecule

in ethene (above), one of the bonds between the carbon atoms is sigma, the other two are pi.

the pi bond is weaker than the sigma bond increasing the reactivity of alkenes compared with alkanes and why alkenes readily undergo addition reactions

Bond Strength and Length

bond length is the distance between two nuclei of the two atoms that are covalently bonded together

the strength is measured in terms of the amount of energy required to break one mole of the bond in the gaseous state

for bonds that are of similar nature, generally, the shorter the bond, the stronger this is a result of the increased electrostatic attraction between the two nuclei and the electrons in the overlapping atomic orbitals.

Electronegativity and Bond Polarity

electronegativity is the ability of an atom to attract a bonding pair of electrons in a covalent bond

if two atoms of the same element are bonded together by the overlap of atomic orbitals, the distribution of electron density between the two nuclei will be symmetrical, as each atom has an identical ability to attract the electrons.

if the two atoms bonded together are from elements of different electronegativities the distribution of electron density won't be symmetrical about the two nuclei

the contour lines will be more closely spaced near to the chlorine atom as it has the higher electronegativity and acquires a slight negative charge and the H will gain a slight positive charge. this is a polar covalent bond (a type of covalent bond between two atoms where the bonding electrons are unequally distributed. because of this, one atom carries a slight negative charge and the other a slight positive charge)

there is a continuum of bonding type. from 100% covalent to 100% ionic. a polar covalent bond will be somewhere in between with some degree of ionic character

Bonding in Discrete Molecules

A discrete (simple) molecule is an electrically neutral group of two or more atoms held together by chemical bonds

covalent and polar covalent bonding in discrete molecules can be shown by dot-and-cross diagrams

the 'octet rule' that in order to form a stable compound, the outer shell of each atom must have 8 electrons. this is not always true.

A displayed (full structural) formula shows each bonding pair as a line drawn between the two atoms involved

Dative Covalent Bonds

formed when an empty orbital of one atom overlaps with an orbital containing a non-bonding pair (lone pair) of electrons of another atom

in the gas phase, just above its boiling temperature, aluminum chloride exists as Al2Cl6 molecules, one chlorine atom from each molecule acts as a bridge connecting the two molecules with dative covalent bonds

Shapes of Molecules and Ions

the electron pair repulsion theory (EPR) / valence shell electron pair repulsion theory (VSEPR) states that:

the shape of a molecule is caused by repulsion between the pairs of electrons, both bond and lone pairs that surround the central atom

the electron pairs arrange themselves around the central atom so that repulsion between them is at a minimum

lone-lone > lone-bond > bond-bond (pair repulsion)

SHAPES/ANGLES

Non-polar & Polar Molecules

the drift of bonded electrons towards the more electronegative element results in a separation of charge called a dipole (when two charges of equal magnitude but opposite signs are separated by a small distance).

each of the bonds in a molecule has its own dipole associated, the overall dipole depends on its shape. depending on the relative angles between the bonds, the individual dipoles can either reinforce (Polar) or cancel out (Non-polar)

Diatomic Molecules:

e.g Hydrogen and Chlorine are non-polar diatomic molecules as the two atoms have equal electronegativities

BUT the bond in Hydrogen Chloride is polar since Chlorine has a higher electronegativity. as this is the only polar bond in the molecule, the molecule itself is polar

Polyatomic Molecules:

in CO2, both bonds are polar but they cancel each other out

in CHCl3, all bonds are polar and the molecule is symmetrical but the dipoles reinforce each other so the molecule is polar

Intermolecular Interactions

considerably weaker than typical covalent/polar covalent bonds

they are intermolecular as they occur between the molecules (not within)

LONDON FORCES

e.g. two non-polar molecules of nitrogen, A and B

although they are non-polar, their electron density can fluctuate over time. if at any instant the electron density in molecule A becomes unsymmetrical in a molecule, a dipole will be generated (instantaneous dipole). inducing a favourable dipole in molecule B (induced dipole)

the two dipoles will be necessarily arranged so that they will interact favourably with each other, causing the London force of attraction

the attractive force increases with increasing number of electrons in the molecule, as it increases the fluctuation in electron density/dipoles

they also depend on the shape and size of the molecules, the more points of contact, the greater the force

they are always present between molecules, regardless of whether they have a permanent dipole/H-bond

PERMANENT DIPOLES

if the molecules possess permanent dipoles, they will also interact, if they are aligned correctly, then there will be a favourable interaction and they will attract each other, however due to the random movement of molecules, the dipoles are not always aligned to produce a favourable reaction and will repel. the interaction between permanent dipoles < instantaneous-induced

HYDROGEN BOND

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CHEMICAL ENERGETICS

Enthalpy & enthalpy change

Chemical Energy

Kinetic energy- measure of motion of particles in a substance

Potential energy- measure of how strongly these particles interact with each other

Heat Energy

the portion of kinetic and potential energy responsible for temperature of a substance

Enthalpy

measure of the total energy of a system, can't be directly determined BUT can measure the enthalpy change (ΔH) that takes place in a physical/chemical change.

