Module 2-Foundations in Chemistry Flashcards
(86 cards)
1
Q
What is the shape and bond angle of a molecule with 4 bond pairs?
A
tetrahedral shape, 109.5 degree bond angles.
2
Q
Why do lone pairs repel more than bond pairs?
A
It is slightly closer to the central atom so repels more (it also occupies more space).
3
Q
What is the shape and bond angle of a molecule with 3 bond pairs and 1 lone pair?
A
pyramidal shape, 107 degree bond angles.
4
Q
What is the shape and bond angle of a molecule with 2 bond pairs and 2 lone pairs?
A
non-linear shape, 104.5 degree bond angles.
5
Q
What is the shape and bond angle of a molecule with 2 bond pairs?
A
linear shape, 180 degree bond angles.
6
Q
What is the shape and bond angle of a molecule with 3 bond pairs?
A
trigonal planar shape, 120 degree bond angles.
7
Q
What is the shape and bond angle of a molecule with 6 bond pairs?
A
octrahedral shape, 90 degree bond angles.
8
Q
Ammonium and sulphate ions.(1)
A
Have the same number of bonded pairs as a methane molecule so have the same shape and bond angle.
9
Q
Carbonate and nitrate ions.(1)
A
Have 3 regions of electron density surrounding the centre atom so have the trigonal planar shape.
10
Q
What effects the electronegativity of molecules?(3)
A
- the nuclear charges are different
- the atoms are different sizes
- the shared pair of electrons may be closer to one nucleus than the other.
11
Q
What is electronegativity?
A
The attraction of a bonded atom for a pair of electrons in a covalent bond.
12
Q
The larger the Pauling electronegativity value the…
How does the vary across the periodic table?
A
more electronegative the element is.
Electronegativity increases across and up the periodic table.
13
Q
Using difference in Pauling values, how do you know if a compound is ionic or covalent?(3)
A
0=covalent
0-1.8=polar covalent
1.8+=ionic
14
Q
What is a non-polar bond?When will one appear?
3
A
The bonded electron pair is shared equally between the bonded atoms.
Appear when:
-bonded atoms are the same(forming pure covalent bonds)
-have the same/similar electronegativity.
15
Q
What is a polar bond?When will one appear?
3
A
The bonded electron pair is shared unequally between the bonded atoms
Appear when:
-bonded atoms are different
-have different electronegativity values.
16
Q
Due to electronegativity,partial charges come about (delta+&-), what does this form?
A
The separation of opposite charges is called a dipole.
17
Q
What is the difference between a permanent and an induced dipole?
A
Permanent means the partial charge values remain the same at all times
Induced is when the dipole is formed due to the oscillation in the electron cloud forming partial charges.
18
Q
Explain why a water molecule is polar.(3)
A
- the two O-H bonds each have a permanent dipole
- the two dipoles act in different directions but do not exactly oppose one another
- overall, the oxygen end of the molecule is still negative whilst the h end is positive.
19
Q
Explain why carbon dioxide is non-polar despite having polar bonds.(3)
A
- the two C=O bonds each have a permanent dipole
- the 2 dipoles act in opposite directions and exactly oppose one another
- overall, the dipoles cancel out so the dipole is 0.
20
Q
3 types of intermolecular forces, label strongest to weakest.
A
- Induced dipole-dipole interactions (London forces)-weakest
- permanent dipole-dipole interactions
- hydrogen bonding-strongest.
21
Q
Explain how London forces come about.(4)
A
- Oscillation of electron cloud produces a changing dipole in a molecule
- At any instant, an instantaneous dipole will exist but its position will be constantly shifting
- The instantaneous dipole induces a dipole on a neighbouring molecule
- The induced dipole induces further dipoles on neighbouring molecules, which then attract resulting in London forces.
22
Q
What molecules tend to have London forces?(2)
A
Noble gases and Diatomic molecules.
23
Q
What effects the strength of the London forces as you encounter bigger elements,why?(3)
A
- As you go down the group there are more electrons
- this results in a greater electron cloud which results in larger instantaneous and induced dipoles
- this means more attractive forces between the molecules.
24
Q
Hydrogen chloride and fluorine molecules are relatively the same in shape, size and have the same number of electrons. Why do their boiling points drastically differ?(3)
A
- Fluorine molecules are non-polar and only have London forces (diatomic molecule)
- Hydrogen chloride molecules, however, are polar and have London forces and permanent dipole-dipole interactions between the molecules
- Extra energy is needed to break the additional permanent dipole-dipole interactions between hydrogen chloride molecules and so therefore is has a higher boiling point.
25
What is a simple molecular substance composed of?What do they form in a solid state?
Simple molecules which contain small units containing a definite number of atoms with a definite molecular formula e.g. neon, Ne, water, H20.
Simple molecular lattices.
26
What happens when a simple molecular substance is heated?
The weak intermolecular forces break-NOT the strong covalent bonds.
