Radioactivity is defined as the spontaneous breaking up of certain unstable nuclei, accompanies by the emission of radiation

Radioactivity is detected by a Geiger-Muller tube.

α particles (alpha)

Alpha particles are helium nuclei with a positive charge and little penetrating energy

Alpha decay involves:

• 2 protons lost – atomic number -2
• 2 protons + 2 neutrons lost – mass number -4
• New element formed

β particles (beta)

Beta particles are electrons with a negative charde and greater penetrating ability than alpha particles

Beta decay involves:

• Neutron breaks up into 1 proton and 1 electron
• Proton stays – atomic number +1
• Neutron replaced by proton – mass number no change
• New element formed

γ rays (gamma)

Gamma rays are high energy electromagnetic radiation, with greater penetrating power than beta particles. Neutral charge.

• Gamma rays are a form of energy, not particles – no mass or charge
• No new element formed
• Energy emitted

Radioactive Reactions

Nuclear reactions cause elements to change into other elements – changes in nucleus

Chemical reactions involve changes in distribution of electrons forming compounds – no new elements

Radioisotopes are unstable radiosactive isotopes e.g. carbon-14

Half-life of a radioactive isotope is the time take for half of the atoms in a sample of the isotope to decay

Background radiation is the low level of ionising radiation surrounding us – mainly radon gas from rocks and soil

Uses for radioisotopes

1. Archaeology

• Used to determine age of objects comtaining carbon
• Measure of the changed ration between stable carbon-12 andn unstable carbon-14

2. Medicine

• Cobalt-60 gamma rays used in radiotherapy to treat cancer

3. Food Preservation

• Cobalt-60 preserves food by irradiation

See Scientists for information on Marie Curie and Henry Bequerel

It is important to have some knowledge about the major scientists involved in the major discoveries of chemistry. This is often only a short question in Q4 or a 3 marker in a long question, but its an easy way to add marks

Boyle

• An element is a substance that cannot be broken down into simpler materials
• Boyle’s Law

Davy

• Developed electrochemical techniques for breaking down compounds
• Use of electrolysis

Moseley

• Discovery of atomic numbers
• Use of X rays in the discovery of a characteristic positive charge in the nucleus

Dobereiner

• Law of Triads
• Groups of three elements have similar properties

Newlands

• Law of Octaves
• Properties repeat every eight elements

Mendeleev

• Periodic table
• Left gaps for undiscovered elements
• Ordered by atomic weights
• Iodine and tellurium the exception (atomic number later explained this)

Dalton

• First atomic theory based on experimental evidence
• Indirect evidence
• Atoms are indivisible and indestructible
• Atoms of a given element are identical
• Atoms of different elements vary in mass
• Compound contains atoms combined in fixed proportions
• Law of conservation of mass

Crookes

• Cathode rays are particles (moved a small paddle)

Thompson

• Cathode rays have negative charge
• Negatively charged particles are extremely light
• Plum pudding model of the atom – electrons embedded in a positively charged sphere

Stoney

• Suggested name “electron”

Millikan

• Oil drop experiment
• Electrical force used to suspend drops mid air
• Electron charge accurately measured

Rutherford

• Alpha particles fired at gold foil
• Large number not deflected (empty space)
• Very small numberbounced back (concentrated positive charge)
• Discovery of nucleus
• Positive charge confined to nucleus
• Discovery of proton

Chadwick

• Aplha particles fired at beryllium
• Alpha particles >>>knocks out>>> neutrons in Be >>>knocks out>>> protons in parafin >>>detected
• Discovery of neutron

Bequerel

• Study of chemicals which emitted light and X rays, which had been exposed to white light
• Photographic plate in black paper with crystals of uranium – placed in sunlight – image of crystals developed
• Without sunlight, image still developed
• Crystals emitted radiation which caused image
• Radiation also ionised air – air conducts electricity

Curie

• Coined term “radioactivity”
• Discovered that amount of radiation depended on proportion of uranium
• Isolated new elements (polonium and radium) which are more radioactive than uranium
• 1903 Nobel Prize (with Bequerel and her husband) – discovery of radioactivity
• 1911 Nobel Prize – discovery of elements

