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A-level Chemistry <3 > 2 - Atoms, Reactions, Bonding And Structure > Flashcards

2 - Atoms, Reactions, Bonding And Structure Flashcards

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1
Q

Define binary compound

A

A compound containing two elements only

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2
Q

Define Polyatomic ion

A

An ion containing more than one atom

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3
Q

What happens when a ionic substance dissolves in water

A

Positive and negative ions separate and become hydrated.

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4
Q

In reactions involving ionic compounds dissolving in water some ions may not be involved.
What are they called

A

Spectator ions

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5
Q

Acid + hydroxide
Ionic equation

A

H+ (aq) + OH- (aq). —> H2O(l)

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6
Q

Acid carbonate (ionic equation)

A

2 H+ (aq) + CO3 2- (aq). —> H20 (l) + CO2 (g)

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7
Q

Acid + hydrogen carbonate
(Ionic equation)

A

H+ (aq) + HCO3 - (aq) —> H2O(l) + CO2 (g)

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8
Q

Acid + ammonia
(Ionic equation)

A

H+ (aq) + NH3 (aq) —> NH4+ (aq)

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9
Q

Substance + oxygen -?

A

Oxides

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10
Q

Metal + water — ?

A

Metal hydroxide + hydrogen

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11
Q

Metal + acid —

A

Salt + hydrogen

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12
Q

Oxide +acid —

A

Salt + water

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13
Q

Hydroxide + acid —?

A

Salt + water

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14
Q

Carbonate + acid —?

A

Salt + water + carbon dioxide

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15
Q

Ammonia + acid —?

A

Ammonium salt

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16
Q

Metal carbonate on heating —?

A

Metal oxide + carbon dioxide

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17
Q

What is avogadros constant

A

6.02 x 10^23 mol-1

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18
Q

Why was avogadros number chosen

A

So that the mass of one mole of particles of a substance equals the mr in grams.

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19
Q

Define isotopes

A

– Atoms of the same element with different numbers of neutrons and different masses. Isotopes have the same atomic number but a different mass number

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20
Q

Isotopes always have- — - - - - chemical reactions as they have - - - - - electron configuration but physical properties like density may be — - - - -

A

The same
The same
Different

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21
Q

What can make an isotope radioactive

A

Sometimes the extra neutrons make the nucleus big and unstable

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22
Q

Isotopes of an element always have the…..

A

same atomic number but a different mass number due to having extra neutrons in the nuclei.

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23
Q

Relative atomic mass

A

The weighted mean mass of an atom of an element relative to one twelfth of the mass of an atom of carbon-12

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24
Q

Relative isotopic mass

A

the mass of an isotope relative to one twelfth of the mass of an atom of carbon-12

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25
Q

Relative molecular mass

A

compares the mass of a molecule with the mass of an atom of carbon-12

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26
Q

How do u calculate mr

A

Adding the Ar of all the elements

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27
Q

How do you calculate the relative formula mass

A

Adding Ar of all the elements in the empirical formula

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28
Q

Define relative formula lass

A

Compared the mass of a formula unit with the mass of an ato, of carbon 12

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29
Q

What is relative atomic mass calculated from

A

Using the isotopic masses and their abundence

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30
Q

How are isotope abundances found

A

Using a mass spectrometer

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31
Q

What is in the x axes of a mass spectrometer

A

m / z

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32
Q

What is in the y axes of a mass spectrometer

A

Percentage abundance

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33
Q

How to calculate relative atomic mass

A

(Atomic mass 1 x% abundance) + (atomic mass 2 x % abundance )
———————————————————
100

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34
Q

How to use a mass spectrometer

A

1- sample In
2- vaporised then ionised to form positive ions
3- ions accelerated - heavier move slower
4- ions detected on mass spectrum as mass: charge ration m/z - as they reach detector they create a signal - detected so the greater the abundance the larger the signal

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35
Q

How to find relative atomic mass without knowing the % abundance (isotopes)

A

(Mass) (% - written as x) + (mass2) (% - written as 100-x)
———————————————— — —
100

1) expanse + simplify
2- work out x
3 - work out 100 - x

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36
Q

Orbitals definition

A

A region of space where there is a high probability of finding an electron

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37
Q

How many electrons can orbitals have

A

Up to 2 that spin in opposite directions

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38
Q

What does isoelectronic mean

A

Same electronic configuration

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39
Q

Electrons charge

A

-1

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40
Q

Where are electrons located

A

In orbitals
These orbitals take up most the volume of the atom

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41
Q

Where is most the mass of the atom concentrated

A

In the nucleus

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42
Q

What subatomic particles are located in the nucleus

A

Protons and neutrons

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43
Q

Relative mass and charge of proton

A

Rm = 1
Charge = +1

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44
Q

Relative mass and charge of neutron

A

Rm= 1
Charge = 0

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45
Q

Relative mass of an electron

A

1
——
1836 (check)

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46
Q

How do you find number of neutrons

A

Mass number minus atomic number

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47
Q

Do negative ions have more electrons or protons

A

Electrons

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48
Q

What decides the chemical properties of an element?
So how does this effect isotopes?

