Module 3.1 Flashcards
1
Q
Give an outline of the periodic table
A
- Consists of rows (periods- horizontally across)
- Consists of groups
- Arranged by increasing atomic number
2
Q
In the periodic table, elements are arranged by …
A
Increase atomic number
3
Q
What groups are fond in the s block?
A
Groups 1 and 2
4
Q
What groups are found in the p block?
A
Groups 3-8
5
Q
What are the names of the elements found in block d?
A
Transition metals
6
Q
Define the term periodicity
A
Trends that occur in physical and chemical properties as we move across the periods of the periodict table
7
Q
When talking about periodicity, list 5 different trends that we talk about
A
- Ionisation energy
- Melting Points/Boiling Points
- Reactivity
- Atomic Radius
- Electronegativity
8
Q
Do all the elements within a group have similar reactions, why?
A
Elements in the same group have the same number of electrons on their outer shell. Therefore, they have similar chemical properties.
9
Q
What was Döbereiner’s theory about?
A
Triads
10
Q
What was John Newlands’ theory about?
A
Law of octaves
11
Q
What was Mendeleev’s theory about?
A
Gaps, elements arranged in order of increase atomic mass
12
Q
What is the modern day theory of the periodic table?
A
Elements are arranged in order of increasing atomic number
13
Q
Explain Döbereiner’s theory behind his model of the periodic table
A
- Döbereiner grouped similar elements in TRIADS e.g. Li,Na and K.
- He ordered elements by atomic mass.
- The middle element in his triads had SIMILAR properties to the other TWO elements.
14
Q
Explain John Newlands’ theory behind his model of the periodic table
A
- Newlands arranged elements in order of mass. He noticed that every 8th element was SIMILAR.
15
Q
Explain Mendeleev’s theory behind his model of the periodic table
A
- He arranged all known elements by atomic mass
- He left GAPS, where no element fitted the repeating patterns
- Predicted the patterns of missing elements
16
Q
Explain the theory behind the modern day periodic table ( Henry Moseley)
A
- Elements are arranged by increasing atomic number
- There are groups and periods
17
Q
Define first ionisation energy
A
The amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms.
18
Q
Give the equation for the first ionisation energy of oxygen
A
O(g) → O^+ (g) + e^-
19
Q
Give the equation for the first ionisation energy of magnesium
A
Mg(g) → Mg^+ (g) + e^-
20
Q
Give the equation for the first ionisation energy of Sodium
A
Na(g) → Na^+ (g) + e^-
21
Q
When talking about first ionisation energy what must you include?
A
- Charges
- Gaseous state
- The amount of energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms
22
Q
The lower the ionisation energy, the easier/harder it is to form an ion?
A
Easier
23
Q
What factors affect ionisation energy?
A
Nuclear charge
Atomic radius
Shielding
24
Q
How does nuclear charge affect ionisation energy?
A
The higher the nuclear charge, the larger the electrostatic attraction between the nucleus and the outermost electrons. Therefore, more energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so the as nuclear charge increase, so does IE.
Higher nuclear charge = Higher Ionisation energy
25
How does the atomic radius affect ionisation energy?
A increase in atomic radius means that the nuclear attraction between outermost electrons and the nucleus is smaller. Therefore, less energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so as the atomic radius increases, ionisation energy falls.
Higher atomic radius = Lower Ionisation energy
26
How does the shielding affect ionisation energy?
As the number of electrons between the outermost electrons and the nucleus increases (As shielding increases), the outer electrons feel less attraction towards the nucleus. Therefore, less energy is required to remove 1 mole of electrons from 1 mole of gaseous atoms, so as shielding increases, ionisation energy falls.
More shielding = Lower ionisation energy.
27
Does Ionisation energy increase or decrease down the group, and why?
Down the group ionisation energy DECREASES. This is because elements down the group have extra electron shells compared to the ones above. The extra shells mean that the atomic radius is larger, so the outer electrons are further away (increasing shielding) from the nucleus, which greatly reduces their attraction to the nucleus.
28
Even though as you go down the group, the positive charge of the nucleus increases, why does ionisation energy still fall?
The effect caused by the positive charge is OVERRIDDEN by the effect of an increase in atomic radius and extra shielding
29
Does Ionisation energy increase or across the period, and why?
Across the period, ionisation energy INCREASES.
This is because across the period, the positive charge of the nucleus increases. This causes electrons to be pulled closer to the nucleus, making the atomic radius smaller. This means that the outer electrons are more strongly attracted to the nucleus. Therefore, more energy will be needed to ionise the element meaning across the period ionisation energy increases
30
Define the term second ionisation energy
The energy needed to remove an electron from each of one mole of 1+ ions in a gaseous state.
