3.1 Periodicity Flashcards
1
Q
how is the modern periodic table arranged?
A
by atomic number
2
Q
what groups are in s block on the periodic table?
A
1 and 2
3
Q
what groups are in d block on the periodic table?
A
transition metals
4
Q
what groups are in f block on the periodic table?
A
bottom two rows
5
Q
what groups are in p block on the periodic table?
A
3, 4 ,5, 6, 7, 0
6
Q
what is the definition of periodicity?
A
the regular repeating patterns in the physical and chemical properties of elements
7
Q
what is the definition of ionisation?
A
when an atom loses an electron from its outer shell
8
Q
what is the definition for first ionisation energy?
A
the energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions
9
Q
what is the first ionisation energy equation for sodium?
A
Na(g) —> Na+(g) + e-
10
Q
what is the second ionisation energy equation for sodium?
A
Na+(g) —-> Na2+(g) + e-
11
Q
what is the third ionisation energy equation for sodium?
A
Na2+(g) —-> Na3+(g) + e-
12
Q
what can we learn from a successive ionisation energy graph?
A
what group an element is in - where the first big jump us shows number of electrons in outer shell
13
Q
what are the 3 factors that affect ionisation energy?
A
shielding
nuclear charge
atomic radius
14
Q
what is the definition of shielding?
A
the affect of the inner electrons shielding the outer electrons from the affect of the charge on the nucleus
15
Q
what is the trend in shielding across period 3?
A
stays the same across period 3
16
Q
what is the definition of nuclear charge?
A
the positive charge on the nucleus
17
Q
what is the trend in nuclear charge across period 3?
A
increases across period 3
18
Q
how do we measure atomic radius?
A
by measuring the distance between two nuclei of touching atoms and halving the distance
19
Q
what is the trend in atomic radius across period 3?
A
decreases across period 3 - and other periods
20
Q
why does the atomic radius decrease across period 3?
A
the nuclear charge increases, the shielding stays the same, the nuclear attraction increases so the atomic radius decreases
21
Q
what is the general trend of ionisation energy in period 3?
A
increases across period 3
22
Q
explain the general trend of ionisation energy in group 3
A
the nuclear charge increases, the atomic radius decreases, the nuclear attraction increases, the shielding stays the same - it takes more energy to remove the first electron
23
Q
why does first ionisation energy decrease between magnesium and aluminum?
A
aluminum has one electron in a higher subshell (3p), this one electron is removed more easily as it is further away from the nucleus, therefore the first ionisation energy is lower than magnesium
24
Q
why does the first ionisation energy go down between phosphorous and sulphur?
A
sulphur has one 3p orbital that contains a pair of electrons - these paired electrons repel each other, so one of these electrons is easier to remove therefore sulphur has a lower ionisation energy than phosphorous
25
what is the trend for shielding and nuclear charge down group 2?
they both increase down the group
26
what is the trend for atomic radius down group 2?
increases down the group
27
what is the explanation for atomic radius increasing down group 2?
shielding increases, nuclear charge increases but is cancelled out by the extra shielding
so the nuclear attraction decreases and atomic radius gets bugger
28
what is the trend in first ionisation energy down group 2?
decreases down the group
29
what is the explanation for the trend in first ionisation energy down group 2?
decreases down the group because there is an increase in shielding, an increase in nuclear charge, nuclear attraction of the electrons to the nucleus decreases, atomic radius increases so less energy is required to remove the outer electron
30
What is the arrangement of elements in the periodic table?
Elements are arranged in increasing atomic number
## Footnote
This arrangement reflects the periodic law and the structure of the atomic nucleus.
31
What do elements in the same group have in common?
Similar physical and chemical properties
## Footnote
This similarity arises from having similar outer shell electron configurations.
32
How are elements classified in the periodic table?
Elements are classified as s, p, or d block
## Footnote
The classification depends on which orbitals the highest energy electrons are in.
33
What is periodicity in terms of elements?
A repeating pattern across different periods
## Footnote
This pattern includes trends in physical and chemical properties.
34
What is the trend of atomic radius across a period?
Atomic radii decrease from left to right across a period
## Footnote
This decrease is due to increased positive charge attraction from protons.
35
What is the first ionisation energy?
Energy needed to remove an electron from each atom in one mole of gaseous atoms
## Footnote
The process is always considered in the gaseous state.
