Chem Flashcards
(214 cards)
1
Q
Explain why MgO has a higher melting point than NaCl.
A
Mg²⁺ and O²⁻ have higher charges compared to Na⁺ and Cl⁻.
Mg²⁺ and O²⁻ are also smaller ions, allowing stronger electrostatic attraction.
This results in a stronger ionic bond in MgO, requiring more energy to break, leading to a higher melting point
2
Q
Why is Al³⁺ smaller than N³⁻, even though they are isoelectronic?
A
Both Al³⁺ and N³⁻ have 10 electrons, but Al³⁺ has more protons (13) than N³⁻ (7).
The greater nuclear charge in Al³⁺ pulls the electrons closer to the nucleus, making it smaller.
3
Q
Why do ionic compounds conduct electricity when molten but not when solid?
A
In the solid state, ions are fixed in place and cannot move.
In the molten state, the ionic lattice breaks down, and ions are free to move and carry charge, allowing the compound to conduct electricity.
4
Q
Describe how the migration of ions can be demonstrated in electrolysis.
A
In electrolysis, positive ions (cations) migrate towards the negative electrode (cathode), and negative ions (anions) migrate towards the positive electrode (anode).
This movement provides evidence for the existence of ions.
As only ions can move this way in an electrical field
5
Q
Why does ionic radius increase as you go down Group 1?
A
As you go down the group, each element has more electron shells.
This increases the distance between the nucleus and the outermost electron, making the ion larger.
6
Q
What two factors determine the strength of an ionic bond?
A
Ionic charge: Higher charges lead to stronger electrostatic attraction.
Ionic radius: Smaller ions allow stronger attraction because electrons are closer to the nucleus
7
Q
Why are ionic compounds generally soluble in water?
A
Water is a polar solvent with partial positive and negative charges.
The positive end of water molecules attracts the negative ions, and the negative end attracts the positive ions, pulling the ions apart and dissolving the compound.
8
Q
Arrange the following isoelectronic ions in order of increasing size: N³⁻, O²⁻, F⁻.
A
The order is F⁻ < O²⁻ < N³⁻.
Although they all have the same number of electrons (10), N³⁻ has the fewest protons and therefore the weakest attraction between the nucleus and electrons, making it the largest.
9
Q
What is meant by molar volume of gas
A
Space occupied by one mole of a gas at a specific pressure and temperature
10
Q
What is metallic bonding?
A
Metallic bonding is the electrostatic attraction between positively charged metal ions (cations) and a “sea” of delocalized electrons that can move freely throughout the metal structure.
11
Q
Describe the structure of metallic bonds
A
In metals, cations are arranged in a regular lattice surrounded by delocalized electrons. This arrangement allows the cations to stay in fixed positions while electrons move freely, contributing to the metal’s strength and conductivity.
12
Q
What factors affect the strength of metallic bonds?
A
Charge on the metal ions (greater charge means stronger bonds).
Metallic radius (smaller radius leads to stronger bonds).
Structure of the metallic lattice (closer-packed structures have stronger bonds).
13
Q
How does metallic bond strength influence melting points and density?
A
Metals with stronger metallic bonds have higher melting points and greater densities due to tighter packing of ions and more delocalized electrons.
14
Q
Why do Group 1 metals have lower densities and melting points compared to Group 2 metals?
A
Group 1 metals lose one electron (ns¹), resulting in a lower charge and weaker metallic bonds. Group 2 metals lose two electrons (ns²), creating smaller ions with higher charges, leading to stronger bonds, higher density, and higher melting points.
15
Q
How does the melting temperature of metals change across a period and down a group, and why?
A
Across a period, melting temperature increases due to smaller radii and more delocalized electrons, resulting in stronger bonds. Down a group, melting temperature decreases as larger radii and electron shielding reduce the attraction between ions and delocalized electrons.
16
Q
What makes metals malleable, and how does their structure support this property?
A
Metals are malleable because layers of metal ions can slide over each other without breaking bonds, as delocalized electrons prevent strong repulsive forces between the ions in adjacent layers.
17
Q
Why do metals have low electronegativity
A
Metals are more likely to loose electrons to form positive ions ,so are unlikely to attract electrons
18
Q
What two factors determine the strength of ionic bonds?
A
Charge on the ions: Higher charges increase the attraction between ions.
Ionic radius: Smaller ions have a stronger attraction due to closer proximity, leading to stronger ionic bonds.
19
Q
How does ionic radius change down a group and across a period?
A
Down a Group: Ionic radius increases as additional electron shells are added.
Across a Period (Isoelectronic Series): For ions with the same electron configuration, radius decreases with increasing nuclear charge, as greater attraction pulls electrons closer.
20
Q
Explain the trend in ionic radius for the isoelectronic series from
N
3
−
N
3−
to
A
l
3
+
A
In the series
N
3
−
,
O
2
−
,
F
−
,
N
a
+
,
M
g
2
+
,
A
l
3
+
N
3−
,O
2−
,F
−
,Na
+
,Mg
2+
,Al
3+
, all ions have the same electron configuration (2,8). As nuclear charge increases from
N
3
−
N
3−
to
A
l
3
+
Al
3+
, the radius decreases due to greater attraction pulling electrons closer to the nucleus.
