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
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
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
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
5
Q
What is the molecular/parent ion
A
the peak with the highest mass caused by the ion formed from the whole molecule
6
Q
Where to find the molecular ion
A
the largest peak at the end of the m/z graph
7
Q
Which metals atom emit coloured flames
A
lithium
sodium
potassium
calcium
strontium
barium
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
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
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
11
Q
Reasons for TREND 1
A
-increase in nuclear charge
-similar shielding and distance
-attraction between outer e- and nucleus
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)
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
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)
15
Q
how is the RMM calculated using the graph
A
the mass value of the molecular/parent ion
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
17
Q
What is ionic bonding
A
ionic bonding is the strong electrostatic attraction between oppositely charged ions
18
Q
what affects ionic bonding
A
-ionic radius
-ionic charge
19
Q
How to answer ionic bond strength questions
A
R - radius (ionic)
C - charge (ionic)
A - [a]energy
20
Q
define isoelectronic ions
A
same electron structure
21
Q
trend for ionic radius down a group
A
increase in ionic radius down a group
more shells
increase in electron density
22
Q
trend for ionic radius arcoss a period of isoelectronic ions
A
-ionic radius decrease
-more protons, attract the elctrons closer
23
Q
What are ionic compounds arranged in
A
they are arranged in giant ionic lattices
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.
25
what does electron density maps show
it shows the likelihood of finding an electron in the region
26
ionic compunds at room temperature and why
- most ionic compounds are solids
- not enough energy to overcome the strong electrostatic forces of attraction between the oppositely charged ions
27
what will many ionic compunds dissolve into and why
- 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
28
what are the factors of how soluble an ionic compound is
- the relative strength of the electrostatic forces of attraction with the lattice
- the attractions between the ions and the polar molecules
29
what gives the evidence of the existence of ions
- electrolysis
- positive ions are attracted to the negative electrode
- negative ions are attracted to the positive electrode
30
Why does it matter if an electron is by it-self or paired in an electron shell?
- 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
31
Why is the first ionisation energy of a group 3 less than group 2?
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.
32
Why is the first ionisation energy of group 6 less than that of group 5?
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.
33
What are orbitals?
Orbitals are a region of space within an atom where there is a probability of finding an electron
34
What is the shape of a P orbital
‘Dumbbell’ shaped
Can appear in groups of 3
35
What is periodicity?
Trends in the properties of elements repeat in successive periods of the periodic table
36
what determine same chemical properties in elements
no. of electrons in outer shells
37
properties of isotopes
similar chemical
different physical
38
What is covalent bonding
the strong electrostatic attraction between the nuclei and the shared pair of electrons between them
39
what is bond length
the internuclear distance of two covalently bonded atom
40
what is the relationship between bond length and strength
longer the length, weaker the strength
shorter the length, stronger the strength
41
why are triple bonds are the strongest
-short bond length
-large electron density between the two nuclei
-increase in attraction between the electrons and nuclei
42
order of repulsion of electron pairs
lone-lone > lone-bonded > bonded-bonded
43
Possible shapes the molecules can form and their angles and give examples
- 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
How to tell how many electrons is used in bonding and structure of a compund
the oxidation state
45
Why isn't BeCl2 ionic bonded
- 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
what is electronegativity
the ability of an atom to attract the bonding electrons in a COVALENT bond
47
What are the 4 types of structure?
Simple molecular
Giant covalent lattice
Giant ionic lattice
Giant metallic lattice
48
Why does NaF have a high melting point?
Giant ionic lattice structure
Many strong electrostatic forces in all directions]
Which require lots of energy to break
49
How many electrons do sulphur, chlorine and phosphorous all have the ability to accommodate in their valence shell?
18
50
What is the definition of dative bonding?
A dative bond is a covalent bond where only one of the bonded atoms donates both electrons being shared
51
Why is X-ray diffraction useful?
It allows bond lengths and spacing between ions or structures of crystals to be investigated
52
What substances have simple molecular structures?
Most non-metal elements (except group 8)
Most non-metal compounds
(- Molecules)
53
What substances have giant covalent structures?
Diamond
Graphite
Silicon dioxide
(- Atoms)
54
difference between simple molecular and giant covalent
simple covalent molecules have a small and fixed number of atoms, while giant structures have large and variable numbers of atoms.