36
What factors affect ionisation energy?
1. The attraction of the nucleus
2. The distance of the electrons from the nucleus
3. Shielding of the attraction of the nucleus
## Footnote
Each of these factors influences how tightly electrons are held by the nucleus.
37
Why are successive ionisation energies always larger?
The ion formed is smaller than the atom, resulting in a greater proton to electron ratio
## Footnote
This leads to stronger attraction between the nucleus and remaining electrons.
38
What does a big jump between successive ionisation energies indicate?
It indicates a change in the electron shell being removed
## Footnote
For example, if there's a large jump between the 2nd and 3rd ionisation energies, it suggests that the 3rd electron is removed from an inner shell.
39
Fill in the blank: The attraction of the nucleus is greater when there are more ______.
protons
## Footnote
A greater number of protons increases the positive charge, enhancing attraction.
40
True or False: The first ionisation energy can vary based on the physical state of the atom.
False
## Footnote
The first ionisation energy is always considered for gaseous atoms.
41
What does a large ionisation energy suggest about an element's group?
It suggests the element is in Group 2 if there is a large jump between the 2nd and 3rd ionisation energies
## Footnote
This is due to the removal of an electron from a shell closer to the nucleus.
42
List the elements in Period 2.
Li, Be, B, C, N, O, F, Ne
## Footnote
These elements display trends in properties across the period.
43
List the elements in Period 3.
Na, Mg, Al, Si, S, Cl, Ar
## Footnote
Similar to Period 2, these elements also exhibit periodic trends.
44
What is periodicity in relation to ionisation energy?
A repeating pattern across a period.
45
What useful information does the pattern in first ionisation energy provide?
Information about electronic structure.
46
Why does Helium have the largest first ionisation energy?
Its first electron is closest to the nucleus with no shielding effects.
47
Why do first ionisation energies decrease down a group?
Outer electrons are further from the nucleus and more shielded.
48
What generally happens to first ionisation energy across a period?
It increases.
49
Why does first ionisation energy increase across a period?
Number of protons increases, enhancing nuclear attraction.
50
Why does Na have a much lower first ionisation energy than Neon?
Na's outer electron is in a 3s shell, further from the nucleus and more shielded.
51
What causes the small drop in ionisation energy from Mg to Al?
Al begins to fill a 3p subshell, which is easier to remove compared to 3s.
52
Why is there a small drop in ionisation energy from P to S?
The 4th electron in sulfur causes slight repulsion in the 3p orbital.
53
What are the three factors that control ionisation energy?
Nuclear charge, distance from nucleus, and shielding effect.
54
Fill in the blank: The first ionisation energy is the energy required to remove an electron from a _______.
neutral atom.
55
True or False: The number of protons in an atom affects its ionisation energy.
True.
56
Fill in the blank: Electrons in the same orbital must have _______ spins.
opposite.
57
What happens when the second electron is added to a 3p orbital?
There is slight repulsion between the two negatively charged electrons.
58
What is metallic bonding?
The electrostatic force of attraction between the positive metal ions and the delocalised electrons
## Footnote
This type of bonding is characteristic of metals.
59
What are the three main factors that affect the strength of metallic bonding?
* Number of protons/Strength of nuclear attraction
* Number of delocalised electrons per atom
* Size of ion
## Footnote
Each factor contributes to the overall strength of the metallic bond.
60
How does the number of protons affect metallic bonding?
The more protons, the stronger the bond
## Footnote
Protons contribute to the nuclear attraction in metallic bonding.
61
How does the number of delocalised electrons per atom affect metallic bonding?
The more delocalised electrons, the stronger the bond
## Footnote
Delocalised electrons are the outer shell electrons that participate in bonding.
62
How does the size of the ion affect metallic bonding?
The smaller the ion, the stronger the bond
## Footnote
Smaller ions can get closer to the delocalised electrons, increasing attraction.
63
Which metal has stronger metallic bonding, Mg or Na?
Mg
## Footnote
Mg has more electrons in the outer shell and a smaller ion size, leading to stronger bonding.
64
What is a macromolecular structure?
Giant molecular structures
## Footnote
Examples include diamond and graphite.
65
Describe the arrangement of carbon atoms in diamond.
Tetrahedral arrangement with 4 covalent bonds per atom
## Footnote
This structure contributes to diamond's hardness.