Intramolecular Bonds the bond within the molecule that holds the atoms together

Intermolecular Forces forces between molecules

Types of Intermolecular Forces

1. Van der Waals Forces

• Weak attractive forces caused by the movement of electrons within a molecule
• Randomly moving electrons may at one point be nearer to one atom than another
• Temporary dipole formed
• If two molecules with opoositely charged temporary dipoles are near each other an attractive force will exist E.g. H(-)H(+) — H(-)H(+)
• Sometimes a temporary dipole in one molecule will induce a dipole in another
• Results in greater boiling point
• Greater number of electrons = greater number of temporary dipoles = greater boiling point
• Occurs in non-polar and polar molecules

2. Dipole – Dipole Interactions

• Negative end of one dipole is attracted to the positive end of another
• Permanent forces (due to polarity)

• Stronger than van der Waals forces, weaker than hydrogen bonding
• Results in greater boiling point
• Occurs in polar molecules

3. Hydrogen Bonding

• Occurs when hydrogen is bonded to highly electronegative atoms (O, N, F)
• Strong – Requires more energy to break than regular covalent bonds – Higher boiling point
• Oxygen (-) has an attraction to neighbouring hydrogen (+) molecules, and vice versa

• Covalent bond > Hydrogen bond > Dipole-dipole interactions > van der Waals forces
• H2S should have a higher boiling point (greater molecular mass) but the H-S bond is much less polar than the O-H bond
• Hydrogen bonding only occurs between hydrogen and small atoms (O, N, F) because the charge is much more concentrated and therefore, more effective

Linear Shape

• BeH2
• 180°
• 2 bonding pairs

Trigonal Planar

• BF3
• 120°
• 3 bonding pairs

Tetrahedral

• CH4
• 109.5°
• 4 bonding pairs

V-Shaped

• H2O
• 104.5º
• 2 bonding pairs and 2 lone pairs

Pyramidal

• NH3
• 107º
• 3 bonding pairs and 1 lone pair

Electron Pair Repulsion Theory the electron pairs in the outer shell of the central atom repela each other and end up as far apart as is geometrically possible

Since lone pairs are closer to the nucleus of the central atom, they are closer together, so their mutual repulsion is greater than that between bond pairs >>>> distorted shape

Order of Strength of Repulsions

1. lone pair:lone pair
2. lone pair:bond pair
3. bond pair: bond pair

Symmetry and Polarity

Symmetrical Atoms

• Linear
• Tetrahedral
• Trigonal Planar
• Non- polar (even if bonds are polar)

E.g. BF3

Boron (central atom) is slightly positive >>> centre of positive charge is here

Each flourine is slightly negative >>> centre of positive charge is the central point between these atoms

Centre of positive coincides with centre of negative >>> non-polar

Non-Symmetrical Atoms

• V-Shaped
• Pyramidal
• Polar

E.g. NH3

Nitrogen (central atom) is slightly negative >>> centre of negative charge will be at apex of pyramid

Hydrogen atoms areslightly positive >>> centre of positive charge at base of pyramid

Centre of negative does not coincide with centre of positive >>> polar

Catalysts work by providing an alternative reaction route with a lower activation energy.

Intermediate Formation Theory of Catalysis

The reactant molecules and the catalyst form and unstable intermediate complex that breaks up to form products and regenerate the catalyst

How does the reaction of aqueous cobalt (II) chloride between H2O2 and potassium sodium tartrate give evidence for the intermediate formation theory of catalysis?

1. Initial solution is pink
2. During reaction, there is a colour change from pink to green [intermediate complex formed]
3. Frothing and bubbling reaction [fast rate of reaction - products formed]
4. Reaction returns to pink [catalyst regenerated]

Surface Adsorption Theory of Catalysis

The reactant molecules adsorb onto solid catalyst where the greater local concentration leads to a quick reaction – bonds formed must be strong enough to adsorb and increase concentration, but weak enough to decompose quickly and form products

How does the oxidation of methanol using a hot platinum catalyst provide evidence for the surface adsorption theory of catalysis?

