Unit 1: Chemical Change And Structure Flashcards
1
Q
What is a catalyst?
A
A substance that increases the rate of a chemical reaction
2
Q
What happens to a catalyst after a reaction occurs?
A
It remains the same
3
Q
Name the two types of catalyst
A
Biological catalyst and chemical catalyst
4
Q
Define a chemical and a biological catalyst
A
A chemical catalyst is elements or substances added to induce a reaction
A biological catalyst is naturally occuring
5
Q
Define a heterogenous catalyst
A
It’s a different state from the reactants
6
Q
Define a homogenous catalyst
A
It’s in the same state as the reactants
7
Q
What do catalysts do?
A
They improve geometry and more successful collisions with lower activation energy
8
Q
Name the 4 types of bonding amongst the first 20 elements
A
Metallic
Covalent network
Covalent molecular
Monatomic
9
Q
Which group have monatomic bonding and what is it?
A
Group 8 have monatomic bonding and it means they consist of single, unbounded atoms
10
Q
Name the 7 elements amongst the first 20 have metallic bonding
A
Lithium
Beryllium
Sodium
Magnesium
Aluminium
Potassium
Calcium
11
Q
What are covalent networks and what’s their state? (What do covalent networks consist of?
A
They consist of thousands of atoms joined together with covalent bonds and they are solids
12
Q
What’s the structure of diamond? (Arrangement)
A
Tetrahedral arrangement, all four electrons used to make bonds
13
Q
Whats the structure of graphite? (Arrangement)
A
3 covalent bonds arranged in hexagonal rings, 4th electron is delocalised in layers
14
Q
Name 4 of the 20 elements that are diatomic
A
Nitrogen
Oxygen
Fluorine
Chlorine
15
Q
What are the diatomic molecules at room temperature? And what’s their size?
A
Gases and small
16
Q
What are the diatomic molecules in the first 20 elements examples of?
A
Covalent molecular molecules
17
Q
What are the 3 elements in the first 20 that are covalent molecular solids?
A
Carbon
Phosphorus
Sulphur
18
Q
What forces do the 3 elements that are covalent molecular solids have? What’s their state at room temperature?
A
Some weak attraction forces, and can be solid at room temperature
19
Q
How was the periodic tables invented? (Order)
A
Invented in order of increasing atomic mass
20
Q
What does melting point and boiling point depend on?
A
The strength of forces between the particles
21
Q
Why are elements on the left of the periodic tables generally stronger?
A
They have stronger intermolecular forces
22
Q
What happens to the melting point and boiling point as you go down group 1?
A
It decreases
23
Q
What’s the density of a substance? Calculation?
A
Mass per unit volume, in g/cm3
24
Q
Define covalent radius
A
Half the distance between the centres, nuclei, of 2 bonded atoms
25
Other name for covalent radius?
Atomic size
26
Here’s covalent radius measured? Unit?
Measured in picometers
27
What happens as you go across a period?
As we go across a period the nuclear charge and number of outer electrons increases
28
What happens as you go down a group?
As we go down a group, the number of electron shells or energy levels increases, but the number of outer electrons stays the same
29
In terms of the trends in atomic size: what happens as you go across a period?
The atomic size decreases as nuclear charge increases and attracts the outer electrons closer to the nucleus
30
In terms of trends in atomic size: what happens as you go down a group?
As you go down a group the atomic size increases as an extra electron shell is added
31
Define the term “First ionisation energy”
The amount of energy required to remove one mole of electrons from one mole of atoms in the gaseous state
32
What does the first ionisation of magnesium look like?
Mg (g) ———> Mg+(g) + e-
33
Which elements are involved in metallic bonding?
Metals
34
What happens to first ionisation energy as you go down a group?
It decreases
35
What happens to the first ionisation energy as you go across a period?
It increases
36
Define “second ionisation energy”
The amount of energy required to remove one mole of electrons from one mole of gaseous 1+ ions
37
Show the second ionisation energy formula for magnesium
Mg (g)+ ——> Mg(g) 2+ + e-
38
Why’s the 2nd ionisation energy of an element always greater than the 1st?