21
Q
Why does
M
g
O
MgO have a higher melting point than
N
a
C
l
NaCl?
A
O
MgO has a higher melting point because
M
g
2
+
Mg
2+
and
O
2
−
O
2−
ions have higher charges compared to
N
a
+
Na
+
and
C
l
−
Cl
−
, leading to stronger electrostatic attractions. Additionally,
M
g
2
+
Mg
2+
and
O
2
−
O
2−
ions are smaller, increasing the attraction.
22
Q
How do charge and ionic radius affect bond strength in ionic compounds?
A
Smaller ions with higher charges can pack closer together, increasing electrostatic attraction, leading to stronger ionic bonds. For example,
M
g
O
MgO has stronger bonds than
N
a
C
l
NaCl due to smaller ion size and higher charge.
23
Q
Why do ionic compounds conduct electricity only when molten or dissolved in water?
A
In solid form, ions in ionic compounds are fixed in place by strong bonds and cannot move. When molten or dissolved, the bonds are broken, allowing ions to move freely and conduct electricity.
24
Q
Describe the experiment using copper chromate to demonstrate ion migration.
A
Copper chromate (
C
u
C
r
O
4
CuCrO
4
) dissolved in dilute
H
C
l
HCl is layered in a U-tube with electrodes connected. When an electric current is applied,
C
u
2
+
Cu
2+
ions migrate to the cathode (turning the solution blue-green), and
C
r
O
4
2
−
CrO
4
2−
ions move to the anode (turning the solution yellow), proving the ionic nature as only ions migrate in an electrical feild
25
How does water’s polarity affect the solubility of ionic compounds?
Water's polarity, with partially positive hydrogen and partially negative oxygen, attracts and separates ionic compounds. This interaction breaks the ionic bonds, causing ions to disperse and dissolve in water.
26
What are electron density maps, and how do they indicate bonding type?
Electron density maps show regions in a molecule where electrons are most likely found. In ionic compounds like
N
a
C
l
NaCl, contour lines don’t overlap, indicating no shared electron density between
N
a
+
Na
+
and
C
l
−
Cl
−
, which confirms ionic bonding,rather than covelant bonding
27
What is a covalent bond, and between which types of elements does it typically form?
A covalent bond is a strong electrostatic attraction between the nuclei of two atoms and a shared pair of electrons. It typically forms between non-metal atoms with an electronegativity of less than about 1.5 units.
28
What is a dative covalent (coordinate) bond, and how does it differ from a standard covalent bond?
A dative covalent bond forms when one atom provides both electrons for a shared pair. It behaves like a regular covalent bond but requires an electron-deficient atom to accept the lone pair from the donor atom.
29
Describe the Octet Rule in covalent bonding
The Octet Rule states that atoms bond covalently to achieve the electron configuration of a noble gas, typically eight electrons in their outer shell, leading to more stability.
30
Explain the concept of orbital overlap in covalent bonding, and name the types of overlaps.
Covalent bonding involves the overlap of atomic orbitals each containing a single unpaired electron. Types of overlaps include:
s-s overlap (e.g.,
H
2
H
2
)
s-p overlap (e.g.,
H
C
l
HCl)
p-p overlap (e.g., between two
p
p-orbitals in different atoms).
31
What is a sigma (σ) bond, and why is it strong?
A sigma (σ) bond is formed by "head-on" overlap of orbitals along the line between the centers of bonded atoms. It is the strongest type of covalent bond because the electron density is concentrated directly between the nuclei.
32
Why does carbon form four covalent bonds instead of the expected two?
Carbon undergoes electron promotion, where one electron from the 2s orbital is promoted to the empty 2p orbital, giving carbon four unpaired electrons. This allows carbon to form four covalent bonds (e.g.,
C
H
4
CH
4
).
33
Why is electron promotion in carbon energetically favorable?
The energy required for electron promotion is small and is more than compensated by the energy released when forming four bonds, making it favorable for carbon to form four bonds instead of two.
34
How does phosphorus expand its octet to form five covalent bonds?
Phosphorus can promote an electron from the 3s to the empty 3d orbital, allowing it to have five unpaired electrons. This enables phosphorus to form compounds like
P
C
l
5
PCl
5
with five bonds.
35
Why can phosphorus form
P
C
l
5
PCl
5
but not
P
I
5
PI
5
?
The
P
−
I
P−I bond is weaker than the
P
−
C
l
P−Cl bond because iodine is larger and forms weaker bonds with phosphorus, releasing less energy, which is insufficient to compensate for the energy required to promote an electron.
36
Describe the difference between a single bond, double bond, and triple bond in terms of sigma (σ) and pi (π) bonds.
Single bond: Consists of one sigma (σ) bond from head-on orbital overlap.
Double bond: Consists of one sigma (σ) and one pi (π) bond; the pi bond is formed by sideways overlap of p orbitals.
Triple bond: Consists of one sigma (σ) and two pi (π) bonds.
37
Explain the concept of a sigma (σ) bond and a pi (π) bond.
