[Year 1] EL Flashcards
(178 cards)
1
Q
What are the relative atomic masses and charges of the components of the atom?
A
Relative Charge
Proton: +1
Neutron: 0
Electron: -1
Relative Mass
Proton: 1
Neutron: 1
Electron: 1/2000
2
Q
Summarise the atomic model through time.
A
Dalton: Atoms are indivisible spheres.
Thomson: Discovered electrons - Positive sphere with negative charge scattered within. [Plum Pudding]
Rutherford: Discovered nucleus - Positive nucleus with a negative cloud which was mainly empty space.
Bohr: Fixed energy shells around a positive nucleus
Quantum: Subshells (s-, p-, d-, f-orbitals)
3
Q
Sumarise the Geigar, Marsden and ruthurfod experiment.
A
- Fired alpha particles at thin gold foil.
- Most went through.
- Small number deflected back.
4
Q
Sumarise Bohrs exeriment.
A
- Fired EM radiation.
- Was absorbed by electrons and was excited.
- When it returned to its shells, radiation was emitted.
5
Q
How were elements made?
A
Fusion reactions (in stars).
6
Q
What is Fusion?
A
The forcing together of 2 nuclei to make heavier nuclei and thus a new element.
7
Q
What conditions do fusion reactions require and why?
A
Very high temperature and pressure, to overcome the repulsive force of fusing 2 positive nuclei together.
8
Q
Show the fusion of Hydrogen to form Helium.
A
²₁H + ¹₁H → ³₂He
Proton number defines elements.
9
Q
where were heavier elements formed?
A
In larger stars (at even higher temperature and pressure)
10
Q
How did elements from stars get to earth?
A
Supernovas (that created earth).
11
Q
What are the different sub-shells?
How many orbitals does it have?
A
S - has 1 orbital
P - has 3 orbitals
D - has 5 orbitals
F - has 7 orbitals
12
Q
How many electrons can each orbital hold?
A
2.
Therfore, the…
…S subshell holds 2 electrons.
…P subshell holds 6 electrons.
…D subshell holds 10 electrons.
…F subshell holds 14 electrons.
13
Q
What is a principal quantum number?
A
The shell number.
14
Q
Relate the distance from the nucleus to the shell number’s energy?
A
- Higher shell number is further from the nucleus.
- So it has a higher energy level.
15
Q
What is the shape of the S orbital?
A
Spherical.
16
Q
What is the shape of the P orbital?
A
Shaped likes dumbells.
Px, Py, and Pz orbitals are 90⁰ from each other.
17
Q
What is spin-pairing?
A
When 2 electrons occupy 1 orbital they ‘spin’ in opposite direction.
18
Q
What does this electron configuration tell us?
1s²
A
1s²
1 : shell number.
s : subshell.
²: number of electrons.
19
Q
What is the electron configuration for Iron?
A
1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
20
Q
When filling electron configurations what rules must be followed?
A
- Fill from the lowest energy level upwards.
- Fill orbitals singularly before pairing.
- 4s fills before 3d
- remove from 4s before 3d
21
Q
Why do electrons fill orbitals singularly before pairing?
A
Electron repulsion.
22
Q
How do we show the electron configuration of ions?
A
Add or remove electrons from the highest energy level.
23
Q
What is the electron configuration for a calcium ion?
A
1s² 2s² 2p⁶ 3s² 3p⁶
24
Q
How are shorthand electron configuration shown?
A
Use the closest noble gas symbol.
e.g.
K = 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹ = [Ar] 4s¹
25
What is the electron configuration for an Iron (3+) ion?
1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵
26
What is an ionic compound?
Oppositely charged ions held together by electrostatic attractions.
27
What ions do these groups form?
```
Group 1?
Group 2?
Group 3?
Group 5?
Group 6?
Group 7?
