foundation Flashcards
1
Q
Relative isotopic mass
A
The mass of iostope relative to 1/12 of carbon-12 atom
2
Q
Relative atomic mass
A
The weighted mean mass of an atom relative to 1/12 of a carbon-12 atom
3
Q
Steps of mass spectroscopy
A
- sample is vapouried
- sample is ionised to 1+ charge
- ions are accelerated by a magnetic field
4
Q
Aim of mass spectroscopy
A
To find the abundancy of each isotope
5
Q
What does a mass spectrometry graph show?
A
Each peak represents the abundancy of a isotope.
6
Q
Isotope
A
Same no. of proton but different no. of neturons
7
Q
Charge of NO2 ion
A
1-
8
Q
Charge of NO3 ion
A
1-
9
Q
Charge of HCO3 ion
A
1-
10
Q
Charge of MnO4 ion
A
1-
11
Q
Charge of Ag ion
A
1+
12
Q
Charge of SO3 ion
A
2-
13
Q
charge of CrO7 ion
A
2-
14
Q
Charge of Zn ion
A
2+
15
Q
Charge of NH4 ion
A
1+
16
Q
Charge of PO4 ion
A
3-
17
Q
What does avogradro constant a value of?
A
The no. of atoms / molecules per moles
18
Q
Equation of amount in gas (Room temp. and pressure)
A
volume (dm3) = moles * 24
19
Q
Ideal gas equation
A
pv = nGT
20
Q
Precentage yield eq.
A
actual/theoritical * 100 = %
21
Q
Atom economy eq.
A
useful/total * 100 = %
22
Q
Benefits of high atom economy
A
↓waste produced : ↑ product formed
Greater profit margin (as cost is reduced)
23
Q
What may cause actual yield < theoritical yield?
A
- side reaction
- incompletion of reaction
- product loss during purification
24
Q
dm3 –> cm3 conversion
A
x 1000
25
K to celcius
K = C + 273
26
Difference between strong and weak acids
Strong acid fully disociates
Weak acid partially disociates
[in solution]
27
Examples of strong acids
1. HCl
2. HNO3
3. H2SO4
28
Examples of weak acids
Any carboxylic acids E.g. HCOOH, H3CCOOH
29
Products of acid and metal reaction
salt + H2
30
Products of acid and base
salt and H20
31
Example of base
Metal oxides, hydroxides, carbonate(have special reactions)
32
Products of Acid and carbonate
salt + H2O + CO2
33
Observation of strong acid reaction
1. effervescence
2. metal dissolves quicker
3. ↑ exothermic
34
Volumetric analyisis
The use of volume of conc. of a standardised solution to determine the conc. of unknown solution.
35
Steps to making standardised solution
1. weigh out precise amount of solid
2. Dissolve solid to small volume of H2O
3. Transfer it to a volumetric flask
4. Add water until the bottom of meniscus is at the scratch mark
36
Use of burrette
To deliver precise volume of solution (WITH known conc.) until point of neutralisation
37
Steps of titration
1. meausre out a known volume of the solution with unknown conc. and transfer to conical flask.
2. Add indicator to the conical flask
3. Add the known solution to burrette
4. turn the burrette on and slow down as it approaches the end point.
5. Take repeats until concordent
38
Define end point
The point where the indicator show the amount of solution required to react completely has been added.
39
Define concordant
when results are 0.1 cm3 apart
40
Resolution of result in titration
0.05 cm3
41
Ox. state of any element
0
42
Ox. state of Group 1 + 2
Group 1 = +1
Group 2 = +2
43
Ox. state of ions
same as their charge
44
Sum of ox. state of any compound
0
45
Sum of ox. state of any polyatomic ion
their charge
46
Rules for assigning oxidisation state
greater electronegativity gets first pick in ox. no.
47
What does roman numerials represents ?
the oxidation state of atom in ion
48
exception in ox no. for hydrogen
-1 for hydrogen in hydrides
49
2 exception in ox. no. for oxygen
+2 in F2O
-1 in H2O2
50
Define disproportionation reaction
Where a species is both oxidised and reduced
51
2 E.g. of disporportionation reaction
1. Chorine with water
2. Chorine with sodium hydroxide
52
4 types of subshells ( and the no. of orbital it holds)
s - (1 orbital)
p - (3 orbital)
d - (5 orbital)
f - (7 orbital)
53
no. of electron each orbital can hold
2 electrons
54
What does electron from the same orbital must have
opposite spin
55
shape of s and p orbital
s - sphere
p - dumbell
56
Ground state
most stable electronic configuration with least energy
57
Order of filling orbital
1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,4f
58
3 Rules for electronic configuration
1. fill from the lowest energy first
2. 4s is filled before 3d
3. else 3d must be half-filled OR fully filled
59
2 common exception for electronic config.
Chromium and Copper(for not filling 4s completely)
60
electronic config of Chromium (24e)
1s[2],2s[2],2p[6],3s[2],3p[6],4s[1],3d[5]
(1/2 filled 3d and 1 in 4s)
61
electronic config of Copper (29e)
1s[2],2s[2],2p[6],3s[2],3p[6],4s[1],3d[10](fully filled 3d and 1 in 4s)
62
Ionic bonding
The electostatic attraction between the metal cation and the anions
63
Covalent bonding
The electrostatic attaction between the positive nuclei and the shared pair of electron
64
Structure of ionic bonding
a strong solid regular, repeating crystalline structure, formed by the oppositvely-charged ions strongly attracting in all directly.