ENTHALPY CHANGE is the heat energy that is transferred from the system to the surroundings at constant pressure

exothermic heat energy is transferred from the system to the surroundings. e.g. HCl(aq) + NaOH(aq) --> NaCl(aq) + H2O(l)..........ΔH = -57.1kJ mol-1

endothermic heat energy is transferred from the surroundings to the system.

e.g. C6H8O7(aq) + 3NaHCO3(s) --> C6H5O73-(aq) + 3Na+(aq) + 3CO2(g)+ 3H2O(l)

the 'system' refers to the reaction mixture, everything outside the system is called the 'surroundings'

any reaction that has to be continually heated in order for it to take place is ENDOTHERMIC

Standard Conditions

100 kPa pressure and a stated temperature, most commonly 298 K

under these conditions, the enthalpy change measured is called the 'standard enthalpy change' given the symbol ΔH°

Standard enthalpy change of reaction ΔrH°

the enthalpy change measured at 100 kPa and at a stated temp. usually 298 K, when the number of moles of substances in the equation as written react

for the reaction:

N2(g) + 3H2(g) --> 2NH3(g) ΔrH°= -92kJmol-1

for the reaction when written as:

1/2N2(g) + 1 1/2H2(g) --> NH3(g) ΔrH°= -46kJ mol-1

in each case, 'per mole' refers to one mole of equation and not one mole of any reactant or product

Enthalpy level diagrams

Change in enthalpy ΔH is given by:

ΔH = H(products) - H(reactants)

exothermic- H(r) > H(p) so ΔH is negative

endothermic- H(r) < H(p) so ΔH is positive

Constructing enthalpy level diagrams:

only need to label horizontal axis (Enthalpy, H)

formulae for both reactants and products should be given, including state symbols

values for ΔH should be given, including the correct sign

it isn't essential to show the activation energy in enthalpy level diagrams BUT should be shown in enthalpy profile diagram

Standard enthalpy change of combustion

standard enthalpy change of combustion, (ΔcH°) is the enthalpy change measured at 100 kPa and a stated temperature, usually 298 K, when one mole of substance is completely burned in oxygen

combustion of Hydrogen:

H2(g) + 1/2 O2(g) --> H2O(l) represents ΔcH°

or

2H2(g) + O2(g) --> 2H2O(l) represents 2x ΔcH°

Experimental determination of enthalpy change in combustion of a liquid

a known mass of the liquid is burned and the heat energy produced is used to heat a known volume of water

a spirit burner containing the liquid under test is weighed

a known volume of water is added to a copper can

the temperature of water is measured

the burner is lit

the mixture is constantly stirred with a thermometer

when the temp. has reached approximately 20°C above its initial temperature, the flame is extinguished and the burner is immediately reweighed

the final temp is measured

calculating enthalpy change of combustion

calculate the heat energy, Q, transferred to the water using the equation Q=mcΔT, where m is the mass of the water and c is the specific heat capacity of water. if m is in grams, then c is quoted as 4.18Jg-1K-1 (assuming the density of water is 1gcm-3, then m=100.0g) the temperature change ΔT has the same value in K as it does in °C so, Q = 100.0g x 4.18Jg-1K-1 x +24.5 K = +10241J = +10.24kJ

calculate the amount, n, of ethanol burned n (C2H5OH) = 0.42g/46.0 gmol-1 = 9.13x10-3 mol

calculate ΔcH using the equation ΔH = Q/n ΔcH°= - +10.24kJ/ 9.13x10-3 mol = 1120 kJmol-1 (3 s.f)

possible source of error/assuptions made

some of the heat energy produced in burning is transferred to the air/copper can and not the water

some of the ethanol may not burn completely to form carbon dioxide and water (incomplete combustion=soot in can)

the conditions are not standard. e.g. water vapour, not liquid water is produced

as the experiment takes a long time, not all of the heat energy transferred from the water to the surroundings is compensated for

