Structure & Reactions Flashcards
1
Q
Examples of slow everyday reactions
A
- Milk going sour
- Iron rusting
- Bread going mouldy
2
Q
Examples of fast everyday reactions
A
- Fireworks
- Popcorn
- Starting a car
3
Q
How do you increase the rate of reaction ~ collision theory
A
DECREASE particle size
INCREASE concentration
INCREASE temperature
ADD a suitable catalyst
4
Q
What is needed for a chemical reaction to occur?
A
- The reacting particles must collide together
- Collisions must have sufficient energy to overcome the energy barrier
- The reacting particles must have the correct collision geometry
5
Q
Why does decreasing particle size speed up a reaction?
A
The smaller the particle size, the larger the surface area.
The more surfaces there are, the more collisions can take place, therefore increasing the reaction rate
6
Q
Why does increasing concentration speed up the rate of reaction?
A
The higher the concentration, the higher the no.of molecules there are present and available for collisions.
Therefore collisions occur more frequently and the reaction goes faster
7
Q
Why does increasing temperature speed up the rate of reaction?
A
At higher temperatures more particles have beget equal to the activation energy.
The activation energy is the amount of energy required for a successful collision to take place.
Therefore more successful collisions take place and the reaction goes faster
8
Q
Why do catalysts speed up reaction rate?
A
They lower the activation energy.
That is they reduce the amount of energy particles need in order to achieve successful collisions.
Therefore successful collisions occur more easily, more frequently, and the reaction goes faster.
9
Q
Formula for average rate of reaction
A
Average rate = change in quantity/ change in time
10
Q
Formula for relative rate of reaction
A
Rate = 1/t
11
Q
What is enthalpy (H)
A
A measure of the energy stored in a chemical
12
Q
What is an exothermic reaction?
A
A reaction in which heat is given off/ lost
13
Q
What is an endothermic reaction?
A
A reaction in which chemicals absorb energy/ gain energy
14
Q
What is a reaction profile?
A
Diagrams showing the difference in the enthalpy the reactants have at the start compared with the enthalpy the products have at the end
15
Q
Formula for enthalpy change
A
Hproducts- Hreactants
16
Q
What are the units for enthalpy change?
A
kJmol-1
17
Q
What is activation energy (Ea)?
A
The minimum kinetic energy that colliding molecules must have in order for a reaction to occur
18
Q
What is an activated complex?
A
An unstable intermediate in which old bonds are partially broken and new bonds are partially formed
Eg. The reaction between hydrogen and bromine
19
Q
Why do reacting particles need the correct collision geometry to collide successfully?
A
If particles do not collide with the correct collision geometry the activated complex cannot form
20
Q
What is temperature?
A
Temperature is a measure of the average kinetic energy of the particles in a substance
21
Q
What is an energy distribution diagram?
A
Diagrams which show the effect that temperature has on the number of successful collisions and hence reaction rate
22
Q
Where on an energy distribution diagram can the particles with sufficient energy for successful collision be found?
A
On the right side of the activation energy line
23
Q
What difference does a small increase in temperature make to the number of particles which have energy equal to or greater than the required Ea?
A
It increases the no.of particles which have energy equal to or greater than the Ea
24
Q
What effect does a small increase in temperature have on the no.of successful collisions?
A
It increases the no.of successful collisions as more particles have the required activation energy
25
What are catalysts?
Catalysts speed up the rate of a chemical reaction and can be recovered after the reaction
26
What are homogeneous catalysts?
Catalysts which are in the same state as the reactants
27
What are heterogenous catalysts?
Catalysts which are in a different state to the reactants
28
How do catalysts work?
Catalysts speed up chemical reactions by providing an alternative reaction pathway which has a lower activation energy
29
What are the three stages involved with catalysts?
Adsorption
Reaction
Desorption
30
What is adsorption?
Reactants form a temporary bond with the catalyst.
This weakens the bonds within the molecules.
Thus, the activation energy is lowered
31
What is the reaction stage of a catalyst?
The molecules react on the catalyst surface.
The collision geometry is more favourable since one of the molecules is fixed
32
What is desorption?
The product molecules leave the catalyst and the vacant site can be occupied by another reaction molecule.
In other words, catalysts can be reused
33
What type of diagrams can catalysts be shown on?
Potential energy diagrams
Energy distribution diagrams
34
Who constructed the periodic table?
Dimitri Mendeleev
35
What group number are the Noble Gases?
0/8
36
What group number are the Halogens?
7
37
What group number are the Alkaline Earth Metals?
2
38
What group number are the Alkali Metals?
1
39
Why do elements in the same group have similar chemical properties?
They have the same number of outer electrons
40
What is bonding?
A term that describes how atoms join together
41
What is structure?
Describes how the atoms in the element are arranged
42
What are properties of an element?
The characteristics of the substance resulting from its bonding and structure, whether physical or chemical
43
What type of structures do metallic substances form?
Lattice structures
44
What allows the outer electrons of metallic atoms to be delocalised?