A sigma (σ) bond forms through head-on overlap of orbitals directly between two nuclei, while a pi (π) bond forms from the sideways overlap of p orbitals above and below the bond axis, only occurring in double or triple bonds.
38
Why is a pi (π) bond generally weaker than a sigma (σ) bond?
A pi (π) bond is weaker because it involves the sideways overlap of p orbitals, resulting in electron density that is less directly between the nuclei compared to the head-on overlap in sigma (σ) bonds.
39
Can atoms rotate freely around a double or triple bond? Why or why not?
No, atoms cannot rotate freely around double or triple bonds because the pi (π) bonds restrict rotation. Breaking a pi bond would be required to allow rotation, which requires energy.
40
No, atoms cannot rotate freely around double or triple bonds because the pi (π) bonds restrict rotation. Breaking a pi bond would be required to allow rotation, which requires energy.
At high temperatures, the pi (π) bond may temporarily break, allowing rotation around the double bond. Once cooled, the structure typically returns to its original state with the pi bond intact.
41
Which bonds allow rotation
Single as only made of a sigma bond,but not double or triple as pi bond restricts rotation
42
Why are geometric isomers stable at room temperature
Geometric isomers have restricted rotation around double bonds due to pi (π) bonding, making the arrangement of atoms fixed and stable at room temperature.
43
How is covalent bond strength measured?
Covalent bond strength is measured by the amount of energy required to break the bond in a gaseous molecule.
44
What factors influence covalent bond strength?
The sum of the atomic radii of the bonded atoms, Smaller atoms and multiple shared electron pairs lead to stronger bonds.
45
Why does diamond have a high melting point?
Diamond has a giant covalent structure where each carbon atom is covalently bonded to four others in a 3D lattice, requiring a large amount of energy to break all covalent bonds.
46
Describe the bonding in silicon dioxide (SiO₂).
Each silicon atom is covalently bonded to four oxygen atoms, and each oxygen atom is covalently bonded to two silicon atoms, forming a 3D lattice similar to diamond.
47
How do the melting points of diamond and SiO₂ compare, and why?
Diamond has a higher melting point than SiO₂ due to the strong C-C bonds in its structure, which are stronger than the Si-O bonds in silicon dioxide.
48
Why is the F-F bond weaker than expected for such a small atom?
The F-F bond is weaker than expected because the small size of fluorine atoms brings lone pairs close together, leading to significant repulsion between these electrons. This repulsion weakens the bond, making F-F less strong compared to other single bonds between small atoms.
49
What is bond energy?
Bond energy is the amount of energy required to break one mole of a bond in a gaseous molecule. It reflects the bond's strength and is influenced by factors like the size of the atoms and the number of shared electron pairs.
50
What determines the strength of ion-ion interactions in ionic compounds?
The strength depends on the charge of the ions (higher charge = stronger attraction) and the ionic radius (smaller ions = stronger attraction). These are non-directional forces that contribute to the reason ionic compounds form crystal lattice structures as there and no specific angles
51
How does bond length affect the strength of covalent bonds?
Shorter bonds are generally stronger because the nuclei of the atoms are closer, increasing electrostatic attraction. Longer bonds are weaker due to increased distance.
52
Explain why smaller metal ions form stronger bonds in a metallic structure compared to larger metal ions.
Smaller ions have more delocalised electrons which is closer to positive charged nucleus ,more compact lattice so electrostatic force of attraction is greater .
53
Describe ion-dipole interactions and provide an example where this interaction is significant.
Ion-dipole interactions occur between a charged ion and a polar molecule. For example, the attraction between Na
+
+
ions and water molecules is crucial for dissolving salts in water.
54
What causes permanent dipole-dipole forces, and between which types of molecules do they occur?
They are intermolecular forces that occur between polar molecules as they contain atoms of different electronegativities which leads to a distribution of electrons densities
55
How do London forces arise, and why are they stronger in larger atoms or molecules?
London forces arise from temporary charge imbalances within a molecule, creating an instantaneous dipole that induces a dipole in a nearby molecule. They are stronger in larger atoms or molecules because they have more electrons, which are more easily polarized.
56
Explain ion-water (polar) interactions
Water is a polar molecule positive metal ions would be attracted to slightly negative oxygen in water as one of the lone pairs in oxygen forms a dative covalent bond with empty orbital in metal. and negative ions would be attracted to positive hydrogen in water ,the water makes ions hydrated
57
What are London forces?
London forces are weak, temporary attractions between molecules, caused by constantly shifting temporary dipoles due to random electron movement.
58
How does the number of electrons in a molecule affect the strength of London forces?
Molecules with more electrons have stronger London forces, as seen in iodine (solid) compared to chlorine (gas) at room temperature.
59
Why does it take more energy to overcome dipole-dipole forces compared to London forces?
Dipole-dipole interactions are generally stronger because they are permanent, while London forces are temporary and weaker.
60
How does molecular shape and contact points influence London forces
Molecules with more contact points experience stronger London forces. For example, long molecules like pentane have stronger forces than spherical molecules like neopentane.