31
List reasons why ionisation ionisation energies higher than the one before?
- Less repulsion (Greater nuclear attraction (due to a fall in atomic radius))
- Positive nuclear charge outweighing the effect of the electron every time the an electron is removed
- Atomic radius decrease
32
Explain why successive ionisation ionisation energies higher than the one before?
- Less repulsion (remaining electrons drawn slightly closer to the nucleus)
- Positive nuclear charge outweighs the negative effect every time an electron is removed
- As the distance of each electron from the nucleus decreases slightly, the nuclear attraction increases. Therefore, more energy is needed to remove each successive electron
33
Define Metallic bonding
The strong electrostatic attraction between positive ions and delocalised electrons.
34
Give an overview of metallic bonding
- Sea of delocalised electrons
- Giant lattice of positive metal ions
- Strong electrostatic forces between positive ions and sea of free electrons.
35
List the properties of metallic bonding
- High melting point/High boiling point
- Can conduct electricity (when solid AND when molten
- Malleable
- Insoluble
36
Do metals have a high melting/boiling point or a low one? Why?
Compounds with metallic bonding have a HIGH melting/boiling point.
This is because the strong electrostatic attractions between the positive ions and the delocalised electrons require a lot of energy to overcome.
37
Can metals conduct electricity? Why?
Compounds with metallic bonding can conduct electricity when solid AND when molten. This is because the DELOCALISED ELECTRONS can move and carry charge.
38
Are metals malleable? Why?
Yes there are.
This is because as there are no bonds holding specific ions together, the metal ions can slide past each other when the structure is pulled.
39
Are metals soluble or insoluble?
Insoluble.
Metals are insoluble except in LIQUID METALS. This is because of the strength of metallic bonds.
40
What is the structure formed by metals?
Giant metallic lattice
41
What is meant by a giant metallic lattice?
A lattice of cations fixed in position surrounded by a 'sea' of delocalised electrons
42
What must you do when drawing metallic lattices?
- Balance the charges over the whole structure (e..g. if its a 2+ metal, have 2 delocalised electrons per 2+ ion)
- Draw at-least 3 rows
- Show charges on ions
- Show the delocalised electrons
43
Define covalent bonding
(insert)
44
What is the name of the structure that diamond, graphite, graphene and silicon form?
Giant COVALENT lattices
45
What is meant by a giant covalent lattice?
Huge networks of covalently bonded atoms
46
Outline the structure of diamond
A form of carbon atoms where each atom forms 4 other carbon atoms around it.
47
List the properties of diamond
- Very high melting point
- Can't conduct electricity
- Very hard
- Insoluble
48
Explain why diamond has a very high melting point?
The extremely strong covalent bonds require a lot of energy to melt. As there are many strong covalent bonds in diamond, a lot (state the temperature, which will probably be given in the question) of energy is needed to melt the strcuture.
49
Explain why diamond can't conduct electricity
All the outer electrons are held up in localised bonds
50
Outline the structure of graphite
Graphite is a form of carbon which forms a giant covalent lattice where carbon atoms are arranged in sheets.
51
List the properties of graphite
- Very high melting point
- Can't conduct electricity
- Very soft
- Insoluble
52
Explain why graphite has a very high melting point
The extremely strong covalent bonds in the sheets require a lot of energy (state the temperature, which will probably be given in the question) to melt.
53
Explain why graphite can conduct electricity
The delocalised electrons in graphite can move along the sheets and carry charge
54
Explain why graphite is soft
The sheets can slide over each other
55
Explain why graphite is insoluble
The covalent bonds in the sheets are too strong to break
56
Outline the structure of graphene
Graphene is a sheet of carbon atoms joined together in hexagons. The sheet is just one atom thick, making it a 2D compound
57
List the properties of graphene
- Very high melting point
- Conducts electricity
- Insoluble
58
Explain why graphene has a very high melting point
The extremely strong covalent bonds in the sheets require a lot of energy (state the temperature, which will probably be given in the question) to melt.
59
Explain why graphene can conduct electricity
The delocalised electrons in graphene are free to move along the sheet
60
What is better at conducting electricity, graphite or graphene?
Graphene
61
Explain why graphene is insoluble?
The covalent bonds in the sheets are too strong to break
62
When asked about the melting/boiling point of simple covalent structures, what must you talk about?
Weak induced dipole-dipole forces
63
Why do simple covalent structures have low melting points/boiling points?
When melted/boiled, the weak induced dipole-dipole forces are broken. These induced dipole-dipole forces are weak and can easily be overcome, so elements with this type of structure tend to have low melting/boiling points.