They have low melting and boiling points.
27
Explain the solubility of simple molecular substances.(4)
non-polar substances:tend to be soluble in non-polar solvents and insoluble in polar substances as the intermolecular forces are too hard to break
polar substances:solubility depends on the strength of the dipole (may dissolve if strong enough to attract eachother)
28
Why don't simple covalent structures conduct electricity?(1)
The have no mobile charge carriers so cannot complete an electrical circuit.
29
When do hydrogen bonds form?(2)
- An electronegative atom with a lone pair of electrons is present: oxygen, nitrogen, fluorine.
- A hydrogen atom attached to an electronegative atom:H-O, H-N, H-F.
30
Why is ice less dense than water?(3)
- H bonds hold water molecules apart in an open lattice structure
- The water molecules in ice are further apart than in water...
- As they have two lone pairs and two h atoms which form four h bonds in a tetrahedral shape with a bond angle close to 180 degrees.
- Therefore it is less dense and floats on water.
31
List the max number of electrons in the first 4 shells?
1-2
2-8
3-18
4-32. formula 2n(squared) gives the max number.
32
Atomic orbitals make up electron shells, what is an atomic orbital?
A region around the nucleus that can hold up to 2 electrons.
33
S-orbitals.(4)
- sphere shape
- can hold up to 2 electrons
- each shell from n=1 contains 1 s orbital
- the greater the number of the shell the greater the radius of the s-orbital.
34
P-orbitals.(4)
- dumb-bell shape
- 3 separate p-orbitals (perpendicular to one another) which can all contain 2 electrons
- each shell from n=2 contains 3 p-orbitals
- the greater the shell number, the greater the distance the p-orbital is from the nucleus.
35
D and F orbitals.(2)
- every shell from n=3 contains 5 d-orbitals
| - every shell from n=4 contains 7 f-orbitals.
36
Electrons pair with opposite spins, explain.(2)
- electrons are negatively charged and repel one another
- 2 electrons in an orbital therefore must have an opposite spin to counteract the repulsion between the negative charge of the 2 electrons.
37
Orbitals with the same energy (in the same sub-shell) are occupied singly first, why?(1)
-Because they have the same energy, one electron has to occupy each orbital before pairing starts to prevent repulsion between paired electrons
38
Define ionic bonding.
The electrostatic forces of attraction between positive and negative ions. It holds together cations and anions in ionic compounds.
39
Common:
- cations
- anions
```
Cations:
-metal ions
-ammonium ions
Anions:
-non-metal ions
-polyatomic ions
```
40
Melting and boiling points of ionic compounds.(2)
- most are solid at room temp and have high melting and boiling points
- a lot of energy is required to overcome the strong electrostatic forces of attraction between the oppositely charged ions within the lattice(this increases if the charges on the ions increase).
41
Solubility of ionic compounds.(2)
- Mainly dissolve in polar solvents, e.g. water, as this breaks down the lattice structure.
- If there are bigger charges on the ions then this may not be the case as the electrostatic forces of attraction may be too great.
42
Solubility in ionic compounds requires 2 main process what are they and why are they important?(3)
-ionic lattice to be broken down
-water molecules must attract and surround the ions.
Therefore, this solubility depends on the strengths of attraction between ions and water molecules and so predictions are hard to make.
43
Why do ionic compounds only conduct electricity when liquid or dissolved in water?(2)
As when they are solid there are no mobile charge carriers as the ions re in a fixed giant lattice structure position
Whereas when they are liquid the solid lattice breaks down meaning the ions are now free to move and act as these mobile charge carriers.
44
Define covalent bonding.
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
45
Common examples of when covalent bonding occurs are:___.(3)
- non-metallic elements, e.g. H2 and O2
- Compounds of non-metallic elements, e.g. H2O and CO2
- polyatomic ions, e.g. NH4+.
46
```
How many bonds do these elements form:
carbon
nitrogen
oxygen
hydrogen.(4)
```
- 4
- 3
- 2
- 1.
47
What is different about the bonding in boron (specifically BF3), what does this prove?(3)
- Boron has 3 electrons in its outer shell and forms covalent compounds
- Boron trifluoride only has 6 outer electrons
- Showing that predictions for bonding cannot be based solely on the noble gas electron structure.
48
Why can fluorides of phosphorus, sulphur, and chlorine form multiple compounds?(2)
for elements in period 2 n=2 so can only hold 8 electrons in the outer shell but for n=3 they can hold 18 electrons so they are all available for bonding so many possible combinations.
49
What is the expansion of the octet, use an example in your explanation.(2)
- SF6, six unpaired electrons from sulphur are paired so the outer shell now contains far more than the nearest noble gas Argon.
- Known as the expansion of the octet and is possible only from the n=3 sub-shell, when a d-sub-shell becomes available for the expansion.
50
What are dative/coordinate bonds?What do you use to show a dative bond in dot and cross diagrams?
