DEFINITIONS Flashcards
1
Q
Relative Isotopic Mass
A
the mass of an atom of an isotope compared with 1/12th of the mass of an atom of carbon-12.
2
Q
Relative Atomic Mass
A
is the weighted mean mass of an atom of an element compared with 1/12th of the mass of an atom of carbon-12.
3
Q
general formula
A
the simplest algebraic formula of a member of a homologous series
4
Q
Empirical Formula
A
the simplest whole number ratio of atoms of each element present in a compound
5
Q
molecular formula
A
actual number of atoms of each element in a molecule
6
Q
structural formula
A
minimal detail that shows the arrangement of atoms in a molecule in space
7
Q
Homologous Series
A
a series of organic compounds containing the same functional group
with successive members differing by -CH2
.
8
Q
carbocation
A
are positively charged carbon atoms with only three covalent bonds instead of four
- primary - least stable - carbon with + charge attached to only one alkyl group
- secondary
- tertiary - most stable
9
Q
radical
A
a species with an unpaired electron
10
Q
orbital
A
region that can hold up to two electrons, with opposite spins
11
Q
Ionic Bonding
A
an electrostatic attraction between positive and negative ions
12
Q
Covalent bonding
A
the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
13
Q
Dative covalent bond
A
a shared pair of electrons in which the bonded pair has been provided by one of the bonding atoms only; also called a co-ordinate bond.
14
Q
Electronegativity
A
the attraction of a bonded atom for the pair of electrons in a covalent bond.
15
Q
First ionisation energy
A
the energy required to remove one electron from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
16
Q
Metallic bonding
A
strong electrostatic attraction between cations and delocalised electrons.
17
Q
Disproportionation
A
is a redox reaction in which the same element is both oxidised and reduced.
18
Q
Hess’s Law
A
the enthalpy change for a reaction that is independent of the route taken.
19
Q
heterogenous catalyst
A
the catalyst is in a different phase to the reactants
20
Q
homogenous catalyst
A
the catalyst is in the same phase as the reactants
21
Q
Adsorption
A
where something sticks to a surface
- One or more of the reactants are adsorbed on to the surface of the catalyst at active sites.
22
Q
electrophile
A
electron pair acceptor
23
Q
nucleophile
A
electron pair donor
24
Q
Structural isomers
A
compounds that have the same molecular formula but different structural formulae
25
Stereoisomers
compounds with the same structural formula but a different arrangement in space.
26
Brønsted Lowry base
species that accepts a proton
27
Brønsted Lowry acid
species that donates a proton
28
Buffer solution
a system that minimises pH changes when small amounts of an acid or a base are added
29
Pi bond
π
A bond formed by the sideways overlap of p orbitals with the electron density above and below the sigma bond.
30
Sigma bond
σ
A bond formed by the overlap of each bonding atom consisting of two electrons and with the electron density centred around a line directly between the nuclei of the two atoms.
31
E/Z isomerism
a type of stereoisomerism caused by the restricted rotation of π bonds. Two
different groups must be attached to each carbon atom of the C=C group
E = opposite so like \ or /
Z = same like _ or -
32
cis-trans isomerism
type of E/Z isomerism in which the two substituent groups attached
to the cabon atoms are the same
cis = opposite so / or \
trans = same so _ or -
33
average bond enthalpy
Breaking of one mole of bonds
In gaseous molecules
34
Enthalpy change of combustion
enthalpy change when one mole of a substance reacts completely with oxygen under standard conditions and standard states
35
Enthalpy change of neutralisation
nthalpy change that accompanies the reaction of an acid by a base to form one mole of water under standard conditions and standard states
36
Enthalpy change of formation
enthalpy change when one mole of a compound is formed from its element under standard conditions and standard states
37
Enthalpy change of solution
the enthalpy change when one mole of a solute dissolves in a solvent under standard conditions.
