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AS Level Chemistry > The Atom > Flashcards

The Atom Flashcards

0
Q

What does the Mass number tell us about an isotope?

A

The number of protons and neutrons in the nucleus

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1
Q

What does the Atomic number tell us?

A

The number of protons in the nucleus

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2
Q

Where are protons and neutrons found?

A

In the nucleus

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3
Q

How would we calculate the number of neutrons in a isotope?

A

Mass number - Atomic number

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4
Q

Give the relative charge and relative mass of protons, neutrons and electrons

A

Proton- charge +1, mass 1
Neutrons- charge 0, mass 1
Electrons- charge -1, mass 1/1840

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5
Q

What did Ernest Rutherford discover?

A

Found that most of the mass and all of the positive charge was in a tiny central nucleus

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6
Q

What did Henri Becquerel discover?

A

Radioactivity, showing that particles come from inside the atom, therefore atoms are divisible

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7
Q

What did JJ Thompson discover?

A

The electron, and that is negatively charged, that all electrons from an atom are the same.

Suggested the precedes of protons to balance charge, suggested plum pudding model

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8
Q

What did John Dalton discover?

A

Suggested elements were made of invisible atoms,

Same elements = same mass

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9
Q

What did Robert Boyle discover?

A

Proposed there ewe substances that could be made simpler- chemical elements

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10
Q

electrons are arranged in energy levels and sub-levels, what letters are given to the sub-levels?

A

s,p,d,f

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11
Q

write an equation for the first ionisation energy of sodium?

A

Na(g) –> Na+(g) +e-

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12
Q

what is the electronic configuration for vanadium? (periodic table needed)

A

1s2, 2s2, 2p6, 3s2, 3p6,3d3, 4s2

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13
Q

which sub-shell fills first- 4s or 3d?

A

4s

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14
Q

what is the maximum number of electrons held in each of the s, p ,d sub-shells?

A

s-2
p-6
d-10

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15
Q

z for krypton is 36, what is the electronic configuration? (periodic table needed)

A

1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6

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16
Q

there are two isotopes of chlorine, Cl-35 (75%) and Cl-37 (25%), show that the R.A.M of chlorine is 35.5

A

(35 x 75) + (37 x 25)
____________________
100
=35.5

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17
Q

how do we work out the relative atomic mass of an element?

A

add up (mass x abundance) for each isotope and then divide by total abundance
(mass x Abundance
__________________
total abundance

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18
Q

what happens to first ionisation energies as we move left to right along the period? why?

A

it increases

because nuclear charge increases but shielding remains the same

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19
Q

what is an atomic orbital?

A

a volume f space that an electron (or pairs of electrons) fill

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20
Q

what is an energy level?

A

“a shelll”

a collection of electrons with similar energy. the first ‘shell’ contains 2, the second 8 and the third 18 electrons.

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21
Q

when moving right to left across the period 3, what exceptions are there for ionisation energy?

A

Auminium- because the electron is being moved from3p, which has a slightly higher energy level than 3s.

Sulfur- because there are two electrons in one of the 3p orbitals and so there is repulsion between them

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22
Q

what can we say about the spin of pairs of electrons?

A

paired electrons will have opposite spin

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23
Q

what is the rule for electron pairing in orbitals?

A

electrons will fill in singly within the same sub-shell, before they pair (because they repulse each other)

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24
Q

write the equation for the second and third ionisation energies of sodium

A

Na+(g) –> NA2+(g) + e-

NA2+(g) –> NA3+(g) +e-

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25
Q

define first ionisation energy

A

the energy required to remove one mole of electrons from one mole of atoms, in the gaseous state

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26
Q

how are electrons removed from atoms?

A

ioniation, the atoms are hit with a beam of electrons from an electron gun

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27
Q

what is an isotope

A

atoms of the same element, with the same number of protons and electrons, but a different number of netrons

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28
Q

what does letter A stand for?

A

mass number

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29
Q

what does the letter Z stand for?

A

atomic number

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30
Q

define second ionisation energy

A

the energy required to remove one mole of electrons from one mole of atoms, with one positive charge, in the gaseous state

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31
Q

when thinking about electrons what should we thnk

A 
S- shielding
A- attraction (to)
N- necleus
D- distance (of outer shell)
E- elctrons
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32
Q

what does ionic bonding happen between

A

metal and non-metal

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33
Q

what properties do ionic compounds have

A

melting points
strong force between oppositely charged ions
molten when dissolved

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34
Q

what is covalent bonding

A

sharing pairs of electrons between atoms

35
Q

what is another name for dative covalent bonding

A

co-ordinate bonding

36
Q

what is dative covalent bonding

A

a covalent bond where both electrons in the bond have been contributed by one species

37
Q

what happens in dative covalent bonding

A

an atom that is electron deficient, as it doesn’t have a full outer shell, the atom that is donating the electrons has a pair of electrons not being used, called a lone pair

38
Q

what crystal structures conduct electricity? why?

A

al metallic, graphite

because they have delocalised electrons that are free to moe through the structure

39
Q

what effect do pairs of electrons both bonding and non-bonding have on each other

A

they repel

40
Q

what shape and bond angle would a molecule with only two bonding pairs of electrons in the outer shall have?

A

Linear

180’

41
Q

what shape and bond angle would a molecule with only three bonding pairs of electrons in the outer shall have?

A

trigonal planar

120’

42
Q

what shape and bond angle would a molecule with only four bonding pairs of electrons in the outer shall have?