A covalent bond is one in which the shared pair of electrons have been supplied by one of the bonding atoms only (i.e. a lone pair).
An arrow from the element donation to the one it is donating to.
51
What is a strong acid?
An acid which releases all its hydrogen atoms into the solution as H+ ions and completely dissociates in aqueous solution.
52
What is a weak acid?-most organic acids are weak acids.
An acid which only releases some of its available hydrogen atoms into the solution as H+ ions and only partially dissociates in aqueous solution.
53
What is an alkali?base?
A base neutralises an acid to form a salt.
| A base that dissolves in water releasing OH- ions into the solution.
54
What happens during neutralisation?
the H+ ions in the acid are replaced by metal or ammonium ions.
55
The dissociation of sulphuric acid.(3)
- has multiple h ions
- when mixed with wtaer each molecule just releases one of its 2 h ions
- this results in hs04 which acts as a weak acid only partially dissociating.
56
When is the oxidation number always 0?How do you always have to write oxidation numbers?
In elements.
| With the charge BEFORE to sign.
57
OILRIG.What does this mean in terms of the oxidation number?
Oxidation is loss (of electrons)-oxidation number increases, Reduction is gain (of electrons)-oxidation number decreases.
58
What is the avogadro constant?
6.02x10^23-amount of particles in one mole of a substance.
59
Equation linking mass(m), moles(n) and Mr(M).
m=nxM
60
Assumptions made when determining the mass and formula of a hydrate salt?(2)
- That all the water has been lost (heat to a constant mass to ensure all the water is gone)
- That there is no further decomposition (most salts decompose when heated)
61
Equation linking moles, concentration and volume.
n=cxv.
62
What is a standard solution?
A solution with a known concentration.
63
How to calculate the volume of a gas?
number of moles x 24(000-in cm3)dm3=volume inn cm3/dm3.
64
What assumptions are made on the ideal gas equation?(4)
- random motion
- elastic conditions
- negligible size
- no intermolecular forces.
65
What is the ideal gas equation?give units.
pV=nRT
```
p=pressure in Pa
V=volume in m3
n=number of moles
R=constant(8.31mol-1K-1)
T=tmeperature in Kelvin.
```
66
Predict the conditions of pressure and temp that cause the ideal gas equation to break down and explain.(4)
High pressure-has molecules are close together and the volume of the molecules becomes significant compared with the volume of the container
Low temperatures-gas molecules slow down and have less energy than higher temperatures. Intermolecular forces may then become significant.
67
Reasons why theoretical yield is never achieved.(3)
- the reaction may not have finished
- other reactions may have taken place
- purification of the product may result in loss of some product.
68
How to calculate % yield.
Actual yield/theoretical yield.
69
What is a limiting reagent?
The reactant that is no in excess as it will all be used up and hence stop the reaction.
70
What is atom economy?
The measure of how well atoms have been utilised.
71
What shows that a reaction has a high atom economy?(2)
- produce a large proportion of desired products and few waste products
- are important for sustainability as they make the best use of resources.
72
How to work out atom economy:
Sum of molar masses of desired products/sum of molar masses of all products
Using a BALANCED EQUATION.
73
What are isotopes?(3)
Atoms of the same element with different numbers of neutrons and different masses but the same number of protons (and electrons meaning they react the same).
74
Define relative isotopic mass.
The mass of an isotope relative to 1/12th of the mass of an atom of carbon-12(the standard).
75
Define relative atomic mass.
The weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12.
76
Describe the process of mass spectrometry.(4)
- a sample is placed in the mass spectrometer
- it is then vapourised and ionised to form positive ions
- the ions are then accelerated. Heavier ions move slower and are harder to deflect so ions of each isotope are seperated
- the ions are detected on a mass spectrum as a mass:charge ratio m/z. Each ion reaching the detector adds to the signals so the greater the abundance the larger the signal.
VIADD!
77
Charges of Silver and Zinc.
Ag+
| Zn2+
78
What is a binary compound?
A compound containing only 2 elements.
79
Formulas of manganate (VII) and dichromate(VI).
MnO4(-). CrO7(2-).
80
What is the formula unit?
The formula work out for ionic compounds from the the ionic charges used in equations.
81
Which bond is NOT regarded as polar?
C-H.
82
What is interesting about the oxygen in H2O2(hydrogen peroxide)?
It contains a 1- charge!
83
What is the electron configuration of Cr, Cu?why?
[Ar] 4s13d5
[Ar] 4s13d10.
It is more stable as it limits repulsion.
84
When determining oxidation numbers (particularly in compounds such as OF2) what do you have to do?
Always go for the most electronegative element.
85
What does the melting and boilin point of an ionic compound depend on?(2)
Charge of the ions (would result in a greater electrostatic attraction).
Size of the ions.
86
Why is a covalent bond described as localised?
Because unlike ionic compounds the attraction is only between the electron pair and the nuclei.