61
How does molecular shape and contact points influence London forces
Molecules with more contact points experience stronger London forces. For example, long molecules like pentane have stronger forces than spherical molecules like neopentane.
62
Why do noble gases have low boiling and melting points?
Noble gases have fewer electrons and lack bonding electrons, resulting in very weak London forces between atoms.
63
What is hydrogen bonding and what elements are typically involved?
Hydrogen bonding is a strong type of intermolecular force, occurring between a δ⁺ hydrogen in one molecule and a lone pair of electrons on a small electronegative atom (usually fluorine, oxygen, or nitrogen) in another molecule.
64
Why can't hydrogen bond with chloride ions despite both having δ⁻ regions?
Chloride ions are too large, which makes hydrogen bonds ineffective, even though chlorine has a δ⁻ charge.
65
Why does water (H₂O) have a higher boiling point than hydrogen fluoride (HF)?
Water forms more hydrogen bonds per molecule due to two lone pairs on oxygen, making it stronger overall in terms of hydrogen bonding than HF
66
What happens when we heat a substance that has covalent bonds
Molecules move faster and intermolecular forces may break,its not the covalent bond that breaks, as intermolecular forces are much weaker than covalent bonds
67
When will an ionic compound conduct electricity
When molten/liquid or when dissolved in water
68
When talking about metallic bonds or ionic bonds what do you have to make sure to mention in questions
Ionic radius , charge on cation
69
What is electronegativity
The power of an atom to attract the shared pair of electrons in a covalent bond
70
What is the most electronegative element
Fluorine (4.0), small radius
71
Define bond polarity
When 2 atoms in a covalent bond have different electronegativities the bond electrons are attracted to the atom with a higher electronegativity creating a polar bond , this results in the more electronegative atom becoming delta negative and the less electronegative becomes delta positive
72
What does a molecules polarity depend on
It’s 3D structure and the polarity of its bonds
73
Define dipole moment
A measure of positive and negative charges in a molecule
74
Why do some molecules that have polar bonds are known as non polar overall
Their dipole moments cancel out so the molecule overall is non polar
75
How can i test for polarity
Run liquid near a charged rod or balloon ,polar liquids are attracted to charge as they have a dipole moment,non polar liquids are not effected
76
I placed an unknown substance by a charged rod and the substance was attracted to the rod ,what can u tell me about the substance
It is polar as it reacts to charge due to its dipole moment
77
What must the difference in electronegativity be for non polar covalent bonds,polar covalent bonds and ionic bonds
Non polar covalent, 0-0.4
Polar covalent, 0.5-1.7
Ionic, 1.7 and above
78
Why won’t a covalent bond with a difference in electronegativity of 0.4 dissolve in water?
Because its non polar covalent so won’t dissolve in water which is polar
79
The difference in electronegativity for a bond is 1.8 ,will this molecule dissolve in water?explain answer.
Yes as this is an ionic bond, and ionic bonds dissolve in water as water is polar and interacts with the positive and negative charges present in ionic bonds causing ionic bonds to dissolve in water
80
Define polymeric
Long chains of molecules bonded covalently with weak forces between chains
81
What is the structure of diamond
Giant covalent , has a continuous network of atoms covalently bonded in a lattice, each carbon atom forms four covalent bonds in a tetrahedral arrangement,creating a very hard 3D lattice, high melt/boil points due to strong covalent bonding
82
Explain the structure of graphite
Structure , each carbon atom bonds to 3 others in hexagonal layers, with a fourth delocalised electron ,that allows electrical conductivity, properties -layers slide over each other ,so its soft and useful as lubricant
83
Why would graphite make a good lubricant
Because hexagonal layers in gains covalent lattice can slide over each other as held by weak London forces making it soft and ideal as a lubricant ,it can also remain stable at high temps
84
Describe fullerenes
Form spherical shapes with hexagonal and pentagonal rings ,they can trap atoms or ions in structure , so important in medical application
85
Name a fullerene and explain its structure and properties
Carbon nanotube -cylindrical structure of carbon atoms in hexagonal arrangement,known for strength and electrical conductivity ,can be used in drug delivery , a cluster of nanotubes is very very strong ,they have delocalised electrons
86
Explain on reason why ice is less dense than water
Open structure, more space between molecules less dense ,where as in water molecules ore compact
87
Explain silicons structure
They bond a giant covalent structure ,each silicon is bonded to another 4 silicon atoms covalently in a tetrahedral structure ,creating a strong lattice ,which contributes to high melt/boil points
88
Explain the structure of SIC. (Silicon carbide)
Each silicon atom is surrounded by four carbon atoms tetrahedrally,which in turn are surrounded by four silicon atoms ,creating a giant covalent lattice network of alternating carbon and silicon atoms
89
Explain SIO2 quartz
Giant covalent lattice where each silicon atom is bonded to four oxygen atoms, and each oxygen atom is bonded to 2 silicon atoms
90
As you go down the group of noble gases what happens to the boiling point
It increases as atoms get larger and have more electrons so Greater London forces so more energy is required to break bonds , hence higher melt/boil
91
What do branches do to the boiling point of alkanes
Branches cause a lower boiling point compared to straight alkanes that have the same molecular formula, as branching reduces surface contact between molecules reducing London forces, so less energy required to overcome bonding.