Standard Enthalpy Change of Neutralisation

standard enthalpy change of neutralisation, (ΔnH°), is the energy change measured at 100kPa and at a stated temperature, usually 298K, when one mole of water is produced by the neutralisation of an acid with an alkali.

e.g. neutralisation of hydrochloric acid:

HCl(aq) + NaOH(aq) --> NaCl(aq) + H2O(l)

making the assumption that strong acids and alkalis are fully ionized in aqueous solution, then the reaction between them is simplified to:

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

experimental determination of enthalpy change of neutralisation

wear safety glasses & lab coat

using a pipette fitted with a safety filter, place 25.0cm3 of 1.00 moldm-3 acid into an expanded polystyrene cup

measure the temperature of the acid

using a pipette fitted with a safety filter, place 25.0cm3 of the alkali (usually dilute sodium hydroxide of a concentration slightly greater than 1.00 mold-3 to make sure all the acid is neutralised) into a beaker

measure the temperature of the alkali

add the alkali to the acid, stir with the thermometer and measure the maximum temperature reached

possible sources of error/assumptions made

answers should be given to only 2 s.f as ΔT is given to only 2 s.f

uncertainties of measurements involved with the use of the pipette and thermometer

some heat energy will be transferred to the thermometer and polystyrene cup

Standard Enthalpy Change of Formation & Hess's Law

standard enthalpy change of formation, (ΔfH°), is the enthalpy change measured at 100kPa and a specified temperature, usually 298K, when one mole of a substance is formed from its elements in their standard states

e.g the standard enthalpy change of formation of a gaseous carbon dioxide is the enthalpy change for the reaction:

C(s,graphite) + O2(g) --> CO2(g) ΔfH°= -394kJmol-1

the standard enthalpy change of formation of liquid ethanol is the enthalpy change for the reaction:

2C(s,graphite) + 3H2(g) +1/2O2(g) --> C2H5OH(l) ΔfH°= -278kJmol-1

Hess's Law

most standard enthalpy changes of formation cannot be determined experimentally

e.g the enthalpy change for the reaction

C(s) + 1/2O2(g) --> CO(g)

is impossible to determine directly so we use Hess's Law (stating that the enthalpy change of a reaction is independent of the path taken in converting reactants into products, provided initial and final conditions are the same in each case) allowing us to calculate the standard enthalpy change of formation of carbon monoxide

the enthalpy changes of combustion of carbon and carbon monoxide are:

C(s, graphite) + O2(g) --> CO2(g) ΔcH°= -394kJmol-1

CO(g) + 1/2 O2(g) --> CO2(g) ΔcH°= -283kJmol-1

there are two ways to use this data to calculate the standard enthalpy change of formation of carbon monoxide:

SUBTRACT EQUATIONS reverse the second equation and then add to the first: CO2(g) --> CO(g) + 1/2 O2(g) ΔH°= +283kJmol-1 C(s, graphite) + O2(g) --> CO2(g) ΔH°= -394kJmol-1 adding the two equations gives: C(s, graphite) + 1/2 O2(g) --> CO(g) ΔH°= +283 + (-394) = -111 kJmol-1 so ΔfH°= -111 kJmol-1

construct an enthalpy cycle using Hess's Law (paper)

Bond Enthalpy & Mean Bond Enthalpy

bond enthalpy, (ΔBH°), is the enthalpy change when one mole of a bond in the gaseous state is broken

for a diatomic (XY) molecule, the bond enthalpy is the enthalpy change for the following reaction:

XY(g) --> X(g) +Y(g)

Cl2(g) --> 2Cl(g) ΔBH°= +243kJmol-1

HCl(g) --> H(g) + Cl(g) ΔBH°= +432kJmol-1

or polyatomic molecules, each bond has to be considered separately, for methane, there are 4 separate bond enthalpies:

CH4(g) --> CH3(g) + H(g) ΔBH°= +423kJmol-1

CH3(g) --> CH2(g) + H(g) ΔBH°= +480kJmol-1

CH2(g) --> CH(g) + H(g) ΔBH°= +425kJmol-1

CH(g) --> C(g) + H(g) ΔBH°= +335kJmol-1

mean bond enthalpy is the enthalpy change when one mole of a bond, averaged out over many different molecules, is broken

the mean bond enthalpy for the C-H bond/ E(C-H) in methane is approximately +416 kJmol-1 [1/4(423+480+425+335)= 415.75]