Overall Flashcards
(177 cards)
1
Q
What are isotopes?
A
Isotopes are atoms of the same element with different numbers of neutrons and different masses
2
Q
Define relative atomic mass (Ar)
A
RAM is the weighted mean mass of an atom of the element relative to one twelfth the mass of one atom of the C-12 isotope, which has a mass of exactly 12
3
Q
Define relative isotopic mass
A
The relative isotopic mass is the mass of an atom of a particular isotope relative to one twelfth the mass of one atom of the C-12 isotope, which has a mass of exactly 12
4
Q
How do you calculate relative atomic mass?
A
((% abundance x mass) + (% abundance x mass))
/100
5
Q
What are the two main uses of mass spectrometry?
A
(i) The determination of relative isotopic masses and relative abundances of the isotopes
(ii) Calculation of the relative atomic mass of an element from the relative abundances of its isotopes
6
Q
How does the mass spectrometer show this?
A
The mass spectrometer provides a trace that shows the mass of each isotope and its relative abundance.
7
Q
What are the five stages, in order, of mass spectrometry?
A
1) Vaporization
2) Ionisation
3) Acceleration
4) Deflection through a magnetic field
5) Detection
8
Q
How are the particles deflected in mass spectrometry?
A
Electrons are attracted to a positively charged plate. They have a large deflection because of their small mass. Neutrons do not change path in a magnetic field. Protons attracted to a negatively charged plate, but their deflection is smaller because their mass is larger.
9
Q
What happens when an organic compound is placed in the mass spectrometer?
A
The organic compound loses an electron and forms a positive ion, called the molecular ion.
10
Q
What does the mass spectrometer detect?
A
A mass spectrometer detects the mass-to-charge ratio (m/z) of the molecular ion, which gives the molecular mass of the ion.
11
Q
How do you calculate a mass-to-charge ratio?
A
(Relative mass of ion) / (Relative charge on ion)
12
Q
What relationship exists between the time of flight of the sample (t) and the mass (m) of the sample?
A
The time of flight of the sample (t) is proportional to the square root of the mass (m) of the sample.
13
Q
What is the mole number in terms of the C-12 isotope?
A
The mole contains the same number of particles as exactly 12 grams of carbon-12. 1 mole of any substance contains the same number of particles.
14
Q
What does Avogadro’s constant (Na) show?
A
Avogadro’s constant is a measure of the number of particles in 1 mole of any substance
6.02 x 10^23 mol-1
15
Q
What is the molar mass of an element?
What are the units?
A
The molar mass of an element is its atomic mass in grams per mole.
Units: g mol-1
16
Q
What is the equation for calculating the number of particles there are in a sample?
A
Moles = Na
Mass = Mr x Mole
17
Q
How is the number of moles (n) calculated?
A
n = mass in g (m) / molar mass in g mol-1 (M)
Molar mass is equal to formula mass
18
Q
What is the molar gas volume?
What are the units?
A
The molar gas volume is the gas volume per mole.
Units: dm^3mol-1
19
Q
What volume does one mole of gas at standard conditions occupy?
A
One mole of any gas at room temperature (240C) and 1 atmospheric pressure has a volume of 24dm3
20
Q
How do you calculate the molar volume of a gas?
A
n= v/V or v=nV
Where n = number of moles
v = volume of gas
V= molar volume (24dm3)
21
Q
What is the empirical formula?
A
The simplest whole number ratio of atoms of each element present in a compound.
22
Q
What is the molecular formula?
A
The number and type of atoms of each element in a molecule.
23
Q
If a salt is hydrated, what does it contain?
A
Water
24
Q
If a salt is anhydrous, what has been removed?
A
Water
25
Define the term 'water of crystallisation'
Water molecules form an essential part of the crystal structure of the compound. % water of crystallisation shows how much of the hydrated salt is water
26
How do you determine the % water of crystallisation?
A known mass of the hydrated salt is heated gently in a crucible until the mass becomes constant despite further heating. The mass of the anhydrous salt is measured, and from this the percentage of water present in the hydrated salt can be calculated.
27
What is the Ideal Gas equation?
pV=nRT
28
What is the standard measure for pressure in the Ideal Gas equation?
The standard unit for pressure is the Pascal (Pa).
| In the Ideal Gas equation, the standard form is 1 atmospheric pressure, which equals 100 x 103 Pa.
29
What is the standard measure for volume in the Ideal Gas equation?
The standard unit for volume is m3.
m^3 = 1000dm^3 = 10^6cm^3
30
In the Ideal Gas equation, what does n stand for?
