3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3.

As you go across period 3, nature of oxides changes from basic to acidic.

Oxide	Acid or base
Na2O	Basic Giant Ionic
MgO
Al2O3 - giant ionic	Amphoteric
SiO2 - giant covalent
P4O10	Acidic Molecular covalent
SO3/SO2
Cl2O7

Amphoteric: 

• posses properties of both acids and bases.
• when it reacts with an acid, it behaves as a base, and vice-versa. 
• tested with aluminum oxide
• When aluminum oxide reacts with an acid such as hydrochloric, aluminum oxide acts as a base.
• Al₂O₃ (s) + 6HCl (aq) → 2AlCl₃ (aq) + 3H₂O (l)
• Now, when aluminum oxide reacts with a base such as Sodium hydroxide, it acts as an acid.
• Al₂O₃ (s) + 2NaOH (aq) + 3H₂O (l) → 2NaAl(OH)₄

Reactions between Na₂O, MgO, P₄O₁₀  and SO₃  and water.

• Na₂O (s) + H₂O (l) →2NaOH (aq)
• Sodium oxide + water → Sodium hydroxide
• MgO (s) + H₂O (l)→ Mg(OH)₂ (aq)
• Magnesium oxide + water → Magnesium oxide
• P₄O₁₀ (s) + 6H₂O (l) → 4H₃PO₄
• Phosphorous (v) oxide + water → Phosphoric (V) acid
• SO₃(l) + H₂O (l) → H₂SO₄(aq)
• Sulfur trioxide + water → Sulfuric (VI) acid.

Across a period 3, there is a change from metallic to non-metallic. 

• Left. ex) Sodium, Magnesium and Aluminium all have the typical properties of metallic substances (E.g. good conductors of electricity), 
• Righ. ex) such as Chlorine and Sulfur are non-metals.

Metals will react to form ionic compounds. 

Non-metals will react to form covalent compounds. 

3.3.1 Discuss the similarities and differences in the chemical properties of elements in the same group.

Group 1

• Group 1 metals = Alkali Metals
• share similarities and differences in their chemical properties.
• reactive, especially with air and water 
• often stored in paraffin to avoid contact with air
• react quite vigorously with water
• ex) lithium with water will most likely lead to a slow reaction
• ex) caesium with water is almost an explosion
• reactivity increases down the group because of decreasing electrostatic attraction to outer shells, therefore electrons can be lost easily.
• good reducing agents because they donate electrons
• the reaction is exothermic (releases energy) 
• metal hydroxide and hydrogen
• first two float
• might ignite and produce a flame (ex) potassium = violet flame 

Alkali Metals reactions with water

• A = generic alkali
• 2A (s) + 2H2O (l) → 2AOH (aq) + H2(g)
• Alkali metal + Water →Metal hydroxide + hydrogen
• We can also write this reaction in its ionic form, as alkali metals form ionic compounds:
• 2A (s) + 2H2O (l) →2A+ (aq) + 2OH- (aq) + H2 (g)

Group 7: Halogens

• 7 outer electrons
• reactivity decreases down the group because the distance between the nucleus and outer-electron increases as you go down the group
• are good oxidizing agents because they accept electrons easily

Halogens reaction with alkalis

• A = generic alkali
• Ha = generic halogen
• Halogen reaction with alkali's
• 2A (s) + HA2→2AH2 (s) = 2KBr +Cl2 → 2KCl +Br2 
• chlorine = gaseous state
• bromine = liquid state
• iodine = solid state 
• after reaction known as a halide
• produce a salt
• we often call this “Displacement reactions”

Group 0/8 : Nobel Gases

• colorless gases
• monoatomic 
• very unreactive because they do not gain or lose electrons
• stable octect (full outer shell)

3.2.4 Compare the relative electronegativity values of two or more elements based on their positions in the Periodic Table.

The most electronegativie element is on the top right of the Periodic Table and the least electronegative element is on the bottom left.

• Increases from left to right across a period because of increase in nuclear charge, resulting in an increased attraction between the nucleus and the bond electrons. 
• Electronegativity decreases down a group. The bond electrons are furthest from the nucleas and so there is reduced attraction. 

• Fluorine has the highest value of 4.0.
• Ceasium has the lowest value of 0.7.

3.2.2 Describe and explain the trends in atomic radii, ionic radii, first ionization energies, electronegativities and melting points for the alkali metals (Li to Cs) and the halogens (F to I). 3.3.3 Describe and explain the trends in atomic radii, ionic radii, first ionization energies and electronegativities for elements across Period 3.

Effective nuclear charge

• Nuclear charge - Number of protons an atom contains. 
• 
  
• The inner electrons reduce the attraction of the nucleus for the outer electrons, therefore the outer electrons are shielded from the nucleus and repelled by the inner electrons. 
• 
  
• The effective charge experienced by the outer electrons is less than the full nuclear charge. 

