Module 2 Flashcards
1
Q
Relative isotopic mass
A
Mass of an isotope relative to 1/12 mass of an atom of carbon 12
2
Q
Relative atomic mass
A
Weighted mean mass of an atom relative to 1/12 of an atom of carbon-12
3
Q
Relative molecular mass
A
simple molecules
4
Q
Relative formula mass
A
Giant ionic compounds
5
Q
Isotopes
A
Atom of same element with different number of neutrons and different mass.
Isoelectronic
Different physical properties
6
Q
Avogadros constant
A
6.02 x 10 ^23 mol-1
7
Q
Isoelectronic
A
Same electron structure
8
Q
Molar mass
A
Mass in g of one mole of a substance
9
Q
mass=
A
Mr moles
10
Q
Number of particles =
A
moles x avagadros constant
11
Q
Empirical formula
A
Simplest whole number ration of atoms of each element present in a compound
12
Q
Water of crystallisation
A
Water that is chemically bonded into a crystalline structure
13
Q
Calculating the formula of a hydrated salt
A
1) moles of anhydrous and H2O
2) calculate ratio of amounts and formulas
13
Q
Check water of crystallisation has been removed by …..
A
Heating to a constant mass
14
Q
Concentration
A
The amount (in mol) of a dissolved substance in 1 dm3 of solution
14
Q
Moles =
A
Concentration x volume
15
Q
Molar gas volume increases as
A
Temperature increases
16
Q
Molar gas volume decreases as
A
Pressure increases
17
Q
Ideal gas equation
A
pV= nRT
T in Kelvin (+273)
V (m3)
R = 8.314 Jmol-1K-1
17
Q
Calculating moles of a gas (RTP)
A
n = vol/24 (dm3)
18
Q
Ideal gas equation Conversions
A
kPa-> Pa X1000
cm3-> m3 X10^-6
19
Q
Stoichiometry
A
Ratio of moles in a chemical reaction
20
Q
Why is a high atom economy good?
A
-efficient
-produce little waste
-less raw materials used
-sustainable
20
Q
Atom economy=
A
(Sum of Mr of desired products/ sum of molar mass of all products) x100
21
Percentage yield=
(Actual/ theoretical) x100
22
acids
Release H+ ions when dissolved in water
Proton donor
22
Strong acids
Fully dissociate when dissolved in water
23
Weak acids
Partially dissociate when dissolved in water
24
Bases
Accept H+ ions
Proton acceptors
25
Examples of bases
-metal oxides
-metal hydroxides
-metal carbonates
-alkalis
26
alkalis
Dissolve in water releasing OH- ions
27
Neutralisation
Reaction between acid and base to produce a salt
28
Neutralisation by carbonate
Acid + carbonate -> salt + water + carbon dioxide
29
Neutralisation by metal oxide
Acid + metal oxide -> salt + water
30
Neutralisation by alkali
Acid + alkali -> salt + water
30
Acid + ammonia
Ammonium salt
31
Standard solution
Solution of known concentration
32
Preparing standard solution of NaOH
-mass (e.g. 1.00 g) weighed out and added to beaker
-dissolve NaOH with distilled water use stirring rod
-pour into 250 cm3 volumetric flask
-rinse beaker with distilled water, wash rinsings into flask
-add distilled water until bottom of meniscus on graduation line
-stopper, invert, thoroughly mix
33
How to find mass needed to prepare standard solution
Find moles (n = cv)
Find mass (m = Mn)
33
Oxidation number ions
Monatomic- oxidation no= charge
33
Carrying out acid-base titration
(Finding conc H2SO4 reacting with 25cm3 NaOH)
-Pipette, add 25.0 cm3 NaOH into conical flask
-white tile, few drops of indicator (phenolphthalein or methyl orange)
-burette, sulfuric acid, initial reading, nearest 0.05 cm3
- add sulfuric acid, swirl
-colour change, final reading
- final- Initial = trial titre
-repeat, dropwise near end point until concordant results (within 0.1)
-mean using concordant results
-calculations
34
Ethanoic acid
CH3COOH
35
Oxidation number elements
0
35
Titration calculation finding concentration
Balance equation
Mol known
Mol unknown
Conc unknown
36
Oxidation no Compounds
Halides -1
H +1
O -2
37
Oxidation no Exceptions
Metal hydride H-1
F2O O +2
Peroxide (H2O2) O -1
38
Oxidation no compounds + polyatomic ions
Sum of oxidation numbers in compound = 0
Sum of oxidation no in polyatomic ion = overall charge
39
How are oxidation numbers shown in elements with variable oxidation numbers?
Roman numerals
E.g. sodium chlorate (I) cl +1
40
Oil rig
Oxidation is loss of electrons
Reduction is gain of electrons
41
Redox in terms of hydrogen
Oxidation is loss of hydrogen
42
Redox in terms of oxygen
Oxidation is gain of oxygen
43
Redox in terms of oxidation no
Oxidation is increase in oxidation number
44
What are the 4 types of orbitals?
S,p,d,f
45
Shape of s orbital
Spherical
46
Shape of p orbital
Dumbell
(3D axis)
47
What is an orbital?