92
What would I draw for skeletal formula
Carbon skeleton but don’t draw carbon atoms or hydrogen atoms ,include other atoms though
93
Why is c3h6 not empirical
Not the simplest ratio of atoms of each element present in compound
94
What the start of the name be fir molecules with these number of carbons
Meth-one
Eth-two
Prop-three
But-four
Pent-five
Hex-six
Hept-seven
Oct-eight
Non-nine
Dect -ten
95
Homologous series
Families of organic compounds with the same functional group and same general formula,all with same chemical properties ,each member differs by ch2 ,there is a gradual trend in physical properties
96
What is the importance of homologous series
Homologous series help chemists predict properties and reactions of organic compounds based on known patterns.
They simplify the study of organic chemistry by grouping compounds with similar properties together.
97
What homologous series have no double carbon bonds
Alkanes
98
If you are given a table of results and asked to find the mean titre what results to we take into accounts
Only the concordat ones which are in 0.1 of each other
99
The student rinses the burette with distilled water before adding in sodium hydroxide solution what effected will the have on titre
It will increase the titre as sodium hydroxide will be more dilute
100
What’s the importance of percentage yield
Idea of getting as much product in a reaction as possible
101
What type of reaction occurs when ammonia gas reacts with hydrogen chloride gas?
Acid base, as ammonia has the ability to except the H+ from hydrogen chloride to become NH4+
102
What makes something a base
If it has the ability to accept an H+ ion
103
Explain why both water and carbon dioxide molecules have polar bonds but only
water is a polar molecule.
Oxygen is more electronegative than hydrogen and carbon.
This creates polar bonds: oxygen δ⁻, carbon/hydrogen δ⁺.
CO₂ is linear and symmetrical, so dipoles cancel.
Water is bent (V-shaped), so dipoles do not cancel.
104
How can the strength of a covalent bond be demonstrated
Through the high melting points of giant covalent structures such as diamond and graphite ,they contain high boiling points due to many covalent bond so lots of energy required to break bonds
105
Why might a lower boiling temperature for water over hydrogen sulphide be assumed even though wrong
Because hydrogen sulphide has more electrons so is more polarisable than water so may be assumed that due to London forces being stronger in hydrogen sulphide it would have a higher boiling temperature than water, however water has hydrogen bonds which are stronger than London forces meaning water has a higher boiling point
106
List the dif in electronegativity and type of molecule formed
0-0.4-non polar covalent,no dipole
0.5-1.7 polar covalent,permanent dipole
1.7 and above -ionic
107
Explain why 2,2-dimethylpropane has a much lower boiling temperature than its isomer
pentane.
Detailed descriptions of the forces involved are not required.
(2-2 shows its branched) branching results in weaker forces as there are less points of contact
108
This question is about some halogens and their compounds.
The intermolecular attractions between halogen molecules are London forces.
(i) Describe how London forces form between halogen molecules.
Random movement of electrons forms instantaneous dipole which induces dipole on neighbouring molecule
109
Explain why hydrogen bonding causes ice to be less dense than liquid water.
More open structure ,due to 3D lattice ,the hydrogen bonds are longer than the covalent bonds
110
Why is ethanol soluble in water
Ethanol is soluble in water because the hydroxyl group (-OH) in ethanol can form hydrogen bonds with water molecules. Hydrogen bonding is a strong intermolecular force, and water is highly polar, which allows it to form hydrogen bonds with ethanol. These interactions make ethanol highly compatible with water, increasing its solubility.
111
Why is cl2 got a lower boiling point than silicon
-silicon giant covalent structure
Cl2-simple molecular
-even though London forces are greater in cl2,the strong covalent bonds in si are stronger than London forces
-more energy required to break strong covalent bonds in si than London forces in cl2
112
Why does more lone pairs decrease bond angle
More lone pairs means more repulsion from the element with the lone pairs which reduces bond angle
113
State the type of bond that joins the two AlCl3 molecules together.
Dative covalent bond
114
Describe the bonding in platinum.
Metal cations in sea of delocalised electrons ,strong electrostatic attraction between cations and delocalised electrons
115
Describe the bonding in the element chromium and use your answer to justify why it has
such a high melting temperature.
Regular arrangement of positive ions,in sea of delocalised electrons , strong forces of attraction between cations and electrons, a lot of heat energy required to break metallic bonds
116
Empirical formula
Simplest whole number ratio of atoms of each element in a compound
117
Molecular formula
Number of atoms of each element present in one molecule of a substance
118
Structural formula
Least detail of how atoms are arranged in molecule
119
Displayed formula
Shows all covalent bonds present in a molecule
120
Homologous series
Families of organic compounds with same functional group and general formula same chemical properties ,gradual trend in physical properties ,each member differ by CH2
121
How do alkanes differ from other homologous series
Have no functional group and
122
If molecules are being reduced what side of half equation are electrons on
LEFT
123
What do oxidation number tell us
How many electrons are being gained/lost in reaction
124
What oxidation number to elements have
0
125
In ionic compounds what do overall oxidation numbers have to add to
Overall ionic charge
126
What is a redox reaction
When oxidation and reduction both occur in reaction
127
2 reasons Why are alkanes very unreactive
Because they are non polar as H and C have similar electronegativities ,bonds are also very strong in alkanes so unreactive
128
What is a hydrocarbon
Compound made of hydrogen and carbon only
129
How is a fraction extracted from crude oil?
