HL CHEM 2: Atomic Structure BY HEIMAN KWOK 12N03S 22.8.13.

HL CHEM 2: Atomic Structure

BY HEIMAN KWOK 12N03S22.8.13

2.1 THE ATOM

RELATIVE ATOMIC MASS; RELATIVE CHARGE

LOCATED IN THE NUCLEUS

Atomic Number and Mass Number• Atomic Number (Z) = # of protons in its

nucleus (number of electrons at ground state); defines the element + its position in the periodic table

• Mass Number (A) = sum of # of protons and # of neutrons in its nucleus

Isotopes and Relative Atomic Mass

• Isotopes: atoms of the same element (same atomic number/ number of protons) but with different number of neutrons (different mass number)

• Relative Atomic Mass: average mass of an element compared to 1/12g of Carbon-12

Comparing Isotopes

Chemical Properties: SAME Physical Properties: DIFFERENT • Density • Rate of Diffusion • Boiling Point • Melting Point

Carbon-14 Radioactive Dating – archaeological objects

• Carbon-14 has too many neutrons (too stable) • It emits beta particles (neutron changes to

proton and an electron; electron is ejected from the atom) to reduce the neutron-to-proton ratio

• Relative abundance of Carbon-14 falls owing to nuclear decay after organisms die

• Rate of decay is measure by its half-life

Radioactive Dating

• (living things contain) carbon (14) • C14 is a radioactive isotope / beta emitter • (radio)activity decreases (over time) • (estimate) half-lives (since material was alive) • compare activity (of sample now with living

tissue) / ratio of C14 to C 12 is fixed in living material

Cobalt-60 in Radiotherapy

• Treats localized solid tumors by damaging genetic material inside a cell by knocking off electrons, making it impossible for the cell to grow

• Although kills both normal and cancerous cells, normal cells are able to recover if the treatment is carefully controlled

• Colbalt-60 commonly used as it emits very penetrating gamma radiation when its protons and neutrons change their relative positions in the nucleus

Iodine-131, Iodine-125 Medical Tracer

• Isotopes (Radio) have the same chemical properties hence they play the same role in the body, yet by detecting radiation, their positions can be monitored

• Sodium Iodide (Iodine-131) – emits beta and gamma rays to diagnose and treat thyroid cancer

• Iodine-125 – longer half life of 80 days allowing low levels of beta radiation to be emitted over an extended period – treatment of prostate cancer

Other Uses of Radioactivity

• Works by injecting a radioisotope, then detecting it with a G.M counter – chosen isotope must have a short half-life but remains long enough to be detected

1. Radioactive Tracers – finding leaks from a pipeline, Fertiliser, human body

3. Sterilising – kill bacteria, moulds and insects in food – prolongs shelf life, preservation; sterilizes medical equipment (eg. Plastic Syringes and packaged food)

Other Uses of Radioactivity

4. Quality Control – measuring how much ionizing radiation passes through then adjusting the thickness accordingly 5. Smoke Detection – Weak source emits alpha particles which ionizes the air and conducts electricity and a small current flows; when there is smoke, it absorbs the alpha particles, current reduces and the alarm sounds

Dangers of Ionising Radiation and Unstable Isotopes

• Mutations in living organisms • Damage to cell and tissue • Problems arising the disposal of radioactive

waste – difficult to dispose hence ^^• How associated risks can be reduced – using

thick sheets of lead/ thick lead container HENCE need for close international cooperation to ensure high safety standards

2.2 THE MASS SPECTROMETER

1) Vaporization - Mass Spectrometer

• A vaporised sample must be used • Sample heated to a gaseous state – so atoms are

separate and free to move independently

• MS is a high vacuum (maintained by slowly transferring sample via airlock) – so ions have a free path and do not collide with air molecules

• MS is at a very low pressure – prevent collisions and to avoid false readings due to presence of other particles

2) Ionisation – Mass Spectrometer

• Bombardment by controlled beam of high-energy (speed) electrons from an electrically heated metal coil (electron gun)

