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Chemistry A Level > Module 2 > Flashcards

Module 2 Flashcards

1
Q

Isotope definition

A

Atoms of the same element with different numbers of neutrons and different masses

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2
Q

Cations

A

Positive ions

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3
Q

Anions

A

Negative ions

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4
Q

Relative isotopic mass

A

Is the mass of an isotope relative to 1/12 th of the mass of an atom of carbon-12

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5
Q

Relative atomic mass

A

The weighted mean mass of an atom of an element relative to 1/12 of the mass of an atom of carbon-12

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6
Q

Mass to charge ratio (m/z) equation

A

Relative mass of ion / relative charge of ion

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7
Q

Relative atomic mass equation

A

(Abundance x mass) (abundance x mass) … / 100

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8
Q

Avogadros constant

A

6l02x10^23, the number of particles in each mole of carbon-12

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9
Q

Mol equation

A

Mol= mass/ mr

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10
Q

Molecular formula definition

A

The number of atoms of each element in a molecule

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11
Q

Empirical formula

A

The simplest whole number ratio of atoms of each element in a compound

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12
Q

mol equation ( from c and v )

A

Mol= conc x vol

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13
Q

Standard solution

A

A solution of known concentration. They are made by dissolving an exact mass of the solute in a solvent and making up the solution to an exact volume

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14
Q

Mol equation at rtp

A

Mol= vol (dm3) / 24

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15
Q

Ideal gas equation

A

pV= nRT

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16
Q

Value for the gas constant

A

8.314 j mol k

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17
Q

How to convert from cm3 to m3

A

X10^-6

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18
Q

How to convert from c to k

A

+ 273

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19
Q

Percentage yield

A

Actual yield / theoretical yield x 100

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20
Q

Theoretical yield

A

The maximum possible amount of product

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21
Q

Why isn’t the theoretical yield always reached

A

The reaction may not have been completely reacted
Side reactions may have taken place
Purification of the product may result in loss of some product

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22
Q

Actual yield

A

Obtained from a reaction is usually lower then the theoretical yield

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23
Q

Limiting reagent

A

The reactant that’s not in excess and when it’s completely used up first it stops the reaction

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24
Q

Atom economy

A

Sum of mr of desired product / sum of mr of all products X 100

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25
Q

Oxidation number for element s

A

0

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26
Q

Oxidation number of H in hydrides (eg NaH, CaH2)

A

-1

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27
Q

Oxidation number of O in peroxides

A

-1

28
Q

Oxidation number of O bonded to F

A

+ 2

29
Q

Reduction in terms of e-

A

Gain of electrons

30
Q

Oxidation in terms of e-

A

Loss of electrons

31
Q

Reduction in terms of oxidation number

A

Decrease in oxidation number

32
Q

Oxidation in terms of oxidation number

A

Increase in oxidation number

33
Q

Metal + acid——>

A

Salt hydrogen

34
Q

How many electrons can an orbital hold

A

1 or 2 electrons no more

35
Q

S-orbital shape

A

Sphere

36
Q

P- orbital shape

A

Dumbbell

37
Q

How many orbital in p

A

3

38
Q

How many orbitals are in d and f

A

D- 5
F- 7

39
Q

Why does the 4s sub shell fill before the 3d sub shell

A

It has a lower energy level

40
Q

Ionic bonding definition

A

The electrostatic attraction between positive and negative ions

41
Q

Ionic bond boiling and melting points

A

Almost all ionic compounds are solids at rtp. Most ionic compounds have high melting and boiling pts

42
Q

Solubility in an ionic bond

A

Many ionic compounds dissolve in polar solvents such as water. Water breaks down the lattice and surround each ion in solution. If the compound is made up of ions with large charges the ionic attraction may be too strong so this compound will then not be very soluble

43
Q

Electrical conductivity in ionic bond

A

In the solid state an ionic compound doesn’t conduct electricity. But once melted or dissolved it will

44
Q

Covalent bonding

A

The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms. Occur in non-metallic compounds

45
Q

Are covalent bonds localised

A

Yes, the attraction is solely between the nuclei and shared electrons , this can result in a small unit called a molecule

46
Q

Double covalent bond

A

The electrostatic attraction between 2 shared pairs of electrons and the nuclei of the bonding atoksn

47
Q

Dative cove lent bonds ( or coordinate bonds )

A

A covalent bond in which the shared pair of electrons has been supplied by one of the bonding atoms only

48
Q

How to measure covalent bond strength

A

Average bond enthalpy. The larger the value of the average bond enthalpy the stronger the covalent bond

49
Q

4 bonding pairs and 0 lone pairs , shape + bond angle

A

Tetrahedral , 109.5

50
Q

3 bonding pairs and 1 lone pair , shape and bond angle

A

Pyramidal , 107

51
Q

2 bonding pairs and 2 lone pairs, shape and bond angle

A

Non-linear , 104.5

52
Q

2 bonding regions , shape and angle

A

Liner , 180

53
Q

3 bonding regions, shape and region

A

Trigonal planar , 120

54
Q

6 bonding regions, shape and angle

A

Octahedral , 90

55
Q

Electronegativity definition

A

The attraction of a bonded atom for the pair of electrons in a covalent bond

56
Q

How is electronegativity measured

A

The Pauling scale

57
Q

Which element is the most electronegative

A

F

58
Q

How are electrons shared in non-polar bonds

A

The electron pair will be shared equally . This is a pure covalent bond

59
Q

How are electrons shared in polar bonds

A

The electron pair is shared unequally between the bonded atoms. Eg the bonded electrons in HCl will be closer to the Cl

60
Q

Intermolecular forces , definition and 3 main categories

A

They are the weak interactions between dipoles of different molecules.
Induced dipole-dipole interactions ( London forces )
Permanent dipole-dipole interactions
Hydrogen bonding

61
Q

Induced dipole- dipole interactions strength

A

They are weak intermolecular forces that exist between all molecules, polar or non-polar. The larger the number of electrons mean larger induced dipoles, more energy is then needed to overcome the intermolecular forces, increasing the boiling point

62
Q

How are London forces formed

A

-Movement of electrons produces a changing dipole in a molecule
-at any instant, an instantaneous dipole will exist but it’s position is constantly shifting
-the instantaneous dipole induced a dipole on a neighbouring molecule
-the induced dipole induced further dipoles on neighbouring molecules, which then attracts one another

63
Q

Properties of simple molecular substances

A

Low melting point and boiling point
Solubility- the solubility of non-polar substances is easier to predict, simple molecular substances tend to be insoluble in polar solvents
Electrical conductivity- there are no mobile charged particles in simple molecular structures so there is nothing to complete an electrical circuit

64
Q

Hydrogen bonds

A

A special type of permanent dipole-dipole interaction found between molecules containing and O, N , F

65
Q

Anomalous properties of water

A

Ice is less dense than water
Water has a relatively high melting and boiling point
Relatively high surface tension
Viscosity

66
Q

Why is ice less dense than water

A

Hydrogen bonds hold the water molecules apart in an open lattice structure. The water molecules in ice are further apart than in water, so solid ice is less dense than liquid water and floats

67
Q

Why does water have a relatively high melting and boiling point

A

Hydrogen bonds are extra forces over and above London forces . An appreciable amount of energy is needed to break the hydrogen bonds in water. Without hydrogen bonds water would have a boiling point of -75c

Chemistry A Level (12 decks)

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