- In the 2nd ionisation energy of negative electrons are being removed from + ions rather than = ones
- in the positive ion there’s a greater attraction for the electron so more energy is needed to remove the 2nd mole of electrons
- successive ionisation energies increases as the atom becomes more positive
- there’s a large jump in ionisation energy when the electron to be removed comes from a new shell, closer to the nucleus
39
Define electronegativity
It’s a measure of an atoms attraction for the shared pair of electrons in a bond
40
Trends in terms of electronegativity: what happens across a period?
Electronegativity increases, the charge in the nucleus increases
41
Trends in terms of electronegativity: what happens down a group?
As we go down a group, electronegativity decreases, atoms have a bigger radius
42
What is covalent bonding a result of?
What do atoms share?
Two positive nuclei being held together by their common attraction for the shared pair of electrons
Atoms share electrons
43
What’s ionic bonding and what’s its force of attraction?
It’s the electrostatic force of attraction between the oppositely charged ions
Strong forces of attraction
44
What type of ions is ionic bonding between? What do ionic compounds form?
It’s between positive and negative ions
Ionic compounds form lattice structures of oppositely charged ions
45
How do polar covalent bonds form?
What’s it sometimes called?
When distribution of electrons is unequal a polar covalent bond forms (this is sometimes called a dipole)
46
What is there in polar covalent bonds? (Electronegativity)
There is a difference in electronegativity
47
What charge does the less electronegative atom develop? Polar covalent bonds
A slightly positive charge and vice versa for the more electronegative atom
48
Where are electrons pulled in polar covalent bonds?
Closest to the atom with the greatest electronegativity
49
How does pure covalent bonding occur?
When the elements are the same as one another or the electronegativities are equal or almost equal
50
What’s the distribution of electrons like in pure covalent bonds?
Bonding electrons are equally shared
51
What are pure covalent and ionic bonding considered to be on the opposite of?
What lays between?
The bonding continuum, with polar covalent bonding laying between the two extremes
52
What happens as the difference in electronegativity between two atoms occurs?
The more polar the bond will be, upto 1.9
53
If the different in electronegativity is 2, what does this mean for the movement of bonding electrons?
The movement of bonding electrons from the element with the lower electronegativity is complete resulting in the formation of ions
54
What type of bond is it if the difference in electronegativity is 0->0.4?
Pure covalent bonds
55
What type of bond is it if the difference in electronegativity is 0.5->1.9?
Polar covalent bonds
56
What type of bonding would it be if the different in electronegativity is 2->>>>
Ionic bonding
57
As a general rule, what’s the electronegativity difference in non polar covalent bonds?
<0.5
58
As a general rule, what’s the difference of electronegativity in polar covalent bonds?
>0.5
59
As a general rule, what’s the difference in electronegativity of ionic bonds?
>1.7
60
What’s the melting point of ionic lattices and covalent networks?
High
61
What’s the melting point of covalent molecular?
Low
62
What’s the melting point of a metallic lattice?
Can vary, usually high
63
How do ionic lattices conduct electricity? In which states?
In liquid and dissolved states. Not solid
64
How do covalent networks conduct electricity? In which states?
They don’t. Only as graphite
65
How does covalent molecular conduct electricity? In which state?
It doesn’t
66
How do metallic lattices conduct electricity? In which state?
They conduct in all states
67
What type of bond is this?
H - Br
2.2 - 2.8
Polar covalent
68
What type of bond is this?
N —- Cl
3.0 - 3.0
Pure covalent (non polar)
69
What must you take into consideration when there’s a molecule containing more than two atoms? (Polarity in molecules)
What does the determine?
- all polar bonds within the molecule
- the shape of the molecules
It’s to determine whether it has a permanent dipole ion (an overall polar structure)
70
Name the two types of covalent molecule (electronegativity)
Polar molecules and non polar molecules
71
Define polar molecules
Theses are molecules with overall polarity, one end is clearly more negative than the other
72
Define non polar molecules
These are molecules with no overall polarity
73
Define a permanent dipole
There is a positive and negative end of the molecule
74
If a molecule contains polar bonds, what determines whether a molecule has an overall polarity?