• Collide with electrons in the particle, knocking them out to form positive ions

• REASON 1) ions accelerated by electric field and 2) ions deflected by magnetic field

• X (g) + e- X+ (g) + 2e-

3) Acceleration – Mass Spectrometer

• Acceleration by passing through an electric field in an electrostatic analyser created by oppositely charged plates with high potential difference

• so all ions have the same kinetic energy repelled out of a hole in a negatively charged plate

4) Deflection – Mass Spectrometer

• The positive vaporized ions are deflected by the passing through a magnetic field; Deflection depends on

1. Mass/charge (m/z or m/e) ratio – the greater the charge, lower the mass, the greater the deflection

2. Strength of the magnetic field 3. Momentum and velocity of the ions –

controlled by strength of the electric field

5) Detection – Mass Spectrometer

• Ions are detected by production and conversion into an electric current

• Ratio of intensity of the peak in the spectrum is equal to the ratio of ions in the sample

• REMARK: Other ions are either deflected too much or too little to have collided into the walls of the MS – picked up electrons and neutralised before reaching the detection plate

Applications of the Mass Spectrometer

1. Determine the mass and relative abundance of different isotopes

2. Identify and quantify a new and (un)known compound – even on different planets

3. Detect and identify use of steroids in athletes 4. Determine if honey is altered with corn syrup 5. Monitor breath of patients by

anaesthesiologists in surgery

Mass Spectra

Results of the analysis are presented as a ^ to provide us the with following information1. No. of isotopes - no of peaks 2. Relative isotopic mass (1/12th of a 12C) – m/z3. Relative abundance of each isotope – height

of the peak = relative atomic mass of the element

Interpreting the Mass Spectra

2.3 ELECTRON ARRANGEMENT

Continuous Electromagnetic Spectrum• Al EM waves travel at the same

speed in a vacuum/ free space (without a medium) at 300,000 km/s (300,000,000 m/s)

• All EM waves are Transverse Waves

• As we go along the EM Spectrum; Wavelength decreases and Frequency increases

Difference between Continuous and Line Spectrum

• Continuous Spectrum – shows an unbroken sequence of all frequencies/ wavelengths/ colors such as the spectrum of visible light

• Line Spectrum (Discontinuous Spectrum) – energy of the atom is quantised, only certain discrete frequencies of light are shown.

• It is produced by excited electrons as they fall back to a lower energy level

• unique to each element – as they all have slightly different energy levels and electron shells

Absorption Line Spectra

• White light passed through a gas • Atom moves from ground state to excited

state – energy absorbed • (result of electron transitions)

• NOTE THAT: Absorption + Emission = Continuous EM SPEC

Hydrogen Absorption Line Spectrum

Emission Line Spectra

• High Voltage applied to the gas • Electron falls from excited state to ground

state – emits radiation – gives out on packet of energy (quantum or photon)

• (result of electron transitions)

Hydrogen Emission Line Spectrum

Planck’s Equation

• Energy of the photon of light emitted = energy change in the atom = Planck’s Constant x frequency

• CONCLUSION: energy depends on the frequency of the wave

Bohr’s Model – Reminders

• Electrons have more energy further away from the nucleus

• Electrons in the small shell have the same amount of energy

• If outermost electron orbit is full – the atom is stable

• Valance electrons = electrons in the outer most energy level

The Hydrogen Spectrum

• When an atom absorbs energy, an electron moves from ground state (lower energy level, n=1) to an higher energy level

• The excited state is produced – unstable, electron soon falls back to the ground state – this emits energy

• When the electron drops back into a lower shell – they emit energy of a particular wavelength, gives out energy and causes a line on the spectrum (emission line spectrum) – flame test

The Hydrogen Spectrum• Energy given out by the

electron (shown on the emission line spectrum) depends on the distance it drops from one shell to another