66
Describe the arrangement of carbon atoms in graphite.
Planar arrangement with 3 covalent bonds per atom in each layer; delocalised electrons between layers
## Footnote
This allows layers to slide over each other, making graphite slippery.
67
Why do macromolecular structures like diamond and graphite have very high melting points?
Because of strong covalent forces in the giant structure
## Footnote
It takes a lot of energy to break the many strong covalent bonds.
68
What type of bonding is described as the electrostatic force of attraction between positive metal ions and delocalised electrons?
Metallic bonding
## Footnote
This is a characteristic feature of metals.
69
What are some examples of macromolecular structures?
* Diamond
* Graphite
* Silicon dioxide
## Footnote
These structures are known for their strong covalent bonds.
70
Fill in the blank: _____ is a type of bonding where a shared pair of electrons is involved.
Covalent
## Footnote
Covalent bonding occurs between non-metals.
71
What is the structure of metallic bonding called?
Giant metallic lattice
## Footnote
This lattice structure is characteristic of metals.
72
What type of bonding is present in macromolecular structures?
Covalent bonding
## Footnote
Macromolecular structures have many strong covalent bonds that require a lot of energy to break.
73
How do the boiling and melting points of macromolecular substances compare?
High
## Footnote
Due to many strong covalent bonds in their structure.
74
What is the solubility of macromolecular substances in water?
Insoluble
## Footnote
Macromolecular substances do not dissolve in water.
75
What is the conductivity of macromolecular substances when solid?
Poor
## Footnote
Electrons are localised and cannot move.
76
What is the conductivity of graphite?
Good
## Footnote
Graphite has free delocalised electrons between layers, allowing for conductivity.
77
What is the general description of giant metallic structures?
Shiny metal, malleable
## Footnote
The identical positive ions allow planes of ions to slide over each other easily.
78
What type of bonding is present in giant metallic structures?
Metallic bonding
## Footnote
Strong electrostatic forces between positive ions and a sea of delocalised electrons.
79
What is the solubility of giant metallic substances in water?
Insoluble
## Footnote
Giant metallic structures do not dissolve in water.
80
How does the conductivity of giant metallic substances compare when molten?
Good
## Footnote
Delocalised electrons can move through the structure when molten.
81
What is the melting and boiling point trend for Na, Mg, and Al?
High melting and boiling points
## Footnote
Due to strong metallic bonding that increases with more electrons released to the sea of electrons.
82
What is the melting point of Si and why?
Very high
## Footnote
Si is macromolecular with many strong covalent bonds requiring a lot of energy to break.
83
What type of forces are present in simple molecular substances like Cl2 and Sg?
Weak London forces
## Footnote
These weak forces require little energy to break, resulting in low melting and boiling points.
84
Why does Se have a higher melting point than P4?
More electrons
## Footnote
Se has stronger London forces due to having more electrons.
85
What type of forces are present in monoatomic gases like Ar?
Weak London forces
## Footnote
These forces result in very low melting points.
86
What trend is observed in period 2 for Li and Be?
Metallic bonding with high melting points
## Footnote
Both elements exhibit strong metallic bonding.
87
What type of bonding is observed in B and C?
Macromolecular
## Footnote
They have very high melting points due to strong covalent bonds.
88
What type of substances are N2 and O2?
Molecular gases
## Footnote
They have low melting points due to small London forces.
89
What is the melting point of Ne?
Very low
## Footnote
Ne is a monoatomic gas with weak London forces.
90
What happens to the atomic radius as one goes down Group 2?
Atomic radius increases down the Group.
## Footnote
Atoms have more shells of electrons making the atom bigger.
91
What is the outer shell electron configuration for Group 2 metals?
Outer shell s² electron configuration.
## Footnote
This configuration is common among all Group 2 metals.
92
How do melting points change down Group 2?
Melting points decrease down the group.
## Footnote
This is due to weakening metallic bonding as atomic size increases.
93
What causes the weakening of electrostatic attractive forces in Group 2 metals?
The distance between positive ions and delocalized electrons increases.
## Footnote
This results in weaker attractive forces.
94
What happens to ionization energies as one moves down Group 2?
First and second ionization energies decrease down the group.
## Footnote
This is due to outermost electrons being further from the nucleus and more shielded.
95
Define first ionization energy.