```
```
Group 1: 1+ ions
Group 2: 2+ ions
Group 3: 3+ ions
Group 5: 3- ions
Group 6: 2- ions
Group 7: 1- ions
```
28
What ions do these molecules form?
```
Hydroxide?
Nitrate?
Ammonium?
Sulfate?
Carbonate?
```
```
Hydroxide: OH⁻
Nitrate: NO₃⁻
Ammonium: NH₄⁺
Sulfate: SO₄²⁻
Carbonate: CO₃²⁻
```
29
Show the formula for the bonding of calcium ions and nitrate ions?
Ca⁺ + NO₃⁻ → Ca(NO₃)₂
30
What is the structure of ionic compounds, like sodium chloride?
Giant Ionic Structure
- Regular structure
- Cubic shape
- Giant repeating pattern
31
Why do most ionic compounds dissolve in water?
- Water molecules are polar.
| - So the δ⁻ oxygen can attract the positive charge in an ion to break up its structure.
32
When do ionic compounds conduct? Why?
- When molten or dissolved in solution.'
| - As ions are free to move around
33
Why do ionic compounds have high melting points?
- There are many, strong electrostatic forces between oppositely charges ions.
- Lots of energy is needed to overcome these forces.
34
What is a covalent bond?
The sharing of outer electrons in order for atoms to obtain a full shell.
35
What force keeps a covalent bond?
An electrostatic attraction force between the shared electron and the positive nucleus.
36
How are covalent bond represented in diagrams?
- Dot and cross.
| - Lines
37
What is a dative covalent/coordinate bond?
Where one atom donate 2 electrons to an atom or ion to form a bond.
E.g. NH₃ + H⁺ → NH₄
has a dative covalent bond.
38
How are dative covalent/coordinate bond shown in a diagram
- Dot and cross (but with only dots or crosses)
- An arrow towards the atom with no electron sharing
Eg. towards H in NH₄
39
What bond does carbon monoxide have?
- a double covalent bond.
| - and a dative covalent bond.
40
What are examples of giant covalent structures?
Graphite.
Diamond.
Silicon(IV) Dioxide.
41
What is the structure of graphite?
Layers of graphene with carbon bonded 3 times.
Therefore there is one delocalised electron.
42
Describe the qualities of graphite.
- Very high melting point, due to strong covalent bonds.
- Conducts electricity, due to delocalised electrons that can carry a charge.
- Layers slide over each other easily, due to weak forces between layers.
- Low density, due to layers being far apart in comparison to bond length.
- Insoluble, due to strong covalent bonds. These are too strong to break.
43
What is the structure of diamond?
- Tetrahedral shape with carbon bonded 4 times.
| - Therefore there are no delocalised electrons.
44
Describe the qualities of diamond/silicon(IV) dioxide.
- Heat conductive, due to the tightly-packed rigid arrangement.
- Diamond can be cut into gemstones, unlike graphite.
- Very high melting point, due to many strong covalent bonds. Also very hard.
- Doesn't conduct electricity, due to no delocalised electrons to carry a charge. (and not when liquid as difficult to melt, and normally sublime)
- Insoluble, due to strong covalent bonds. These are too strong to break.
45
What is a metallic bond?
Positive metal ions that donate electrons to form a 'sea' of delocalised electrons.
46
What force keeps a metallic bond together?
An electrostatic attraction between positive metal ions and negative delocalised electrons.
47
How can we work out what giant metalic lattice structure has a higher melting point?
If its atom donates more electrons to the delocalised system, the structure will have a higher melting point.
Mg can donate 2 electrons, whereas Na can only donate 1. Thus Mh has a higher melting point than Na.
48
Describe the qualities of a giant metalic lattice structure.
- Good thermal conductors, due to the delocalised electrons, can transfer kinetic energy.
- Good electrical conductors, due to the delocalised electrons are mobile and can carry a current.
- High melting points, due to strong electrostatic attractions.
- Insoluble, due to strong covalent bonds. These are too strong to break.
