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as-level chemistry > Foundations in Chemistry > Flashcards

Foundations in Chemistry Flashcards

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1
Q

What are isotopes?

A

Atoms of the same element either a different number of neutrons and different masses.

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2
Q

What is relative atomic mass?

A

Weighted mean mass compared with 1/12th mass of carbon-12

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3
Q

What is relative isotopic mass?

A

Mass compared with 1/12th mass of carbon-12

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4
Q

What is a mole?

A

Unit for amount of substance

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5
Q

What is Avogardo’s constant?

A

The number of particles per mole, 6.02 x 10^23

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6
Q

Definition for empirical formula

A

simplest whole number ratio of atoms of each present in a compound

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7
Q

Definition for molecular formula

A

Number and type of atoms of each element in a molecule

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8
Q

Calculation for moles

A

Moles = mass / RFM

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9
Q

Calculation for concentration

A

Conc = mole / vol (dm^3

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10
Q

Equation for volume of gas at room temp and pressure

A

Volume = mole x 24 (dm^3)

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11
Q

Ideal gas equation with units

A

Pressure x Volume = Moles x Gas constant x Temperature

Pressure = Pa Volume = m^3 n = moles (mol) R= 8.314 T= Kelvin

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12
Q

Equation for Atom Economy

A

% atom economy = Mr of desired product / Mr of reactants * 100

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13
Q

What is meant by one mole of substance?

A

The amount of substance which contains as many particles as there are carbon atoms in 12g of 12 carbon atoms

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14
Q

What are acids?

A

Release H+ ion in aqueous solution

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15
Q

What are alkalis?

A

Release OH- ions in aqueous solution

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16
Q

What makes an acid weak or strong?

A

A strong acid will more readily dissociate and will become fully ionised in water.
A strong acid will fully dissociate while a weak acid will only partially dissociate

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17
Q

What is neutralisation?

A
  • H+ and OH- to form H2O
  • Acid with bases to form salts
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18
Q

What is a standard solution?

A

A solution for analysis whose concentrations are known presisely

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19
Q

How is a standard solution made?

A
  • Weigh out a precise amount of solid
  • Add to a small volume of water and pre-dissolve the solid
  • Transfer to a volumetric flask
  • Make up to the scratch mark with more water
  • Add stopper and mix the contents
  • Rinse the beaker with distilled water and add the rinsings to the flask.
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20
Q

What are the steps in a titration?

A
  • Measuring a known volume (usually 20 or 25 cm3) of one of the solutions with a volumetric pipette and placing it into a conical flask
  • The other solution is placed in the burette
  • A few drops of the indicator are added to the solution in the conical flask
  • The tap on the burette is carefully opened and the solution added, portion by portion, to the conical flask until the indicator starts to change colour
  • As you start getting near to the end point, the flow of the burette should be slowed right down so that the solution is added dropwise
  • You should be able to close the tap on the burette after one drop has caused the colour change
  • Multiple runs are carried out until concordant results are obtained
  • Concordant results are within 0.1 cm3 of each other
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21
Q

What are atomic orbitals?

A

An orbital is a region around the nucleus which can hold a maximum of 2 electrons, with opposite spin.

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22
Q

What are the shapes of s and p orbitals?

A

S -> spherical orbital
P -> dumb bell orbital
Other orbitals are more complex

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23
Q

What are the rules for filling orbitals?

A
  • Lowest energy filled first
  • Each orbital can hold 2 electrons
  • Orbitals of the same energy are filled singly before pairing up
  • Fill the 4s sub shell before the 3d sub shell
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24
Q

Show the max configuration you have to know for a-level?

A

1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^2 4p^6 4d^10 4f^14

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25
Q

What are the exceptions and what are their configurations?

A

Cr -> 1s^2 2s^2 2p^6 3s^2 3p^6 3d^5 4s^1
Cu -> 1s^2 2s^2 2p^6 3s^2 3p^6 3d^10 4s^1

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26
Q

What is ionisation energy?

A

Energy needed to remove one electron from each atom in one mole of gaseous atoms to form 1+ ions

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27
Q

Why does the IE increase when removing electrons from the same shell?

