Unit 1 - Chemical Changes

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Unit 1 - Chemical Changes and Structure
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Reaction Rates
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Reaction Rates

• During the course of a chemical reaction,
  reactants are being converted into products.
• Measurement of the rate of reaction involves
  measuring the change in the amount of a
  reactant or product in a certain time.
• The rate of reaction changes as it progresses,
  being relatively fast at the start and slowing
  towards the end.
• What is being measured is the average rate over
  the time interval chosen.
• Reactions can be followed by measuring changes in
  concentration, mass and volume.

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Nat 5

• Where property mass/volume/concentration
• The above is used when there is no change in
• mass/volume/concentration measured, for example
• during a colour change reaction.

Higher
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Collision Theory

• A chemical reaction can only occur if there is a
• successful collision between reactant molecules.
• From national 5 we know that we can speed up a
• chemical reaction by
• Decreasing particle size (increasing surface
  area)
• Increasing concentration (of reactant)
• Increasing temperature
• Adding a catalyst

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Collision Theory Particle Size

• The smaller the particle size, the higher the
  surface area.
• The higher the surface area, the greater the
  number of collisions that can occur at any one
  time.
• The greater the number of collisions, the faster
  the reaction.
• Therefore the smaller the particle size, the
  faster the reaction rate.

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Collision Theory Concentration

• The higher the concentration, the higher the
  number of particles.
• The higher the number of particles, the greater
  the chance of collisions that can occur.
• The greater the number of collisions, the faster
  the reaction.
• Therefore the higher the concentration, the
  faster the reaction rate.

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Collision Theory Temperature

• The higher the temperature, the higher the energy
  the particles have.
• The higher the energy, the faster the particles
  move.
• The faster the particles move, the greater the
  chance that they can collide with sufficient
  energy (activation energy) to be successful.
• The greater the number of collisions, the faster
  the reaction.
• Therefore the higher the temperature, the faster
  the reaction rate.

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Collision Theory Catalyst

• A catalyst speeds up a chemical reaction by
  lowering the activation energy. (i.e. catalyst
  provides another easier route)
• The lower the activation energy, the greater the
  chance of successful collisions.
• The more collisions in a period of time, the
  faster the reaction rate.

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Catalysts

• A catalyst is a substance which speeds up a
  chemical reaction without getting used up or
  changed itself.
• There are two main categories of catalyst
• a) Heterogeneous
• b) Homogenous.

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Heterogeneous Catalysts

• Heterogeneous catalysts have active sites on
  their surface.
• Reactant molecules form weak bonds with the
  surface in a
• process called adsorption.
• At the same time bonds with the adsorbed reactant
  
• molecules are weakened.
• The reactant molecules are also held at a
  favourable angle
• for a collision with another reactant molecule to
  occur.
• The product molecules then leave the active site
  in a stage
• called desorption. The active site is then
  available again.

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surface
Active site
Desorption product molecules formed.
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• Unfortunately unwanted substances can often be
• adsorbed onto the active sites thus making them
• unavailable for the normal reactants. (Example
• lead in petrol.)
• When this happens the catalyst is said to be
• poisoned.
• Sometimes it is not possible to regenerate a
• poisoned catalyst and it must be
  replaced/renewed.
• This adds to industrys costs so every effort is
• made to remove any impurities from reactants that
  
• might poison a catalyst.

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Reactant 1
Reactant 2
Poison blocking active site
Youll need 3 colours when drawing this diagram
(underneath - if possible the previous
catalyst note)
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Homogeneous Catalysts

• A catalyst that is in the same state as the
  reactants is
• said to be a homogeneous catalyst.
• The catalyst forms an intermediate compound with
  one of
• the reactants. (this intermediate compound later
• decomposes to reform the catalyst.)
• For example (using Reactants A and B)

Reactant A Catalyst Intermediate Interme
diate Reactant B Product
Catalyst
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Potential Energy Diagrams
See jotter for labelled diagrams

• Labels include
• Exothermic or Endothermic
• Activated complex
• Enthalpy change
• Reaction pathway
• Potential Energy (KJ)
• Activation Energy

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Bonding, Structure and Properties of Elements
He
S
Fe
C
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Summary of Bonding types in first 20 elements.
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• Inter means in between.
• In other words an INTERmolecular bonds means
  bonds in
• between the molecules.
• Intra means within.
• In other words an INTRAmolecular bond means bonds
  
• within the molecule.