- Malleable and ductile, due to ion layers being able to slide when hit with a force, but still retain an attraction between ions and delocalised electrons.
49
Why do molecules have a specific shape with specific angles?
Because bonds repel each other so that electrons are as far apart as possible.
Lone pairs repel more than bond pairs.
50
How do lone pairs change the shape of molecules?
Lone pairs will push bond pairs closer together as they repel more.
For every lone pair, the bond angle reduces by 2.5°
(with some exceptions).
51
What method can be used to work out bond angles?
- Draw dot an cross.
- Count bond pairs and lone pairs.
- Link total pairs to a shape.
- Adjust according to lone pairs.
```
E.g.
Water:
BP: 2
LP: 2
Total: 4
```
Because total is 4, it's based on tetrahedral.
Tetrahedral = 109.5°
- 2 lone pairs = 109.5° - (2 x 2.5°)
= 104.5°
52
What is the bond angle of a linear-shaped molecule?
180°
53
What is the bond angle of a Trigonal Planar-shaped molecule?
120°
54
What is the bond angle of a Tetrahedral-shaped molecule?
109.5° (3D)
55
What is the bond angle of a Trignal Bipyramidal-shaped molecule?
120° AND 90°
| like trigonal planar but with a perpendicular line through its middle
56
What is the bond angle of an Octahedral-shaped molecule?
90°
57
What is the bond angle of a Pyramidal-shaped molecule?
107°
58
What is the bond angle of a Bent-shaped molecule?
104.5°
59
What is the bond angle of a Square planar-shaped molecule?
90°
60
What shape do you get with:
BP = 2
LP = 0
Linear
61
What shape do you get with:
BP = 3
LP = 0
Trigonal Planar
62
What shape do you get with:
BP = 4
LP = 0
Tetrahedral
63
What shape do you get with:
BP = 5
LP = 0
Trigonal Bipyrimidal
64
What shape do you get with:
BP = 6
LP = 0
Octahedral
65
What shape do you get with:
BP = 3
LP = 1
Pyramidal
66
What shape do you get with:
BP = 2
LP = 2
Bent
67
What shape do you get with:
BP = 3
LP = 2
Trigonal Planar
| lone pairs repel equally from opposite sides
68
What shape do you get with:
BP = 4
LP = 2
Square Planar
| lone pairs repel equally from opposite sides
69
Describe the qualities of simple covalent molecules.
- Liquid or gas at room temperature, due to the low melting point. (Iodine is solid due id-id to crystal structure).
- Doesn't conduct electricity, due to no delocalised electrons.
- Solubility depends on the polarity of the molecule.
Polar dissolves well in polar. Non-polar don't.
- Low boiling point due to weak id-id forces.
70
What does the EM spectrum show?
The types of radiation at a different frequency.
71
List the EM spectrum.
```
Radio
Micro
Infra Red
Visible
Ultra Violet
X-Rays
Gamma
```
72
What is the trend in Energy on the EM spectrum?
Energy increases across:
Radio = lowest
Gamma = highest
73
What is the trend in Frequency on the EM spectrum?
Frequency increases across:
Radio = lowest
Gamma = highest
74
What is the trend in Wavelength on the EM spectrum?
Wavelength decrease across:
Radio = longest
Gamma = shortest
75
What is line spectra?
A way of identifying elements and in evidence for energy level shells.
76
What is the energy level closest to the nucleus known as?
The ground state.
77
Can electrons be between energy levels?
No, they have discrete energy values and electrons can only move from one shell to another.
78
What happens when an electron absorbs energy?
Providing it has enough energy, it is excited to a higher energy level.
79
What happens when an electron returns to a lower energy level?
Energy is released.
(and is a form of energy found on the EM spectrum).
80
What does an emission spectrum show?
The frequency of light given out when an electron moves down energy levels.
We see this as a coloured line on a black background.
