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Lower Sixth Chemistry > NEEEEEEEED TO KNOW > Flashcards

NEEEEEEEED TO KNOW Flashcards

1
Q

What is the relative atomic mass

A

Relative atomic mass (Ar) is the weighted mean of the atoms of a normal sample relative to 1/12 the mass of the C-12 isotope

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2
Q

What is the relative isotopic mass

A

Relative isotopic mass is the mass of an isotope of the element relative to 1/12 the mass of the C-12 isotope

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3
Q

What is relative molecular mass

A

Relative molecular mass (Mr) is the mass of a molecule relative to 1/12 the mass of the C-12 isotope

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4
Q

What is the relative formula mass

A

Relative formula mass is the sum of all the atomic masses of all the atoms in a particular formula

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5
Q

What is the molecular/parent ion

A

the peak with the highest mass caused by the ion formed from the whole molecule

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6
Q

Where to find the molecular ion

A

the largest peak at the end of the m/z graph

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7
Q

Which metals atom emit coloured flames

A

lithium
sodium
potassium
calcium
strontium
barium

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8
Q

Definition of first ionisation energy

A

the energy required to remove one electron from each atom in one mole of gaseous atoms producing one mole of gaseous ions with one positive charge

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9
Q

What does the magnitude of the first ionisation energy tell us

A

the magnitude tells us about the force of attraction between the nucleus and the outer electrons

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10
Q

trends for first ionisation energies across a period

A

TREND 1: general increase
TREND 2: little drop from group 2 to 3
TREND 3: little drop from group 5 to 6
TREND 4: big drop from group 0 to 1/ one period to another

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11
Q

Reasons for TREND 1

A

-increase in nuclear charge
-similar shielding and distance
-attraction between outer e- and nucleus

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12
Q

Reasons for TREND 2

A

-distance increase from s to p orbital
-same shielding
-attraction between outer e- and nucleus decrease
-(despite higher nuclear charge)

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13
Q

Reasons for TREND 3

A

-outer electron is group 6 is sharing a p orbital
-repulsion between the electrons in the p orbital
-attraction between outer e- and nucleus decrease

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14
Q

Reasons for TREND 4

A

-large increase in distance and shielding, into a new electron shell
-attraction between outer e- and nucleus decrease
-(despite higher nuclear charge)

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15
Q

how is the RMM calculated using the graph

A

the mass value of the molecular/parent ion

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16
Q

how to answer an ionisation energy question

A

PSDSS
P-protons (no. of)/nuclear charge
S-shielding(no. of)
D-distance from nucleus
S-spin pair repulsion
S-strength of attraction

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17
Q

What is ionic bonding

A

ionic bonding is the strong electrostatic attraction between oppositely charged ions

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18
Q

what affects ionic bonding

A

-ionic radius
-ionic charge

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19
Q

How to answer ionic bond strength questions

A

R - radius (ionic)
C - charge (ionic)
A - [a]energy

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20
Q

define isoelectronic ions

A

same electron structure

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21
Q

trend for ionic radius down a group

A

increase in ionic radius down a group
more shells
increase in electron density

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22
Q

trend for ionic radius arcoss a period of isoelectronic ions

A

-ionic radius decrease
-more protons, attract the elctrons closer

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23
Q

What are ionic compounds arranged in

A

they are arranged in giant ionic lattices

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24
Q

why are anions bigger than than cations

A

cations have the same amount protons than its parent ion but less electrons . therefore the attraction of electrons to nucleus is more, so a smaller size.
there is more electron-electron expulsion and lower nuclear charge per electron. the electrons push each other away and make the anion bigger.