Energy needed to remove an electron from each atom in one mole of gaseous atoms.
## Footnote
Represented by the equation: H(g) → H⁺(g) + e⁻.
96
What is the second ionization energy?
Enthalpy change when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge.
## Footnote
Represented by the equation: Ti⁺(g) → Ti²⁺(g) + e⁻.
97
How does reactivity change in Group 2 metals?
Reactivity increases down the group.
## Footnote
This is due to increased atomic radii and more shielding.
98
What occurs when Group 2 metals react with oxygen?
They burn in oxygen.
## Footnote
Mg burns with a bright white flame producing MgO.
99
What is the reaction equation for magnesium burning in oxygen?
2Mg + O₂ → 2MgO.
## Footnote
MgO is a white solid with a high melting point due to ionic bonding.
100
What happens to magnesium ribbon when exposed to oxygen?
It often has a thin layer of magnesium oxide on it.
## Footnote
This layer forms slowly and needs to be cleaned before reactions.
101
What is the significance of cleaning magnesium ribbon before reactions?
An uncleaned Mg ribbon would give a false result in reaction rate tests.
## Footnote
Both Mg and MgO would react but at different rates.
102
What is the reaction equation for magnesium reacting with hydrochloric acid?
Mg + 2HCl → MgCl₂ + H₂.
## Footnote
This illustrates the reaction of magnesium with acid.
103
What is the reaction equation for magnesium oxide reacting with hydrochloric acid?
MgO + 2HCl → MgCl₂ + H₂O.
## Footnote
This shows the reaction of magnesium oxide with acid.
104
Name three properties of metals
- strong metallic bonds
- high electrical conductivity
- high melting and boiling points
105
Describe the melting and boiling points of metals
melting point depends on the strength of the metallic bond, strong attraction results in more metals having high melting and boiling points
106
Are metals soluble?
No, most metals do not dissolve
107
Describe giant covalent structure
Simple molecular lattice structure - held together by weak intermolecular forces, so have low melting and boiling points
Have a giant covalent lattice for giant structures
108
Describe giant covalent compounds properties
High melting and boiling point - lots of energy needed to break the bonds
solubility - insoluble
electrical conductivity - do not conduct electricity other than graphite and graphene
109
What is the trend in reactivity down group 7?
Reactivity decreases
110
Why does boiling point increase down group 7?
more electrons
more shells
stronger London forces
more energy required to break the intermolecular forces
so boiling point increases
111
What colour does aqueous chlorine go in water in a displacement reaction?
pale green
112
What colour does aqueous bromine go in water in a displacement reaction?
orange
113
What colour does aqueous iodine go in water in a displacement reaction?
brown
114
What colour does aqueous chlorine go in cyclohexane solution in a displacement reaction?
pale green
115
What colour does aqueous bromine go in cyclohexane solution in a displacement reaction?
orange
116
What colour does aqueous iodine go in cyclohexane solution in a displacement reaction?
violet
117
Graph for melting point across period 3
118
why do the melting points and boiling points rise across the three metals in period 3?
The change on the metal ions increase from +1 to +3
This means the number of delocalised electrons increases, so the strength of the metallic bond increases
So more energy is needed to break the stronger metallic bonds so the melting points and the boiling points increase
119
Why does silicon have the highest melting point on the period 3 graph of melting and boiling points?
Silicon is a metalloid
It has a giant covalent structure exactly the same as carbon in diamond
Silicon has a high melting point due to the strong covalent bonds that need breaking in order to melt it - this requires a lot of energy to break
120
Why do the last 4 elements across period 3 have the lowest melting points?
They are non metals
P, S + Cl exist as simple molecules with strong covalent bonds between their atoms
Their melting points are low because when they are melted or boiled it is their London Forces that are broken - these are very weak bonds so little energy is needed to overcome them
121
When does the strength of London Forces increase?
Increases with the number of electrons
122
Phosphorous molecule
P4
123
Sulphur molecule
S8
124
Chlorine molecule
Cl2
125
Argon molecule
A (monatomic)
126
What is the reactivity trend down group 2?
Reactivity increases down the group
127
Why does reactivity increase down group 2?