Element	Na	Mg	Al	Si
Nuclear Charge	11	12	13	14
Electron Arrangement	2, 8, 1	2, 8, 2	2, 8, 3	2, 8, 4

From left to right, across a periods, one proton is added to the nucleus, there fore its added to the outer most shell, the effective charge increases with nuclear charge because there is not change in the number of inner electrons.

Element	Nuclear charge	Electron arrangement
Li	3	2, 1
Na	11	2, 8, 1
K	19	2, 8, 8, 1

As we descend the group, the increase in the nuclear charge is largely balanced by the increase in the number of inner electrons, both increase by eight between successive elements. The effective nuclear charge stays the same down the group. 

3.2.1 Define the terms first ionization energy and electronegativity.

First ionization energy - The first ionization energy is the energy required to remove the most loosely held electron from one mole of gaseous atoms to produce 1 mole of gaseous ions each with a charge of 1+. This is more easily seen in symbol terms. 

In simpler terms the first ionization energy is the energy required to remove the first electron from an atom, thus producing an ion with a 1+ charge. 

Electronegativity - Electronegativity is the ability of an atom to attract electrons to itself in a covalent bond. The Pauling scale is the most commonly used. Fluorine (the most electronegative element) is assigned a value of 4.0, and values range down to caesium and francium which are the least electronegative at 0.7.

definitions from www.chemguide.co.uk/

3.1.4 Apply the relationship between the number of electrons in the highest occupied energy level for an element and its position in the Periodic Table.

The block structure of the Periodic Table is based on the electron sub-levels of the atom.

The position of an element in the Periodic Table is based on the sub levels of the highest-energy electron in the ground-state atom.

Element	Period	Group	Electron Arrangement	Electron configuration
Helium	1	0	2	1s2
Lithium	2	1	2,1	1s2
Carbon	2	4	2,4	1s22s22p2
Aluminum	3	3	2,8,3	1s22s22p63s23p1
Chlorine	3	7	2,8,7	1s22s22p63s23p5
Potassium	4	1	2,8,8,1	1s22s22p63s23p64s1
Calcium	4	2	2,8,8,2	1s22s22p63s23p64s2

The number of electrons in the outer shell of elements with higher atomic numbers can be deduced from the group number of the element.

3.1.3 Apply the relationship in between the electron arrangement of elements and their position in the Periodic Table up to Z = 20.

The position of an element is related to the electron arrangement in the atom.

• The period number represents how many shells it has. 
• The group number represents how many valence electron(s) the atom contains.

example) Sodium

• period 3 - sodium has 3 occupied energy levels, (shells).
• group 1 - one electron in the outer shell. 

3.1.2 Distinguish between the terms group and period.

groups - columns of the periodic table

periods - rows of the periodic table 

Groups are numbered from 1 to 7. The gap between group 2 and group 3 is filled by transition elements from the fourth period and onwards. 

3.1.1 Describe the arrangement of elements in the Periodic Table in order of increasing atomic number.

Elements are placed in order of increasing atomic number (Z). Which is the number of protons in the nucleus of the atom. 

2.3.4 Deduce the electron arrangement for the atoms and ions up to Z=20.

Element	Electronic arrangement	Element	Electronic arrangement
Hydrogen	1	Sodium	2,8,1
Helium	2	Magnesium	2,8,2
Lithium	2,1	Aluminum	2,8,3
Beryllium	2,2	Silicon	2,8,4
Boron	2,3	Phosphorus	2,8,5
Carbon	2,4	Sulfur	2,8,6
Nitrogen	2,5	Chlorine	2,8,7
Oxygen	2,6	Argon	2,8,8
Fluorine	2,7	Potassium	2,8,8,1
Neon	2,8	Calcium	2,8,8,2

2.3.3 Explain how the lines in the emission spectrum of hydrogen are related to electron energy levels.

The movement of electrons between the shells is called electron transitions.

The emission and absorption spectra are both the result of electron transitions, they can be used like bar codes to identify the different elements.

When electron transitions take place the energy emitted can be detected and its wavelength measured. This provides information about the relative energies of the energy shell.

One packet of energy or photon is released for each electron transition.

image and the following  information from http://eilat.sci.brooklyn.cuny.edu/cis1_5/Old%20HWs/HW2d_C.htm

Electrons in their shells can receive energy in the form of heat or electricity and jump to higher energy shells (promotion). They cannot remain at these higher levels (excited state) for very long and soon fall back to their original shell (or other shells). When they fall back (relax) they have to lose the energy difference between the two shells. his loss of energy is performed by releasing electromagnetic energy in the form of infrared, visible light or ultraviolet radiation.

In the hydrogen atom (the simplest case with only one electron to 'jump' between shells) the energy emitted appears in several series of lines each series corresponding to electrons falling back to different levels. This is shown in the diagram below.

The Lyman series corresponds to transitions between the higher shells and the lowest shell (ground state).

The energy shells are usually given a letter 'n' to describe the specific energy level. The lowest level is n=1 the second level is n=2 etc.

Transitions from higher shells (n >2) to n=2 produce radiation in the visible region of the spectrum and we can actually see it by splitting the light using a prism or diffraction grating and projecting it onto a screen. 