The number of moles
31
What is the standard measure for temperature in the Ideal Gas equation?
Standard: 298 Kelvins
32
How do you convert Celsius to Kelvin?
Celsius + 273 = Kelvins
33
In the Ideal Gas equation, what does R stand for?
R is the gas constant.
| It has a value of 8.31441 J K^-1 mol^-1
34
How do you calculate the molecular formula from the empirical formula?
The molecular formula shows how many times the empirical formula goes into the molar mass.
(molar mass / empirical formula)
Multiply the empirical formula by the answer to find the molecular formula.
35
What do ionic equations show?
Ionic equations show only the reacting particles (and the products they form.) Ions that don’t get involved in the reaction - the spectator ions - aren't written.
36
What is the formula of the nitrate ion?
NO3 ^-
37
What is the formula of the carbonate ion?
CO3 ^2-
38
What is the formula of the sulphate ion?
SO4 ^2-
39
What is the formula of the hydroxide ion?
OH ^-
40
What is the formula of the ammonium ion?
NH4 ^+
41
What is the formula of the zinc ion?
Zn ^2+
42
What is the formula of the silver ion?
Ag+
43
What is an acid?
An acid is a proton donor. When mixed with water, they release hydrogen ions - H+.
44
What is a base?
A base is a proton acceptor. They want to grab H+ ions.
45
What is an alkali?
An alkali is a base that is soluble in water. Alkalis produce OH- ions in an aqueous solution.
46
Name three common acids and give their formula
Hydrochloric acid - HCl
Sulfuric acid - H2SO4
Nitric acid - HNO3
Ethanoic acid - CH3COOH
47
Name three common bases and give their formula
Sodium hydroxide - NaOH
Potassium hydroxide - KOH
Ammonia - NH3
48
What kind of reaction occurs between an acid + water or a base + water?
Reversible
49
What happens when strong acids are reacted with water? Which arrow is used for the reaction equation?
For strong acids, nearly all of the acid will dissociate in water, and lots of the H+ ions are released. As the equilibrium lies very far to the right, a forward arrow is used.
50
What happens when strong bases are reacted with water? Which arrow is used for the reaction equation?
For strong bases, nearly all of the base will dissociate in water, and lots of the OH-ions are released. As the equilibrium lies very far to the right, a forward arrow is used.
51
Which arrow is used when weak acids and bases react with water? Why?
A reversible arrow is used, as the backward reaction is favoured and only a few H+ or OH- ions are released.
52
What is the general equation for a neutralisation reaction between an acid and a base?
Acid + Base -> Salt + Water
The hydrogen ions released by the acid and the hydroxide ions released by the alkali combine to form water:
H+(aq) + OH-(aq) ≈ H2O(l)
The salt is produced when hydrogen ions in the acid are replaced by metal ions from the alkali.
53
Metal + Acid = ?
Metal Salt + Hydrogen
54
Metal Oxide + Acid = ?
Salt + Water
55
Metal Hydroxide + Acid = ?
Salt + Water
56
Metal Carbonate + Acid = ?
Metal Salt + Carbon Dioxide + Water
57
When ammonia reacts with acid, what is produced?
An ammonium salt (aqueous)
58
Which indicator is used when adding acid to alkali and what colour does it turn?
Methyl orange turns from yellow to red when adding acid to alkali
59
Which indicator is used when adding alkali to acid and what colour does it turn?
Phenolphthalein indicator turns from pink to colourless when adding acid to alkali
60
What is a diprotic acid?
A diprotic acid can donate two protons (i.e. H2SO4).
61
If you are completing a titration calculation, what do you have to consider when using a diprotic acid?
You need to double the number of moles of a base to neutralise a diprotic acid.
62
What is the molar volume equation?
Moles = concentration x (vol (cm ^3) / 1000)
63
How do you calculate percentage yield?
Percentage yield =
(Actual yield / Theoretical yield) x 100
64
Define 'atom economy.'
Atom economy is a measure of the proportion of reactant atoms that become part of the desired product.
65
How do you calculate atom economy?
(Molecular mass of desired products / Sum of molecular mass) x 100
66
What is an addition reaction?
In an addition reaction, the reactants combine to form a single product. The atom economy for addition reactions is always 100% since no atoms are wasted.
67
What is a substitution reaction?
A substitution reaction is one where some atoms from one reactant are swapped with atoms from another reactant. This type of reaction always results in at least two products.
68
Give two environmental benefits of having a high atom economy.