64
When asked why do other simple molecular structures have higher melting/boiling points than other simple molecular structures, what do you say?
X has a great melting/boiling point than why because it has more atoms. More atoms in a molecule mean stronger induced dipole-dipole forces. This results in a higher boiling point and melting point as more energy is required to overcome these 'stronger' induced dipole dipole forces.
65
Explain the variation in melting points across Periods 2 and 3 for metals ONLY
For the metals, the melting point increases across the period because the metallic bonds get stronger, the ionic radius decreases and the number of delocalised electrons increases
66
Explain the variation in melting points across Periods 2 and 3 for giant covalent elements ONLY
These elements have giant covalent lattice structures linking all their atoms together. Aa lot of energy is needed to break these bonds
67
Explain the variation in melting points across Periods 2 and 3 for simple molecular elements ONLY
The elements that form simple molecular structures have only weak intermolecular forces. These IMFs, require little energy to overcome, so they have low melting points and boiling points
68
Explain the variation in melting points across Periods 2 and 3 for group 8 elements ONLY
The noble gases have the lowest melting point in their periods. This is because they are held together by the weakest forces.
69
How many electrons do group 2 elements have in their outer shell?
2
70
Down group 2, what is the trend for reactivity?
Down group 2 reactivity increases.
71
Why does reactivity increase down group 2?
As you go down group 2, the IE decreases. This is due to the increasing atomic radius and shielding effect. When group 2 elements react, they lose their electrons. The lower, the ionisation energy, the easier it is to lose electrons. Therefore, the easier it is to lose electrons, the more reactive the metals. This means that reactivity increases down the group
72
When group 2 elements react with water, what is produced?
A metal hydroxide (salt) and hydrogen
73
Ca (s) + 2H2O (l) →
Ca(OH)2 + H2
74
When group 2 elements react with oxygen, what is produced?
A metal oxide
75
2Ca(s) + O2(g) →
2CaO (s)
76
When group 2 elements react with a dilute acid, what is produced?
A salt and hydrogen
77
Ca(s) + 2HCl (aq) →
CaCl2 (aq) + H2(g)
78
CaO(s) + H2O (l) →
Ca(OH)2 (aq) + 2OH^-
79
What is calcium hydroxide used for?
To neutralise acidic soils
80
How does calcium hydroxide neutralise acidic soils
It reacts with acids in soil (Remember salt + water will be formed)
81
How does calcium hydroxide help farmers?
It neutralises acidic soils - increasing : crop health, plant life and crop production
82
What is Magnesium Hydroxide used for?
Treating Indigestion
83
Give a reaction that shows how Magnesium hydroxide treats indigestion
Mg(OH)2 (s) + 2HCl (aq) → MgCl2 (aq) + 2H2O (l)
84
What is calcium carbonate used for?
Treating Indigestion
85
Give a reaction that shows how Calcium hydroxide treats indigestion
CaCO3 (s) + 2HCl (aq) → CaCl2 (aq) + H2O (l) + CO2 (g)
86
Give the formula for fluorine, chlorine, bromine and iodine
F2, Cl2, Br2 and I2
87
What state is fluorine, chlorine, bromine and Iodine under standard conditions?
```
Fluorine = Gas
Chlorine = Gas
Bromine = Liquid
Iodine = Solid
```
88
Halogens exist as ________ molecules
Diatomic e.g. Br2
89
Why do halogens exist as diatomic molecules
They share an electron to give them both a full shell.
90
Down group 7, the boiling point of halogens increase because?
As we go down the group, the atomic radius increases. An increase in atomic radius makes the number and strength of induced dipole-dipole forces increase. Therefore, more energy is required to break the weak bonds and separate the molecule apart
91
How many electrons do group 7 elements have in their outer shell?
5
92
How do halogens react?
By gaining an electron in their outer shell to form 1- ions.
93
Do halogens displace less reactive halide ions?
Yes they do
94
What is the most reactive halogen?
Fluorine
95
What is the least reactive halogen?
Iodine
96
Give the full and ionic equation for the reaction between Br2 + 2KI
Full reaction = Br2 (aq) + 2KI (aq) → 2KBr (aq) + I2 (aq)
Ionic reaction = Br2(aq) + 2I^- (aq) → 2Br^- (aq) + I2 (aq)
97
When Bromine reacts with Potassium Iodide what happens and why?