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25
Q

what does electron density maps show

A

it shows the likelihood of finding an electron in the region

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26
Q

ionic compunds at room temperature and why

A
  • most ionic compounds are solids
  • not enough energy to overcome the strong electrostatic forces of attraction between the oppositely charged ions
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27
Q

what will many ionic compunds dissolve into and why

A
  • polar solvents like water
  • δ+ end of the molecule surround the negative anions
  • δ- end of the molecule surround the positive cations
  • greater the ionic charge, the less soluble an ionic compound is
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28
Q

what are the factors of how soluble an ionic compound is

A
  • the relative strength of the electrostatic forces of attraction with the lattice
  • the attractions between the ions and the polar molecules
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29
Q

what gives the evidence of the existence of ions

A
  • electrolysis
  • positive ions are attracted to the negative electrode
  • negative ions are attracted to the positive electrode
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30
Q

Why does it matter if an electron is by it-self or paired in an electron shell?

A
  • Two electrons in the same orbital experience a bit of repulsion from each other
  • This offsets the attraction of the nucleus, so it is easier to lose an electron
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31
Q

Why is the first ionisation energy of a group 3 less than group 2?

A

The first ionisation energy of group 3 is less than that of group 2 because removing an electron from a higher energy orbital requires less energy. The electron removed from group 3 is in a p sub shell, which has a higher energy than the s electron in group 2.

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32
Q

Why is the first ionisation energy of group 6 less than that of group 5?

A

The highest energy electron in both group 5 and 6 is the p sub-level, and group 6 has a higher nuclear charge. However, group 6 has an electron that is spin-pairing, leading to repulsion of the two electrons. making it easier to remove.

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33
Q

What are orbitals?

A

Orbitals are a region of space within an atom where there is a probability of finding an electron

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34
Q

What is the shape of a P orbital

A

‘Dumbbell’ shaped
Can appear in groups of 3

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35
Q

What is periodicity?

A

Trends in the properties of elements repeat in successive periods of the periodic table

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36
Q

what determine same chemical properties in elements

A

no. of electrons in outer shells

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37
Q

properties of isotopes

A

similar chemical
different physical

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38
Q

What is covalent bonding

A

the strong electrostatic attraction between the nuclei and the shared pair of electrons between them

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39
Q

what is bond length

A

the internuclear distance of two covalently bonded atom

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40
Q

what is the relationship between bond length and strength

A

longer the length, weaker the strength
shorter the length, stronger the strength

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41
Q

why are triple bonds are the strongest

A

-short bond length
-large electron density between the two nuclei
-increase in attraction between the electrons and nuclei

42
Q

order of repulsion of electron pairs

A

lone-lone > lone-bonded > bonded-bonded

43
Q

Possible shapes the molecules can form and their angles and give examples

A
  • linear, 180, CO2
  • bent/v-shaped, 104.5, H2O
  • trigonal planar, 120, BF3
  • trigonal pyramidal, 107, NH3
  • trigonal BIpyramidal, 90, 120, PCl5
  • tetrahedral, 109.5, CH4
  • octahedral, 90, SF6
44
Q

How to tell how many electrons is used in bonding and structure of a compund

A

the oxidation state

45
Q

Why isn’t BeCl2 ionic bonded

A
  • Be2+ is so small, 4 protons and 2 electrons
  • highly polarising and strong attractions
  • Cl- is big
  • Be2+ distorts the e- cloud of Cl- so much towards itself that it covalently bonds
  • incomplete octet, could accept 2 lone pairs of electrons from another molecule datively
46
Q

what is electronegativity

A

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

47
Q

What are the 4 types of structure?

A

Simple molecular
Giant covalent lattice
Giant ionic lattice
Giant metallic lattice

48
Q

Why does NaF have a high melting point?

A

Giant ionic lattice structure
Many strong electrostatic forces in all directions]
Which require lots of energy to break

49
Q

How many electrons do sulphur, chlorine and phosphorous all have the ability to accommodate in their valence shell?

A

18

50
Q

What is the definition of dative bonding?

A

A dative bond is a covalent bond where only one of the bonded atoms donates both electrons being shared

51
Q

Why is X-ray diffraction useful?

A

It allows bond lengths and spacing between ions or structures of crystals to be investigated

52
Q

What substances have simple molecular structures?

A

Most non-metal elements (except group 8)
Most non-metal compounds
(- Molecules)

53
Q

What substances have giant covalent structures?