oil is heated and vaporised,gas rises up column ,condense at dif BP ,longer chains will condense at bottom which is hottest part of column as more energy required to break strong London forces in larger molecules ,shorter chains at top where its cooler less energy to break London forces in small molecules ,similar molecules condense together
130
Why do we use cracking for longer chains
Longer chains are less flammable so less useful so we use cracking to break C-C bonds to make shorter chain alkanes so more flammable more useful to meet fuel demand
131
What bonds break when cracking and when vapourising
Vapour-breaks intermolecular forces not bonds
Cracking breaks covalent bonds
132
What does cracking produce
Alkane and alkene
133
What do we do with alkene left over from cracking
Used as Polymer
134
Catalytic converters?
These remove CO, NOx and unburned hydrocarbons (e.g. octane,
C8H18) from the exhaust gases, turning them into ‘harmless’ CO2, N2
and H20 releasing them into atmosphere
135
Pollutant environmental consequences
Nitrogen oxides (formed when N2 in the air reacts at
the high temperatures and spark in the engine)
NO is toxic and can form smog
NO2 is toxic and acidic and forms acid rain
Carbon monoxide toxic to humans reduce oxygen transport
Carbon dioxide Contributes towards global warming
Unburnt hydrocarbons (not all fuel burns in the
engine)
Contributes towards formation of smog
Soot/particulates Global dimming and respiratory problemst
136
Describe structure of catalytic converter how its structure helps it to convert pollutants
Converters have a ceramic
honeycomb coated with a thin
layer of catalyst metals
platinum, palladium, rhodium
– to give a large surface area so pollutants can be converted quickly
137
What is a disproportionation reaction
When single element undergoes oxidation and reduction same time in reaction
138
What is reforming
The process of turning straight hydrocarbons into branched chains alkenes and cyclic hydrocarbons for efficient combustion
139
What is reforming
The process of turning straight hydrocarbons into branched chains alkenes and cyclic hydrocarbons for efficient combustion
140
What are the limitations of free radical substitution
Impurities -due to multiple termination steps
Further substitution -more than one halogen is substance onto a carbon
Substitution of dif c atoms-the halogen gets substituted onto dif c atoms
141
What is a free radical
Reactive specie that has an unpaired electron
142
Oxidising and reducing agent
Oxidising agent-gain electrons become reduced they cause other substances to be oxidised
Reducing agent-loose electrons and become oxidised cause other substances to be reduced
143
Dif between branched and straight hydrocarbons
Straight-have no side groups but both can have double bonds
144
Advantages of biofuels
Considered carbon neutral -as plants grow produce o2 which equals amount of c02 released when burned.
Biodiesel and biogas-reduce waste in landfill as waste used to produce them.
Biofuel-provide money for less developed countries who have space for crops required
Short time to produce
Renewable
145
Disadvantages of biofuels
Cost of converting engines and machinery to run of biofuel rather than petroleum/diesel
Many developed countries don’t have space for enough crop to reach demand as space needed for food
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Main 3 biofuels
Biodiesel -made from refining renewable fats and oils
Bio ethanol -made by fermentation
Bio gas-made when organic waste breaks down
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A sample of copper contains the two isotopes 63Cu and 65Cu only. It has a relative
atomic mass, Ar, less than 64. The mass spectrum of this sample shows major peaks with
m/z values of 63 and 65, respectively
Suggests that 63cu is more abundant than 65cu Ar will be skewed towards 63CU
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63cu2+ has smaller peak than 63cu+ why
More energy required to remove second electron
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In mass spec how are ions detected and how is abundance measured
Positively charged ions are deflected to detector
Causing current to flow
Size of current at detector depends on number of ions hitting it
Allowing mass spectrometer to measure relative abundance of each isotope
150
State the difference, if any, in the chemical properties of isotopes of the same
element
Isotopes have same protons so same electron configuration so no difference in chemical properties
151
Define the term atomic number of an element.
number of protons in one atom or nucleus
152
Why is it necessary to ionise atoms before acceleration?
Ions must be charged to be accelerated by electric field
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What deflects the ions?
Magnetic field
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What is adjusted in order to direct ions of different mass to charge ratio onto the
detector?
Magnetic field
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What does mass spectrometer measure
Relative abundance of isotopes
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During mass spec Why do the Kr+ ions from this sample of krypton separate into four paths?
Different isotopes
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What deflects the moving ions round a curved path?
Magnet
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How can we reduced deflection of ions in mass spectrometer
Reduce magnetic field
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What factor, other than the mass to charge ratio of an ionised particle, determines how
much that particle is deflected in a magnetic field of a given strength?