A

The number and arrangement of eelctromsn

They have the same chemical properties as they still have the same electron configuration

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49
Q

What did some ancient Greeks use to think about atoms?

A

All matter was made from indivisible particles

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50
Q

At the start of the 19th century how did John dalton describe atoms

A

Sold spheres
And different types of spheres make up differnt elements

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51
Q

What did JJ Thomson prove + how

A

Atoms weren’t sold and indivisible.
His measurements of charge and mass showed that an atom must contain even smaller , negatively charged particles (he called them ‘corpuscles’ - we call them electrons)

The new model was known as the plumb pudding model - a positively charged sphere with negativley charged electrons embedded in

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52
Q

When did JJ Thomson do his experiments

A

1897

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53
Q

When did John dalton describe atoms as solid spheres

A

19th century

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54
Q

When did Ernest Rutherford do the gold foil experiment and with who?

A

1909
His students - hans Geiger and Ernest marsden

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55
Q

What was the gold foil experiment
And what were they expecting to happen becuase of the plumb pudding model

A

They fired alpha particles (which are Moseley postives charge) as an extremely thin sheet of gold.

Expected most of the alpha particles to be deflected very slightly by the positive pudding that made up most of the atom

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56
Q

What actually happened in the gold foil experiment

A

Most the alpha particles passed straight through the gold atoms and a very small number where deflected backwards - proved the plum pudding model was wrong

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57
Q

What did Rutherford replace the plumb pudding model with

A

The nuclear model

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58
Q

What was Rutherfords nuclear atom

A

1) a tiny positively charged nucleus at the centre of the atom, where most of the atoms mass is concentrated
2) the nucleus is surrounded by a cloud of negative electrons
3) most of the atom is empty space

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59
Q

What did Henry Moseley discover

A

The charge of the nucleus increased from one element to another in units of one.
Which led to him discovering it contained positively charged particles called protons.

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60
Q

Why did James Chadwick continue to experiment and What did James Chadwick discover

A

There was a problem with the previous model - the nuclei of the atoms was heavier than it would be a if it just contained protons.

James predicted that there was another subatomic particle in the nucleus wuth mass but no charge - he discovered the neutron

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61
Q

Why did Neil bohr continue to experiment

A

Scientists realised that electrons in a cloud around the nucleus of an atom would spiral down into the nucleus causing the atom to collapse

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62
Q

What 4 basic principles where in Niel bonds new model of the atom

A

1) electrons can only exist in fixed orbits or shells, and not anywhere in between

2) each shell has a fixed npenergy

3) when an electron moves between shells electromagnetic radiation is emitted or absorbed

4) becasue the energy of shells is fixed, the radiation will have a fixed frequency

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63
Q

How did Bohr mod well explain why some elements are inert (noble gasses)

A

The shells of an atom can only hold fixed number of electron,and an elements reactivity is due to its electrons.
When an atom has full shells of electrons it is stable and doesn’t react

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64
Q

What does the most accurate model of the atom we know today involve

Inc positive and negatives

A

Complicated quantum mechanics.
You can never know where an electron is or which direction it is goin inn at any moment, but you can say how likely it is to be at any particular point in the atom,m

P- more accurate and explains some observations that Bohr model can’t
N-more complicated visually

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65
Q

When is relative molecular mass used

A

When referring to simple molecules

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66
Q

How do you find relative molecular mass

A

Add up the relative atomic mass values of all the atom is in the molecule

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67
Q

When is relative formula mass used

A

For compounds that are ionic (or covalent)

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68
Q

How do you find the relative formula mass

A

Add up the relative atomic masses of all the ions in the formula unit

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69
Q

What is a mass spectra produced by

A

A mass spectrometer

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70
Q

What is a mass spectrometer

A

A device which are used to find out what samples are made up of by measuring the masses of their components

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71
Q

What can a mass spectra tell us

A

The relative isotopic masses and abundances of differnt elements

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72
Q

What is on the y-axis for a mass spectra

A

The abundance of ions as a %
- for an element the height of each peak gives the relative isotopic abundance

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73
Q

What is on the x-axis for a mass spectra

A

Units given as ‘mass/charge’ ration
(m /Z)

Since the charge on most ions is +1 you can often assume this axis is the relative isotopic mass

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74
Q

What can the mass spectra be used to work out

A

The relative atomic masses of differnt element

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75
Q

How to calculate relative atomic masses using a mass spectra

A

1) multiply each relative isotopic mass by its relative isotopic abundance and add up the results