A

tetrahedral

109.5’

43
Q

what shape and bond angle would a molecule with only five bonding pairs of electrons in the outer shall have?

A

trigonal bipyramid

90’ and 120’

44
Q

what shape and bond angle would a molecule with only six bonding pairs of electrons in the outer shall have?

A

octahedral

90’

45
Q

what shape and bond angle would a molecule with three bonding pairs and one lone pair have?

A

triangular pyramid

107’

46
Q

what shape and bond angle would a molecule with two bonding pairs and two lone pairs have

A

v-shape

104.5’

47
Q

what shape and bond angle would a molecule with four bonding pairs and two lone pairs have

A

square planar

90’

48
Q

what repels more, bonding pairs or lone pairs

A

lone pairs

49
Q

what is ionic bonding

A

the electrostatic attraction between oppositely charged ions in a lattice

50
Q

what is an ion

A

a charged particle

51
Q

what charge do the ions of elements the groups 1,2,3,6,7 have?

A 
1= +1
2= +2
3= +3
6= -2
7=-1
52
Q

what is metallic bonding

A

attraction between delocalised electrons and positive ions, in a lattice

53
Q

what is electro-negativity

A

the power of an atom to withdraw the electron density from a covalent bond

54
Q

what is meant by a polar bond

A

where the charge in a molecule is not symmetrical, so one area is slightly more positive and another slightly more negative

55
Q

what can we say about the polarity of bonds in elements

A

they are not polar

56
Q

what can we say about the polarity of bonds between different elements

A

they will be polar to different extents depending on the elements

57
Q

starting with the weakest, name the three types of intermolecular attraction

A

van der walls
permanent dipole-dipole
hydrogen bonding

58
Q

what is a single covalent bond

A

a shared pair of electrons

59
Q

what is a double covalent bond

A

two shared pairs of electrons

60
Q

how does the strength of a covalent bond compare with the strength of a hydrogen bond

A

covalent bonds are stronger

61
Q

what causes van der waals force

A

caused by instantaneous dipoles. these occur because in any instant electrons are not spread evenly and more will be in one area than another. this will, in that instant induce a dipole in its neighbour, leading attraction

62
Q

where do we find van der waals forces and what affects its strength

A

found in all atoms and molecules. the more electrons present, the bigger the force

63
Q

what causes permanent dipole- dipole force

A

attraction between molecules with permanent dipoles

64
Q

what effects the strength of permanent dipole-dipole force

A

the bigger he difference in electro-negativity between a bonding pair, the greater the dipole and so the greater the force

65
Q

what causes hydrogen bonding

A

when hydrogen is covalently bonded to NOF. the hydrogen can be shared between this and a lone pair on another molecule

66
Q

why does temperature not increase while a substance is melting or boiling

A

he energy is absorbed as the bond weakens

67
Q

what is the energy required to weaken the forces of attraction in a solid enough for a solid to turn into a liquid called

A

enthalpy change to fusion

68
Q

what are the four types of crystal structure

A

ionic, metallic, giant covelent , molecular

69
Q

what are ionic crystals held together by

A

intermolecular forces (van der waals, hydrogen, dipole-dipole) hold molecules together. covalent bonds hold atoms within the molecules together

70
Q

what are metallic crystals held together by

A

attractions between positive metal ions and a negative ‘sea of delocalised electrons

71
Q

what are giant covalent (macromolecular) crystals held together by

A

covalent bonds

72
Q

what type of structure do the following crystals have? NaCl, Mg, Diamond, Graphite Iodine, Ice

A 
NaCl- ionic
Mg- metallic
Diamond- Giant Covalent
Graphite- Giant Covalent
Iodine- Molecular
Ice- molecular
73
Q

what is electron repulsion theory

A

each pair of electrons around an atom will repel all other electron pairs
the pairs of electrons will take positions as far away from each other as possible

74
Q

what does macromolecular mean

A

a macromolecular structure is one in which large numbers of atoms are linked in the regular three-dimensional arrangement by covalent bonds

75
Q

why is the second ionisation energy of silicon is lower than the second ionisation energy of aluminium

A

Electron in Silicon is removed from a higher energy orbital/ more shielded

76
Q

Predict the element in period 3 that has the highest second ionisation energy

A

Na

Electron is removed from 2p

77
Q

explain why the ionisation energy of every element is endothermic

A

Energy needed to overcome attraction between electron and nucleus

78
Q

Suggest why the electronegativity of the elements increase from lithium to fluorine

A

The attraction to the nucleus increases
The distance of the outer electron increases
But shielding stays the same

79
Q

Explain how the current is generated in a mass spectrometer

A

Electrons transfer at the detector to the positive ion

80
Q

Identify the period two element, from carbon to fluorine that has the largest atomic radius. Explain answer

A 
Carbon
> fewest protons
> smallest nuclear charge 
> least attraction between protons and electrons 
> weakest nuclear attraction 
> similar shielding
81
Q

State one reason why the first ionisation energy of rubidium is lower than the first ionisation energy of sodium

A

> Further from nucleus
from higher energy level
more shielding

82
Q

Why is there no hydrogen bonding between phosphine molecules

A

There is not a great enough difference in electronegativity

83
Q

In terms of structure and bonding, why does graphene have a high melting point

A

Covalent bonds are hard to break

84
Q

Why can a meal be hammered into different shapes

A

The layers can slide over each other making it malleable

85
Q

Suggest one reason why particles with the same mass and velocity can be deflected by different amounts in the same magnetic field

A

Different charges

AS Level Chemistry (7 decks)

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