1. Series of mild explosions and glowing platinum [fast exothermic rate of reaction]
2. H atoms are removed more quickly by reactant adsorbing to surface of catalyst which weakens and breaks bonds.

Catalytic Poisons

When the active sites of the catalyst are blocked by substances bonding to it more strongly than the reactants, the catalyst is poisoned

e.g. Lead, arsenic, sulfur

A catalyst is a substance that alters the rate of a chemical reaction but is not consumed in the reaction

Heterogeneous catalysis involves reactants in different physical states i.e. liquid reacting with solid e.g.  MnO2 on H2O2

Homogeneous catalysis involves reactants in same physical states i.e. both in aqueous solution e.g. potassium iodide on H2O2

Enzymes are homogeneous biological catalysts e.g. amylase on starch

Autocatalysis occurs when the product of a reaction increases the reaction rate i.e. reaction makes its own catalysts e.g. reduction of manganate (VII) ions with Fe 2+

Activation Energy

Activation energy is the minimum energy required by particles colliding to cause a reaction

Exothermic reactions give out heat [Energy of: products < reactants ]

Endothermic reactions take in heat [Energy of: products > reactants ]

Average kinetic energy of particles is directly proportional to the temperature – greater the energy, greater the speed, greater the reaction rate. This means:

1. the number of collisions per second increases
2. each collision is more energetic and a higher proportion of collisions has the necessary activation energy

The second factor is more significant

Pollution and Catalytic Converters

Engines produce harmful CO, NO, NO2 and hydrocarbons.

Catalytic converters (e.g. palladium and platinum) speed up reactions to reduce harmful emissions

e.g. 2CO + O2 >>> 2CO2

 2CO + 2NO >>> 2CO2 + N2

This is an example of heterogenous catalysis.

Rate of Reaction is the change in concentration in unit time of any one reactant or product

Average Rate – Change in Concentration

                     Time taken

Instantaneous Rate of Reaction is the rate at a particular point in time during the reaction

To calculate the instantaneous rate of reaction

1. Draw a tangent to the curve
2. The tangent is the hypotenuse to a right angled triangle
3. θ = angle by tangent to the horizontal axis
4. Slope = tan θ = 18.5/1.2 = 15.4 cm³/min

On graphs:

• The reaction time is inversely proportional to concentration (x axis = time) i.e. the shorter the time, the greater the concentration
• When using the reciprocal (1/time) the reaction time is directly proportional  to concentration i.e. the greater the time, the greater the concentration

Factors Affecting Reaction Rate

1. Temperature

• Increase in temperature – increase in reaction rate
• In some cases, an increase of 10 K can up to double reaction rate

2. Concentration

• Increase in concentration – increase in reaction rate

3. Particle Size

• Smaller particles – increase in reaction rate
• Finely divided solids – Greater surface area for reaction to occur over

Dust Explosions

The factors for a dust explosion to occur are:

1. Finely divided combustible particles
2. An enclosed space
3. Enough oxygen to sustain combustion
4. A spark to ignite combustion

4. Nature of Reactants

• Covalent bonds – slow reaction
• Bonds must be broken before reaction can occur – Rate will depend on bonds (Single bonds easier to break than double bonds)
• Ionic bonds – Quick in aqueous solution
• Ions pulled apart by water – instant reaction

5. Presence of Catalysts

• Can speed up or slow down reaction
• Catalyts work by lowering the energy required for a reaction to occur

Oxidation Number the charge an atom appears to have when electrons are distributed according to certain rules

Rules for assigning charges

• Free elements: 0 e.g. O2 – o
• Sum of oxidation numbers: 0 e.g. H2O – o
• Oxidation numbers equals charge on ion e.g. Fe3+ – +3
• Sum of oxidation numbers in a complex ion equal to charge e.g. NO3- = +5 +3(-2) = -1
• Hydrogen: +1 except hydrides: -1
• Oxygen: -2 except H2O2 – O: -1   and   F2O – O:+2
• Group 1: +1, Group 2: +2…etc.
• Halogens: -1 (when bonded to less electronegative) e.g. NaCl
• Transition metal may have many oxidation numbers

What is the oxidation number of C in C6H12O6?