Spatial arrangement
75
If a molecule contains polar bonds, what determines whether a molecule has an overall polarity?
Spatial arrangement
76
What’s the structure of CO2? Is it polar? Why?
O=C=O
It’s non polar because it’s linear so both ends have a slightly negative charge
77
Generally are tetrahedral molecules polar or non polar?
Non polar
78
What are intermolecular forces also known as?
van der waals forces
79
Where are intermolecular forces present?
Between all molecules
80
What are the three types of van der waals forces?
1) London dispersion forces (LDFs)
2) Permanent dipole to permanent dipole interaction
3) hydrogen bonding
81
What break when you melt or boil a substance?
The van der waals forces, NOT the covalent bonds within the molecule
82
Is it possible for a molecule to have more than one type of van der waals force?
Yes
83
LDFs: what do electrons do and what does this mean?
Electrons move around atoms randomly which means that more often than not there’s more electrons on one side of the atom than the other
84
Where do LDFs exist?
Between all atoms and molecules
85
Why do LDFs arise?
Because electrons move around the atoms
86
Which molecules are LDFs the main force?
Non- polar molecules
87
What does the uneven distribution of LDFs cause?
A slightly negative and a slightly positive charge on either side of the atom, known as a temporary dipole
88
What is an induced dipole and how does it happen? (LDFs)
When another atoms is nearby, the electrons of that atom are repelled by the electron density of the temporary dipole, pushing them to the other side of that atom, creating a second dipole. This is called an induced dipole.
89
Define an LDF
The forces of attraction between temporary dipoles and induced dipoles caused by movement of electrons in atoms and molecules
90
What does the strength of LDFs depend on?
The size of the molecule/atom
91
What does the strength of LDFs depend on?
The size of the molecule/atom
92
How are larger dipoles established?
Larger atoms and molecules have more electrons. This leads to larger dipoles being established.
93
What is a permanent dipole to permanent dipole interaction and how does it occur?
When two polar molecules are near each other there’s an attraction between the negative end of one of the molecules and positive end of the other. This is called a permanent dipole to permanent dipole interaction.
94
What’s stronger? LDFs or permanent dipole to permanent dipole interaction?
Permanent dipole to permanent dipole interaction
95
Rank all the van der waals forces by strength
1 = strongest
1 - hydrogen bonds
2- permanent dipole to permanent dipole interaction
3 - London dispersion forces (LDFs)
96
Where do hydrogen bonds occur? Van der waals forces
Between molecules where there is a very high electronegativity
97
For a molecule to have hydrogen bonding, it must have one or more of which bonds?
Nitrogen (N-H)
Oxygen (O-H) All highly polar
Fluorine (F-H)
NOF bonding****
98
Where do permanent dipole to permanent dipole interactions occur?
Between polar molecules
99
Do polar molecules react to an electric field?
Yes
100
What would happen if you help a charged rod near a stream of a polar liquid?
The liquid will be attracted or repelled
101
What is viscosity?
The measure of how thick a liquid is
102
Do viscous liquids have strong intermolecular forces?
Yes
103
What is an oxidising agent?
A substance that accepts electrons, whilst being reduced
104
What does an oxidising agent assist?
Oxidation
105
Which elements tend to be oxidising agents?
Those with high electronegativities tend to form ions by gaining electrons, so act as oxidising agents
106
What’s a reducing agent?
A substance that donates electrons, whilst it is oxidised
107
What does a reducing agent assist?
Reduction
108
Which elements tend to be reducing agents?
Those which have low electronegativities tend to form ions by losing electrons, so act as reducing agents
109
Which elements tend to be reducing agents?
Those which have low electronegativities tend to form ions by losing electrons, so act as reducing agents
110
Which group have the strongest reducing agents?
Group 1
111
Which group have the strongest oxidising agents?