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Types of Bonding in elements

• There are 3 types
• Metallic Bonding (intramolecular)
• Covalent Bonding (intramolecular)
• Van der Waals/London Forces (intermolecular)

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Metallic Bonding

• Metallic bonding unsurprisingly only appears
  in metal elements.
• Metallic bonding occurs between (positively
  charged) metal ions and delocalised outer shell
  electrons.
• delocalised means the electrons are common to
  all of the ions (i.e. they move from one to
  another)

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• The movement of delocalised electrons allow
  metal elements
• to conduct electricity

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Covalent Bonding

• Covalent bonding occurs between two non metal
  atoms.
• Covalent bonds are held together through the
  attraction between the positively charged nucleus
  of one atom and the negatively charged outer
  electrons of the other atom.
• Outer electrons are shared in covalent bonding.

whiteboard example
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Van der Waals Bonding

• Van der Waals bonding is weak bonding which
  occurs BETWEEN molecules.
• London forces are temporary dipole to temporary
  dipole attractions.
• Temporary dipoles occur when electrons lie
  slightly closer to one atom than the other. This
  means for a short time one of the atoms is
  slightly negative and the other is slightly
  positive (i.e. electrons not shared equally)
• London forces (and other intermolecular forces)
  are useful when explaining patterns in the
  periodic table e.g. melting/boiling points.

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Bonding in Specific Groups

• Groups 1, 2 and 3
• All elements in groups 1, 2 and 3 have strong
  metallic bonds holding them together.
• Metallic bonds allows metals to be shaped (i.e.
  malleable and ductile)
• Metals have high melting/boiling pts due to
  strong metallic bonds
• Boron is the only exception as it has very
  complex bonding. B12 is almost as hard as
  diamond. This suggests a covalent network
  structure.

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The 3 Structures of Carbon (group 4)

• Each atom covalently bonds
• to 4 other atoms.
• This means covalent bonds must
• be broken to melt/ boil very
• high m.pt/b.pt values.
• No free electrons no conduction.
• Tunnels between atoms allow
• light through transparent
• structure.

Diamond
(covalent network)
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• Each atom forms 3 covalent bonds and its last
  valence electron becomes delocalised.
• As the delocalised electrons are only held weakly
  they can flow i.e. graphite conducts electricity.
• The delocalised orbitals sit between the layers -
  as a result there are 3 strong covalent bonds
  WITHIN the layers but only weak interaction
  BETWEEN the layers.
• Due to these weak interactions, graphite is flaky
  as the layers can be easily separated.
• Graphite layers are offset (i.e. not above each
  other) - light cant travel through it meaning it
  is not transparent.

Graphite
(covalent network)
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• 3. Buckminsterfullerene (aka Bucky Ball)

• The fullerenes, despite being large molecules,
  are discrete covalent molecules.
• The smallest of fullerenes is a molecule known as
  Buckminsterfullerene (C60).
• This is a spherical molecule containing 5 and 6
  membered carbon rings.
• The properties are still being researched so the
  full applications are still unknown.

'Bucky Ball'
(covalent molecule)
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• Group 5 (nitrogen and phosphorus)

• Nitrogen atoms form diatomic molecules with a
  triple covalent bond.
• This means that nitrogen only has London forces
  between the molecules.
• London forces are easily broken and as a result
  nitrogen has a low boiling pt. This is why
  nitrogen is a gas at room temperature.
• Phosphorus forms tetrahedral P4 molecules which
  are larger than N2 molecules and as a result it
  has stronger London forces between its molecules.
  This stronger attraction means phosphorus has a
  higher boiling pt and is a solid at room temp.

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• Group 6 (oxygen and sulphur)

• Oxygen atoms form diatomic molecules with a
  double covalent bond.
• This means that oxygen only has London forces
  interaction between the molecules.
• London forces are easily broken and as a result
  oxygen has a low boiling pt. Hence oxygen is gas
  _at_ room temp.
• Sulphur forms 8 membered rings. The London forces
  between the molecules are strong enough in
  sulphur to make it solid at room temperature.

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• Group 7 (the halogens)
• The halogens form diatomic molecules (i.e. they
  bond with themselves)
• As with oxygen and nitrogen this results in the
  halogens only having London forces with each
  other molecule.
• Fluorine and chlorine are very volatile (and
  therefore reactive) gases due to these weak
  intermolecular forces.
• Group 8 (the noble gases)
• As the noble gases have a stable outer electron
  shell they do not form bonds. As a result they
  remain monoatomic.