C6H12O6

6x+12(+1)+6(-1)=0

6x+12-12=0

x=0

Oxidation is an increase in oxidation number

Reduction is a decrease in oxidation number

What is oxidised and what is reduced in ZnS + 2O2 >>> ZnSO4

ZnS + 2O2 >>> ZnSO4

+2 -2  2(0)>>>+2 +6 4(-2)

S: -2>>-6 oxidised

O: 0>>-2 reduced

Balance the equation MnO4- + Fe2+ + H+ >>> Mn2+ + Fe3+ +H2O

MnO4- + Fe2+ + H+ >>> Mn2+ + Fe3+ +H2O

+7 4(-2)  +2      +1       +2     +3     2(+1)-2

Mn: +7 >> +2 : +5 e-

Fe: +2 >> +3 : -1e-

Mn + 5e- >> Mn

5Fe – 5e- >> Fe

Mn + 5Fe    Ratio 1:5

MnO4- + 5Fe2+ + H+ >>> Mn2+ + 5Fe3+ + H2O

MnO4- + 5Fe2+ + 8H+ >>> Mn2+ + 5Fe3+ + 4H2O

This is the theory section of acid-base titrations.
Experiments make up the majority of this section.

1 molar = 1 mole per litre e.g. 1 mole Na2CO3 = 106g/L : 0.1 molar = 1.6g/L

Concentration is the amount of solute in a specified amount of solution e.g. moles per litre or grams per litre

Ways of Expressing Concentration

1. Percentage weight per volume (w/v) e.g. 3% NaCl solution = 3g NaCl in 100cm³ solution
2. Percentage volume per volume (v/v) e.g. 3% alcohol solution = 3cm³ alcohol in 100cm³ solution
3. Percentage weight per weight (w/w) e.g. 3% sugar solution = 3g sugar in 100g solution
4. Parts per million (p.p.m.) e.g. 2 p.p.m. solution has 2mg substance per litre

Standard solution a solution whose concentration is accurately known

Primary standard a water-soluble substance that is stable and available in pure form

Formula for Titration Problems

V1 x M1 = V2 x M2
    n1          n2                           Where V=volume, M=molarity, n=moles of solution present

If 20cm³ of 0.3 molar NaOH are neutralised by 25cm³ of H2SO4 solution, find the concentration of H2SO4 in (i) moles/L (ii)g/L according to the equation 2NaOH + H2SO4 >>> Na2SO4 + 2H2O

V1=25cm³  
M1=M1  n1=1    
V2=20cm³ 
M2=0.3  n2=2

25 x M1 = 20 x 0.3
   1            2

M1 = 0.12 molar

(i) 0.12 moles/litre    
(ii) 0.12 x 98 g = 11.76 g/litre

Equipment:

• 0.02M potassium iodate solution
• 0.5M potassium iodide solution
• 1M sulfuric acid
• Sodium thiosulphate solution
• Starch indicator solution
• Deionised water
• Pipette and filler
• Burette
• Funnel
• COnical flask
• Ahite tile
• Graduated cylinder
• Retort stand and clamp
• Beakers

Experiment:

1. Wash pipette, burette and conical flask with deionised water
2. Rinse pipette with potassium iodate and burette with iodine thiosulphate
3. Pipette 25 cm³ potassium iodate into conical flask, with 20cm³ sulphuric acid and 10cm³ potassium iodide
4. Fill burette with sodium thiosulphate. Titrate against iodine solution.
5. When colour of solution fades to pale yellow, add starch indicator. Blue-black colour appears.
6. Continue titration until colour changes from blue-black to colourless.
7. Carry out two more accurate titrations
8. Calculate concentration of thiosulphate solution.