Group 7
112
Example of a strong (strongest) reducing agent
Carbon monoxide
113
Example of oxidising agent
Hydrogen peroxide H2O2
114
Uses for oxidising agents
They can inactivate viruses and their ideal for bleach due to the oxidation process being an effective means of breaking down coloured compounds
115
Where are the strongest oxidising agents placed in the electrochemical series?
Bottom left hand column of the electrochemical series
116
Where are the strongest reducing agents placed in the electrochemical series?
The top of the right hand column
117
What’s the balanced redox equation for this?
Al(s) ——-> Al3+ (aq) + 3e-
2H+ (aq) + e- ——-> H2 (g)
6H+ + 2AL ——> 2AL3+ + 3H2
118
What’s always cancelled out in redox equations?
Electrons
119
What’s miscibility?
The ability for a liquid to “dissolve” another liquid
120
What aren’t immiscible liquids miscible?
They have a boundary between them
Example: water and hexane
Water is polar, hexane is not
121
Define solvent
The liquid in which the substance dissolves
122
What group and one element make covalent network structures?
Group 4 and boron
123
What is thermochemistry?
The study in changes in heat energy which occur during a chemical reaction
124
What’s an endothermic reaction?
The reaction takes in heat
125
What’s an exothermic reaction?
A reaction that releases heat into the atmosphere
126
Which has higher products energy after a reaction? Endothermic or exothermic
Endothermic because it takes IN heat
127
What is an activated complex?
An unstable arrangement of atoms formats as react at bonds are broken and product bonds are formed.
128
When does an activated complex occur?
At the maximum point of the potential energy diagram
129
What does it mean if an atom has a symmetrical shape? Polarity…
It means there’s no polarity, therefore it’s not permanent
130
How does sulphur exist?
S8
131
How does phosphorus exist?
P5
132
Define the term “nuclear charge”
The force of attraction between the nucleus and outer electrons
133
Name all the diatomic elements
Chlorine
Hydrogen
Iodine
Nitrogen
Fluorine
Oxygen
Bromine
134
Alkali metals: what happens as you go down a group?
- Reactivity increases
- melting point and boiling point decreases
135
Alkali metals: what happens when they react with water? Final product?
React with water to release hydrogen and form hydroxides which dissolve in water making alkaline solutions
136
Alkali metals: density?
Low density, lithium, sodium and potassium float on water
137
Transition metals: what do they form and are they catalysts?
They’re catalysts and they form coloured compounds
138
Transition metals: compared with group 1, melting point, strength, density, reactivity?
Melting point - higher minus per hey
Strength - stronger
Density - harder
Reactivity - less reactive and don’t react as vigorously with water or oxygen
139
Transition metals: properties (similar/different) etc
Similar properties and some special ones because a lower energy level is being filled in the atoms of elements
140
Noble gases: 3 features (bonding, reactivity, colour)
Bonding - monatomic
Reactivity - unreactive
Colour - colourless
141
Halogens: trends down a group? (Mp/bp, reactivity)
Reactivity decreases down the group
Melting point and boiling point increase
142
Halogens: what two things can they form?
1) ionic salts with metals
2) simple covalent molecules with other non-metals
143
What happens when you put a more reactive halogen and a less reactive halogen in an aqueous solution?
The more reactive halogen can displace the less reactive one
144
What’s a fullerene? Bonding
Discrete covalent bond
145
Covalent networks, strength?
What three elements can form these?
Very strong forces between bonds
3 elements: carbon, boron, silicon
146
Define a covalent bond
(What creates it?)
When 2 positive nuclei are held together by their common attraction to a shared pair of electrons
147
Name the 4 factors of reaction rate
Temperature
Concentration
Particle size
Use of a catalyst
148
Name 5 signs that a chemical reaction has occurred
1 - colour change
2 - smell
3 - change in pH
4 - temperature change
5 - fizzing
149
What must happen for a reaction to occur?
Particles must collide with a minimum energy to react, called the activation energy
150
How does pressure affect reaction rate?
Increasing the pressure in a reaction involving gases, reduces the volume for particles to move around in. This results in more collisions, therefore increasing the reaction rate.
