Foundations in chemistry Flashcards
1
Q
Relative isotopic mass
A
The mass of an atom of an isotope in relation to 1/12th of the mass of carbon-12
2
Q
Relative atomic mass
A
The weighted mass of an atom of an element compared to 1/12th of the mass of carbon-12.
3
Q
How can the percentage abundance of an isotope be found? What stage?
A
Using a mass spectrometer,
- ionisation
- accelerator
- deflection
- detection
relative abundances are recorded as peaks on a graph and percentage can be found from these results.
x-axis is atomic mass over charge (m/z) however on most ions in this case the charge is +1 so the x-axis scale is equal to atomic mass.
4
Q
1st ionisation energy
A
Energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
5
Q
factors that affect ionisation energy.
A
- atomic radius (larger atomic radius = smaller ionisation energy)
- atomic charge/nuclear attraction (larger atomic charge=greater attraction of electrons = greater ionisation energy)
- electron shielding (inner electrons shield outer electrons from the positive attraction of the nucleus)
6
Q
Explain the increase in ionisation energies from 1st–>2nd->3rd
Why are some differences much greater than others?
A
The second electron is being removed from a species with a positive charge. The electrons are pulled closer to the nucleus as more electrons are removed.
Ca –> Ca+ + e-
Ca+ –> Ca2+ + e-
etc.
often there is a large jump between particular ionisation energies in elements. For example in group 2 elements the jump between 2nd and 3rd ionisation energies is larger due to the 3rd electron being in a different shell closer to the nucleus. Atomic radius and electron shielding are smaller values and therefore ionisation energies are larger.
7
Q
Photo electron spectrum?
Why does peak high not matter?
A
graph demonstrates the energy required to remove electrons from different energy levels in a atom.
Peak height does not matter as it is relative to the highest peak and it is not an absolute value.
8
Q
How many orbitals and electrons in each type of subshell?
A
s-subshell = 1 obital containing two electrons (groups 1/2) p-subshell = 3 orbitals each containing 2 electrons therefore with a maximum of 6 electrons. (groups 3-8) d-subshell = 5 orbitals, maximum of 10 electrons (transition metals)
9
Q
order of energy levels (subshells)?
A
1s2s2p3s3p4s3d4p
electrons are always lost from orbital furthest from the nucleus but do not always fill in this order.
10
Q
Oxidation
A
Is the loss of electrons
Increase in oxidisation number
Mg –> mg^2+ + 2e-
11
Q
Reduction
A
Is the gain of electrons
decrease in oxidisation number
Cl2 + 2e- —> 2Cl-
12
Q
Oxidising agent
A
Is a reagent that oxidises (takes electrons from) another species
13
Q
Reducing agent
A
Is a reagent that reduces (gives electrons to) another species
14
Q
Redox
A
Redox reactions are reactions in which both oxidation and reduction happen
15
Q
Oxidation number rules
A
- Uncombined elements have an oxidation number of 0
- group I metals have and oxidation number of +1
- group II metals have an oxidation number of +2
- Flourine toms have an oxidation number of -1
- Oxygen atoms have an oxidation number of -2 except in peroxides when it is -2
- hydrogen has an oxidation number of +1 except in hydrides wen it is - 1
16
Q
Hydrochloric acid
A
HCl
17
Q
Sulfuric acid
A
H2SO4
18
Q
Nitric acid
A
HNO3
19
Q
Phosphoric acid
A
H3PO4
20
Q
ethanoic acid
A
CH3COOH
21
Q
Acid definition
A
An acid is a proton donor, and release H+ ions in aqueous solution, when dissociated
22
Q
weak acids
A
Acid that do not dissociate completely, at any one time some of the aid molecules have split into H+ and an appropriate anion and some are still associated.
The weaker the acid is the lower proportion of the acid is dissociated, the less h+ ions are in solution .
normally carboxylic acids
23
Q
Strong acids
A
Are always fully dissociated in solution.
Both weak and strong acids can be diluted, but it is the percentage of acid that is dissociated that defines weak or strong acids
normally mineral acids
24
Q
Bases
A
Bases are proton acceptors
25
Sodium hydroxide
NaOH b
26
Ammonia
NH3 b
27
Potassium oxide
K2O b
28
Calcium Hydroxide
Ca(OH)2 b
29
Copper oxide
CuO b
30
Water as an acid or base
as a weak acid:
H2O H+ + OH- (H+ ion in aqueous solution)
as a base
H+ + H2O -----> H3O+ (proton acceptar)
31
Alkali
A soluble base that release OH- ions in aqueous solution
eg. KOH Potassium hydroxide. white solid
32
Ammonium
NH4+
33
Ammonia as a alkali
NH3 + H2O ----> NH4+ + OH-
| reacts with water to form OH- ions and therefor is an alkali
34
Alkali and base relationship
All alkalis are bases but not all bases are alkali.