38
Enthalpy change of hydration
the enthalpy change when the separate gaseous ions interact with polar water molecules to form hydrated aqueous ions
39
Lattice enthalpy
enthalpy change that accompanies the formation of one mole of an ionic compound from its gaseous ions under standard conditions
40
Entropy
measure of the dispersal of energy in a system which is greater, the more disordered a system.
41
Spontaneous
energetically feasible (i.e. whether the reaction is able to happen i.e. delta G < 0)
42
Oxidising agent
takes electrons from the species being oxidised, it contains the species being reduced.
43
Reducing agent
adds electrons to the species being reduced, the reducing agent contains the species that is being oxidised
44
Polar bond
covalent bond in which there is an unequal share of the electrons between the 2 atoms due to the differing electronegativities of the atoms involved. One atom will have a partial positive charge while the other will have a partial negative charge.
45
Aromatic
organic compound with benzene ring
46
Dipole
partial charge on an atom which is caused by the differing electronegativities of
atoms in a covalent bond.
47
Alicyclic
an aliphatic compound that is arranged in non-aromatic rings (with or without side
chains)
48
Aliphatic
a compound containing carbon and hydrogen atoms joined in straight or branched
chains or in non-aromatic rings.
49
Heterolytic Fission
when a covalent bond breaks, one bonding atom receives both electrons from the bonded pair
- different products +ve ion and one -ve ion
50
homolytic fission
each of the covalently bonded atoms takes one of the shared electrons, forms 2 radicals
- same products
51
Le Chatelier’s Principal
if a system at equilibrium is subjected to a small change, the equilibrium tends to shift so as to minimise the effect of the change
52
chiral centre
an atom which is bonded to four different substituents / groups
53
optical isomers
- octahedral
- bidentate ligand
- non superimosable mirror images
54
Test for Alkene
Use bromine water
+ = decolourises brown to clear duh
55
Test for halide ions
use AgNO3 and ethanol
positive test:
colourless (no ppt) - F-
white ppt = Cl-
Cream ppt = Br-
Yellow ppt = I-
children will cry yearly
56
Test for Carbonyl
| C=O and 2 R group of C
Use Brady's reagent
+ = yellow/orange ppt
57
Test for Aldehyde and ketone
| aka fehlings solution
add 2-4 dinitrophenylhydrazine
+ = blue solution to red ppt shows aldehyde or ketone present C=O present
58
Test for Aldehyde
| H - C = O and R group of C
use Tollens reagent
+ = silver mirror
- = ketone has no change
59
Test for Carboxylic acid
| HO- C = O and R group of C
Use sodium carbonate and test gas for limewater
+ = bubbling and white ppt formed with limewater
60
Test for Alcohol
Use potassium dichromate
+ = orange to green IF primary alcohol
- = no change as tertiary alcohol
61
Test for ammonium ions
use sodium hydroxide solution dropwise and gently heat then use red litmus paper
+ = red litmus turns blue
62
Test for transition metal ions
use aq ammonia or sodium hydroxide dropwise
Cu(II) = turns blue ppt
Fe(II) or Cr(III) = turns green ppt
Mn(II) = turns brown ppt
63
Test for sulfate ions
add dilute HCl then barium chloride
+ = white ppt formed
64
Test for carbonate
| do this FIRST
add dilute HNO3 and + test is bubbling
65
NO radicacls catalyse breakdown of ozone
Propogation step 1:
NO● + O₃ ➜ NO₂● + O₂
Propogation step 2:
NO₂● + O ➜ NO + O₂
66
K<1
indicates that there are more reactants than products at equilibrium
67
k>1
indicates that there are more products than reactants at equilibrium
68
K=1
neither reactants nor products are favored
the concentration of reactants and of products are equal at equilibrium
69
bidentate ligand
2 lone pairs are donated to form 2 dative covalent bonds
70
coordination number
the number of dative covalent bonds formed between ligands and a metal ion centre
71
Why is Sc and Zn not transition elements?