They’re held loosely
45
What does metallic bonding consist of?
The atoms losing their outer electrons to a common ‘pool’ of delocalised electrons.
As each atom has lost one or more electrons, the atoms become positively charged ions.
The charged metal ions are now attracted to the pool of electrons
46
Why do Metals conduct electricity?
As the electrons are free to move
47
Why are Metals shiny/ have a lustre?
When light is shone onto a metal, the delocalised electrons absorb the energy from the light and then re-emit the light which gives the metal its characteristic shiny appearance
48
Why are most metallic substances solid at room temperature?
Metallic bonding is strong
Due to the energy required to break these bonds and change them into a liquid
49
What may covalent substances take the form of?
* Discrete covalent molecules
| * Covalent molecules
50
What is the trend between the properties of covalent molecular substances?
Low melting and boiling points
51
Why do discrete covalent molecular substances have low melting and boiling points?
Due to the fact that molecules of these substances are held together by weak intermolecular bonds called London Dispersion Forces.
As these intermolecular forces are weak it doesn’t require a lot of energy to overcome them
52
What structure do sulphur atoms form?
Closed, eight membered, puckered rings
53
What structure do phosphorus atoms form?
Tetrahedral molecules
54
What are Fullerenes?
Discrete, covalent bonded molecules of carbon
Spherical
Buckminsterfullerene
55
What are covalent networks?
Huge complex structures of thousands of atoms covalently bonded together
56
What are the four covalent network structures?
* Carbon Diamond
* Carbon graphite
* Silicon
* Boron
57
Why is carbon diamond exceptionally hard and ridged?
There are no discrete molecules; all carbon atoms are joined together in the structure
58
Describe the structure of a carbon diamond network
Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement
59
Why does carbon diamond not conduct electricity?
There are no free electrons
60
Why is carbon diamond transparent?
Due to the way the structure is arranged, ‘light tunnels’ form between the atoms
61
Describe the structure of carbon graphite
Graphite has a structure based on three covalent bonds from each atom, forming layers of hexagonal rings
Each carbon atom has a fourth unpaired electron
This causes a pool of electrons
The result is strong covalent bonding within the layers but only weak interactions between the layers
62
Why does graphite conduct?
Fourth unpaired electron
Pool of electrons
63
What is silicon oxide?
Sand
64
How common is silicon?
Second most common element found on earth
65
Example of a silicon compound
Silicon oxide - sand
66
What is silicon used for?
In electronics in the form of silicon chips
67
Example of a use of boron
Pyrex glassware - boron oxide
Strength and ability to withstand high temperatures
68
What makes boron almost as hard as diamond?
It’s covalent network structure
69
What’s the trend between the properties of covalent networks?
Exceptionally high melting and boiling points
70
Why do covalent networks have high melting and boiling points?
The strong covalent bonds within the structure require high quantities of energy to overcome
71
What are monatomic elements?
Single atoms
The Noble gases
72
What is the covalent radius?
Size of an atom
The covalent radius of an atom is taken to be half the distance between the nucleus of 2 of its bonded atoms
73
Atomic radius ______ decreases as you go across a period
Decreases
74
Atomic radius ______ as you go down a group
Increases
75
Explain why atomic radii increases on going down a group
* although the no.of protons and therefore the nuclear charge increases, its effects are overpowered by the no. Of electron layers increasing
* every additional electron layer ‘shields’ the outer electrons from the positive nucleus - they stop them getting as close to the nucleus so the protons pull is weaker
* This results in atomic size increasing
* this ‘sheilding’ effect is called screening
76
Explain why atomic radii decreases on going across a period
* As you go across a period, the atomic number increases so the nuclear charge increases
* That then causes a greater attraction for the electrons and they are pulled in closer
* This results in the atomic size decreasing
77
What is the covalent bond a result of?
Two positive nuclei being held together by their common attraction for the shared pair of electrons
78
What is electronegativity?
A measure of the attraction an atom involved in a bond has for the electrons of the bond
79
As you go down a group, electronegativity ______
Decreases
80
As you go across a period, electronegativity _____
Increases
81
Explain why electronegativity across a period increases
Because nuclear charge increases each time we move across a period the attraction between the electrons and the nucleus gets stronger.
In addition, each time the nuclear charge increases the distance between the outer electrons and the nucleus decreases.
Therefore, the attraction between the outer electrons and the nucleus becomes stronger.
Increased nuclear charge and decreased distance from the nucleus means that electronegativity across a period increases.
82
Explain why electronegativity decreases down a group.
Because an additional electron cloud is added to an atom each time we move down a group the distance between the outer electrons and the nucleus increases.
Therefore, the attraction between the electrons and the nucleus becomes weaker.
In addition, each time an electron cloud is added it acts like a screen between the nucleus and outer electrons and weakens the attraction between the electrons and the nucleus.
Increased distance from the nucleus and screening means that electronegativity decreases down a group.
83
What is the first ionisation energy of an element?