55
electronegativity down a group
- decreases
- more shells, more shielding
- increase in atomic radius
- negligible increase in nuclear charge
- less attraction between nucleus and outer bonding electron
56
electronegativity across a period
- increases
- increase in nuclear charge
- decrease in atomic radius
- greater attraction between nucleus and outer bonding electron
57
how does the difference in electronegativity dictate the type of bond
- 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
Steps for working out the shape of molecules
Number of valence electrons on central atom
Number of atoms bonded to central atom
Number of lone pairs
Bonding pairs + lone pairs
59
what are intermolecular forces and types
weak forces between molecules
- london forces, induced dipole-dipole forces
- permanent dipole-dipole forces
- hydrogen bonding
60
ranking of all bonding forces from stongest to weakest
- metallic bonding
- ionic bonding
- covalent bonding
- hydrogen bonding
- p d-d forces
- london forces, instantaneous d-d forces
61
what force occurs in all molecules
london forces
62
how does london forces occur
- 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
how does permanent d-d forces occur
- 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
how does hydrogen bonding occur
- 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
why is H2O is require mor energy than HF and NH3
- can form two hydrogen bonds, twice as many
66
How to answer difference in boiling/melting temperature
- no. of electrons - for LFs
- electronegativity
- intermolecular forces on each molecule
- [number hydrogen bonds form]
-energy required [lower/higher than]
67
density of ice vs water
- 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
What are the properties of molecular substances?
- 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
Describe the properties of graphite
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
Describe the properties of diamond
Allotrope of carbon
Each C atom is covalently bonded to 4 other C atoms
Tetrahedral arrangement
Insoluble in water: No interactions
71
How do we test for polar molecules?
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
What does oxidation mean?
Loss of electrons
Higher oxidation number
73
What does reduction mean?
Gain of electrons
Decreased oxidation number
74
What is an oxidising agent?
Something that causes oxidation by taking away electrons
75
What is a reducing agent?
Something that causes reduction by giving electrons
76
What is a disproportionation reaction?
A reaction where one element in one species is both oxidised and reduced
77
How to write half equations
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
Desribe and explain the trend in boiling temperatures of alkanes with increasing chain length
- as chain increase, boiling point increase
- more points of contact/ surface area with adjacent molecules
- stronger LFs between adjacent molecules
79
Describe and explain the effect of branching in the carbon chain on the boiling temperatures of alkanes
- as branching increases, the boiling point decreases
- less points of contact with adjacent molecules
- weaker LFs
80
the relatively low volatility (higher boiling temperatures) of alcohols compared to alkanes with a similar number of electrons
- alkanes only LFs
- alcohols form hydrogen bonds
- require more energy
81
Describe and explain the trend in boiling points in the first four hydrogen halides
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
why do some ionic compunds do not dissolve into water
the electrostatic between the ions is too great for water molecules to overcome
83
what is hydration of ions
when + end of water (hydrogens) surround anions and - end of water (oxygen) surround cations and pull the ionic lattice apart
84
why can alcohols dissolve in water
they can form hydrogen bonds with water
85
why can't halogenoalkanes can't dissolve in water
unable to form hydrogen bonds despite being polar and unable to interact with water molecules
86
why can't water dissolve not ionic or able to form h bonds
- 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
what can non-polar molecules can dissolve into
- non-polar, non-aqueous solvents
- similar forces between molecules
- able to interact with each other freely
88
what is metallic bonding
the strong electrostatic attraction between metal ions and the delocalised ions
89
reactivity of group 2
- reducing strength (strength to give away e-) increases down the group (reactivity increases)
90
why reducing strength increases down group 2
- sum of 1st & 2nd ionization energy decrease
- strength of metallic bonding decreases
91
solubility of hydroxides of G2
more soluble down the group
92
solubilty of sulphates of G2
more insoluble down the group
93
thermal stability trend for G2
more stable down the group
decompose slower
higher temp
94
G2 + oxygen
metal oxide
ionic
white solid
high MP
95
G2 + chlorine
metal chloride
ionic
white solid
high-ish MP
96
G2 + water
metal hydroxide + hydrogen
metal disappears
effervescence
increase in solubility of salt
97
G2 oxide + water
G2 hydroxide
Be(OH)2 no reaction
solid disappears
98
G2 oxide + acid
salt + water
watch out for salt solubility
99
G2 hydroxide + acid
salt + water
watch out for salt solubility
100
why calcium carbonates decompose when heated
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
trend in thermal stability in group 1+2 and why
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
What G1/G2 carbonates decompose into
Li + G2 decompose
Li2CO3[s] > Li2O[s] + CO2[g]