1. Few waste products: Waste products need to be disposed of and can be harmful to the environment.
2. More sustainable: Many raw materials are in limited supply, so using them efficiently makes them last as long as possible.
69
Give two economical benefits of having a high atom economy.
1. It costs money to separate and dispose of waste products. Less waste = less cost.
2. Reactant chemicals are costly. The more which can be transferred to useful products, the better.
70
What is an oxidation number?
Oxidation numbers tell you how many electrons an atom has donated or accepted to form an ion, or to form part of a compound.
71
What is the oxidation number of all uncombined elements?
0
72
The oxidation number of a monatomic ion is always the same as its ________.
Charge
73
For molecular ions, the sum of the oxidation ions is always equal to the ______ _________.
Overall charge
74
Oxygen nearly always has an oxidation number of ____
| Exceptions: peroxides and molecular oxygen.
-2
75
In peroxides, oxygen has an oxidation number of ____.
-1
76
Hydrogen usually has an oxidation number of ____
| Exceptions: metal hydrides and molecular hydrogen.
+ 1
77
In metal hydrides, hydrogen has an oxidation number of ___.
- 1
78
Roman numerals on an ion tell you the oxidation number. What oxidation number would be assigned to iron in iron (II) sulfate?
+ 2
79
What occurs in a redox reaction?
Electrons are transferred.
80
A loss of electrons is called _______.
Oxidation
81
A gain of electrons is called _______.
Reduction
82
What happens to an oxidising agent?
It accepts electrons and gets reduced. The oxidation number decreases by 1 for each electron gained.
83
What happens to a reducing agent?
It loses electrons and gets oxidised. The oxidation number increases by 1 for each electron lost.
84
Metals are _______ when they react with acids.
Oxidised
| This produces positive metal ions (in salts), and their oxidation number increases.
85
What is the relationship between the quantum number and the number of electrons the shell can hold?
2n ^2
86
How many electrons can the first four shells hold?
1st: 2
2nd: 8
3rd: 18
4th: 32
87
How do you define an atomic orbital?
A region in space around the nucleus that can hold up to two electrons with opposite spin.
88
What is the shape of an s orbital?
Spherical
89
In which quantum shells do s orbitals appear?
There is an s orbital in every quantum shell.
90
What is the shape of a p orbital?
Dumbbell shaped. There are three p orbitals and they are at right angles to one another.
91
In which quantum shells do p orbitals appear?
All quantum shells apart from the first
92
How many electrons in total can the p orbitals in one shell hold?
6
93
How many electrons in total can d orbitals in one shell hold?
10
94
Which three rules apply when filling energy levels?
1. Electrons enter the lowest energy level available
2. Two electrons in the same orbital must have opposite spin
3. Electrons will try and remain unpaired to reduce electrostatic repulsion and increase stability.
95
What is an ionic bond?
An electrostatic attraction between positive and negative ions in a giant lattice.
96
Describe a giant ionic lattice structure.
Oppositely charged ions are strongly attracted in all directions.
97
Do ionic compounds conduct electricity?
Only when molten or in solution, as the charged ions are free to move.
98
How is the ionic structure linked to boiling point?
Ionic compounds have high melting and boiling points. It takes a lot of energy to overcome the strong electrostatic forces of attraction.
99
Are ionic substances soluble in water?
Most ionic compounds will dissolve in water, as water molecules are polar and are attracted to the charged ions. An ionic substance only dissolves if more energy is released by separating the bonds then breaking the bonds.
100
What is a covalent bond?
A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
101
What is a dative (coordinate) bond?
A dative bond occurs when both of the electrons in a single bond come from one of the atoms. The bond length and strength is identical to any other covalent bond.
102
What does average bond enthalpy measure?
Average bond enthalpy measures the energy required to break a covalent bond. The stronger a bond = the greater the average bond enthalpy.
103
What is electronegativity
The power of an atom to attract the electrons in a covalent bond.
104
What are the three factors which affect electronegativity?
1. Atomic radius (distance from nucleus.)
2. Shielding
3. Nuclear charge
105
Which element is the most electronegative?
Fluorine - it has an electronegativity of 4.0.
106
What is the link between electronegativity and bonding?
Compounds made up of elements with a large difference in electronegativity will tend to be ionic in character. If they are similar in electronegativity they are likely to be covalent in character.
107
Molecular shape is determined by the electrostatic repulsion of electrons round the central atom. Order these angles, from smallest to largest bond angle:
- lone pair/ lone pair angles
- lone pair/bonding pair angles
- bonding pair/bonding pair angles
Smallest
bonding pair/bonding pair angles
lone pair/bonding pair angles
lone pair/lone pair angles
Biggest
108
A molecule with two bonded pairs has a _____ shape. The bond angle is _______.
