energy matters unit 1. reaction rates from standard grade you should remember that a reaction can be...
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Energy matters
Unit 1
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Reaction rates
• From standard grade you should remember that a reaction can be speeded up by;
• Decreasing particle size
• Increasing concentration
• Increasing temperature
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Following the course of a reaction
• In general to measure the rate of a reaction we must choose some measurable quantity which changes as the reaction proceeds.
e.g mass of reactants in a flask, volume of gas produced, colour intensity, concentration of reagent
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Following the course of a reaction
• If we react marble chips (Calcium carbonate) with hydrochloric acid we can monitor the course of the reaction.
CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)
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Marble chips & acid
• As we are producing a gas, it will escape from the vessel causing the total mass to drop.
• If we measure this change in mass over a fixed period of time we can calculate the rate of the reaction.
Marble chips
Balance
Cotton woolHCl(aq)
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Time (s) Total mass of flask (g)
Decrease in mass (g) or Mass of CO2
produced
0
30
60
90
120
150
180
210
240
270
300
149.00
147.75
147.08
146.60
146.26
145.94
145.68
145.48
145.32
145.19
145.08
-
1.25
1.92
2.40
2.76
3.06
3.32
3.52
3.68
3.81
3.92
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Decrease in mass (g) -
00.5
11.5
22.5
33.5
44.5
0 50 100 150 200 250 300 350
Time (seconds)
Decre
ase in
mass
(g)
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Average rate of reaction
• It is difficult to measure the actual rate at any one instant since the rate is always changing.
• We can calculate average rate over a certain period of time.
Average reaction rate = Change in mass of product Time taken for change
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Example
• Calculate average rate of reaction between 30 and 60 seconds.
Average reaction rate = Change in mass of product
Time taken for change
Average reaction rate = 1.92 - 1.25
30
Average reaction rate = 0.022gs-1
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Collision theory
• For a chemical reaction to occur, the reactants must collide.
• Any factor that increases the number of collisions per second is likely to increase reaction rate.
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Particle size
• More collisions occur if the particle size of a solid reactant is decreased, since its overall surface area is increased.
• Powdered marble (calcium carbonate) reacts much faster than marble chips.
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Concentration
• If concentration is increased, there are more reactant particles.
• The more particles there are in one space, the more collisions.
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Raising temperature
• Raising the temperature at which a reaction takes place does more than merely raise the number of collisions.
• Temperature is a measure of the average kinetic energy of particles in a substance.
• Therefore at higher temperatures, particles have greater kinetic energy and they collide with more force.
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Collisions
• Not all collisions cause a reaction to occur e.g. nitrogen & oxygen particles in the air.
• The colliding particles must have a minimum amount of kinetic energy for a reaction to occur.
• This minimum kinetic energy is called the Activation energyActivation energy (EA)
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Activation energy
• Activation energy required varies from one reaction to another.
• If the activation energy of a reaction is high, only few particles will have enough energy to successfully collide.
• Conversely, a reaction with low activation energy will be very fast.
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Kinetic energy• At a given temperature (T1) individual molecules of
a gas have widely different kinetic energies.
• Most molecules will have energy near to the average energy but some will be well below average, and some well above.
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Activation energy
• The shaded area represents the all of the molecules which have kinetic energy greater than the activation energy.
• The shaded area represents the portion of molecules that will react
EA
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Temperature• Distribution of energy changes when the temperature
changes.• A small rise from T1 to T2 considerably increases the
number of particles capable of reacting.• Hence increasing the reaction rate.
T2
Kinetic energy
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Catalysts
• Substance that alters rate of reaction without being used up.
• Homogeneous catalyst: Same state as the reactants.
• Heterogeneous catalyst: Different state as the reactants.
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Heterogeneous catalyst
• The catalyst has a large surface area.• Catalysis occurs at certain points on the
catalyst called ‘active sites’. • At these sites reactant molecules are
adsorbed onto the surface of the catalyst.• At least 1 reactant is held in place on active
site, making collision more likely.
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Catalyst poisoning
• Occurs when reactants or impurities become preferentially adsorbed or even permanently attached to the catalyst surface.
• Hence reducing number of active sites and therefore rendering the catalyst as useless.
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Catalytic converters• Petrol engine cars must now be fitted
with a catalytic converter.
• The contains a honeycomb network of platinum, converting harmful gases into less harmful ones.
CO, NOx, O2
CO2
H2OO2
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Industrial catalystCatalystCatalyst ProcessProcess ImportanceImportance
Vanadium(v) oxide
Contact Manufacture of H2SO4
Iron Haber Manufacture of ammonia
Platinum Oxidation of ammonia
Manufacture of nitric acid
Nickel Hydrogenation Manufacture of margarine
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Enzymes
• Biological catalyst.
• Examples of enzymes:– Amylase, catalyses the hydrolysis of starch.– Catalase, catalyses the decomposition of
hydrogen peroxide. Catalyase is found in the blood, preventing build up of hydrogen peroxide in the body.