Shielding increases
nuclear charge increases
nuclear attraction decreases
atomic radius increases
so it is easier to lose electrons
128
What are the observations of a group 2 metal reacting with water?
fizzing / bubbling
metal disappears
heat given out (exothermic)
129
What happens to the pH when a group 2 metal reacts with water?
pH increases
130
Equation for metal + water
metal + water --> metal hydroxide + hydrogen
this is a redox reaction
131
What are the observations of a group 2 metal reacting with an acid?
Bubbling / fizzing
metal disappears
132
What happens to the pH when a group 2 metal reacts with acid?
pH increases
133
Equation for metal + acid
metal + acid ---> salt + hydrogen
this is a redox reaction
134
What are the observations of a group 2 metal reacting with oxygen?
bright white light
colour change
135
What happens to the pH when a group 2 metal reacts with oxygen?
pH increases
136
Equation for metal + oxygen
metal + oxygen ---> metal oxide
this is a redox reaction
137
What are the observations of a group 2 metal carbonate reacting with acid?
fizzing / bubbling
metal carbonate disappears
138
What happens to the pH when a group 2 metal carbonate reacts with an acid?
pH increases
139
Equation for a metal carbonate reacting with an acid
metal carbonate + acid ---> salt + carbon dioxide
this is not a redox reaction
140
What is produced when an oxide of group 2 reacts with water?
Hydroxide ions are released forming a metal hydroxide which is alkaline
141
What is the trend in solubility of group 2 metal oxides?
solubility increases down the group
pH increases
142
What are some uses of group 2 compounds?
antacids for treating indigestion - magnesium hydroxide and calcium carbonate
barium sulphate used in barium meals
calcium hydroxide used to neutralise soil in agriculture
143
What is the trend in reactivity down group 7?
decreases
144
Why does reactivity decrease down group 7?
more electrons
more shielding
decrease in nuclear attraction
larger atomic radius
nuclear charge increases
it is harder to gain an electron
145
What is the trend in melting points and boiling points down group 7?
they increase
146
Why do the melting and boiling points increase down group 7?
more electrons
stronger London forces
more energy needed to break the forces
147
What is the trend in electronegativity down group 7?
decreases down the group
148
Why does electronegativity decrease down group 7?
due to an increase in atomic radius
149
What is a displacement reaction?
when a more reactive halogen displaces a less reactive halogen from its compound
150
What is the ionic equation for chlorine displacing iodine in a compound?
2I-(aq) + Cl2(aq) ---> 2Cl-(aq) + I2(aq)
151
What is the ionic equation for chlorine displacing bromine in a compound?
2Br-(aq) + Cl2(aq) ---> 2Cl-(aq) + Br2(aq)
152
What is the ionic equation for bromine displacing iodine in a compound?
2I-(aq) + Br2(aq) ---> 2Br-(aq) + I2(aq)
153
How many displacement reaction does chlorine carry out out of the 3 halogens?
2 out of 3
154
How many displacement reaction does bromine carry out out of the 3 halogens?
1 out of 3
155
How many displacement reaction does iodine carry out out of the 3 halogens?
0 out of 3
156
What acts as the reducing agent in a reaction?
the substance loosing electrons - therefore the substance that is oxidised
157
What acts as the oxidising agent in a reaction?
the substance that is gaining electrons - therefore the substance that is reduced
158
What is the trend of the halogens as oxidising agents down the group?
their ability as oxidising agents decreases down a group
159
Why does the halogens oxidising ability decrease down the group?
more shielding
increased nuclear charge
a decreased nuclear attraction
so more difficult to accept electrons
160
What is the trend in the reducing ability of halide ions?
the ability of the halides as reducing agents increases down the group
161
Why does the ability of halide ions to be reducing agents increase down the group?
more shielding
increased nuclear charge
decreased nuclear attraction
easier to lose / donate electrons
162
Why are some people against the chlorination of water?
lack of individual freedom - you cannot opt out of it
163
What two acids are formed when small, amounts of chlorine is added to water?
hydrochloric acid and chloric acid
Cl2 + 2H2O --->HClO + HCl
this is a disproportionation reaction as chlorine is oxidised and reduced
164
What are the conditions needed for the reaction with chlorine and water?
sodium hydroxide has to be COLD AND AQEOUS
165
What is the method for testing for carbonate ions?
1. Add dilute nitric acid to the solid or solution
2. Bubble any gas made through limewater - observation will be fizzing and limewater turns cloudy if it is a positive test
166
What is the method for testing for sulphate ions in a stand alone test?