81
Why are emission spectra unique to different elements?
- Every element has a different electron configuration.
- This means it will absorb and emit different frequencies of radiation.
- Hence the unique emission spectra.
82
What does an absorption spectrum show?
The missing frequencies of radiation when passing through the sample, as they have been absorbed by the electrons).
We see this as a black line on a coloured spectrum.
83
When do we get an absorption spectra?
When EM radiation is passed through an element in the gas state.
The electrons absorb specific frequencies that correspond to the difference in energy between the discrete shells. As E=hf.
84
What does a line spectra show?
The different regions of the EM spectrum. (on a graph)
85
What do we call a group of lines on a line spectra?
How are these created?
A series.
When electrons move to the same energy level from different ones.
86
On a line spectra, what happens to the lines as the energy and frequency increase?
The lines get closer.
87
What do arrows show on a line spectra?
The potential path an excited electron could follow to different energy levels.
88
What part of the spectrum is emitted or absorbed when the arrows go to the ground state?
Is it emitted or absorbed?
Electrons fall to the ground state (n=1).
Thus energy is emitted in the UV part of the spectrum.
89
What is the trend between the distance between bands and the width of band on an absorption spectra?
Wider bands with larger distances between other bands fall from lower energy levels (e.g. n=2 to n=1 OR n=3 to n=2) .
90
What part of the spectrum is emitted or absorbed when the arrows go to the second energy level?
Is it emitted or absorbed?
Electrons fall to the second energy level (n=2).
Thus energy is emitted in the visible part of the spectrum.
91
What part of the spectrum is emitted or absorbed when the arrows go to the third energy level?
Is it emitted or absorbed?
Electrons fall to the third energy level (n=3).
Thus energy is emitted in the IR part of the spectrum.
92
How is emission spectra evidence to prove the existence of quantum shells?
The defined lines prove electrons exist in quantum shells.
They can never exist between them.
This proves the shells have fixed energy.
Which is why the radiation emitted will have a fixed frequency.
Wich means EM radiation is absorbed to move electrons to higher quantum shells and emitted when they drop.
Giving our defined lines.
93
How is energy and frequency linked?
ΔE = hν
ΔE: energy difference between shells. (J)
h: Planck's constant. (JHz⁻¹)
ν: Frequency (Hz or s⁻¹)
94
How is the speed of light linked to frequency?
c = νλ
c: speed of EM radiation/light through a vacuum (ms⁻¹)
ν: Frequency (Hz or s⁻¹)
λ: Wavelength (m)
95
How is energy and wavelength linked?
ν= c/λ
...then use ν in...
ΔE = hν
Do these separately to avoid mistakes and to get all the marks (watch out for rounding errors).
96
How can we test for cations in a solid sample?
(cations = posative ions)
Flame tests.
Look at colours emitted and correspond it to an element. OR look though spectra scope and observe bands on emission spectra.
97
Describe a test for cations.
1. Dip nichrome wire in concentrated HCl
2. Dip into the solid sample
3. Place loop into blue Bunsen flame and observe colour
98
List the colours of flames for each cation (that should be known).
```
Li⁺: Crimson.
Na⁺: Yellow-Orange.
K⁺: Lilac.
Ca²⁺: Dark Red.
Ba²⁺: Green.
Cu²⁺: Green-Blue.
```
99
Describe how mass spectrometers work.
(HINT: 5 steps)
NOT ON SPEC
1. Vaporisation- vaporise sample so it can travel through the mass spectrometer.
2. Ionisation - Bombard sample with high energy electrons to ionise the atom to its +1 charge. Allowing it to be accelerated and detected.
3. Acceleration- Pass through an electric field. Low mass/charge ration will accelerate quicker.
4. Ion Drift - Particles travel though with a constant speed and KE. They drift through and particles with lower m/z ratio will travel faster.
5. Detection - Ions are detected as an electrical current when it hits a plate. Particles with lower m/z reach the detector first as they are fastest.