• Drops to n=1 – most energy – Ultra-violet • Drops to n=2 – visible light (what we can see!)• Drops to n=3,4,5 – least energy – Infrared

• n=∞ - electron leaves the atom – atom is positively ionised

Q: Emission Spectrum of Hydrogen

Outline how this spectrum is related to the energy levels in the hydrogen atom:• Series of lines/lines - lines up with the concentric energy

levels (n=1 - 5)• Electron transition between higher energy level to lower

energy level OR transition into 1st energy level causes UV series / transition into 2nd energy level causes visible series / transition into 3rd energy level causes infrared series;

• Convergence at higher frequency/energy/short wavelength;

Ionisation Energy

• Ionisation energy = energy needed to remove an electron from the ground state of each atom in a mole of gaseous atom, ions or molecules

Red, Green, Blue-purple,

Pink and Black

12.1 ELECTRON CONFIGURATION

Writing about electrons

• Electron arrangement: up to 20 (Ca) 2,8,8,2

• Electron configuration: up to 52 (Xe - Xenon)• Eg. (1 = energy level, s = orbital shape type, 2

= no of electrons)

Heisenberg’s Uncertainty Principle

• FUNDAMENTAL PROBLEM: Bohr’s Atomic Model assumes that one can precisely describe the electron’s trajectory – yet it is impossible as any attempt to measure its position will disturb its motion – focusing radiation to locate the electron gives the electron a ‘random kick’ which sends it hurtling off in a random direction

Orbital (sub-shell) rules

• ^ Define: orbital – the region in space where there is a 95% probability of finding the electron

• In each every level, there are 4 types of sub-shells – s, p, d and f (in order of ascending energy)

• Each orbital can hold a max of 2 electrons as long as they have opposite spin (Pauli exclusion principal)

Orbital (sub-shell) rules

• Electrons enter the lowest available energy level (Aufbau principle)

• When in orbitals of equal energy, electrons will try to remain unpaired (Hund’s rule of maximum multiplicity)

• Within each energy level, max 1 s, 3 p, 5 d and 7 f

Orbital shapes – learn to draw• Orbitals have different shape depending on

which sub-shell they are in – so they all fit spatially around the nucleus, no crash

• S-orbital – spherical path ---------- >• P-orbital – dumbbell path

Orbital shapes – recognise

• 5 d orbitals – double dumbbells

• 7 f orbitals – very complex

Glitch in the system

• Orbitals overlap – as energy levels converge as one gets further away from the nucleus

• First ‘glitch’ in the periodic table: 4s is physically further away from the nucleus but has lower energy than 3d yet 4s is still part of the 4th energy level – electron fill 4s before 3d• Electron Configuration of Ti (Z = 40)

Filling Orbitals

• Way to remember which orbital to fill in first

• Ie. 4s is filled in before 3d

• EXTRA: [Ar] is Zn shortened way for electric configurations after Argon

Special Cases – Cr and Cu

• Chromium Cr

• Copper Cu

Arrangement of six unpaired electrons has a lower energy (and requires less) than if two electrons (in the 4s orbital) are paired (repelling each other) – more stable as 3d sub-shell is symmetrical

Ionised electronic configurations

• Electrons are removed first from the occupied orbital farthest away from the nucleus as it is easier and requires less energy

• Ions of s and p block elements are ‘isoelectronic’ (same electronic configuration) with a noble gas but contain a different number of protons and are charged

• Na • Na+ • Ne

Ionised electronic configurations

• Initially (when filling) 4s is lower in energy than 3d

• Yet when removing electrons (to become positively charged) 4s is higher in energy – hence its electrons are removed first

• Ti

First Ionisation Energy

• Minimum energy (in kJ mol⁻¹) per mole needed to remove the first electron from a neutral atom; in the gaseous state;

- REMEMBER STATE SYMBOLS AND + AND -

• Periodicity: repeating pattern of (physical and chemical) properties

Size of first ionisation energy

• Size of first ionisation energy depends on: 1. Nuclear Charge – greater no. of protons, greater

the size of charge, increase in IE 2. Distance – greater the distance between the

nucleus and the outermost electrons, the less the attractive force, decrease in IE