35
Salts
Salts are an ionic compound formed when H+ ions in acids are replaced by metal or ammonium ions
36
Carbonate ion
CO3 2- H2CO3
37
Silver ion
Ag+
38
Hydrogen carbonate ion
HCO3 -
39
Reactions of acids
Acid + Metal hydroxide ----> salt + water
Acid + Metal Oxide -----> salt + water
Acid + metal carbonate -----> carbon dioxide + salt + water
Acid + Reactive metal -----> salt + hydrogen
The acid dictates what salt is formed
40
Cu(NO3)2 (aq) Observations
Green clear solution
41
Ionic compound properties
- Hard crystalline structure
- High melting points and solid at room temp (strong electrostatic forces require large amounts of energy to break)
- Soluble in water (polarity of water pulls molecules apart)
42
Ions that form soluble compounds
group 1
ammonium
nitrate
hydrogen carbonate
halides except when combined with Ag+ Pb2+ Hg2+
Sulfate except when combined with Ag+ Pb2+ Ca2+ Sr2+ Ba2+
43
Ions that form insoluble compounds
Oxide except with group 1 or Ca/Ba 2+
Carbonate except group 1 or ammonium
hydroxide except group 1 ammonium or Ba/Ca/Mg/Sr 2+
sulfide except group 1 or ammonium
Chromate Cr2O7 2- except group 1 ammonium or Ca/Mg 2+
44
silver chloride
white precipitate
| soluble in dilute and conc ammonia solution
45
silver bromide
cream precipitate
| soluble in conc ammonia
46
silver iodide
yellow precipitate
| not soluble in dilute or conc ammonia
47
test for CO3 2-
Add nitric acid
| effervescence and gas produced turns lime water cloudy when bubbled through
48
test for SO4 2-
Add Ba+ ions and a white precipitate formed
49
test for Cl-
add aq silver nitrate (Ag+ ions) white recipitate forms and dissolves in dilute ammonia solution
50
test for Br-
add aq silver nitrate (Ag+ ions) cream precipitate forms and dissolves in conc ammonia solution
51
test for I-
add aq silver nitrate (Ag+ ions) yellow precipitate forms and does not dissolves in dilute or conc ammonia solution
52
Test for NH4+
add sodium hydroxide solution and warm mixture
| strong smell of ammonia released. and damp red litmus turns blue
53
ionic bonding defnition
Electrostatic attraction between positive and negative ions
| containing a metal (which forms positive ion) and a non-metal (negative ion)
54
covelant bonding definition
A strong electrostatic attraction between a shared pair of negative electrons (bonded pair) and the positive nuclei of the bonded atoms
called a MOLECULE
55
A lone pair
A pair of electrons in the outer shell not used in bonding
56
Dative covalent bond
Is a bond where the shared pair of electrons has been provided by only one of the two bonded atoms.
57
average bond enthalpy
Is the average enthalpy change which takes place when breaking, by homolytic fission, one mole of a given type of bond in the molecules of a gaseous species
stronger bonds have higher bond enthalpies
58
metals structure type / explaination
giant continuous structure
because of metals relativly low ionisation energies electrons in valance or outer most shells are held losely, in metallic solids each atom loses at least 1 electron forming a sea of delocalised negative electrons which are attracted to the now positive metal ions and hold the structure together
59
Metallic bonding definition
The electrostatic attraction between positive metal ions and delocalised negative electrons
strong electrostatic attraction results in high melting and boiling point
60
why are metals always good conductors
delocalised electrons are free to move and can carry charge in metal solids
when molten, both positive ions and delocalised electrons are free to move and can be carriers of charge
61
Why are metals insoluble in water or polar solvents
some metals react to make soluble products however in most the electrostatic attraction is strong and requires alot of energy to break
62
why do group 1 have lower melting points than group 3 in the same period?
More electrons are donated to create 1+ /2+ /3+ ions depending on the number of electrons in the valent shell
63
malleable/ ductile definition and reason
malleable- easy to bend into different shapes
ductile- easy to draw into a wire
metals have these properties as layers of electrons easily slide over each other
64
alloys why?
They are used to modify metals properties, become less malleable/ductile due to different sized molecules so layers can't slide as easily
65
why do gaint ionic lattices exist? What are the properties?
continuous lattice arrangement of + and - ions in all 3 dimensions
HIgh melting/boiling points due to strong electrostatic attraction which requires a lot of energy to overcome
Non-conductor as a solid - ions are fixed in a lattice so can not move and carry charge
Conductor when molten- ions are no longer fixed so they can carry charge
Often good solubility in water /polar solvents - oxygen attracted to positive ions and hydrogen attracted to negative ions this breaks down lattice and dissolves it
brittleness- slight movement of layers will result in negative being next to negative causing repulsion and cracking
66
Simple molecular structure info
Covalent compound.
Small units containing definite number of atoms with a definite molecular formula
Low melting and boiling points due to weak intermolecular forces which do not require large aounts of energy to break.