100
What can be plotted from the data in a mass spectrometer?
What does this show?
%Abundance - m/z graph
OR
rel.Abundance - m/z graph
(m/z is the same as the isotopic mass if a 1+ charge is used.)
Shows the Abundance of isotopes in the sample.
101
How do we work out the RAM from the data from a mass spectra graph?
RAM = Σ(Abundance x m/z) / Total Abundance
102
What is an ion?
Ions have a different number of electrons to protons.
103
Wha is an isotope?
Isotopes have a different number of neutrons but the same number of protons.
104
Define the Relative atomic mass (Ar)
The weighted mean mass of an atom of an element, compared to ¹/₁₂ᵗʰ of the mass of an atom of carbon-12.
105
Define the Relative molecular/formula mass (Mr)
The mean mass of a molecule compared to ¹/₁₂ᵗʰ of the mass of an atom of carbon-12.
106
Define the Relative isotopic mass?
The mass of an atom of an isotope compared to ¹/₁₂ᵗʰ of the mass of an atom of carbon-12.
107
How many atoms are in one mole of the atom?
Avogadro's number of atoms
=
6.02x10²³
108
How do we calculate the number of particles in a substance if we know the number of moles?
Number of particles = Avogadro's number x Number of Moles
109
How can we work out the number of moles?
Moles = Mass / Mᵣ (OR Aᵣ)
Moles = Concentration (moldm⁻³) x Volume (dm³)
110
How do we convert cm³ to dm³?
÷ 1000
111
What do ionic equations show?
The ions that are formed in solution and which particles are reacting.
112
For the following equation show the full ionic equation and then the simplest ionic equation:
H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O
H₂SO₄ + 2KOH → K₂SO₄ + 2H₂O
2H⁺ + SO₄²⁻ + 2K⁺ + 2OH⁻ → 2K⁺ + SO₄²⁻ + 2H₂O [water isn't an ion so its left as is]
[cancle any ions that appear on both side]
2H⁺ + 2OH⁻ → 2H₂O
(Charges must balance)
113
What are the ions not involved in the ionic equation called?
Spectator ions.
114
What is an empirical formula?
The simplest whole-number ratio of elements in a compound.
115
How would you work out this question:
How much CaO can be made when 34g of Ca is burnt completely in oxygen?
1. Write out a balanced equation.
2Ca(s) + O₂(g) → 2CaO(s)
2. Work out Mr/Ar.
2Ca = 40.1x2 = 80.2
2CaO = (40.1 + 16)2 =112.2
3. Use Moles = mass/Mr equation to work out moles for the given mass.
Moles = 34/80.2 = 0.42mol
4. Use the ratio from the equation to work out the mole of the product.
Ca:CaO = 2:2 = 1:1 = 0.42:0.42
5. Use mass = mole x Mr.
mass = 0.42 x 112.2 = 47.6g
116
What is an emirical formula?
The simplest whole number ration of elecment in a compound.
117
How would you work out this question:
A comound contains 23.3% Mg, 30.7% S, and 46.0% O. What is the empirical formula of this compound?
1. Write out elements involved.
Mg S O
2. Write % as masses.
23. 3g 30.7g 46.0g
3. Find Mr of each.
24. 3 32.1 16.0
4. Do Moles = mass/Mr.
0. 96 0.96 2.88
5. Divide by smalles number of moles and write as ratio.
0.96:0.96:2.88
1 : 1 : 3
6. Write final formula.
MgSO₃
118
How do you work out the molecular formula using the empirial formula?
1. Work out the Mr of the emiricle formula.
CH₂O = 12.0 + 2(1.0) + 16.0 = 30
2. Divide by the Mr of the molecular formula.
Mr of molecule = 180
180/30 = 6
3. Use the number to multipy by all the atoms in the empirical formula.
Cx6 H₂x6 Ox6
C₆H₁₂O₆
119
How would you work out this question:
1.88g of hydrated CaSO₄.xH₂O was heated until there was no more water of crystallisation left in the sample. The mass of this anhydrous compound is 1.13g. Work out the value of 'x'.