Speed
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In a question if bromine is oxidised and reduced however not 2 bromine products are formed only one is it a disproportination reaction
No
161
Explain why alkanes in crude oil can be separated by fractional distillation
Alkanes have different boiling points depending on the lengths of the chain so can be separated as they condense at different boiling points
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What can oh groups form and what does this mean in terms of energy
OH groups can form hydrogen bonds so more energy required to break intermolecular forces
163
Explain the two major reasons for cracking hydrocarbons
Shorter chain alkenes and alkanes formed
Alkenes can be used to make polymers
Alkanes can be used as fuels to meet high demand
Shorter chain alkanes are more useful as More flammable so are high in demand
164
Some isomer of hexane have lower boiling points why
They are branched so have fewer contact points weaker London forces less energy to overcome intermolecular forces
165
In the initiation step, chlorine molecules are converted into radicals.
Identify the type of bond broken and the type of bond fission occurring in this step
Sigma bond breaks
And type of bond fission in homolytic
166
Describe the stages by which the reaction of NO and CO turns into co2 and n2 occurs in a catalytic converter
Absorption of gases to catalytic surface
Weakening of bonds on catalytic surface
Release or products from catalytic surface
167
Deduce two possible reasons why the density of iron (7.86 g cm−3) is much greater
than the density of graphite (2.2 to 2.8 g cm−3).
Iron atoms have a greater mass than carbon atoms
Irons atoms pack closer together in graphite compares to carbon atoms
168
Describe the key feature of the bonding of the carbon atoms in graphite that results in it
being an electrical conductor.
Delocalised electrons within layer to carry current each carbon is covalently bonded to three other carbons
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Graphene, graphite and diamond are all forms of solid carbon.
Explain, in terms of structure and bonding, why graphene and graphite are good
electrical conductors but diamond is a poor electrical conductor
Graphene-single layer of hexagonal sheets and has delocalised electrons to carry charge
Graphite -has layers and each carbon is bonded to 3 others ,so has delocalised electrons to carry charge ,in diamond each carbon atom is bonded to four other carbon atoms in tetrahedral structure ,diamond has no delocalised electrons to carry charge.
170
Explain why the melting temperature of silicon(IV) oxide is much higher than that of iodine,
even though the bonding in both is covalent.
Silicon oxide is giant covalent so has Many strong covalent bonds whereas iodide is simple covalent weak London forces ,more energy required to break stronger bonds in silicon oxide hence higher boiling point
171
Explain why diamond has a much higher melting temperature than iodine
Diamond -giant covalent lattice structure every carbon bods to 4 carbons in tetrahedral arrangement held by strong covalent bonds,iodine is simple covalent,weak intermolecular forces ,more energy required to break stronger bonds in covalent bonds in diamond compared to weak intermolecular forces in iodine
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What is homolytic fission
When a covalent bond breaks into two and each atom receives electron from the bond (initiation step)
173
Relative mass of electron
1/1840
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How can mass spec identify elements
By looking at spectrum can identify elements by unique isotopes and relative abundance
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What does Relative atomic mass take into account
Weighted mass of element taking into account isotopes and relative abundance
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What do i write if asks for species for a peak on mass spec
If asked to give the species for a peak
in a mass spectrum then give charge
and mass number e.g. 24Mg+
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R.A.M formula
Sum of (isotopic mass x percentage abundance)/100
178
What are fragments ,why may there be smaller peaks ,why are fragmentation patterns useful
During ionisation molecules can break down into smaller fragments ,these fragments produce peaks with lower m/z ratio,these fragmentation patterns help identify structure of molecule,can tell what parents of molecule are stable ,shows clues for bonding and functional groups ,helps identify M+ peak
179
What does the tallest peak refer to and what does the peak with largest m/z ratio refer to
Tallest-most abundant ion ,the taller means easiest to form, most commonly formed fragment and most stable during fragmentation
Molecular ion peak-largest m/z ratio,shows mr (relative molecular mass)represents the entire molecule with one electron removed
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Successive ionisation energy
Energy required to remove and electron from an atom or ion after one or more electrons have already been removed
181
Why are successive ionisation energy’s always larger
The ion is being removed from a more positively charged ion so the electrons tic attraction between the nucleus and outer electrons increases and shielding decreases so electrons are more bound by nucleus harder to remove electron so more energy would be required
182
Why does boron have lower first ionisation energy than beryllium if its further to the right of the periodic table
Beryllium’s outermost electron is in its 2s sub shell which is closer to the nucleus and more tightly bound compared to boron which has outermost electron in 2p sub shell which is at a higher energy level and further away from the nucleus and therefore experiences more shielding by the full s sub shell so less energy required to remove electron from boron
183
Why does oxygen have lower first ionisation energy than nitrogen even though oxygen is closer to right
In nitrogen electrons in 2p sub shell are all unpaired so experience less repulsion so harder to remove electrons tan in oxygen which has a pair of electrons in one of the p orbitals,these experiencing more repulsion making it easier to remove less energy required to do so
184
What does a large jump in ionisation energy indicate
Removal of electron from closer shell
185
Why is second initiation energy so much greater than first for sodium
After the one electron is removed from 3s sub shell to remove another one would change to 2p shell which is full ,and has more effective nuclear charge
186
If there is big jump in ionisation energy between 2nd electron removed and 3rd what group is element in
Group 2 as the removal of one more electron shows change in sub shell from shell closer to nucleus hence why more energy needed to remove electron
187
Why is there a drop in ionisation energy for mg-al and for p-s
Mg-al the outset electron for mg is in 3s sub shell whereas for mg it has started to fill 3p sub shell which is further from nucleus and not as bound, so less energy required to remove electron from 3p than full 3s sub shell
P-s =sulphur has 4 electrons in 3p orbital whereas p has only 3 so wont experience as much repulsion as in s which has one pair of electrons in 3p shell so easier to remove electron due to repulsion
188
What do orbitals represent
The mathematical probability of finding an electron within certain distances of the nucleus ,they all have approximate 3D shape ,can never 100 percent know where electron is and can’t draw orbital shape precisely
189
How can i tell what a translation metal is based on configuration
A transition metal is an element that:
Has a partially filled d-orbital (d¹ to d⁹) in either its atom or ion form.
