Transcript Document

Energy matters
Unit 1
Reaction rates
• From standard grade you should
remember that a reaction can be
speeded up by;
• Decreasing particle size
• Increasing concentration
• Increasing temperature
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
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)
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.
HCl(aq)
Cotton wool
Marble chips
Balance
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
Decrease in mass (g)
Decrease in mass (g) 4.5
4
3.5
3
2.5
2
1.5
1
0.5
0
0
50
100
150
Time (seconds)
200
250
300
350
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.
Change in mass of product
Average reaction rate = Time taken for change
Example
• Calculate average rate of reaction between
30 and 60 seconds.
Average reaction rate =
Change in mass of product
Time taken for change
1.92 - 1.25
Average reaction rate =
30
Average reaction rate = 0.022gs-1
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.
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.
Concentration
• If concentration is increased, there are more
reactant particles.
• The more particles there are in one space,
the more collisions.
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.
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 energy (EA)
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.
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.
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
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
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.
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.
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.
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.
CO2
H2O
O2
CO, NOx, O2
Industrial catalyst
Catalyst
Process
Importance
Vanadium(v) oxide
Contact
Manufacture of
H2SO4
Iron
Haber
Platinum
Oxidation of
ammonia
Hydrogenation
Manufacture of
ammonia
Manufacture of
nitric acid
Manufacture of
margarine
Nickel
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.
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.
Industrial enzymes
Enzyme
Process
Lipase
Enhance flavour of
cheese, ice-cream &
chocolate
Rennin
Cheese production
Protease
Tenderising meat
Amylase
Desizing (removing
starch from fabric)
Potential energy
• Potential energy is the energy
possessed by the reactants.
• In an exothermic reaction, the products
have less potential energy than
reactants.
Potential energy
• In an endothermic
reaction, the opposite
is true. Reactants
must absorb energy
from their
surroundings.
• Products have more
energy than the
reactants.
Enthalpy
• The difference in potential energy between
reactant and product is called the enthalpy
change (ΔH)
• Enthalpy changes are normally quoted in kJ
mol-1
Activation energy
H
• The rate of reaction
depends on the height of
the Ea barrier.
• Rate of reaction does
not depend on the
enthalpy change ( H )
Catalyst
• Catalysts provide alternative reaction
pathways.
• Thus lowering the activation energy.
Energy
Reaction pathway
Activated complex
• When reactants change into products, they pass
through a very unstable state known as the activated
complex. (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.
Activated complex
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.
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.
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.
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.
First ionisation energy
• Energy required to remove an electron from
one mole of free atoms in a gaseous state.
• K(g)
K+(g) +
e-
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
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
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 eNa e-
K
e-
Bonding, structure
and properties of
compounds
 Metallic Bonding
 Covalent bonding
 Polar covalent bonding
 Ionic bonding
Metallic Lattice
Covalent molecular
Covalent
network
Carbon atoms
Covalent bonds
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.
Polar covalent bonding
• 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
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.
Ionic bonding
Summary
Intermolecular
forces of
attraction
Covalent molecular
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.
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).
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.
Symmetry
• Some molecules have a symmetrical
arrangement of polar bonds.
• This cancels out the polarity over the molecule
as a whole.
Polar or Non-polar?
Boiling point
• Polar molecules have higher boiling
points than non-polar molecules with
similar molecular mass.
Hydrogen bonds
• Bonds consisting of a hydrogen atom bonded to an
atom of a strongly electronegative element such as
fluorine, oxygen or nitrogen.
Water molecules
Ice
Glycerol
Sulphuric Acid
Phosphoric
Acid
Covalent molecular
Covalent
network
Carbon atoms
Covalent bonds
Diamond
Fullerenes
• Discrete covalently
bonded molecules
• Consisting of
pentagonal &
hexagonal panels.
Graphite
Bonding, structure &
properties of elements
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.
Metallic Bonding
1 exception: Boron
• Structure made up of
B12 groups,
interbonded with
other groups.
• This results in an
element almost as
hard as diamond.
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.
Diamond
Graphite
Fullerenes
• Discrete covalently
bonded molecules
• Consisting of
pentagonal &
hexagonal panels.
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.
Group 6
• Oxygen: 2 unpaired electrons, therefore
forms 2 pure covalent bonds.