1. Write out molecules involved.
CaSO₄ H₂O
2. Write down the mass of each molecule.
1. 13g 0.75g
3. Do Moles = mass/Mr.
0. 0083 0.0416
4. Divide both by smallest number of moles.
1 5
5. Thus x=5
120
What is the formula for % yeild?
%Yield = Actual / Theoretical x100
121
What is the theoretial yeild?
The amount of product produced assuming NO products are lost and ALL reactantsd react fully.
122
What can effect %Yeild?
- Evaporation occurs in liquids.
- Transferring liquids from one container to another as some liquid will always be left in the container or you can spill the liquid when transferring.
Filtering the solution, because liquid will be dissolved into the filter paper or solid left behind.
Reactants may not react to make the product, especially if the reaction is reversible.
123
What is a standard solution?
A solution, used in titrations, with a known concentration.
124
Describe the method in making a standard solution.
Weigh out the amount of solid precisely using a balance and plastic/glass weighing boat.
Transfer solid from weighing boat to a beaker.
Wash any solid left behind into the beaker using DI water.
Dissolve solid fully using DI water. Stir to ensure the solid is dissolved fully.
Transfer solution to volumetric flask. Use a funnel to avoid spliiage adn rinse the bealker AND glass rod into the flask to make sure most is transfered.
Use DI water to fill to the graduation line. Donnot go above this line. You can use a pipette as you near the line.
Cover and inver the flask a few times. This is to ensure your solution is thoroughly mixed and ready to use.
Use mass and volume to work out the concentation of the solution.
125
How would you work out this question:
What mass of solid NaOH is reuired to make 250cm³ of 0.75mol dm⁻³ NaOH solution?
1. Work out moles via Moles = Concentration(mol dm⁻³) x Volume(dm³).
Moles = 0.75 x 250x10⁻³ = 0.1875 mol
2. Work out mass via mass = Moles x Mr
mass = 0.1875 x 40 = 7.5g
126
How would you work out this question:
If we wanted to make a 250cm³ solution of 0.75mol dm⁻³ HCl from 2.5moldm⁻³ HCl, the volume of more concentrated HCl would be-
Volume to use = Final con/initial con x volume required
Volume to use = 0.75/2.5 x 250 =75cm³ of conc HCl needed
127
What can titration be used to work out?
The concentration of an acid or alkali.
128
What does concordant mean in titrations?
Within 0.10cm³ of each other.
129
What are indicatior that can be used in a titration?
Phenolphthalein
Acid to Alkili
Colourless to Pink.
Methyl Orange
Acid to Alkili
Yellow to Red.
130
How would you work out this question:
18.3cm³ of 0.25mol dm⁻³ HCl was required to neutralise 25cm³ of KOH. Calculate the concentration of KOH.
1. Write out balence equation.
HCl + KOH → KCl + H₂O
2. work out moles of known via Moles = Conc x Vol
Moles = 0.25 x 18.3x10⁻³ = 4.58x10⁻³ mol
3. Use ration to work out moles on unknow.
HCl:KOH
1:1
4.58x10⁻³:4.58x10⁻³
4. Work out conc of unknown via Conc = Moles / Vol
Conc = 4.58x10⁻³ /25x10⁻³ = 0.18mol dm⁻³
131
What is the conversion of dm³ to cm³?
x1000
132
Describe the trend in melting points of the first 3 elements across period 3.
First 3 element are metrals and thus have matlaic bonding.
This increases across as the metal ions have an increasing posative charge, increasing the number of delocalised electrons and smaller ionic radius.
Thius mean s a stronger metalic bond is possible.
SMAE FOR P2
133
Describe the peak in melting point of period 3 elements.