1. Add dilute hydrochloric acid
2. Follow with aqueous barium chloride solution - positive result is a white precipitate formed
167
What is the method for testing for sulphate ions in a series of tests?
1. Add nitric acid and aqueous barium nitrate
positive test is a white precipitate is formed (solid barium sulphate)
168
What is the test for halide ions?
1. Add dilute nitric acid
2. Followed by aqueous silver nitrate solution
positive tests -
Cl- = white precipitate
Br- = cream precipitate
I- = yellow precipitate
169
What is a further test to test for chloride ions?
Add aqueous dilute ammonia
only chloride ions will redissolve and will turn solution colourless
bromine and iodide do not react with aqueous ammonia
170
What is a further test to test for bromide ions?
Add concentrated ammonia
only bromide ions will dissolve and become colourless
iodide ions do not react
171
What is the order you should complete a series of tests in?
CaSH
carbonate
sulphate
halide
172
What is the test for ammonium ions?
1. Add aqueous sodium hydroxide
2. Gently warm with a bunsen
3. Hold damp pH paper above test tube
positive result - will turn blue
half equation - NH4+(aq) + OH-(aq) ----> NH3(g) + H2O(l)
173
174
What is produced when magnesium reacts with steam?
Magnesium oxide and hydrogen
## Footnote
The reaction is represented as: Mg (s) + H2O (g) → MgO (s) + H2 (g)
175
How do the group 2 metals react with cold water?
With increasing vigour down the group to form hydroxides
## Footnote
Examples include:
* Ca + 2 H2O (l) → Ca(OH)2 (aq) + H2 (g)
* Sr + 2 H2O (l) → Sr(OH)2 (aq) + H2 (g)
* Ba + 2 H2O (l) → Ba(OH)2 (aq) + H2 (g)
176
What do group 2 metals form when they react with acids?
A salt and hydrogen
## Footnote
Examples include:
* Ca + 2 HCl (aq) → CaCl2 (aq) + H2 (g)
* Sr + 2 HNO3 (aq) → Sr(NO3)2 (aq) + H2 (g)
* Mg + H2SO4 (aq) → MgSO4 (aq) + H2 (g)
177
What is the product when magnesium reacts with warm water?
Magnesium hydroxide and hydrogen
## Footnote
The reaction is represented as: Mg + 2 H2O → Mg(OH)2 + H2. This reaction is slower than with steam.
178
What observable signs indicate a reaction of group 2 metals with water?
Fizzing, metal dissolving, solution heating up, and white precipitate appearing
## Footnote
These effects become more vigorous down the group.
179
What happens when barium metal reacts with sulfuric acid?
It reacts slowly due to barium sulfate forming a barrier
## Footnote
The reaction is represented as: Ba + H2SO4 → BaSO4 + H2.
180
How do group 2 oxides react with water?
To form hydroxides of varying solubility
## Footnote
Examples include:
* CaO (s) + H2O (l) → Ca(OH)2 (aq) pH 12
* MgO (s) + H2O (l) → Mg(OH)2 (s) pH 9
181
What is the solubility characteristic of magnesium hydroxide?
Slightly soluble in water
## Footnote
This results in a pH of around 9 due to fewer free OH ions.
182
What is magnesium hydroxide used for in medicine?
To neutralise excess acid in the stomach and to treat constipation
## Footnote
The reaction is: Mg(OH)2 + 2 HCl → MgCl2 + 2 H2O.
183
What is lime water and its significance?
An aqueous solution of calcium hydroxide used to test for carbon dioxide
## Footnote
The reaction is: Ca(OH)2 (aq) + CO2 (g) → CaCO3 (s) + H2O (l). It turns cloudy when CO2 is present.
184
What is the effect of excess calcium hydroxide in soil?
Soils can become too alkaline to sustain crop growth
## Footnote
Calcium hydroxide is used to neutralise acidic soils.
185
True or False: Magnesium hydroxide is highly soluble in water.
False
## Footnote
It is classed as partially soluble.
186
Fill in the blank: Group 2 oxides are ______ as they accept H+ ions.
basic
187
What are the diatomic molecules of halogens?
Fluorine (F2), Chlorine (Cl2), Bromine (Br2), Iodine (I2)
## Footnote
All halogens exist as diatomic molecules.
188
What is the physical state and color of Fluorine?