This means the d-orbital cannot be empty like D0 or full like D10 can’t be D5 4p as 4p comes after d5,
190
Group 2 metals similar to 1
Explain dif in reactivity how mg reacts vs BA
Alkaline metals
Increase reactivity down group -more shield,larger radius
Last electron on s orbital
Density and melting point decreases -metallic bond weakens as size increases and increase distance from nucleus and delocalised electrons electrostatic force weakens less energy to break metallic bonds
1st Ionisation energy decreases,more likely to loose electron rather than gain,despite increase nuclear charge.
Increase solubility down group
Ph increases down group
Hydroxides of group 2 - Charge density decreases down group so solubility increases due to lower attraction between hydroxide and larger 2+ ions,the ions split up easier ,will result on more OH- ions in water so PH increases
BA more reactive than MG both in same group
Group 2 metals burn in oxygen MG burns with bright white flame
BA burns with an apple green flame
Be won’t react with water ,Mg-reacts slowly with cold water but quickly in steam
BA-vigorous reacts with cold water
Oxides of group 2 are ionic solids except BeO which has covalent character
Group 2 and 1 sulphates- solubility decreases down group
191
Why is it important to ionise isotopes in mass spec
So that they can b accelerated by electrical field, when the ions hit detector the can produce a current,so can be detected
192
How will i know which isotope hits the detector first
The one with the lowest m/z ratio
193
How does mass spec separate species into dif curves
Ions are accelerated by electrical field and ions are deflected by magnetic field based on m/z ratio the ones with lower m/z ratio are deflected more,this separates the ions and curves them on to dif paths
194
How does mass spec separate species into dif peaks based on m/z ratio
During ionization (often by electron impact), the molecule absorbs energy and becomes unstable, breaking into smaller charged fragments.
These fragments have distinct
m/z values, producing different peaks in the mass spectrum
The fragments are deflected based on m/z value
195
How can strength of covalent bond be demonstrated
By the high melting points of covalent structures like diamond and graphite where a lot of heat is required to break the many strong covalent bonds
196
Electronegativity is measured on what scale
Pauling scale which ranges from 0-4
197
What is a symmetrical molecule
Where no loan pairs all bonds are the same,these bonds won’t be polar even if individual bonds are as dipole moment cancels out
198
Why is ammonium chloride ionic
Ammonium ion is a positive charge cl is negative
199
What does it mean if shape is planar
It’s 2d no 3D shape like tetrahedral
200
Why is an na+ ion smaller than f-
they are isoelectronic so have same configuration however NA has more protons than F- so stringer pull on outer electrons from nucleus making na smaller
201
What is an isomer
Molecules with the same molecular formula but different arrangement of atoms within the molecule
202
Structural isomer
Have same molecular formulae but different structural arrangement of atoms within can be straight branched etc but will always have same molecular formula
203
Position isomer
Same functional group but in different position of chain e.g propane-1-ol and propane-2ol but have same molecular formula
204
Functional group isomers
Same molecular formula but atoms are arranged differently
205
Sterioisomers
Same structural formula but different spatial arrangement of atoms
206
What is E-Z isomerism
In molecules with double carbon bond the pi bond restricts rotation this leads to two different spatial arrangement depending on how the groups are positioned around the double bond
207
E
The 2 higher priority groups are on opposite sides of double carbon bond
208
Z
The 2 priority groups are on the same side of the double carbon bond
209
Cis ,trans
Cis means same side ,trans means opposite sides ,only works when where hydrogen bonds are there to compare to ,u can’t prioritise groups like in e-z
210
Homolytic fission
Each atom gets one electron from covalent bond that breaks to form two free radicals
211
Heterolytic fission
One atoms gets both electrons
212
213
What does the ionisation energy of boron go against the general trend for period two elements
In boron electron is removed from 2p subshell unlike Be and Li where it’s removed from the 2S ,the 2p sub shell is further away from the nucleus and therefore there is more shielding blocking the electron from the nucleus force so it’s easier to loose electron that’s why there is a drop in ionisation energy of boron which goes against the trend that ionisation energy increases from left to right.
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