Region around Nucleus where an electron is most likely to be containing 2 electrons with opposite spins
47
Number of electrons in first four shells
2,8,18,32
48
Sub shells
Orbitals within energy level grouped together
49
Metals form ——- ions
Positive (cations)
49
Exceptions- electron structure
Cr and Cu
Less energy to fill 3d and have 1 4s
Or have 5 3d (all occupied singularly) and 1 4s
50
Electron configuration
Electrons fill orbitals in increasing energy
4s lower energy than 3d
Maximum of 2 electrons with opposite spin represents by up and down arrow
Occupied singularly before pairing
51
Non-metals form ——— ions
Negative (anions)
52
Ionic bonding definition
Electrostatic attraction between oppositely charged ions
53
Structure of ionic compounds
Giant ionic lattices, ions fixed in place
54
Properties of ionic compounds
High mp and bp (lots of energy needed to overcome strong ionic bonds)
Soluble in polar solvents (water)
Insoluble in non-polar solvents (hydrocarbons)
55
Strength of ionic bonds
Smaller ions form stronger bonds
Highly charged ions form stronger bonds
56
Conduct because
Moving charged particle (electron, ion)
57
Covalent bond definition
Attraction between positive nuclei and shared pair of electrons
58
Dative covalent bond
Both electrons contributed by one atom
59
Lone pairs
Pair of valence electrons nor bonded to anther atom
Able to form dative covalent bonds with atoms that have vacant orbitals
60
Valence
Outer shell
61
Average bond Enthalpy
Measure of covalent bond strength
Larger= stronger
62
Salt
H+ ion replaced by metal ion or NH4+ ion
63
Displayed formula
Bonds represented by lines
64
Covalent bonding happens between ….
Non-metals
65
Dative covalent bond displayed formula
Arrow
66
Electron deficient
Less than 8 electrons
67
Expanded octet
More than 8 electrons
68
valence shell electron pair repulsion theory
e- repel as far apart as possible
69
How much does each lone pair reduce the angle by??
≈2.5 °
70
LP> LP-BP> BP
repulsion
71
2 bonding regions
2bp, 0lp
linear 180 °
72
3 bonding regions
3bp, olp
trigonal planar 120 °
73
3 bonding regions
2bp, 1lp
non-linear 117.5 °
74
4 electron pairs
4bp, 0lp
tetrahedral 109.5°
75
4 electrons pairs
3bp, 1lp
pyramidal 107°
76
4 electron pairs
2bp, slp
non-linear 104.5°
e.g. H2O
77
5 bonding regions
5bp 0lp
trigonal bipyramidal
120°
90°
77
5 bonding regions
4bp, 1lp
pyramidal 119° 89°
see-saw 119° 89°
78
5 bonding regions
3bp, 2lp
trigonal planar 120 °
T-shaped 89°
79
6 electron pairs
6bp, 0lp
octahedral 90°
80
VSEPR
valance shell electron pair repulsion theory
81
6 electron pairs
4bp, 2lp
square planar 90°
81
6 electron pairs
5bp, 1lp
square pyramid 89°
82
method for determining shapes of elements
eg BrF3
central atom Br
outer e- 7
e- gained from bonds 3
e- gained from change 0
total e- 10
bp 3
lp 2
83
shapes of ions; SO4 2-
2 x S=O
2 x S--O
two atoms with single bond O have extra e- in outer shell resulting in 2- charge
tetrahedral shape 109.5
84
what can be said about the bonding affects of single bonds to double bonds?
they are similar
85
covalent bond H-H
attract electrons equally
electron equally shared
86
electronegativity
the ability of an atom to attract the bonding electrons in a covalent bond
87
trends in electronegativity
F most electronegative
decreases down group
increases across period
88
polar bond
pair of e- shared unequally
δ- slightly negative
δ+ slightly positive
89
dipole
separation of partial charges in a molecule
90
polar molecules
polar bonds (regions with different e- densities)
non-symmetrical
dipoles do not cancel
91
non-polar molecules
if polar bonds: symmetrical, dipoles cancel
92
types of intermolecular forces
induced dipole-dipole interactions (London forces)
permanent dipole-dipole interactions
hydrogen bonds
93
london forces
*e- move randomly, at any instant in time, distribution of e- may be uneven
*this creates weak instantaneous dipoles δδ-
*induces dipole in neighbouring molecules
94
permanent dipole-dipole interactions
2 polar molecules with permanent dipoles attract
95
hydrogen bonds
strong permanent dipole-dipole interactions
H and O/F/N
96
hydrogen bond diagrams
dipoles labelled
h in one molecule bonded to lone pair in another
hydrogen bond represented by dotted line
straight line
97
ice less dense than water
regular structure, open lattice
when ice melts, the open lattice structure collapses allowing H2O molecules to move closer together increasing density
98
water expands when it freezes
tetrahedral structure maximising hydrogen bonding
when ice melts, crystal structure melts as water molecules fall into empty spaces
99
H2O higher melting point than expected
hydrogen bonds
100
Aufbau principle
e- enter lowest energy level available
101
Hund's rule
e- prefer to occupy orbitals on their own and can only pair up when no orbital of the same energy is available (bus analogy)
occupy singularly before pairing