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Explaining the Melting and Boiling pts Trend In
small discrete covalent molecules the melting and
boiling points are low. This is because only
weak intermolecular London forces have to be
overcome when boiling or melting. The strong
covalent bonds are left unaffected. In the
covalent network solids (carbon, silicon and
boron) strong covalent bonds MUST be broken when
melting or boiling. Breaking these bonds requires
a lot more energy and therefore we get very high
values. In the metal groups (1, 2, 3) strong
metallic bonds MUST be overcome thus they have
high melting/boiling points.
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Summary of Bonding types in first 20 elements.
Except Buckminsterfullerenes !
H
He
Li
Be
B
C
N
O
F
Ne
Mg
Al
Si
P
S
Cl
Ar
Na
Ca
K
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Bonding, Structure and Properties of Compounds
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Types of Bonding in compounds

• There are 4 main types
• Ionic Bonding intramolecular
• Covalent Network Bonding intramolecular
• Polar/Non Polar Covalent Bonding - intramolecular
• Intermolecular (Van der Waals)
• Hydrogen bonding (present in H2O, NH3 and HF)
• Permanent dipole Permanent dipole interactions
• London forces Temporary dipole interactions

Strongest to Weakest
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Ionic Bonding

• Ionic bonding is an electrostatic attraction
  between the positive ions and negative ions.
• Ionic bonding is related to the
  electronegativities of elements. The greater the
  difference in e.n the less likely the elements
  are to share outer electrons. (electronegativity
  definition and trends are found in later section
  of jotter.)
• Instead the element with the higher e.n value
  will gain the electrons to form a negative ion
  and the element with lower e.n value will lose
  the electrons to form a positive ion.
• Due to the trends of electronegativity, the
  elements that are far apart from one another in
  the periodic table form ionic bonds. (normally
  metal and non metal.) Caesium fluoride is the
  compound with the greatest ionic character.

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Structure
diagram

• Ionic compounds do not form molecules. Instead
  the positive and negative ions come together to
  form lattice structures.
• When the lattice forms, energy is released. This
  is known as lattice energy or enthalpy.
• The overall charge of the lattice must be zero
  and therefore this affects the number of each
  ions we have present.
• In sodium chloride (NaCl) there is an equal
  number of Na and Cl- ions.
• In calcium fluoride (CaF2) there are twice as
  many F- ions than Ca2

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Covalent Compounds

• There are 3 types of covalent bonding in
  compounds (all involving combinations of non
  metals)
• Covalent network structures.
• Polar covalent molecules
• Non Polar covalent molecules

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Covalent Network

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• These covalent network compounds have the same
  properties as covalent network elements.
• Both SiC and SiO2 have very high melting pts. as
  melting requires breaking strong covalent bonds.
• Silicon carbide (structurally similar to diamond)
  has many uses due to its strength, durability
  and low cost.
• Silicon carbide is often referred to as
  carborundum

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Polar Covalent Bonding

• Most covalent compounds are made from atoms with
  slightly different electronegativity (e.n.)
• This difference is not significant enough for one
  of the atoms to fully remove an electron from the
  other. (Approx difference of between 0.5 and 1.6)
• As a result the atom (element) with the higher
  e.n. holds the electron slightly closer to itself
  and therefore becomes slightly negative (?-)
• The atom with the lesser e.n. is therefore
  slightly positive (?) as the electron is sitting
  further away from it.
• Covalent bonds with unequal electron sharing are
  called polar covalent bonds.

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Non Polar Covalent Bonding

• Non polar (or pure) covalent bonding normally
  occurs
• when
• electrons are equally shared between the two
  different atoms. i.e. equal electronegativity.
• E.g. Phosphorus Hydride
• the compound structure is symmetrical and
  therefore charges are overall balanced.
• E.g. CH4 (methane) and CO2 (carbon dioxide)

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• Both polar and non polar?
• Its possible for non polar covalent molecules
• to have individual polar bonds.
• For example Carbon Tetrachloride (CCl4)

diagram from whiteboard
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Non Polar
?
?
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• Summary of electronegativity values bonding
• In general
• If the electronegativity difference (usually
  called ?EN) is less than 0.5, then the bond is
  non polar (pure) covalent.
• If the ?EN is between 0.5 and 1.6, the bond is
  considered polar covalent
• If the ?EN is greater than 2.0, then the bond is
  ionic.

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Properties of Polar/ Non Polar Covalent Bonds

• Boiling Points
• Polar covalent molecules have higher boiling
  points than non polar covalent molecules with a
  similar mass.
• This is because the intermolecular forces are
  stronger (changing from London forces to
  permanent permanent dipole interactions.)
• Permanent dipole to dipole interaction is caused
  via the constant attraction between the ? atoms
  and ?- atoms of neighbouring molecules.

diagrams from white board
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• Solvent Action
• like substances dissolve in like substances
• This means that polar molecules will dissolve in
  polar solutions but not in non polar solutions
  and vice versa.
• This is due to the attraction between ? and ?-
  atoms of the water and the polar substance.
• Ionic compounds dissolve in polar solutions in a
  similar way due to the interaction between the
  ions and the ? and ?- atoms.
• Ions surrounded by a layer of water molecules
  held by electrostatic attraction are said to be
  hydrated.