Not a conductor due to lack of mobile charge carriers
Soluble in polar solvents as all bonds in mixture require little amounts of energy to overcome
67
Giant molecular structure info
Covalent compound
Billions of atoms held together to form a network of strong covalent bonds
High melting points as strong bonds have to be broken which requires a large amount of energy
Non-conductor as no mobile charge carriers
Not soluble in polar solvents as strong bonds require too much energy to overcome
68
VSEPR theory
Valence shell electron repulsion theory states that -electron pairs repel each other and the replusion between lone pair-lone pair > lone pair - bonded pair > bonded pair - bonded pair
69
linear angle
180
70
triagonal planar angle
120
71
tetrahedral angle
109.5
72
pyramidal angle
107
73
v-shaped/bent angle
104.5
74
octahedral angle
90
75
Van der Waals forces are both....
```
London forces( induced dipole-dipole forces) and permanent dipole dipole forces
only occur in simple covalent structures
```
76
How do London forces arise?
electrons in molecules are constantly moving, at any instant the distribution may not by symmetrical. This results in a instantaneous temporary dipole. This dipole induces dipoles in neighbouring molecules and leads to attraction between neighbouring charges in the dipoles. These attractions between molecules are known as London forces and are a type of Van der waals force.
77
What affects the strength of London forces?
The more electrons there are in a molecule the stronger the fluctuations in the electron cloud and the stronger the instantaneous dipole- induced dipole forces.
For molecules with the same number of electrons, the more contact area between dipoles the stronger the induced dipoles develop.
Long thin molecules can line up closer than spherical or branched ones. hence unbranched molecules have stronger dipole forces than branched molecules
78
define electronegativity
The ability of an atoms to attract the bonded electrons in a covalent bond
79
what is a pauling value
It is a measurement of the degrees of electronegativity
80
What elements have the highest pauling value, state their values.
```
Flourine 4.0
Oxygen 3.5
Nitrogen 3.0
chlorine 3.0
greater value = greater electronegativity
```
81
What is a permanant dipole
A small charge difference that does not change across a bond, with partial charges on the bonded atoms (indicated by delta +/- signs) It is a result of bonded atoms having different electronegativities
induced dipole -dipole forces occur in all covalent molecules and permanent dipoles occure in addition to this meaning substances with permanent dipoles have higher melting/boiling points
82
polar covalent bond
A bond with a permanent dipole, having positive and negative partial charges on the bonded atoms
83
Polar molecule
A molecule with an overall dipole having taken into account any dipoles across bonds and the overall shape and symmetry of the molecule
Not all molecules with polar bonds are polar overall
84
What is a Hydrogen bond?
A hydrogen bond is a strong dipole-dipole attraction between an electron deficient hydrogen atom from -NH, OH or HF on one molecule and a lone pair of electrons on a highly electronegative atom N, O or F on a different molecule.
85
How do H bonds arise ( in water)?
O-H bond is polar due to large difference in electronegativity between oxygen and hydrogen, If another water molecule (b) were to approach the existing molecule (a) then the delta positive end of b will be attracted to the delta negative end of a.
One of the lone pairs of electrons on O from a can form a partial dative bond with the delta positive H on b
This very strong dipole-dipole interaction is a hydrogen bond.
86
Draw a hydrogen bond between two water molecules.
Check internet
lone pair drawn on oxygen, relevent delta charges drawn on, straght line of HOH with parrallel line between O of one molecule and H of the other.
87
Properties of water due to H bonds
Hydrogen bonds are responsible for the higher than expected melting and boiling points of water for a molecule of its size.
Ice floats on water, meaning ice is less dense despite in being a solid, this means for a given volume there are fewer molecules of H2O in ice than in water. meaning molecules are less closely packed. This is because when water freezes the molecules are pushed further apart in the structure of ice. making it less dense than water. Hydrogen bonding is responsible for the open structure of ice.
88
define amount of substance
A quantity whose unit is the mole, used as a means of counting any species such as atoms ions or molecules
89
define mole
The amount of any substance containing as many elementary particals as there are in exactly 12g of carbon-12 isotope. That is 6.02 x 10^23 particles
90
define Avogadros constant
The number of atoms per mole of the carbon-12 isoptope
| 6.02 x 10^23
91
Moles, Mass and molecular mass relationship.
number of moles = mass (g) / molercular mass
| n=m/Mr
92
define empirical formula
Simplest whole number ratio of atoms of each element in a compound
93
define molecular formula
The actual number of atoms of each element in a compound
94
Define anhydrous
Does not contain water of crystallisation
95
Define Hydrated
Does contain water of crystalisation
96
Define water of crystallisation
A fixed number of molecules of water incorporated into a lattice structure
(when a substance is crystallised from aqueous solution)
97
Define solute
The material which dissolves
98
Define solvent
The liquid the solute dissolve in
99
Define Solution
The resulting mixture from dissolving a solute in a solvent
100
when something is 1mol/dm^3 it means....
In each dm^3 of solution, 1 mol of the element/compound is present
101
3 main mole equations
No. of moles = mass /molar mass
No. of moles = concentration x volume
No. of moles = volume / 24