Linear
| 180o
109
A molecule with three bonded pairs has a _____ ________ shape. The bond angle is _______.
Trigonal planar
| 120o
110
A molecule with four bonded pairs has a _____ shape. The bond angle is _______.
Tetrahedral
| 109.5o
111
A molecule with five bonded pairs has a _____ _______ shape. The bond angle is _______.
Trigonal bipyramidal
| 90o / 120o
112
A molecule with six bonded pairs has a _____ shape. The bond angle is _______.
Octahedral
| 90o
113
If a molecule contains lone pairs, the angle is based on the number of ______ _____ but the name on the number of ________ ________.
The angle is based on the number of electron pairs.
The name is based on the number of bonding pairs.
114
Each lone pair reduces the bond angle by approximately ____o from the bond angle determined by the number of electron pairs.
2.5o
115
How is a permanent dipole-dipole interaction caused?
In a covalent bond between two atoms of different electronegativities, the bonding electrons are pulled towards the more electronegative atom. This makes the bond polar.
116
Do diatomic gases (i.e. H2) have permanent dipoles?
No - the atoms have equal electronegativities and so the electrons are equally attracted to both nuclei.
117
How is charge arranged in a polar molecule?
The arrangement of polar bonds in a molecule determines whether or not the molecule will have an overall dipole. If charge is arranged unevenly across the molecule, there will be an overall dipole and the molecule will be polar.
118
How is the dipole moment calculated?
Multiplying the size of the charge by the distance between the charge centres.
119
In which atoms and molecules are induced dipole-dipole forces found?
All atoms and molecules
120
How are induced dipole-dipole interactions (London [dispersion]) forces caused?
Electrons in charge clouds are always moving. At any one moment, a temporary dipole is caused as electrons are on one side of the atom. This dipole can cause another temporary (induced) dipole in another atom in a chain. The overall effect is that the atoms are attracted to one another.
121
Do stronger induced dipole-dipole forces increase or decrease boiling point?
They increase boiling point. More energy is needed to overcome stronger intermolecular forces.
122
Which factors increase the strength of induced dipole-dipole forces?
1. Size of electron cloud: larger molecules have larger electron clouds, meaning stronger induced dipole-dipole forces.
2. Surface area: bigger exposed electron cloud = stronger induced dipole-dipole forces
123
Order these forces from strongest to weakest:
- Induced dipole-dipole forces
- Permanent dipole-dipole interactions
- Hydrogen bonding
1. Induced dipole-dipole forces
2. Permanent dipole-dipole interactions
3. Hydrogen bonding
124
Define hydrogen bonding.
Intermolecular bonding between molecules containing fluorine, nitrogen or oxygen and the H atom of -NH, -OH or HF.
125
Why does hydrogen bonding occur?
Hydrogen has a high charge density because it’s so small, and the other elements are very electronegative. The bond is so polarised that a bond forms between the hydrogen of one molecule and a lone pair on the other molecule.
126
Hydrogen bonding _______ boiling points and freezing points
Increases
127
What effect does hydrogen bonding have on water when it freezes?
When ice freezes, in becomes less dense. This is because it contains more hydrogen bonds then liquid water.
128
Explain two properties of simple covalent compounds.
1. Low melting and boiling points: weak intermolecular forces, takes little energy to overcome them
2. Polar molecules are soluble
3. Don't conduct electricity: overall covalent molecules are uncharged
129
How are elements arranged in the periodic table?
1. By increasing atomic (proton) number
2. In periods showing repeating trends in physical and chemical properties (periodicity).
3. In groups with similar chemical properties.
130
What is meant by periodicity?
The study of trends across a period
131
```
What is the name given to the following groups:
Group I (1)
Group II (2)
Group VII (7)
Group 8/0
```
1: Alkali metals
2: Alkaline (earth) metals
7: Halogens
8/0: Noble gases
132
Which groups are in the s block?
I, II and hydrogen
133
Which groups are in the d block?
The transition metals
134
Which groups are in the p block?
III, IV, V, VI, VII, 0
135
Which groups are in the f block?
Lanthanides and actinides
136
Define first ionisation energy.
The energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms
137
Is ionisation and endothermic or exothermic reaction?
Endothermic
138
How would you write an equation for the first ionisation of oxygen?
O (g) --> O+ (g) + e-
The gas state symbol must be used
139
What are the factors affecting ionisation energy?