65
explain ionic structure's high BP and MP
1. no intermolecular forces present
2. the electrostatic attraction is very strong
3. require a lot of energy to overcome
66
explain the conductivity of ionic structures
1. can't conduct in solid lattice
(ions are in fixed position)
2. conducts when molten and in solution
(regular lattice is distrupted and the ions are mobile)
67
Solubility of ionic structure
Dissolves in any polared solvent (i.e. water) as the polared solvent attracts the ions in the lattice to disrupt the regular lattice.
68
A measurement of the strength of covalent bond
Average bond enphalpy
69
Define dative covalent bonds
when a lone pair of electrons from one atom is shared with an electon deficient atom.
70
Three giant covalent structure
1. Diamond
2. Graphite
3. Silica (Silicon dioxide)
71
how many carbons is each carbon bonded to in diamond and graphite [and what does that suggest]
4 atoms in diamond 3 atoms in graphite
1. there is one delocalised e- per Carbon
2. as it's not occupied in a bond
3. hence can conduct.
72
Properties of graphite
1. conduct charge
2. slippery - as weak London forces between layers
3. high MP and BP
73
structure of SiO2
giant covalent lattice
74
shape of molecule
2 bonds and NO lone pairs (with bond angle)
linear (180)
75
shape of molecule with
2 bonds and 1 lone pair
(with bond angle)
non-linear / bent (104.5)
76
shape of molecule
3 bonds and NO lone pair
(with bond angle)
trigonal planar (120)
77
shape of molecule
3 bonds and 1 lone pair
(with bond angle)
pyramidal (107)
78
shape of molecule
4 bonds
(with bond angle)
tetrahedral (109.5)
79
shape of molecule
5 bonds
(with bond angle)
trigonal bipyramidal (120 or 90)
80
shape of molecule
6 bonds
(with bond angle)
octahedral (90)
81
Summarise the valance shell electron repulsion theory (VSEPRT)
Given electron have like charge and repulse each other, lone pair and bonds repel each other and itself.
The order of repulsion strength goes as
1. lone pair - lone pair
2. lone pair - bond pair
3. bond pair - bond pair
82
Electronegativity
The power of an atom to attract the pair of electrons in a covalent bond
83
Nuclear charge
Attraction between the nucleus and electrons
Greater nuclear charge = more electronegative
84
How does electronegativity increase
highhest = top-right
↑ group = ↑ electronegative
↑period = ↓ electronegative
85
Pauling scale
A numerical scale of electronegativity based on bond energies
86
How does distance from the nucleus affect electronegative
Greater distance = lower attraction
87
What affects the bond polarity
Greater the difference in electronegativity between the two atoms, more polared the bond would be
88
How can a molecule be non-polared
When there is no net dipole moment, either when:
1. the electronegativity of the atoms are same
2. dipole moments cancel out by geometry
89
3 types of intermolecular forces
1. temporary dipole-dipole forces
2. Permanent dipole-dipole forces
3. hydrogen bonding
90
What is the general name for dipole-dipole forces
london forces
91
What is the difference between permanent and temporary d-d forces
Permanent - only occurs with polared
Temporary - occurs between any molecules
92
Steps to temporary dipole-dipole forces [6]
1. electron clouds in non-polared molecules are constantly moving randomly
2. when electron clouds are more on one side
3. forms a instantaneous dipole
4. Inducing a dipole on neighbouring molecules
5. so molecules to attract
6. As the clouds are constantly moving, dipole will disappear
93
How to ↑ temporary d-d forces
↑ no of electrons
94
Steps to permanent D-D forces
1. the molecule is naturally have one end more electronegative than the other
2. which causes δ+ end attract δ- end
3. As the dipole is permanent, so does the attraction
95
The three species that bonds with hydrogen molecules AND will form 'hydrogen bond'
Oxygen, Nitrogen and Flourine
96
Two anomulous properties of water
1. Ice is denser to water
2. water's usually high MP and BP
97
Explain water's unusually high boiling and melting point
1. Oxygen is a species that qualifies for hydrogen bonding
2. The strongest type of London forces
3. Requires the lots of energy to overcome
98
Explain why ice floats on water?
1. In ice, the water molecules are held together by hydrogen bonds
2. hydrogen bonds has a slightly longer bond length
3. and molecules are further apart in this lattice than in liquid
99
Three types of covalent structure
non-polared, polared, giant
100
the boiling point of the covalent structure
simple molecular - low
giant - high
101
why does boiling point vary between covalent?
BP(Simple molecular) < BP(gaint covalent)
London forces require ↓ energy to overcome than covalent bonds
102
Solubility of ionic, non-polared, polared, giant ( in both polared and non-polared)
non-polared - non-polared
polared - polared
giant - None
ionic - polared
103
What is the solubility rules of covalent simple molecules
like dissolves like
104
Why does ionic substance dissolves in polared solvent?
polared molecules can used their dipoles (δ+and δ-) to disrupt the regular lattice
surrounding cations and anions.
105
2 states ionic substance conducts in and WHY ?
When molten / in solution
as the ions are not in fixed position of the lattice