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Enzymes continued
• Enzymes are highly specific. • Enzymes work best at their
optimum temperature & pH.• Optimum temperature for
human enzymes will be 37°C.
• Greatly exceeding either of these will result in the protein being denatured.
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Industrial enzymesEnzymeEnzyme ProcessProcess
Lipase Enhance flavour of cheese, ice-cream &
chocolate
Rennin Cheese production
Protease Tenderising meat
Amylase Desizing (removing starch from fabric)
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Enthalpy
From SG: Exothermic reaction
Combustion
Neutralisation
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Potential energy• Potential energy is the energy
possessed by the reactants.
• In an exothermic reaction, the products have less potential energy than reactants.
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Potential energy
• In an endothermic reaction, the opposite is true. Reactants must absorb energy from their surroundings.
• Products have more energy than the reactants.
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Enthalpy
• The difference in potential energy between reactant and product is called the enthalpy enthalpy changechange (ΔH)
• Enthalpy changes are normally quoted in kJ mol-1
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Activation energy
• The rate of reaction depends on the height of the Ea barrier.
• Rate of reaction does not depend on the enthalpy change ( )
HH
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Catalyst• Catalysts provide alternative reaction
pathways.
• Thus lowering the activation energy.
Energy
Reaction pathway
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Activated complex
• When reactants change into products, they pass through a very unstable state known as the activated activated complexcomplex. (Situated at the maximum potential energy).
• The activated complex is a highly energetic arrangement of atoms that exists for a short time.
• The activated complex loses this energy by either forming products or reforming as reactant particles.
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Activated complex
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Patterns in the periodic table
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Density
• The amount of material packed into a given volume.
• Density values are much larger for Solid & liquid elements.
• Density increases down each group.• Across the period from L to R,
density increases towards the centre of the period, then decreases again towards the noble gases.
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Atomic size: Groups• Atomic size is measured in covalent
radius. This is the distance from the nucleus to the outer electrons.
• As you move down a group the atomic radius increases.
• This is due to the increased number of occupied electron shells.
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Atomic radius: Periods• Across a period atomic number and electron number
increase by one.
• Although the number of outer electrons is increasing across the period, the atomic radius decreases.
• This is due to the increasing attraction between the nucleus and the outermost electrons.
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Ionisation energies
• The attraction between the nucleus and the outer electrons means that energy is required to remove electrons from the atom.
• Ionisation energy is a measure of the nuclear attraction for outer electrons.
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First ionisation energy
• Energy required to remove an electron from one mole of free atoms in a gaseous state.
• K(g) K+(g) + e-
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Second ionisation energy
• Energy required to remove an electron from one mole of ions with a charge of 1+ in the gaseous state.
• K+(g) K2+(g) + e
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Third ionisation energy
• Energy required to remove an electron from 1 mole of ions with 2+ charge in the gaseous state.
• K2+ K3+(g) + e
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Ionisation energies• The first ionisation energy decreases as
you go down a group.• This is due to the increasing atomic radius.• As the radius increases, the attraction
between the nucleus and the outermost electrons decreases.
• Screening / Shielding effect.• Therefore the energy required to remove
that electron decreases.
Li
Na
K
e-
e-
e-
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Bonding, structure and properties of
compounds
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Metallic Bonding
Covalent bonding
Polar covalent bonding
Ionic bonding
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Metallic Lattice
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Covalent molecular
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Carbon atoms
Covalent bonds
Covalent network
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Electronegativity
• The greater the difference in electronegativity between two elements, the less likely they are to share electrons and form covalent bonds.
• Caesium fluoride is the compound with the greatest degree of ionic bonding.
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• Formed when atoms of different electronegativities bond to form a covalent compound.
• Bonding electrons are not shared equally.
• The atom with the greater share of electrons becomes slightly negative (δ-)
• The other atom becomes slightly positive (δ+)
• These molecules have a permanent dipole.
Polar covalent bonding
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Polar covalent bonding
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Ionic bonding
• Different elements have different attraction for bonding electrons, (electronegativity values).
• One atom may attract electrons very strongly and another atom may attract them very weekly and lets them go.
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Ionic bonding
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Summary
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Intermolecular forces of attraction
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Covalent molecular
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Intermolecular interactions
• Van der Waal forces are a result of electrostatic attraction between temporary dipoles and induced dipoles caused by movement of electrons in atoms and molecules.
• All covalent molecules interact by van der Waals bonding, as all molecules possess temporary dipoles.
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Halogens
• All halogen have 1 unpaired electron in the outer shell. Therefore form 1 pure covalent bond. E.g. F2, Cl2, Br2, I2
• These molecules interact only weakly by van der Waals’ mechanism, this makes them very volatile. (Fluorine & chlorine are gaseous).
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Permanent dipole
A molecule can be described as polar if it has a permanent dipole. A permanent dipole is due to a difference in electronegativity between the atoms involved in a covalent bond.