A

Diamond
Graphite
Silicon dioxide
(- Atoms)

54
Q

difference between simple molecular and giant covalent

A

simple covalent molecules have a small and fixed number of atoms, while giant structures have large and variable numbers of atoms.

55
Q

electronegativity down a group

A
  • decreases
  • more shells, more shielding
  • increase in atomic radius
  • negligible increase in nuclear charge
  • less attraction between nucleus and outer bonding electron
56
Q

electronegativity across a period

A
  • increases
  • increase in nuclear charge
  • decrease in atomic radius
  • greater attraction between nucleus and outer bonding electron
57
Q

how does the difference in electronegativity dictate the type of bond

A
  • when the difference is more than 1.7, ionic bond
  • when the difference around 0.4 to 1.7, covalent bond, bond will be polar
58
Q

Steps for working out the shape of molecules

A

Number of valence electrons on central atom
Number of atoms bonded to central atom
Number of lone pairs
Bonding pairs + lone pairs

59
Q

what are intermolecular forces and types

A

weak forces between molecules
- london forces, induced dipole-dipole forces
- permanent dipole-dipole forces
- hydrogen bonding

60
Q

ranking of all bonding forces from stongest to weakest

A
  • metallic bonding
  • ionic bonding
  • covalent bonding
  • hydrogen bonding
  • p d-d forces
  • london forces, instantaneous d-d forces
61
Q

what force occurs in all molecules

A

london forces

62
Q

how does london forces occur

A
  • electron cloud constantly moves in non polar molecule
  • instaneous dipole occurs temporarily with one side + and other -
  • it induces a dipole to neighbouring molecules
  • the - of one creates a i d-d force between the + of another and vice versa
  • strength depends on no. of electrons
63
Q

how does permanent d-d forces occur

A
  • formed from polar molecules
  • always a + and - end
  • p d-d forces occurs when two oppositely charged ends of two neigbouring molecules are attracted
64
Q

how does hydrogen bonding occur

A
  • special p d-d force
  • needs O,N, or F bonded to H
  • the bond between is highly polarised can all electron are pulled to the very electronegative elements
  • H becomes so + charged so it can bond with with lone pairs of O, N or F in another molecule
65
Q

why is H2O is require mor energy than HF and NH3

A
  • can form two hydrogen bonds, twice as many
66
Q

How to answer difference in boiling/melting temperature

A
  • no. of electrons - for LFs
  • electronegativity
  • intermolecular forces on each molecule
  • [number hydrogen bonds form]
    -energy required [lower/higher than]
67
Q

density of ice vs water

A
  • lower density than water
  • less packed
  • water molecules are packed in an open lattice
  • the long hydrogen bonds pushes the water molecules futher apart
  • therefore less dense
68
Q

What are the properties of molecular substances?

A
  • Low mp/bp: Little energy required to overcome IM forces
    • Neutral: Cannot carry charge
    • Polar molecules dissolve in polar solvents: PD-PD interactions form
    • Non-polar molecules dissolve in non-polar solvents: London forces can form between the solvent and solute
69
Q

Describe the properties of graphite

A

Allotrope of carbon
Each C atom is covalently bonded to 3 other C atoms
Hexagonal arrangement
Remaining electron is delocalised between layers
Insoluble in water: No interactions

70
Q

Describe the properties of diamond

A

Allotrope of carbon
Each C atom is covalently bonded to 4 other C atoms
Tetrahedral arrangement
Insoluble in water: No interactions

71
Q

How do we test for polar molecules?

A

Diverting streams method:
- Use a burette to create a stream of the liquid you are testing
- Charge a plastic rod
- If the liquid is polar the stream will be deflected

72
Q

What does oxidation mean?

A

Loss of electrons
Higher oxidation number

73
Q

What does reduction mean?

A

Gain of electrons
Decreased oxidation number

74
Q

What is an oxidising agent?

A

Something that causes oxidation by taking away electrons

75
Q

What is a reducing agent?

A

Something that causes reduction by giving electrons

76
Q

What is a disproportionation reaction?