Silocon has the highest melting pont in period 3.
It has a giant covalent structure.
Many strong covalent bonds hold the silicon atoms together (tetrahedral shaped).
A large amount of energy is needed to overcome these strong bonds.
SMAE FOR P2
134
Describe the trend in melting points for the 4th element in period 3.
Phosphosus has the formula P₄.
It has a lower melting point than Silicon due to weaker simple molecular stucture.
The melting point is down to weaker intermolecular forces.
SMAE FOR P2
135
Describe the trend in melting points for the 5th element in period 3.
Sulfur has the formula S₈.
It has a higher melting point due to a larger simple molecular strucute.
It has larger intermolecular forces and hence a higher melting point.
SMAE FOR P2
136
Describe the trend in melting points for the 6th element in period 3.
Chlorine has the formula Cl₂.
It has a lower melting point than phosphorus and sulfur due to a smaller simple molecular strucute.
It has smaller intermolecular forces and hence a lower melting point.
SMAE FOR P2
137
Describe the trend in melting points for the 7th element in period 3.
Argon has the formula Ar.
It has the lowest melting point in period 3 due to existing as individual atoms.
It has smaller intermolecular forces and hence a lower melting point.
SMAE FOR P2
138
What is ionisation energy?
The minimum amount of energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous statr.
139
Show the first ionisation enthalpy of Sodium.
Na(g) → Na⁺(g) + e⁻
140
What is electron shielding in bonding?
The more electron shells between the +tive nulceus and the -tive electrons that is being removed the less enrgy is required.
There is a a weaker attraction.
141
What is atomic size in bonding?
The bigger the atom the further away the outer electron are from the nucleus.
The attractive force between the nucleus and outer electrons reduces -easier to remove eletrons.
142
What is nuclear charge in bonding?
The more protons in the nucleus the bigger the attraction between the nucleus adn outer electrons.
This means more energy required to remove the electron.
143
What is the trend in ionisation enthalpy down a group?
Decrease.
- Atomic size increases
- Sheilding increases
144
What is the trend in ionisation enthalpy across a period (in P2 and 3)?
Increase.
- Increased nuclear charge.
- Similar sheilding
145
Group 2 + Water →
Metal Hydroxide + Hydrogen gas.
146
What is the trend in reactivity down group 2?
It increases (Be no reaction).
- Atomic size increases.
147
What do group 2 oxides look like?
White solids.
148
Why are metal hydroxides alkili?
Oxides react readily with water to make hydroxides whihc dissociate to from OH⁻ ions.
149
What is the trend in alkilinity of metal hydroxides down a group?
They become more strongly alkaline as we go down the grouo as the hydroxides become more soluble.
150
What is the trend in solubiltiy of group 2 hydroxides and carbonates.
Carbonates decrease down.
(double charged become less soluble down a group)
Hydroxides increase down.
(singel charge become more soluble down a group)
151
What happends in the thermal decomposition of group 2 carbonates?
Break down into metal oxides and carbon dioxide.
152
What happends in the volume of CO₂ made from the thermal decomosition of group 2 carbonates as you go down a group.?
As we go down the group carbonates of the same mass produce less CO₂.
153
What is the trend in thermal stability of group 2 carbonates down a group.
Carbonates become more thermally stable down a group.
Due to large electron cloud that can be distorted when nearby positive group 2 metal ions
154
What is the trend in charge density of group 2 ions down a group.
They get lower.
- As charge is always 2+ but atom becomes larger.
- Thus more charge needs to be shared over a larger space.
155
What are the formulas for these cations.
```
Zinc(II)
Lead(II)
Copper(II)
Ammonium
Iron(II)
Iron(III)
```
Zinc(II) = Zn²⁺
Lead(II) = Pb²⁺
Copper(II) = Cu²⁺
Ammonium = NH₄²⁺
Iron(II) = Fe²⁺
Iron(III) = Fe³⁺
156
What are the formulas for these anions.