A
  • When electrons are in the same shell each successive ionisation energy is slightly higher than the one before.
  • This is because there are the same number of protons (and same nuclear charge) attracting fewer electrons.
  • As each electron is removed there is less repulsion between the remaining electrons and so the shell is drawn in closer to the nucleus.
  • So, the electrons are attracted more strongly to the nucleus and more energy is needed to remove them.
28
Q

How is successive ionisation energy evidence for shells? (When removing electrons from a different shell)

A
  • When electrons are being removed from another shell closer to the nucleus there is a large increase in ionisation energy.
  • This is because as the electron being removed is in a new shell much closer to the nucleus, it is more strongly attracted to it and so more energy is needed to remove an electron. There is also less electron shielding.
  • This large increase in IE is evidence that shells exist.
29
Q

Describe a graph to show the first ionisation energies for the elements in period 3

A

Up, up, down, up, up, down, up, up

30
Q

Explain a drop in ionisation energy between group 2 and group 3

A
  • The outer most electrons of the group 3 element is in a p sub shell whereas the group 2 element’s outer electrons is in an s sub shell .
  • The p sub shell is slightly higher in energy than the s-sub shell so less energy is required to remove this electron as it is less strongly attracted to the nucleus
31
Q

Explain a drop in ionisation energy between group 5 and group 6

A
  • The electron being removed from the group 6 element is from an orbital containing 2 electrons (a spin-paired electron).
  • The repulsion between these paired electrons means that less energy is required to remove that electron.
32
Q

Explain the trend in ionisation energy as you go down the group

A
  • number of shells increases
  • shielding increases
  • atomic radius increases
  • first ionisation energy decreases
33
Q

What are all the different shapes of molecules ?

A
  • Linear ( 2bp)
  • Trigonal Planar (3bp)
  • Tetrahedral (4bp)
  • Octahedral (6bp)
  • Pyramidal (3bp, 1lp)
  • Non-linear (2bp, 2lp)
34
Q

What’s the bond angle and shape of a molecule with 2bp?

A

Linear
180

35
Q

What’s the bond angle and shape of a molecule with 3bp?

A

Trigonal Planar
120

36
Q

What’s the bond angle and shape of a molecule with 4bp?

A

Tetrahedral
109.5

37
Q

What’s the bond angle and shape of a molecule with 6bp?

A

Octohedral
90

38
Q

What’s the bond angle and shape of a molecule with 3bp and 1lp?

A

Pyramidal
107

39
Q

What’s the bond angle and shape of a molecule with 2bp and 2lp?

A

non-linear
104.5

40
Q

What is a dative bond?

A

Shared pair of electrons where both of the electrons have been provided by one of the bonding atoms only.

41
Q

Explain why bond angles are different when there is a lone pair?

A

Because lone pairs repel more than bonded pairs and so the angle is different

42
Q

What is electronegativity ?

A

The ability of an atom to attract the bonding electrons in a covalent bond

43
Q

What is a non - polar bond?

A

When the electrons in the bond are evenly distributed between the atoms that make up the bond.

44
Q

What is a permanent dipole?

A

Small charge difference across a bond that results in a difference in the electro-negativities of the bonded atoms. A polar covalent bond has a permanent dipole.

45
Q

What is a polar- covalent bond?

A

The electrons in the bond are not evenly distributed between atoms that make up the bond and this bond has a permanent dipole.

46
Q

What is a polar molecule?

A
  • A molecule that is asymmetrical and has atoms of different electronegativity.
  • As a result, the pull of a one electronegative atom is not cancelled out by the pull of the other electronegative atoms.
47
Q

What is a non-polar molecule?

A

A molecule that is symmetrical and so the pull of a one electronegative atom is cancelled out by the pull of the other electronegative atoms.

48
Q

What are intermolecular forces?

A
  • Individual compounds are attracted to each other by intermolecular forces
  • They occur due to the random constant movements of the electrons within the shells of the atoms in molecules
  • They are weaker than chemical bonds
    there are three types: permanent dipole-dipole induced dipole interactions, permanent dipole - permanent dipole interactions, London forces and hydrogen bonding
49
Q

Explain dipole-induced dipole interactions

A
  • If a molecule has a permanent dipole it has a slightly negative and slightly positive end
  • When it is near other non polar molecules, it is able to cause electrons in the shells in the nearby molecule to shift slightly (by being repelled or attracted by its charged ends)
  • This causes a non-polar molecule to become slightly polar and then an attraction occurs
50
Q

Explain dipole- permanent dipole interactions

A
  • Molecules with permanent dipoles will also be attracted to other molecules with permanent dipoles
  • The negative ends are attracted to the positive ends and vice versa
  • Happens between HCL molecules
51
Q

Explain London forces

A
  • Weak forces between non-polar molecules
    Caused by random movement of electrons in shells
  • Causes unbalanced distribution of charge
  • At any moment, there will be an instanteous dipole across a molecules
  • This instantaneous dipole induces a dipole in neighbouring molecules which do the same for their neighbouring molecules
  • The small induced dipoles attract one another, causing London forces
  • The size of London forces increases with increasing numbers of electrons. The greater the number of electrons, the larger the induced dipoles and the greater the attractive forces between molecules.
52
Q

Explain the effect of London forces on bpt

A
  • When a substance is boiled, heat energy is being used to overcome intermolecular forces between molecules
  • In case of non-polar molecules, there are only London forces to be overcome and as they are so weak they have low boiling points
  • As the number of electrons increases, so does the boiling point. When moving down groups, the bpt becomes greater.
53
Q

What is a hydrogen bond ?