- transition elelements have an ion with an partially-filled d sub-shell
- so Sc and Zn are not transition elements
- Sc3+ = 1s2 2s2 2p6 3s2 3p6
Zn2+ = 1s2 2s2 2p6 3s2 3p6 3d10
- Sc3+ d sub-shell empty
- Zn2+ d sub-shell full
72
Optical isomers
- octahedral
- bidentate ligands
- non superimposable mirror images
73
How does lig sub allow haemoglibin to transport oxygen in blood?
O2/oxygen bonds to Fe2+/Fe(II)
(When required,) O2 substituted OR O2 released
74
In the presence of carbon monoxide, less oxygen is transported in the blood.
Suggest why, in terms of bond strength and stability constants.
- Stability constant value with CO is greater than with
complex in O2
- CO has greater affinity for ion/metal/haemoglobin
75
Config of Nickel
1s2 2s2 2p6 3s2 3p6 3d8 4s2
76
config of chromium
1s2 2s2 2p6 3s2 3p6 3d5 4s1
77
config of copper
1s2 2s2 2p6 3s2 3p6 3d10 4s1
78
# disproportionation reaction
Chlorine and treating water
Cl2 + H2O --> HCl + HClO
| Cl2 = 0 and HCL = -1 and HClO = +1
79
# disproportionation reaction
Chlorine and cold dilute AQ NaOH
| BLEACH!!!
Cl2 + 2NaOH --> NaCl + NaClO + H2O
ionic= Cl2 + 2OH- --> Cl- + ClO- + H2O
| cl2 = 0 and cl- = -1 clo- = +1
80
At room temperature and pressure, the first four members of the alcohol homologous series are liquids whereas the first four members of the alkanes homologous series are
gases.
- Alkanes have london forces
- Alcohols have hydrogen bonds and london
- Hydrogen bonds are stronger
81
A proposed mechanism for this reaction takes place in several steps. Suggest two reasons why it is unlikely that this reaction could take place in one ste
- rate eq does not match in overall eq
- collision unlikely to happen with more than 2 species
82
A small amount of aqueous ammonia, NH3(aq), is added to the buffer solution. Explain, in terms of equilibrium, how the buffer solution would respond to the added NH3(aq).
- NH3 reacts with H+ of acid forming NH4+ and so EQ shifts in opposite direction to form more H+
83
Properties of a transition element
- forms voloured compounds
- has diff oxidation states
- elements and compounds used as catalysts
- element forms AT LEAST 1 stable ion with a partially filled d sub shell
84
[Cu(H₂O)₆]²⁺
Pale blue solution
85
[Cu(H₂O)₆]²⁺ + weak dilute NaOH
Pale blue ppt
## Footnote
[Cu(H₂O)₆]²⁺ + 2OH- -->[Cu(H2O)4(OH)2] + 2H2O
86
[Cu(H₂O)₆]²⁺ + weak NH₃
Pale blue ppt
## Footnote
[Cu(H₂O)₆]²⁺ + 2NH- -->[Cu(H2O)4(OH)2] + 2NH4+
87
[Cu(H₂O)₆]²⁺ + excess NH₃
Dark blue solution
## Footnote
[Cu(H₂O)₆]²⁺ + 4NH3 -->[Cu(NH3)4(H2O)2]2+ + 4H2O
88
Ammonia acts as what before a ligand
BASE
89
[Cu(H₂O)₆]²⁺ + conc HCl
Yellow Solution
| [Cu(H₂O)₆]²⁺ + 4CL-- -->[CuCl4]2- + 6H2O