A

A reaction where one element in one species is both oxidised and reduced

77
Q

How to write half equations

A
  1. Calculate oxidation states on each side of the equation.
  2. Balance the element changing oxidation state.
  3. Sort out Os. For every O gained/lost, add/remove one H2O molecule.
  4. Sort out Hs. For every H gained/lost, add/remove one H+ ion.
  5. Sort out electrons. If the oxidation state becomes more negative then it gains electrons. If the oxidation state becomes more positive then electrons are lost.
  6. Check – if the total electric charge on the left equals that on the right then it is probably correct. If it is not then you know you have gone wrong!
78
Q

Desribe and explain the trend in boiling temperatures of alkanes with increasing chain length

A
  • as chain increase, boiling point increase
  • more points of contact/ surface area with adjacent molecules
  • stronger LFs between adjacent molecules
79
Q

Describe and explain the effect of branching in the carbon chain on the boiling temperatures of alkanes

A
  • as branching increases, the boiling point decreases
  • less points of contact with adjacent molecules
  • weaker LFs
80
Q

the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons

A
  • alkanes only LFs
  • alcohols form hydrogen bonds
  • require more energy
81
Q

Describe and explain the trend in boiling points in the first four hydrogen halides

A

highest to lowest
- HF, HI, HBr, HCl

-HF can form hydrogen bonds
- others can’t, low electronegativity can’t create a sufficient dipole
- increases down the group
- iodine has more electrons, stronger LFs

82
Q

why do some ionic compunds do not dissolve into water

A

the electrostatic between the ions is too great for water molecules to overcome

83
Q

what is hydration of ions

A

when + end of water (hydrogens) surround anions and - end of water (oxygen) surround cations and pull the ionic lattice apart

84
Q

why can alcohols dissolve in water

A

they can form hydrogen bonds with water

85
Q

why can’t halogenoalkanes can’t dissolve in water

A

unable to form hydrogen bonds despite being polar and unable to interact with water molecules

86
Q

why can’t water dissolve not ionic or able to form h bonds

A
  • the strong hydrogen bonds between water resticts the molecules to intersperse, preventing them from moving unless they can form equally strong interactions with water
87
Q

what can non-polar molecules can dissolve into

A
  • non-polar, non-aqueous solvents
  • similar forces between molecules
  • able to interact with each other freely
88
Q

what is metallic bonding

A

the strong electrostatic attraction between metal ions and the delocalised ions

89
Q

reactivity of group 2

A
  • reducing strength (strength to give away e-) increases down the group (reactivity increases)
90
Q

why reducing strength increases down group 2

A
  • sum of 1st & 2nd ionization energy decrease
  • strength of metallic bonding decreases
91
Q

solubility of hydroxides of G2

A

more soluble down the group

92
Q

solubilty of sulphates of G2

A

more insoluble down the group

93
Q

thermal stability trend for G2

A

more stable down the group
decompose slower
higher temp

94
Q

G2 + oxygen

A

metal oxide
ionic
white solid
high MP

95
Q

G2 + chlorine

A

metal chloride
ionic
white solid
high-ish MP

96
Q

G2 + water

A

metal hydroxide + hydrogen
metal disappears
effervescence
increase in solubility of salt

97
Q

G2 oxide + water

A

G2 hydroxide
Be(OH)2 no reaction
solid disappears

98
Q

G2 oxide + acid

A

salt + water
watch out for salt solubility

99
Q

G2 hydroxide + acid

A

salt + water
watch out for salt solubility

100
Q

why calcium carbonates decompose when heated

A

Ca2+ attract some e- density in the CO3 2- (polarises)
lower e- density in CO3 2- weakens C-O bond
C-O bond breaks if theres heat energy

101
Q

trend in thermal stability in group 1+2 and why

A

more stable down the group
- cation radius increases
- less polarisation of carbonate/nitrate ion
- C-O/N-O bonds weakened less
- break slower or needs higher temp

102
Q

What G1/G2 carbonates decompose into

A

Li + G2 decompose
Li2CO3[s] > Li2O[s] + CO2[g]

Lower Sixth Chemistry (14 decks)

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