The Bromine displaces the Iodine, in KI to form KBr. This is because Bromine is MORE REACTIVE than Iodine
98
Give the equation for the reaction between I2 (aq) + 2KF (aq)
I2 (aq) + 2KF (aq) → NO REACTION
Iodine isn't reactive enough to displaces the Fluorine from KF
99
Give the full and ionic equation for the reaction between Cl2 + 2KI
Full reaction = Cl2 (aq) + 2KI (aq) → 2KCl (aq) + I2 (aq)
Ionic reaction = Cl2 (aq) + 2I^- (aq) → 2Cl^- (aq) + I2 (aq)
100
Explain the trend in reactivity of halogens due the decreasing ease of forming 1- ions
As you go down the group, the atomic radius increases meaning that the outermost shell is further away from the nucleus as there are more inner shells. As there are more inner shells between the nucleus of the halogen and the outermost electron shells, shielding increases. . This makes it harder for larger atoms to attract the electron needed to form an ion, so larger atoms are LESS REACTIVE. This means that down the group, halogens becomes less oxidising.
101
Define disproportionation
The oxidation AND reduction of the same element element in a redox reaction
102
Give the equation for the reaction between Chlorine and water, what type of reaction is this?
Cl2 (g) + H2O (l) → HCl (aq) + HClO
Type of reaction = Disproportionation as Chlorine is both reduced and oxidised
103
Why do we react chlorine with water?
To purify water
104
Give the equation for the reaction between Chlorine and cold, dilute aq NaOH, what type of reaction is this?
Cl2 (g) + 2NaOH → NaClO (aq) + NaCl (aq) + H2o (l)
Type of reaction = Disproportionation as Chlorine is both reduced and oxidised
105
Give an analogous reaction between X2 and Sodium Hydroxide, what type of reaction is this?
X2 + 2NaOH (aq) → NaOX (aq) + NaX (aq) + H2O (l)
106
Why is chlorine good?
- It kills disease causing microorganisms
- Purifies water
- Some chlorine remains in water, and prevents reinfection further down the supply
- Prevents the growth of algae, eliminating bad tastes and smells from water.
107
Why is chlorine bad?
- Chlorine can react with organic matter in water such as bits of leaves to form chlorinated hydrocarbons. It has been suggested that some chemicals formed by chlorine reacting with organic matter in water may lead to the the formation of some cancers.
108
Should we chlorinate water, or nah? (Cost-Benefit- Analysis)
In conclusion, the benefits of purifying water outweighs the costs. For example, the increased risk of people getting cancer due to chlorinated hydrocarbons is small compared to the risks from untreated water, such risks include a cholera epidemic, which could kill thousands of people
109
How can we test for Halide ions?
Dissolve the suspected halide in water and add aq solution of SILVER NITRATE. Watch for a coloured precipitate
110
When testing, for halides, if the ppt formed is white, what halide ion is present?
A chloride ion
111
When testing, for halides, if the ppt formed is cream, what halide ion is present?
A bromide ion
112
When testing, for halides, if the ppt formed is yellow, what halide ion is present?
An iodine ion
113
In the halide ion test, if the colour is hard to distinguish, what do we do?
Add dilute ammonia, then add concentrated ammonia
114
Does the white ppt dissolve in dilute ammonia?
Yes
115
Does the cream ppt dissolve in dilute ammonia?
NO
116
Does the cream ppt dissolve in conc ammonia?
Yes
117
Does the yellow ppt dissolve in either dlute or conc ammonia?
NO
118
What is the ionic equation if a chloride ion is present when testing for halides?
Ag+ (aq) + Cl- (aq) → AgCl (s)
119
What is the ionic equation if a bromide ion is present when testing for halides?
Ag+ (aq) +Brl- (aq) → AgBr (s)
120
What is the ionic equation if a iodide ion is present when testing for halides?
Ag+ (aq) + I- (aq) → AgI (s)
121
When testing unknown substances, what order MUST we carry out our tests in to avoid false positives?
1) Carbonate test
2) Sulfate Test
3) Halide Test
122
How can we test for carbonates?
Add HCl to the suspected carbonate and collect any gas formed in the reaction and pass this gas through limewater
123
What is the positive test for carbonates?
Limewater will turn cloudy
124
Give the ionic equation when testing for carbonates
CO3^2- (s) + 2H+ (aq) → CO2 (g) + H2O (l)
125
How can we test for sulfates?
Add dilute HCl and BaCl2 to the suspected solution
126
What is the positive test for sulfates?
A white ppt of BaSO4 will form
127
Give the ionic equation when testing for sulpfates
Ba^2+ (aq) + SO4^2- → BaSO4
128
How can we test for ammonium compounds?
Add NaOH to the suspected ammonium compound and warm gently. Pass the gas produced through RED litmus paper
129
What is the the positive test for ammonium?
Gas produced turns RED litmus paper BLUE
130
What gas is produced when testing for ammonium coimpounds?
Ammonia
