1
Q
Define an isotope
A
sme protons
same electrons
atoms of same element with same number of protons but different number of neutrons
2
Q
Define Relative Atomic Mass
A
The average mass of an atom of an element compared to 1/12th the mass of Carbon-12 atom
3
Q
Define Relative Isotopic Mass
A
The mass of an atom of a particular isotope of an element relative to 1/12th the mass of Carbon-12 atom
- can be found from mass spec
4
Q
Define Weighted Mean Mass
A
The mean mass taking into account the relative abundances of the isotopes
5
Q
What does the atomic number of an atom tell you?
A
Number of protons
6
Q
What does the mass number of an atom tell you?
A
The number of neutrons and protons
7
Q
How do you calculate the RFM given abundances and relative atomic masses?
A
abundance x masss + …
/
100
8
Q
On a mass spectrometer, what do the peaks indicate?
A
Isotopes
9
Q
What happens to all elements when in a mass spectrometer?
A
Develop a 1+ charge
10
Q
How to calculate the RFM from a mass spectrometer, where the relative abundances don’t equal 100?
A
abundance x mass…. + …
/ total relative abundance
11
Q
What peaks should be shown on mass spectrometers with diatomic elements?
A
All possible combinations of diatomic molecules, and all monotomic atoms
12
Q
How many electrons occupy the first shell?
A
2
13
Q
How many electrons occupy the 2nd shell?
A
8
14
Q
How many electrons occupy the 3rd shell?
A
18
15
Q
How many electrons occupy the 4th shell?
A
32
16
Q
How many orbitals in s subshell?
A
1
17
Q
How many orbitals in p subshell?
A
3
18
Q
How many orbitals in d subshell?
A
5
19
Q
How many electrons in ANY orbital?
A
2
20
Q
Define an orbital
A
A 3D region of space around the nucleus, that can hold up to 2 electrons with opposite spins
21
Q
What is the shape of an s orbital?
A
spherical
22
Q
What is the shape of a p orbital?
A
dumbell - at right angles
23
Q
What is the exception to the rule of subshells filling with increasing energy?
A
4s fills before 3d
24
Q
How do electrons fill sub-shell orbitals as you go along the period?
A
Singularly, before doubling up in each orbital
25
Why is an element in the x block?
As its highest energy electrons occupy the x subshell
26
When writing electronic configurations, how should you write it when the 3d subshell is full?
Write it before 4s
27
What electrons are removed when positive ions form from d block elements?
4s are removed before 3d
28
Electronic configuration for chromium
29
Electronic configuration for copper
30
Define Ionic Bonding
The strong electrostatic attraction between positive and negative ions
31
Define an Ionic Lattice
A repeating pattern of oppositely charged ions
32
When drawing ionic lattices, what must be included in each circle?
Both the formula and the charge
33
Why do ionic compounds have high melting/boiling points?
Giant ionic lattice structure has lots of very strong ionic bonds. Lots of energy is required to break all the strong ionic bonds.
34
What ions are always soluble?
* Na+
* K+
* NH4+
* NO3-
35
State and explain the conductivity of solid ionic compounds
Does not conduct. Ions are fixed in position, not mobile and free to move around giant ionic grid-like lattice structure.
not able to move as mobile charge carries
36
State and explain the conductivity of aqueous and molten ionic compounds
Does conduct. Ions aren’t fixed in position, ions are mobile and free to move around the giant ionic grid-like lattice structure.
37
Solubility: How do many ionic compounds dissolve?
They dissolve in polar solvents, like water. Polar water molecules break down the ionic lattice and surround the ions.
38
What are the two main processes required for an ionic compound to dissolve?
1. The ionic lattice must be broken down. 2. Water molecules must attract and surround the ions.
39
What factor makes an ionic compound less soluble in water?
Large charges on the ions because the ionic attraction may be too strong for water to break down the lattice structure.
40
How does the ionic charge generally affect the strength of the attractions in the giant ionic lattice, and thus the solubility?
As ionic charge increases the attractions in the giant ionic lattice have a greater effect, and solubility decreases
41
Define covalent bonding
The strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms
42
Define a Dative Covalent Bond
A covalent bond where both electrons are donated by one atom
43
dative covalent drawn
44
What is used to measure the strength of a covalent bond?
Average bond enthalpy
45
Why does diamond/silicon have a very high melting/boiling point?
Each carbon atom makes 4 strong covalent bonds with other carbon/silicon atoms. Lots of energy is needed to break all the strong covalent bonds in the giant tetrahedral covalent lattice.
46
Why does graphite/graphene have a very high melting/boiling point?
Each carbon atom makes 3 strong covalent bonds with other carbon atoms. Lots of energy is needed to break all the strong covalent bonds in the giant covalent lattice structure.
47
State and explain the conductivity of diamond/silicon
Cannot conduct. Each carbon/silicon atom makes 4 strong covalent bonds with other carbon/silicon atoms. There are no delocalised electrons, or ions mobile and free to move around the giant tetrahedral covalent lattice structure.
48
State and explain the conductivity of graphite/graphene
Does conduct. Each carbon/silicon atom makes three strong covalent bonds with other carbon atoms. There is one delocalised electron per atom mobile and free to move around the giant covalent lattice structure.
49
Why is graphite soft?
Layers of carbon atoms are held together by weak forces of attraction.
Layers are able to slide over each other.
50
Which giant covalent structures are soluble?
None of them
51
Why do simple molecular substances have a low boiling point?
Molecules held together by weak intermolecular forces. Not a lot of energy is needed to overcome the weak intermolecular forces that act between molecules.
52
State and explain the conductivity of simple molecular substances
None conduct as all the molecules are neutral. As there's no delocalised electrons or ions present.