• Intermolecular interactions are weak van
der Waals, therefore volatile & gaseous.
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.
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.
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.
Solvent action
• In general polar solvents dissolve polar
substances and ionic substances.
Non polar solvents…(e.g hexane)
• Dissolve non polar solvents
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.
Avagadro constant (L)
• One mole of any substance contain L,
6.02x1023 formula units.
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
(6.02x1023)
atoms
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 Number of
No. of
Total No. of
substance
molecules
atoms per
atoms
molecule
2g of Hydrogen, H2
L
2
2L
L
3
3L
L
8
8L
18g of Water, H2O
30g of ethane, C2H6
Ionic compounds
• Formula unit consists of a ratio of ions
expressed by ionic formula.
Quantity of substance
No. of
No. of +ve
formula and –ve ions
units
58.5g of Na+Cl-
L
74g of Ca2+ (OH-)2
L
342g of (Al3+)2(SO42-)3
L
LNa+ and LClLCa2+ and 2LOH-
2L Al3+ and 3L
SO42-
Total
No. of
ions
2L
3L
5L
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 1023 molecules
Example 2
• What mass of Nitrogen gas contains
18.06x1022 atoms of Nitrogen?
6.02x1023 molecules of N2
6.02x1023 molecules of N2
1 molecule of N2
18.06x1022 molecules of N2
1 mole
28g
28/L
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
Molar volume
• The volume occupied by one mole of gas at
specific temperature and pressure.
• At room temperature and pressure, i.e. 20°C
and 1 atmosphere pressure, the molar
volume of any gas is approximately 24 litres
mol-1
Molar volume
• Volume of gas changes if the temperature
and/or pressure changes.
• Therefore you must specify at what
temperature and pressure the volume is
being measured.
Mass
Volume
Density
Molar volume
In litres mol-1
GFM
Vmol
Density in g l-1
Density
Example 1
• Calculate the density in g l-1 of the
following gases at room temperature.
The molar volume under these
conditions is 24 litres mol-1.
(a) Neon (b) Ammonia
(a) D = GFM / Vmol
= 20.2 / 24
= 0.84 g l-1
(b) V = GFM / Vmol
= 17 / 24
= 0.71 g l -1
Density (gl-1) Molar volume
(litres mol-1)
0.65
24.6
Gas
Formula
Methane
CH4
Oxygen
O2
1.33
24.0
Nitrogen
N2
1.15
24.3
Carbon
dioxide
Argon
CO2
1.81
24.3
Ar
1.63
24.5
Volume of gas
In litres
V
No. of moles
n
Molar volume
In litres mol-1
Vmol
Example 2
•
The molar volume at 0°C and 1
atmosphere pressure is 22.4 litres mol-1,
Calculate
(a) The volume of 0.025 mol of oxygen
(b) The no. of moles of nitrogen in 4.48 litres under
these conditions.
(a) Volume of oxygen V = n X Vmol = 0.025 X 22.4
= 0.56 litres
= 560cm3
(b) No. of moles of Nitrogen in 4.48
litres under these conditions?
No. of moles of nitrogen, n = V/Vmol
= 4.48/22.4
= 0.2 moles
Example 1
• Calculate (i) the volume of oxygen required for the
complete combustion of 100cm3 of ethane and (ii) the
volume of each product.
2C2H6(g) + 7O2(g)
4CO2(g) + 6H2O(g)
2 mol
2 vol
4 mol
4 vol
6 mol
6 vol
2 vol
3 vol
7 mol
7 vol
Simplified ratio
1 vol
3.5 vol
(i) 100cm3 of ethane requires 350cm3 O2 and (ii)
produces 200cm3 CO2 and 300cm3 H2O.
Example 2
• A mixture of 20cm3 propane and 130cm3 Oxygen was
ignited and allowed to cool. Calculate the volume and
composition of the resulting gaseous mixture.
C3H8(g) + 5O2(g)
3CO2(g) + 4H2O(l)
1 mol
1 vol
3 mol
3 vol
5 mol
5 vol
4 mol
-
20cm3 propane requires 5 x 20 = 100cm3 Oxygen.
Oxygen in excess by 30cm3 (130cm3 – 100cm3)
CO2 formed = 3 x 20 = 60cm3
Resulting gas mixture = 30cm3 O2 and 60cm3 CO2.