```
Sulfate
Carbonate
Hydrogencarbonate
Hydroide
Nitrate
```
Sulfate = SO₄²⁻
Carbonate = CO₃²⁻
Hydrogencarbonate = HCO₃⁻
Hydroide = OH⁻
Nitrate = NO₃⁻
157
When are nitrates soluble and insoluble?
```
SOLUBLE
All Potassium
All Litium
All Ammonium
All Nitrates
All Sodium
[PLANS]
```
INSOLUBLE
158
When are halides soluble and insoluble?
```
SOLUBLE
Most Halides (Chlorides, Bromides, Iodides)
```
INSOLUBLE
Silver Halides
Copper Iodide (white)
Lead Chloride, Bromide (white), Iodide (yellow)
159
When are hydroxides soluble and insoluble?
```
SOLUBLE
Sodium
Calcium
Ammonioum
Lithium
Potassium
Strontium
Barium
[SCALPS+B]
```
INSOLUBLE
Most Hydroxides
160
When are sulfates soluble and insoluble?
SOLUBLE
Most Sulfates
INSOLUBLE
Barium
Calcium
Lead
161
When are carbonates soluble and insoluble?
```
SOLUBLE
Potassium
Ammonimum
Lithium
Sodium
[PALS]
```
INSOLUBLE
Most Carbonates (white)
Copper (blue/green)
Silver (yellow)
162
What is a precipitation reaction?
A reaction that produces an insoluble solid.
163
How can we speperate insoluble salts?
Filtration
- Filter
- Rinse (with DI water)
- Dry
164
How can we make an insoluble salt?
Acid + metals/insoluble bases
165
How can we speperate soluble salts?
Crystalisation
| - Heat salt to evaporate water
166
How can we make an insoluble salt (with alkalis)?
- Titrate with indiactor
- note volume
- re-titrate without indicator for the pre-measured volume.
167
How can we test for cations?
- Flame tests.
| - Using Sodium Hydroxide.
168
How can we use sodim hydroxide to test for cations?
- Add NaOH to unknown solution.
| - Precipitate will form.
169
What colour precipitates form when NaOH reacts with the following cations:
```
Alluminum
Zinc
Silver
Lead(II)
Copper(II)
Calcium
Iron(II)
Iron(III)
```
Alluminum = White (disolve in exess)
Zinc = White (disolve in exess)
Silver = Brown (silver oxide)
Lead(II) = White
Copper(II) = Blue
Calcium = White
Iron(II) = Green
Iron(III) = Brown-Red
170
How can we test for carbonates?
- Add acid
| - Turns limewater cloudy
171
How can we test for sulfates?
- Add HCl
- Add Barium Chloride
- White precipitate forms
172
How can we test for ammonum compounds?
- Add NaOH
- Heat gently
- Gas should produce
- Pass though damp litmus paper
- Should turn blue
173
How can we test for hydroxides?
- Turn red litmus blue
| - further testing needed to confirm
174
How can we test for halides?
- Add dilute nitric acid.
- Add silver nitrates.
```
Chloride = White
Bromide = Cream
Iodide = Yellow
```
Chlorine will disolve in dilute ammonia.
Bromine will disolve in concentrated ammonia.
Iodine will not dissolve.
175
Why do we add nitric acid when testing for halides?
To react with any other anions (e.g carbonates which also form yellow precipitate with silver).
Thus avoiding false results.
176
How can we test for nitrates?
- Add NaOH.
- Add aluminum foil.
- Heat.
- If present, will be reduced by alluminum to ammonium ions
- This will react withh OH to form ammonia gas.
- Which turns red litmus blue
177
Why do we test for ions in a spesific order?
To avoid false positives.
178
In what order should we test for ions?
- Carbontes.
THEN
- Sulfates.
THEN
- Halides.