Very pale yellow gas
## Footnote
Fluorine is highly reactive.
189
What is the physical state and color of Chlorine?
Greenish, reactive gas
## Footnote
Chlorine is poisonous in high concentrations.
190
What is the physical state and color of Bromine?
Red liquid that gives off dense brown/orange poisonous fumes
## Footnote
Bromine is a liquid at room temperature.
191
What is the physical state and color of Iodine?
Shiny grey solid that sublimes to purple gas
## Footnote
Iodine can change from solid to gas without becoming liquid.
192
What trend occurs in melting and boiling points of halogens down the group?
Increase down the group
## Footnote
Larger molecules have more electrons leading to stronger induced dipole-dipole forces.
193
What type of intermolecular forces increase down the group of halogens?
Induced dipole-dipole forces (London forces)
## Footnote
More energy is required to break these forces.
194
What is the electronic configuration of group 7 elements?
Outer shell s2p5
## Footnote
They often react by gaining one electron in redox reactions.
195
What happens in displacement reactions involving halogens?
A more reactive halogen displaces a less reactive halogen from its compound
## Footnote
The reactivity of halogens decreases down the group.
196
Which halogens can Chlorine displace?
Bromide and iodide ions
## Footnote
Bromine can only displace iodide ions.
197
What observation is made when Chlorine is added to potassium chloride?
Very pale green solution, no reaction
## Footnote
Indicates Chlorine's inability to displace itself.
198
What observation is made when Chlorine is added to potassium bromide?
Yellow solution, Cl has displaced Br
## Footnote
The color indicates the presence of free Chlorine.
199
What observation is made when Chlorine is added to potassium iodide?
Brown solution, Cl has displaced I
## Footnote
Indicates Chlorine's higher reactivity compared to Iodine.
200
What observation is made when Bromine is added to potassium iodide?
Purple solution, Br has displaced I
## Footnote
The color indicates the presence of free Bromine.
201
What color does the organic solvent layer show when Chlorine is present?
Colourless
## Footnote
Indicates the presence of free Chlorine.
202
What color does the organic solvent layer show when Bromine is present?
Yellow
## Footnote
Indicates the presence of free Bromine.
203
What color does the organic solvent layer show when Iodine is present?
Purple
## Footnote
Indicates the presence of free Iodine.
204
Why is Chlorine more reactive than Bromine?
Chlorine is smaller and less shielded, allowing it to gain an electron more easily
## Footnote
The outermost shell of chlorine is more attracted to the nucleus.
205
What is the reaction equation for Chlorine displacing Bromine?
Cl (aq) + 2Br- (aq) → 2Cl- (aq) + Br2 (aq)
## Footnote
This reaction shows Chlorine's higher reactivity.
206
What is the reaction equation for Chlorine displacing Iodine?
Cl (aq) + 2I- (aq) → 2Cl- (aq) + I2 (aq)
## Footnote
Demonstrates Chlorine's ability to displace Iodine.
207
What is the reaction equation for Bromine displacing Iodine?
Br2 (aq) + 2I- (aq) → 2Br- (aq) + I2 (aq)
## Footnote
Shows Bromine's capability to displace Iodine.
208
What is disproportionation?
A reaction where an element simultaneously oxidises and reduces.
209
What is the reaction of chlorine with water?
Cl2(g) + H2O(l) → HClO(aq) + HCl(aq)
210
What oxidation states does chlorine change to in the reaction with water?
Changes from 0 in Cl2 to -1 in HCl and +1 in HClO.
211
What happens to a universal indicator when chlorine is added to water?
Turns red due to acidity, then turns colourless as HClO bleaches the colour.
212
What is chlorine used for in water treatment?
To kill bacteria.
213
What are the benefits of using chlorine in water treatment?
Kills bacteria, outweighs risks of toxic effects and chlorinated hydrocarbons.
214
What is the reaction of chlorine with cold dilute NaOH solution?
Cl2(aq) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H2O(l)
215
What is the product of the reaction of chlorine with cold NaOH?
A mixture of NaCl and NaClO (sodium chlorate (I)).
216
What is the use of the mixture of NaCl and NaClO?
Used as bleach and to disinfect/kills bacteria.
217
What happens when hot sodium hydroxide is used with chlorine?
A different disproportionation reaction occurs forming sodium chlorate(V).