151
Define activation energy
The activation energy, Ea, is the minimum kinetic energy required by colliding particles before a reaction may occur
152
What affect does a catalyst have on activation energy?
They lower the required activation energy
153
What is the stored energy that every substance contains called?
Enthalpy
154
What must happen for a successful collision to take place?
The collision geometry must be correct (the reactant molecules have to be facing the right way! So that the activated complex can be formed
155
Define metallic bonding
The electrostatic force of attraction between positively charged ions and delocalised electrons
156
Metal + oxygen —-> ?
Metal oxide
157
Metal + water ——> ?
Metal hydroxide + hydrogen
158
Metal + dilute acid ——-> ?
Salt + hydrogen
159
Name the four states of matter
Liquid. Solid, gas and aqueous
160
What is an atomic number?
Number of protons
161
What’s a mass number?
Protons + neutrons
162
What are atoms made of?
Three subatomic particles
Neutrons, protons and electrons
163
What’s an isotope?
Atoms with the same atomic number but a different mass number
164
Name the four structures of molecules
Linear
Angular
Trigon pyramidal
Tetrahedral
165
An element has the atomic number 11. Using the periodic table you can identify it as?
Sodium
166
Label this nuclide notation
35
17 Cl
35 —-> mass number
17 —-> atomic number
Cl ——€ symbol
167
What’s an ionic lattice?
A regular repeating arrangement of metal and non metal ions which creates compounds with very high melting points
168
The difference between the first ionisation energies of sodium and chlorine is mainly due to the different in?
Number of protons
169
Explain why electronegativity values decrease going down group 7
Covalent radius increases so attraction of the nucleus for the outer electrons decreases
170
Explain fully why the boiling points of the halogens increase going down group 7 (3 marks)
It’s because the intermolecular forces increase going down a group
LDFs are forces between the molecules
The more electrons the stronger the LDFs
171
Explain the decrease in atom size going across the third period from sodium to argon
Increase in the number of protons in the nucleus
172
Silicon nitride has a melting point of 1900°C and does not conduct electricity when molten
Explain fully, in terms of structure and bonding, why silicon nitride has a high melting point (2 marks)
Silicon nitride is a covalent network.
Strong covalent bonds are broken
173
Explain fully why, of these three chlorides, silicon tetrachloride is the most soluble in hexane.
Silicon tetrachloride, phosphorus and sulfur (2 marks)
Silicon tetrachloride and hexane are non-polar.
Silicon tetrachloride is non-polar due to its shape cancelling out
174
The different between the first ionisation energies of sodium and chlorine is mainly due to the difference in the…
Number of protons
175
Which of the following contains pure covalent bonds?
CO2, H2S, PH3, CF4
PH3
176
A student is carrying out a titration. Which of the following would help the student to accurately observe the end-point?
1) repeating the titration
2) using a white tile under the flask
3) rinsing the flask between titrations
4) disregarding the rough titre
Using the white tile under the flask
177
Write the equation for the first ionisation of phosphorus
P(g) ——> P+(g) + e-
178
Explain why the melting point of phosphorus P5, is much higher than that of nitrogen N2. Refer to the intermolecular forces involved. 3 marks
1) stronger intermolecular forces between phosphorus compared to nitrogen
2) LDFs are the intermolecular forces present
3) There are more electrons in P5 compared to N2
179
Why does silicon dioxide have a high melting point? (2 marks)
Silicon dioxide is a covalent network
Strong covalent bonds are broken
180
Explain full why chloromethane has a lower boiling point than water. In your answer you should refer to the intermolecular forces involved
Chloromethane has permanent dipole to permanent dipole interactions whereas water has hydrogen bonding
Hydrogen bonding is stronger than permanent dipole to permanent dipole interactions
181
Titanium chloride is a liquid at room temperature. Suggest the type of bonding and structure present in titanium chloride
Covalent molecular
182
Write the ion-electron equation for the oxidation of magnesium atoms
Mg(s) —-> Mg2+ (aq) + 2e-