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Symmetry
• Some molecules have a symmetrical arrangement of polar bonds.
• This cancels out the polarity over the molecule as a whole.
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Polar or Non-polar?
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Boiling point• Polar molecules have higher boiling
points than non-polar molecules with similar molecular mass.
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Hydrogen bonds• Bonds consisting of a hydrogen atom bonded to an
atom of a strongly electronegative element such as fluorine, oxygen or nitrogen.
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Water molecules
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Ice
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Glycerol Sulphuric Acid
Phosphoric Acid
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Covalent molecular
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Carbon atoms
Covalent bonds
Covalent network
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Diamond
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Fullerenes• Discrete covalently
bonded molecules
• Consisting of pentagonal & hexagonal panels.
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Graphite
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Bonding, structure & properties of elements
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Groups 1,2 & 3• Not enough electrons to achieve full outer shell.
• Elements contribute electrons to a common ‘pool’ of delocalised electrons.
• This binds the resultant positive ions.
• Bonding is less directional, therefore metals are more ductile & malleable.
• Delocalised electrons, therefore conduct electricity.
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Metallic Bonding
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1 exception: Boron• Structure made up of
B12 groups, interbonded with other groups.
• This results in an element almost as hard as diamond.
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Group 4
• Standard structure: Infinite 3D network or lattice, e.g. diamond, silicon.
• Therefore exceptionally hard & rigid.
• No discrete molecules, each atom joined to another.
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Diamond
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Graphite
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Fullerenes• Discrete covalently
bonded molecules
• Consisting of pentagonal & hexagonal panels.
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Phosphorus (group 5)
• Phosphorous bonds to 3 other phosphorous atoms to form tetrahedral P4 molecules.
• Fewer electrons in P4 than S8 make van der Waals forces weaker in phosphorous, therefore lower m.p.
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Group 6
• Oxygen: 2 unpaired electrons, therefore forms 2 pure covalent bonds.
• Intermolecular interactions are weak van der Waals, therefore volatile & gaseous.
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Sulphur
• Sulphur atoms can bond to more than one other sulphur, forming an 8 member ring.
• Van der Waals forces strong enough to make sulphur a solid at room temperature.
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Groups 5, 6 & 7
• Intra molecular forces (bonds within molecules) are covalent.
• Intermolecular forces are very weak van der Waals forces.
• Therefore most elements are volatile even if solid at room temperature.
• This is due to the little energy required to break intermolecular forces in order to melt/boil.
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Bonding in elements: Noble gases
• There are no covalent or ionic bonds between atoms in group 8.
• Uneven distribution of electrons within the atom produce temporary (or transient) dipoles on the atom.
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Solvent action
• In general polar solvents dissolve polar substances and ionic substances.
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Non polar solvents…(e.g hexane)
• Dissolve non polar solvents
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The mole
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The Avagadro constant
• 1 mole of any element contains the same number of atoms.
• This number is known as the Avagadro constant.
• This constant is given the symbol (L) after the first person to calculate a numerical value.
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Avagadro constant (L)
• One mole of any substance contain L, 6.02x1023 formula units.
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Formula units• For metals & monatomic species e.g. Noble
gases, a formula unit is an atom.
• Thus 4g helium
40g of calcium
197g of gold
Contain L Contain L (6.02x10(6.02x102323) ) atomsatoms
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Covalent substances• A formula unit is a molecule• The total number of atoms can be found by
multiplying L by the number of atoms in the molecule.
Quantity of substance
Number of molecules
No. of atoms per molecule
Total No. of atoms
2g of Hydrogen, H2
18g of Water, H2O
30g of ethane, C2H6
L
L
L
2
3
8
2L
3L
8L
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Ionic compounds• Formula unit consists of a ratio of ions
expressed by ionic formula.
Quantity of substance No. of formula
units
No. of +ve and –ve ions
Total No. of ions
58.5g of Na+Cl-
74g of Ca2+ (OH-)2
342g of (Al3+)2(SO42-)3
L
L
L
LNa+ and LCl-
LCa2+ and 2LOH-
2L Al3+ and 3L SO4
2-
2L
3L
5L
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Example 1
• How many molecules are there in 8.8g of CO2?
1 mole of CO2 contains L molecules
44g of CO2 contains L molecules
1g of CO2 contains L/44 molecules
8.8g of CO2 contains L/44 x 8.8 molecules
= 1.204 X 10= 1.204 X 102323 molecules molecules
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Example 2
• What mass of Nitrogen gas contains 18.06x1022 atoms of Nitrogen?
6.02x1023 molecules of N2 1 mole
6.02x1023 molecules of N2 28g
1 molecule of N2 28/L
18.06x1022 molecules of N2 28/L x 18.06x1022
= 8.4g
Therefore 8.4g of N2 gas contains 18.06x1022 molecules
4.2g of N2 gas contains 18.06x1022 atoms