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• Behaviour in electric field
• Copy figure 4.8 on page 49

Viscosity (thickness/ability to pour) Summarise
textbook notes on page 54 (2/3 sentences
max) Miscibility (ability to mix) Summarise
textbook notes on page 57 (2/3 sentences max)
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Physical properties of some hydrides
Boiling Points
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O H
F H
N H

• The above bonds are very polar due to the large
  difference in the electronegativity values.
• This interaction is called hydrogen bonding.
• Hydrogen bonding occurs in any molecule that
  contains any of the above bonds but mainly in
  hydrogen fluoride (HF), water (H2O) and ammonia
  (NH3). These intermolecular forces affect the
  properties of these compounds.
• Hydrogen bonding is stronger than both London
  forces and permanent dipole to dipole
  interactions but weaker than covalent bonding.

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Density of Water
Water is unusual as its solid form (ice) is less
dense than its liquid form (water). This means
that ice floats on water whereas most other
solids sink in their own liquid forms. This
phenomenon is due to the structure and bonding
which takes place between the water
molecules. As water molecules cool, they
contract. However, at 4oC they begin to expand.
This is because of hydrogen bonds between the
water molecules. This decreases the density of
ice (greater volume/same mass) compared to that
of the liquid water. Ice floating is vital to
real life i.e. fish/marine life surviving
under frozen lakes etc (https//www.youtube.com/wa
tch?vT4GCShGvw-M)
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Trends in the Periodic Table
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• Density (measured in g/cm3)
• Across a period (starting from group 1) the
  density increases towards the centre (group 4)
  and then decreases.
• Density tends to increase down a group (as atomic
  number increases.)

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• Atomic size
• Atomic size (or covalent atomic radius) is half
  the distance between the nuclei of two bonded
  atoms.
• Single bond lengths between atoms of different
  elements can be found by adding their individual
  covalent radii.
• e.g. the covalent radii of hydrogen and chlorine
  are 37 and 99 pm so the bond length in HCl is 37
  99 pm 136 pm
• There are two clear trends in the periodic table
• Going across a period covalent radii decreases.
  This is because
• Going down a group covalent radii increases. This
  is because

Good to know but not essential...
the nuclear charge increases but the number of
electron shells stays the same i.e. the outer
electrons are held more tightly, making the atom
smaller.
the number of electron shells increases and
therefore the inner (full) electron shells shield
the outer electrons from the nuclear charge.
(known as shielding effect) i.e. the outer
electron are held less tightly, making the atom
bigger.
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First Ionisation Energy

• Definition
• First ionisation energy is the energy required to
  remove one mole of electrons from an element in
  gaseous state. (example at top of page 11 in data
  book no excuses!)
• Trends
• Across a period increases due to an increased
  nuclear charge holding the outer electrons more
  closely (smaller atoms.) This means you need MORE
  energy to remove a mole of electrons.
• Down a group decreases due to the shielding
  effect (bigger atoms.) This means the outer
  electrons are further away from the nucleus and
  therefore this attraction is less, thus it is
  easier (less energy) to remove one mole of
  electrons.

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Other types of ionisation questions

1. Calculation (whiteboard examples)

• The first ionisation energy of lithium
  520kJmol-1 but its second ionisation energy value
  7298kJmol-1.
• Why is there such a big difference between the
  two
• values?
• Lithium achieves a stable outer electron shell
  (octet) when it loses one electron. Therefore
  first ionisation energy is small. The second
  electron would therefore be removed from an
  stable octet which is unfavourable and requires a
  lot of energy.

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Electronegativity

• Electronegativity is a measure of the tendency of
  an
• atom to attract electrons. (think pulling power)
• Electronegativity is measured on the Pauling
  scale.
• Trends
• Electronegativity increases across a period.
• Electronegativity decreases down a group.
• (Hint Fluorine is the highest value)

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• Melting and Boiling Points
• Generally, the stronger the bond between atoms,
  the higher the energy required to break that
  bond. Melting/boiling points are varied and don't
  generally form a trend across a period however
• Metals generally possess a high melting point.
• Most non-metals possess low melting points.
• Group 4 have the highest values.
• Metal group example
• In group 1 (alkali metals), the melting/boiling
  pts decrease as the atomic number increases. This
  is because there is an decrease in the attraction
  between the particles. (refer to bonding)
• Non metal group example
• In group 7 (halogens) the melting/boiling pts
  increase as the atomic number increases. This is
  because there is an increase in the attraction
  between the particles. (refer to bonding)

refer to bonding
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