The energy required to remove one mole of electrons from one mole of atoms in the gaseous state
84
As you go down a group, ionisation energy ____
Decreases
85
As you go across a period the ionisation energy _____
Increases
86
Explain why ionisation energy decreases as you go down a group
Because an additional electron cloud is added to an atom each time we move down a group the distance between the outer electrons and the nucleus increases.
Therefore, the attraction between the electrons and the nucleus becomes weaker.
In addition, each time an electron cloud is added it acts like a screen between the nucleus and outer electrons and weakens the attraction between the electrons and then nucleus.
Increased distance from the nucleus and screening means that ionisation energy decreases down a group.
87
Explain why ionisation energy increases as you go across a period.
Because nuclear charge increases each time we move across a period the attraction between the electrons and the nucleus gets stronger.
In addition, each time the nuclear charge increases the distance between the outer electrons and the nucleus decreases.
Therefore, the attraction between the outer electrons and the nucleus becomes stronger.
Increased nuclear charge and decreased distance from the nucleus means that ionisation energy across a period increases.
88
Elements with the greatest electronegativity is more likely to ____ electrons to form ____ ions
Gain
Negative
89
Elements with the
| Smallest electronegativity is more likely to ____ electrons to form ____ ions
Lose
Positive
90
Conductivity of ionic substances?
Conducts when molten/ in solution due to lattice breaking allowing electrons to move freely
91
What is a pure covalent bond?
When atoms of the same element form a covalent bond they share the pair of electrons as they have the same electronegativity
92
What type of covalent bond would hydrogen and hydrogen form?
Pure covalent bond
93
Example of pure covalent bond
Hydrogen-hydrogen
94
What’s a polar covalent bond?
When atoms with different electronegativities form bonds and the electrons are pulled closer to the the atom with the highest electronegativity
95
Example of polar covalent bond
Hydrogen fluoride
96
What are Van der Waal’s forces?
The weak forces of attraction between atoms
97
What are the 3 types of Van der Waal’s forces?
* London Dispersion Forces
* Permanent dipole - permanent dipole attractions
* Hydrogen bonding
98
What’s a temporary dipole and why is it temporary?
The entire electron cloud tends to wobble around which means at any one moment the majority of electrons are in one side of the atom, creating a slightly negative side, and the nucleus is on the other side, creating a slightly positive side.
Temporary bc the electron cloud could wobble in another direction a moment later
99
What is the London Dispersion Force?
A temporary, weak electrostatic attraction between the DELTA + end of a temporary dipole of one atom and the DELTA - end of a temporary dipole of another atom.
LDF are the weakest type of Van der Waal’s attraction
100
The _____ the size of an atom, the greater the electron cloud wobble
The greater the electron could wobble the greater the ____ in the DELTA+ and DELTA- ends of the dipole, which results in a greater attraction between the two atoms and ___ LDF
Greater
Difference
More
101
What are permanent dipole- permanent dipole attractions?
Weak electrostatic attractions that exist between polar molecules
102
Why are permanent dipole- permanent dipole attractions stronger than London Dispersion Forces?
Because they occur between permanent dipoles rather than momentary (temporary) dipoles
103
What’s hydrogen bonding?
* strongest Van der Waal’s force
* Weak electrostatic attraction that exists between polar molecules when the molecules contain a hydrogen atom bonded to nitrogen, oxygen, or fluorine
104
Why do N, O, and F form strong hydrogen bonds?
They’re all extremely electronegative compared with hydrogen so the bonds they form are highly polar
105
Example of molecule with hydrogen bonding
Water
106
What’s ionic character?
The bond’s tendency to be ionic
107
Ionic character _____ as the difference in electronegativity increases
Increases
108
The larger the difference in electronegativities between bonded atoms, the more ____ the bond will be
Polar
109
Where do LDF exist?
Between all molecules + atoms, but the only force between non-polar molecules and monatomic elements.
Weakest attractive force.
110
Example of molecules which contain LDF
Between halogen molecules and between bible gas atoms
111
Example of molecule which contains permanent dipole- permanent dipole attractions
Between ICl molecules
112
Describe viscosity
The stronger the intermolecular force of attraction between the molecules the more viscous the liquid will be
113
Why do stronger intermolecular forces result in a more viscous liquid?
Bc the stronger the attraction there is between molecules the more tightly the molecules will be held together and the thicker the liquid will be
114
What’s the solubility of a substance?
It’s ability to dissolve in a liquid
115
What’s the miscibility of a substance?
It’s ability to mix with another substance
116
When are substances said to be immiscible?
If they don’t mix and separate into separate layers
117
What’s the general rule for solubility/miscibility?
Like dissolves like
118
Polar and ionic substances tend to dissolve in ____ solvents
Polar
119
Non polar substances tend to dissolve in _______ substances
Non polar
120
Polarity of water?
Extremely polar
121
Why is ice less dense than water?
* when water freezes the H bonding between molecules forces them into a hexagonal shape
* This makes molecules spread out more than they do as liquid