53
Define Metallic Bonding
Strong electrostatic attraction between a lattice of cations and a sea of negatively charged delocalised electrons
54
Why do metallic compounds have high melting/boiling points?
There are lots of strong metallic bonds in the giant metallic lattice structure. Lots of energy is required to break all the strong metallic bonds.
55
State and explain the conductivity of metals in all states
Conducts in all states as the sea of delocalised electrons are mobile and free to move around the giant metallic lattice structure.
56
When are metallic substances soluble?
Never
57
draw a linear molecule
58
draw a trigonal planar molecule
59
draw a tetrahedral molecule
60
draw an octahedral molecule
61
draw a bent molecule (h20)
62
What are the relative repulsions of lone electron pairs and bonded electron pairs?
Lone pairs repel more than bonded pairs
63
How to write an answer, comparing bond angles?
State the number of bonded and lone pairs of each molecule. Lone pairs repel more strongly than bonded pairs. Bonded electron pairs repel equally.
64
Define electronegativity
The ability of an atom to attract the bonding electrons in a covalent bond
65
What is the most electronegative element?
Fluorine
66
How to determine what element is more electronegative?
Whatever is closest to fluorine in the periodic table
67
In a covalent bond, which atom do the electrons move closer to?
The more electronegative atom
68
How to explain if this molecule is or isn’t polar?
Has polar bonds. Molecule is symmetrical. Dipoles cancel out. Non-polar.
69
How to explain if this molecule is or isn’t polar?
Has polar bonds. Molecule isn’t symmetrical. Dipoles don’t cancel out. Polar.
70
Where do all intermolecular forces act?
Between molecules
71
In what molecules do London forces occur?
In every molecule (NOT giant covalent)
72
Why as molecules get larger, do their boiling points increase?
More electrons. Stronger London forces. More energy required to break London forces.
73
How are London forces induced?
Electrons move randomly in molecule. Creates a temporary dipole in the molecule. Induces temporary dipoles in neighbouring molecules.
74
Describe a permanent dipole force
Polar molecules have dipoles. Dipoles interact to form the dipole-dipole force.
75
Between what molecules does hydrogen bonding occur?
One lone pair of fluorine, oxygen or nitrogen and hydrogen of another molecule
76
What is the relative strength of the different intermolecular forces?
* Hydrogen bonding strongest
* Permanent dipole-dipole forces
* London forces
77
draw a hydrogen bonding diagram for two water molecules
78
What must be included in all hydrogen bonding diagrams?
* Lone pairs
* Dipoles
* Hydrogen bond
79
Describe and explain the anomalous properties of ice - relatively high melting point
Hydrogen bonding is very strong, so lots of energy needed to overcome it
80
Describe and explain the anomalous properties of ice - ice is less dense than water
Water molecules held apart in an open lattice structure by hydrogen bonds
81
State and explain the solubility of non-polar substances in non-polar solvents
Soluble. IMF’s form between molecules in solvent and molecules in solute. Weakens the IMF’s in the simple molecular solvent.
82
State and explain the solubility of polar substances in non-polar solvents
Insoluble.
Attraction between molecules in solvent and ions in ionic lattice not strong enough to break ionic bonds in ionic lattice.
83
State and explain the solubility of polar substances in polar solvents
Soluble. Polar bonds in solute attract polar bonds in solvent.
84
which block does mg belong to
s block
highest energy electrons are in an s sub shell/ orbital
85
CU 2+ ion electron structure
1s2 2s2 2p6 3s2 3p6 3d9
86
Si is
giant covalent
87
p s and cl
simple molecular
88
Different isotopes of antimony have the same chemical properties
same number of electrons in outer shell
89
systematic name for fe2o3
Iron (III) oxide
90
bonding pairs
repel
91
lone pairs
repel more than bonding pairs
92
isotope defintion
same element
same protons and electrons
different neutrons
93
sulfide vs sulfate
ide= s 2-
ate- so4 2-
94
shape around carbon atoms in graphene
trigonal planar
95
which is not a neutration reaction
acid + metal
96
neutralisation reaction
produces water
97
metal oxide + acid
salt and water
98
polar
not symmetrical bonds
99
which substance has London forces in solis state
molecules
- not giant covalent
100
more branching
- decreases Bp
- less surface contact
- weaker London forces
- less energy to break weak London forces
101
lone pairs..
repel more than bonding pairs
102
bent v shape
104.5
103
Si is
giant covalent
104
ive sign
means more electronegative
105
London forces
between moleucles
106
relative atomic mass
the average mass of an atom relative to 1/12 of the mass of an atom of carbon - 12
107
tetrahedral
109.5
4Bp
0LP
108
pyramidal
107
3Bp
1Lp
109
non linear (bent)
104.5
2Bp
2LP
110
trigonal planar
120
3Bp
0LP
111
linear
180
2Bp
0LP
112
octahedral
90 and 180
6Bp
0LP
Chemistry
flashcards
Decks in class (36)
# Cards
-
A1 Atomic structure112
-
Formula17
-
A2 energetics 144
-
A3 inorganics77
-
A4 kinetics 124
-
A5 Equilibrium27
-
A6 kinetics 272
-
acids 232
-
B1 moles, equations59
-
B2 redox and periodicity73
-
B5 Energetics 287
-
A8 - Redox15
-
A9- Transition metals34
-
B3 AlkAnes49
-
AlkEnes46
-
alcohols30
-
B4- haloalkanes23
-
organics 1 synthesis25
-
B4 organic analysis52
-
Benzene59
-
phenol and directing groups43
-
Carbonyl16
-
carboxylic acids9
-
esters6
-
acyl chlorides and acid anhydrides10
-
amines15
-
amino acids and amides8
-
optical isomers7
-
polymers7
-
organic practicals31
-
tests14
-
reaction pathway43
-
HARD31