1. Atomic radius (distance from nucleus.)
2. Shielding
3. Nuclear charge
140
Define second ionisation energy.
The energy required to remove 1 mole of electrons from 1 mole of gaseous singly positive ions to form 1 mole of gaseous +2 ions.
141
Give two reasons why the second I.E. is always greater than the first.
1. The electron is being pulled away from a more positive species.
2. There is less electron shielding as you move down a shell.
142
Why does ionisation energy decrease down a group?
1. Larger atomic radius
| 2. Inner shells provide shielding
143
Why does ionisation energy increase across a period?
1. Nuclear charge increases as the number of protons increases. This decreases the atomic radius and the bonding electrons are more strongly attracted to the nucleus.
2. No extra shielding as the same number of shells are present
144
Group 2 elements form _____ ions.
2+
145
Why does reactivity increase down a group?
Ionisation energy decreases due to increasing atomic radius and shielding. The easier it is to lose electrons, the more reactive the element.
146
What is produced when group 2 elements react with water?
Metal hydroxide + hydrogen
| M(OH)2 + H2
147
What is produced when group 2 elements burn in oxygen?
They form solid white oxides.
| 2M(s) + O2 = 2MO(s)
148
What is produced when group 2 elements react with dilute acid?
Salt + hydrogen
149
The oxides of the group 2 metals can produce metal hydroxides when water is added. The OH- ion makes the solutions strongly alkaline. What is the trend as you go down the group?
The oxides form more strongly alkaline solutions as you go down the group, because the hydroxides get more soluble.
150
Give two uses of group 2 compounds
1. Calcium hydroxide can be used to neutralise acidic fields.
2. Magnesium hydroxide and calcium carbonate can be used to neutralise excess stomach acid - they are antacids.
151
Halogens always exist as ________ molecules.
Diatomic
152
The boiling and melting points ________ down the group due to the ________ strength of the London forces and size and relative mass of the atom.
Increase
| Increasing
153
Define general formula
The simplest algebraic formula of a member of the homologous series
i.e. CnH2n+2
`
154
Define structural formula
The minimal detail that shows the arrangement of atoms in a molecule
i.e. CH3CH2CH2CH3
155
Define displayed formula
Shows the relative positioning of atoms and the number of bonds between them.
H--C(H2)--C(H2)--OH ; with all bonds in brackets shown
156
Define skeletal formula
The simplified organic formula, shown by removing hydrogen from the alkyl chains, leaving just a carbon skeleton and associated functional groups.
157
Define catenation
The ability to form strong covalent bonds to atoms of the same element - i.e. carbon.
158
Define homologous series
A series of organic compounds having the same functional group but with each successive member differing by CH2.
159
Define functional group
A group of atoms responsible for the characteristic reactions of a compound. Molecules with the same functional group are classified into homologous series.
160
Define aliphatic
A compound containing carbon and hydrogen joined together in straight chains, branched chains or non-aromatic rings.
161
Define alicyclic
An aliphatic compound arranged in non-aromatic rings with or without side chains.
162
Define aromatic
A compound containing benzene rings (C6H6).
163
Define saturated
Single carbon-carbon bonds only
164
Define unsaturated
Contains multiple carbon-carbon bonds
165
Alkanes have the ending ____.
- ene
166
Haloalkanes have the suffix's ______, _______, ______ or ________.
- chloro
- bromo
- fluoro
- iodo
167
Alcohols have the ending ____ to replace the last e if one functional group is present. If more than one functional groups are present, it begins with ________.
-ol
| hydroxy-
168
Aldehydes have the ending ____ in the place of the last -e. Aldehydes must always be at the ____ of the chain.
-al
| End
169
Ketones have the ending -one in place of the last -e. If another functional group is present, the functional group goes at the beginning as ____.
Oxo-
170
What is the difference between aldehydes and ketones?
The oxygen is at the end of the chain on an aldehyde, but is in the middle of the chain on a ketone.
171
Carboxylic acids have the ending -oic acid. The _____ functional chain must always be at the end of the chain.
COOH
172
Nitriles have the ending -ile. Does the C of the CN group count as a carbon on the carbon chain?
Yes
173
Amines end in -amine, which becomes ______- if another functional group is present.
Amino-
174
Define structural isomerism
Compounds with the same molecular formula but a different structural formula.
175
Structural isomers have ________ chemical properties.
The same
176
Structural isomers have ________ physical properties.
Differing
177
The greater the degree of branching, the _____ the boiling point. Why?
Lower
| Branching decreases the effectiveness of the intermolecular forces so less energy is needed to separate the molecules.