When an electron falls from a lower to a higher energy level, energy is absorbed and a line in the absorption spectrum is produced. 

When an electron falls from a higher to a lower energy level radiation is given out by the atom and a line in the emission spectrum is produced.

When an atom is at the highest energy n=∞, it is no longer in the atom, it has been ionized. 

Ionization energy - energy needed to remove an electron from the ground state of each atom in a mole of gaseous atoms.

2.3.2 Distinguish between a continuous spectrum and a line spectrum.

Continuous spectrum - produced when white light passes through a spectrum. It shows all the frequencies.

Line spectrum - produced when white light passes through hydrogen gas, it shows selected line frequencies. An absorption spectrum is produced.

• If high voltage is applied to the gas, an emission spectrum is produced.
• Colors in the emission spectra are missing from the absorption spectra.

2.3.1 Describe the electromagnetic spectrum.

• electromagnetic radiation comes in different forms of different energy

• all electromagnetic waves have the same speed but we can tell them apart by their wave lengths 

• different colors of visible light also have different wavelengths 

• frequency - number of waves which pass a particular point in one second

• the shorter the wavelength the higher the frequency 

• speed = frequency x wavelength 

2.2.3 Calculate non-integer relative atomic masses and abundance of isotopes from given data.

Mass Spectra - results of the analysis by the mass spectrometer are presented in the form of mass spectra.


Example)
  The mass spectrum for boron.




Calculations

(11 x 100) + (10 x 23)           1330
-------------------------- =  ---------------- = 13.3
             100                            100

2.2.2 Describe how the mass spectrometer may be used to determine relative atomic mass using the Carbon-12 scale.

The mass spectrometer can be used to measure the mass of individual atoms.

The mass needs to be recorded relative to some agreed standard.

As carbon is a common element (most abundant), and it is easy to transport, store, etc therefore it was chosen as the standard in 1961.

Standard Isotope	Symbol	Relative Atomic Mass
Carbon-12	12      C	12.000
Element	Symbol	Relative Atomic Mass
Carbon	C	12.011
Chlorine	Cl	35.453
Hydrogen	H	1.008
Iron	Fe	55.845

2.2.1 Describe and explain the operation of a mass spectrometer.

The masses of the different isotopes  and their relative abundance can be measured using a mass spectrometer.

5 Basic Operations - V.I.A.D.D

• Vaporization - Vaporized sample is injected so individual atoms can be analysed.

• Ionization - Atoms are hit with electrons which knock out other electrons producing cations.

• Acceleration - Positive ions are attracted to negatively charged plates. They are accelerated by an electrical field and pass through a hole in the plate.

• Deflection - Accelerated cations are deflected by a magnetic field placed at right angles to their path. Ions that are deflected more are the ones with the smaller mass/charge ratio. 

• Detection - cations of a particular mass/charge ratio are detected and a  signal is sent to a recorder. The strength of a signal is a measure of the number of ions with that particular ratio that are being detected. 

image from http://www.wou.edu/las/physci/taylor/gs331/mass_spec.html

2.1.7 Discuss the uses of radioisotopes.

The stability of a nucleus depends on the balance between the number of protons and neutrons. When a nucleus contains too may or too few neutrons, it is radioactive and becomes stable by giving out radiation.

Alpha Particles - emitted by nuclei with too many protons to be stable, (are helium atoms).

Beta Particles - emitted by nuclei with too many neutrons, (are fast moving electrons).

Gamma rays - form of electromagnetic radiation, (are waves).

Radioactive isotopes can be used for;

• preserve food
• detect cracks in structural materials 
• generate energy in nuclear power stations 
• sterilize surgical  instruments in hospitals 

USES OF ISOTOPES 

Carbon-14 used for dating materials

• unstable isotope of Carbon-12
• present in living plants
• continually replenished from carbon present in carbon dioxide in the air  
• when organisms die, no more carbon-14 is absorbed and the levels of carbon-14 fall because of nuclear decay
• this process occur at a regular rate, it can be used to date carbon containing materials
• rate of decay is measured by its half life

Cobalt-60 used in radiotherapy

• treatment of cancer and other diseases with ionization radiation
• treatment damages genetic material inside a cell by knocking off electrons and making it impossible for the cell to grow 
• normal cells are able to recover if treatment is carefully controlled
• cobalt-60 emits very penetrating gamma radiation

Iodine-131 as a medical tracer

• radioisotopes have the same chemical properties as their stable atom, but their positions can be monitored by detecting radiation levels, making them suitable for medical tracers
• emits both beta particles and gamma rays 
• short half life (of 8 days), so can be eliminated from the body quickly

Examples of iodine isotopes 

• compound sodium iodide to investigate the activity of thyroid gland and to diagnose and treat thyroid cancer
• iodine-125 used in treatment of prostrate cancer. Pellets of the isotope are implanted into the gland. (Half life of 80 days).

Dangers 

• living organisms can be affected if they are exposed to uncontrolled radiation, (excessive treatment).