218
What is the reaction of chlorine with hot sodium hydroxide?
3Cl2 + 6NaOH → NaClO + 5NaCl + 3H2O
219
What is the purpose of adding nitric acid in the halide ion test?
To react with any carbonates present to prevent formation of Ag2CO3.
220
What precipitate do chlorides produce in the reaction with silver nitrate?
A white precipitate.
221
What is the reaction of silver nitrate with chlorides?
Ag+(aq) + Cl-(aq) → AgCl(s)
222
What precipitate do bromides produce in the reaction with silver nitrate?
A cream precipitate.
223
What is the reaction of silver nitrate with bromides?
Ag+(aq) + Br-(aq) → AgBr(s)
224
What precipitate do iodides produce in the reaction with silver nitrate?
A pale yellow precipitate.
225
What is the reaction of silver nitrate with iodides?
Ag+(aq) + I-(aq) → AgI(s)
226
What happens to silver chloride when treated with dilute ammonia?
Dissolves to form a complex ion: AgCl(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + Cl-(aq)
227
What happens to silver bromide when treated with concentrated ammonia?
Dissolves to form a complex ion: AgBr(s) + 2NH3(aq) → [Ag(NH3)2]+(aq) + Br-(aq)
228
Does silver iodide react with ammonia?
No, it is too insoluble.
229
How can you test for the presence of a carbonate?
Add any dilute acid and observe effervescence.
## Footnote
Fizzing due to CO2 would be observed if a carbonate was present.
230
What happens when you bubble gas through limewater to test for CO2?
It will turn limewater cloudy.
## Footnote
The reaction is: Ca(OH)2 + CO2 → CaCO3 (s) which is the cloudiness observed.
231
What is the reaction equation for testing carbonates with hydrochloric acid?
2HCl + Na2CO3 → 2NaCl + H2O + CO2.
## Footnote
This reaction produces carbon dioxide, indicating the presence of carbonate.
232
How can you test for the presence of sulfate ions?
Add acidified BaCl2 solution as a reagent.
## Footnote
A white precipitate forms if sulfate ions are present.
233
What is the reaction that occurs when barium chloride is added to a solution containing sulfate ions?
Ba^2+(aq) + SO4^2-(aq) → BaSO4(s).
## Footnote
The formation of BaSO4 indicates the presence of sulfate ions.
234
Why is sulfuric acid not used to acidify the mixture when testing for sulfate ions?
Because it contains sulfate ions which would form a precipitate.
## Footnote
This would lead to false results.
235
What is the procedure for testing halide ions?
Make the test solution acidic with nitric acid, then add silver nitrate solution dropwise.
## Footnote
The nitric acid prevents the formation of insoluble carbonate precipitates.
236
What precipitate is formed when chloride ions are present?
A white precipitate of AgCl is formed.
## Footnote
The reaction is: Ag^+(aq) + Cl^-(aq) → AgCl(s).
237
What precipitate is formed when bromide ions are present?
A cream precipitate of AgBr is formed.
## Footnote
The reaction is: Ag^+(aq) + Br^-(aq) → AgBr(s).
238
What precipitate is formed when iodide ions are present?
A pale yellow precipitate of AgI is formed.
## Footnote
The reaction is: Ag^+(aq) + I^-(aq) → AgI(s).
239
What happens to silver chloride when treated with dilute ammonia?
It dissolves to form a complex ion, producing a colourless solution.
## Footnote
The reaction is: AgCl(s) + 2NH3(aq) → [Ag(NH3)2]^+(aq) + Cl^-(aq).
240
What happens to silver bromide when treated with concentrated ammonia?
It dissolves to form a complex ion, producing a colourless solution.
## Footnote
The reaction is: AgBr(s) + 2NH3(aq) → [Ag(NH3)2]^+(aq) + Br^-(aq).
241
Does silver iodide react with ammonia?
No, it does not react because it is too insoluble.
## Footnote
Silver iodide remains as a solid precipitate.
242
What is the correct sequence of tests for anions?
Carbonate, sulfate, then halide.
## Footnote
This sequence prevents false results from BaCO3 and Ag2SO4.
243
How can you test for ammonium ions (NH4^+)?
By reacting with warm NaOH(aq), forming NH3 gas.
## Footnote
Ammonia gas can be identified by its pungent smell or by turning red litmus paper blue.