A

A hydrogen bond is a strong permanent dipole -permanent dipole attraction between:
- An electron- deficient hydrogen atom on one molecules (0-H, N-H or F-H) and
- A lone pair of electrons on a highly electronegative atom (O, N or F ) on a different molecule

54
Q

What is ionic bonding?

A

Electrostatic force of attraction between positive and negative ions

55
Q

Describe the structure of ionic compounds

A

A giant ionic lattice structure, resulting from positively charged ions strongly attracted in all directions to negatively charged ions.
The positive and negative electrons alternate in this structure.

56
Q

Why do ionic substances have high mpt?

A

The giant ionic lattices are held together by strong electrostic forces. It takes loads of energy to overcome these forces, so their melting and boiling points are very high.

57
Q

Describe and explain the solubility of ionic compounds

A

Water molecules are polar - part of the molecule has a small negative charge, and the other bits gave small positive charges. The water molecules are attracted to the charged ions. They pull the ions away from the lattice and cause it to dissolve.

58
Q

Describe and explain the conductivity of ionic compounds pounds both as solids, liquids and aqueous solutions

A

The ions in a liquid are mobile( and they carry a charge). In a solid they’re fixed in position by the strong ionic bonds.

59
Q

What is a covalent bond

A

Strong electrostatic attraction between shared electrons and nuclei of the bonded atoms

60
Q

What are Giant Covalent Lattices? (Give some examples)

A
  • Giant covalent lattices are huge networks of covalently bonded atoms.
  • Carbon atoms can form this type of structure because they can each form four strong, covalent bonds.
  • Different forms of the same element in the same state are called allotropes. Carbon has several allotropes.
  • These include diamond, graphite and graphene.
61
Q

Describe the structure of diamond and therefore why it has its properties?

A
  • Each carbon atomis covalently bonded to four other carbon atoms to form a tetrahedral shape.
  • Because it has a lot of strong covalent bonds:
  • Diamond has a very high mpt as lots of energy is needed to overcome the strong covalent bonds
  • Diamond is extremely hard for the same reason
  • Vibrations travel easily through the stiff lattice, so its a good thermal conductor
  • It can’t conduct electricity - all the outer electrons are held in localised bonds
  • It won’t dissolve in any solvent.
62
Q

Describe the structure of graphite ?

A
  • The carbon atoms are arranged in sheets of flat hexagons covalently bonded with three bonds each
  • The fourth outer electron of each carbon atom is delocalised between the sheets of hexagons
  • The sheets of hexagons are bonded together by weak induced dipole-dipole forces.
63
Q

Why does graphite have its properties?

A
  • The weak forces between the layers in graphite are easily broken, so the sheets can slide over each other- graphite feels slippery and is used as a dry lubricant and in pencils
  • The ‘delocalised’ electrons in graphite aren’t attached to any particular carbon atom and are free to move along the sheets, so an electric ,current can flow
  • The layers are quite far apart compared to the length of the covalent bonds , so graphite is less dense than diamond and is used to make strong, lightweight sports equipment.
  • Because of the strong covalent bonds in the hexagon sheets, graphite also has a very high melting point
    Like diamond, it is insoluble.
64
Q

What is the structure of graphene and how does it explain its properties?

A
  • Graphene is a sheet of carbon atoms joined together in hexagons.
  • The sheet is just one atom thick, making it a two-dimensional compound.
    It is the best known electrical conductor on earth as its delocalised electrons can also move and carry charge but do it quicker as graphite doesn’t have layers.
  • The delocalised electrons also strengthen the covalent bonds between the carbon atoms. This makes graphene extremely strong.
  • A single layer of graphene is transparent and incredibly light.
  • It has a high mpt and is not soluble in any substance.
65
Q

Describe metallic bonding

A
  • The electrons in the outermost shell of a metal atom are delocalised - the electrons are free to move about the metal.
  • This leaves a positively charged metal cation.
  • The metal cations are electrostatically attracted to the delocalised negative electrons.
  • They form a lattice of closely packed cations in a sea of delocalised electrons.
66
Q

How does metallic bonding explain the properties of metals?

A
  • The number of delocalised electrons per atom affects the melting point. The more there are, the stronger the bonding will be and the higher the melting point.
  • As there are no bonds holding specific ions together, the metal ions can slide past each other when the structure is pulled, so metals are malleable and ductile.
  • The delocalised electrons can pass kinetic energy to each other, making metals good thermal conductors.
  • Metals are good electrical conductors because the delocalised electrons can move and carry charge.
  • Metals are insoluble
67
Q

What is a salt?

A

When the H+ in an acid is replaced by a metal ion OR an ammonium ion OR a + ion

as-level chemistry (4 decks)

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