3. Shielding – repelling force from the filled inner orbitals makes the electrons in the outer orbitals easier to remove, decrease in IE; paired electron is easier to remove due to repelling force

As the charge on the nucleus increases so the energy required to remove the electron increases.As the distance between the outermost electron and the nucleus increases so the energy required to remove it decreases

This graph shows how the first ionisation energy varies moving from element to element in the periodic table. The outermost electron is being removed in each case and so the amount of energy needed to remove it is a function of the force holding the electron in position around the atom.

First Ionisation Energy across a period

• Na 3s¹ to Mg 3s² - up: up nuclear charge • Mg to Al 3s²3p¹ - down: up distance, up shielding > up

nuclear charge • Al to Si 3s²3p² to P 3s²3pᶾ - up: up nuclear charge • P to S 3s²3p⁴ - down: electrons are paired – up shielding

+ repelling force • S to Cl 3s²3p⁵ to Ar 3s²3p⁶ - up: up nuclear charge

• Suggests and amounts for the existence of sub-levels within the main energy levels

Evidence for energy levels and subshells

1. Highest values of IE for noble gases / lowest values of IE for alkali metals;

2. General increase in IE across a period;3. drop in I.E. from Be to B/Mg to Al/Group 2 to

Group 3;4. drop in I.E. from N to O/P to S/Group 5 to

Group 6;

Successive Ionisation Energy

• Successive electrons (are more difficult to remove because each is) taken from more positively charge ion; and experience increased electrostatic attraction (effective nuclear charge) (and decreased shielding)

Large increases – change in sub-shell/ energy level

Example: 2nd to 3rd – 3s to 2p; 8th to 9th – 2p to 2s;

10th to 11th – 2s to 1s• 10th electron comes from 2nd energy level

(n=2) and the 11th electron comes from 1st energy level (n=1)

• Electron in 1st energy level is closer to nucleus; not shielded by inner electron and are exposed to greater effective nuclear charge;

Successive Ionisation Energy Graph

• From the graph one can deduce the existence of energy levels; its period and group by the number of electrons before a big jump in its ionisation energy

Explaining: Sodium Na - 1s22s22p63s1

• First electron easiest to remove as it is the furthest from the nucleus in the outermost energy level, n = 3

• Large increase between 1st and 2nd IE as electron now removed from the second energy level, n = 2;

• Next 8 electrons more difficult to remove as they are in the same energy level (2nd, n=2) and there is an increase in nuclear charge as the ion is increasingly positive

• large increase between 9th and 10th IE as electron now removed from n = 1 OR last two electrons are the most difficult to remove as they are in the innermost/lowest energy level (n=1) and are closest to the nucleus

• Electron 11 also comes from 1s, so shows a small increase;

It gets easier to remove electrons down the group

• Mg outer electron is the 3s sub-shell rather than 2s of Be – which is higher in energy

• 3s is further away from the nucleus – is shielded more and less force of electrostatic attraction from the nucleus to the electron

• 3s is more easily removed

Commenting on trends

• Number of outer electrons? Full or not? • Electron configuration? Electron paired? • Energy level and sub-shell? • All the electrons are unpaired – it requires

more energy to pair up the electrons in S so it has a lower ionisation energy

• - more stable, higher ionisation energy• - less stable, slightly lower lE

Group 2 – ionisation energy

• For elements that reach a filled or half filled sublevel by removing 2 electrons 2nd IE is lower than expected; easier to remove (eg. Ca)

• Makes it easier to achieve a full outer shell • True for 2s – hence alkaline earth metals

(group 2) from 2+ ions

Group 3 – ionisation energy

• S2p1 atoms have a low 3rd IE • 2nd IE and 3rd IE are always higher than 1st IE • Atoms in the Al family from 3+ ions