90
[Fe(H₂O)₆]³⁺
Yellow solution
91
[Fe(H₂O)₆]³⁺ + weak dilute NaOH
Brown ppt
| [Fe(H₂O)₆]³⁺ + 3OH- --> [Fe(H2O)3(OH)3] +3H2O
92
[Fe(H₂O)₆]³⁺ + weak NH₃
Brown ppt
| [Fe(H₂O)₆]³⁺ + 3NH3 --> [Fe(H2O)3(OH)3] +3NH4+
93
[Fe(H₂O)₆]²⁺
Pale Green solution
94
[Fe(H₂O)₆]²⁺ + weak dilute NaOH
Dark green ppt
| [Fe(H₂O)₆]²⁺ +2OH- --> [Fe(H2O)4(OH)2] +2H2O
95
[Fe(H₂O)₆]²⁺ + weak NH₃
Dark green ppt
| [Fe(H₂O)₆]²⁺ +2NH3- --> [Fe(H2O)4(OH)2] +2NH4+
96
[Mn(H₂O)₆]²⁺
Pale pink solution
97
[Mn(H₂O)₆]²⁺ + weak dilute NaOH
Pale brown ppt
| [Mn(H₂O)₆]²⁺ +2OH- --> [Mn(H2O)4(OH)2] + 2H2O
98
[Mn(H₂O)₆]²⁺ + weak NH₃
Pale brown ppt
| [Mn(H₂O)₆]²⁺ +2NH3 --> [Mn(H2O)4(OH)2] + 2NH4+
99
[Cr(H₂O)₆]³⁺
Violet solution
100
[Cr(H₂O)₆]³⁺ + weak dilute NaOH
Dark green ppt
| [Cr(H₂O)₆]³⁺ + 3OH- --> [Cr(H2O)3(OH)3] + 3H2O
101
[Cr(H₂O)₆]³⁺ + excess NaOH
Dark green solution
| [Cr(H₂O)₆]³⁺ + 6OH- --> [Cr(OH)6]3- + 6H2O
102
[Cr(H₂O)₆]³⁺ + weak NH₃
Dark green ppt
| [Cr(H₂O)₆]³⁺ + 3NH3 --> [Cr(H2O)3(OH)3] + 3NH4+
103
[Cr(H₂O)₆]³⁺ + excess NH₃
Purple solution
| [Cr(H₂O)₆]³⁺ + 6NH3 --> [Cr(NH3)6]3+ + 6H2O
104
Ti
+3 = purple
105
V
+2 = purple
+3 = green
+4 = blue
+5 = yellow
106
Cr
+2 = blue
+3 = green
+6 = orange
107
Mn
+2 = pink
+4 = dark pink
+6 = green
+7 = purple
108
Co
+3 = green
109
Ni
+2 = green
110
does pH of a buffer solution change
- pH stays same
- ratio of [HA]/[A-] is same
111
standard conditions for electrode potential
298k or 25 degrees
1mol dm^-3
112
which equation do u flip when making an overall eq for electrode potential eq
THE LESS POSITIVE ONE XOXO
113
Strongest reducing agent
- the one being oxidised
- most negative cell potential
114
Healthy human blood needs to be maintained at a pH of 7.40 for the body to function normally.
Carbonic acid, H2CO3, is a weak acid which, together with hydrogencarbonate ions, HCO3 acts as a buffer to maintain the pH of blood.
The pKa value for the dissociation of carbonic acid is 6.38.
Explain, in terms of equilibrium, how the carbonic acid–hydrogencarbonate mixture acts as a
buffer in the control of blood pH, and calculate the [HCO3−] : [H2CO3] ratio in healthy blood
| 6 marks
H2CO3(aq) ⇌ H+(aq) + HCO3-(aq)
* Addition of H+ causes ⇌ to shift to left
* Addition of OH– causes ⇌ to shift to right
When increasing H+ it will react with HCO3-
When increasing OH- it will react with H2CO3
work out ratio !!
115
Red blood cells contain haemoglobin.
Explain using ligand substitutions:
* how haemoglobin transports oxygen around the body
* why carbon monoxide is toxic.
| 3 marks
- O2 binds with Fe2+ in haemoglobin
- it is replaced y H2O or CO2 when required
- CO forms a stronger biond than O2 so CO toxic