2) divide by the sum of the isotopic abundances

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76
Q

What is avogadro constant

A

6.022 x 10 23^

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77
Q

Formula for finding number of moles from the number of atoms / molecules

A

Nu. Moles= nu. of particles you have / nu.Particles in a moles

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78
Q

T or f
Is molar mass the same as relative molecular mass

A

T

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79
Q

What is molar mass units

A

G mol-1

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80
Q

Moles formula

A

Mass /mr

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81
Q

What is the order of the orbitals

A

1s2 2s2 2p6 3s2 3p6 4s2

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82
Q

How many electrons in energy level 1

A

2

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83
Q

How many electrons in energy level 2

A

8

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84
Q

How many electrons in energy level 3

A

18

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85
Q

How many electrons in energy level 4

A

32

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86
Q

What are the differnt types of orbitals

A

S
P
D
F

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87
Q

What shape is orbital s

A

Spherical

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88
Q

What shape is orbital p

A

Dumb-bell shell

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89
Q

What are the rules of filling orbitals with electornx

A

Aufbau principle
Fill in order of inc energy

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90
Q

What is paulis exclusion principle

A

Within an orbital, electrons pair up with opposite spin so that the atom is as stable as possible.
Electrons in the same orbital must have opposite spin

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91
Q

What are the 3 rules for writing out electron configuration

A

1) lowest energy orbitals filled first
2) one electron occupies each orbital before pairing begins (hands rule)
3) no single orbital holds more than 2 electrons

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92
Q

When doing electron configuration for ions, they will lose electron in reverse order except what

A

4s will empty before 3D

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93
Q

If electrons are unpaired and therfore unbalanced it produces what

A

Natural repulsion between the electrons making the atom very unstable

The electrons may take on differnt a range,ents to improve stability

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94
Q

What is ionic bonding

A

Electrostatic forces of attraction between oppositely charged ions

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95
Q

What 2 elements are exceptions to the rule about electrons filling up orbitals

A

Copper and chromium
As half filled sub shells are more stable

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96
Q

How are water molecules part of the crystalline structure represented by

A

XXX.H20

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97
Q

What are the units for volume

A

1dm 3

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98
Q

Dm3 = xxxxxx cm3

A

1000

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99
Q

Concentration and volume equation

A

Conc. = mole / vol

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100
Q

What are the units for concentration

A

Mol / dm3
M

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101
Q

What does standard solution mean
+how

A

A solution of a known concentration.
An exact mass of solute is dissolved in a solvent and made up to an exact volume of solution

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102
Q

How do you go from mol/dm3 to g/dm3

A

Mol/dm3 x molar mass

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103
Q

What is avogadros law

A

At the same temperature and pressure, equal volumes of differnt gases contain the same number of molecules

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104
Q

What is molar volume

A

The molar gas volume Vm is the volume per mole of gas molecules at a stated temperature and pressure

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105
Q

At room temperature the molar gas volume is

A

24.0 dm3mol-1

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106
Q

Gas volume equation

A

Moles = vol / molar volume(24)

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107
Q

What is the method to carry out an experiment to determine the water crystallisation in hydrated crystals

A

1) weight empty crucible
2) add hydrated salts into weight crucible. Weigh the crucible containing the hydrated salt.
2) using a pipe-clay triangle, surpport the crucible contains the hydrated slat on a tripod. Heat the crucible and contents gently for around a Minuit. Then hear it strongly for another 3 mins
4) leave crucible To cool. Weigh crucible and anhydrous salt.
Calculate
1) calculate amount, in mol, of Anhydrous salt.
2) calculate mass + amount of water
3) find smallest whole number ratio

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108
Q

What are the bonds like in ionic bonding
And how does this effect melting and boiling points

A

Very strong
Attractions between the opposite ions bind the together.
It takes a lot of energy to break these bonds in order for the ions to move apart and change state.

  • so the melting and boiling points of such compounds are high (over 500°c)
  • the melting and boiling points usually increases as the change of the ions increase
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109
Q

Do solid ionic lattice conduct electricty

A

No
They cannot move as they are held tightly

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110
Q

Do liquid / aqueous solutions of ionic lattices conduct electricty

A

Ions are free to move and so can carry a charge

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111
Q

Why are ionic compounds hard

A

The strong attractive forces holding the lattice together

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112
Q

Why are ionic crystals brittle

A

When hit in the correct place ions are lined up alternatively to attract.
A force can push the rows so same ions now repel and the crystal will flake away

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113
Q

Why does it mean that water is a polar solvent

A

The distribution of electrons in the water is uneven giving it slight positive and slight negative chagre at each end

These charges attract the positive and negative ions and surround them according to charge . The ions are now hydrated , and the compound is dissolved

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114
Q

In order for ionic lattice to dissolve what needs to happen

A

Energy must be provided to break down the lattice . This is provided by the hydrates ion.
The water molecule must then attarct and surround the ions

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115
Q

In ionic compounds where ions have larger charges the ionic attraction may be too strong for what?

A

Water to break down the lattice structure

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116
Q

What does solubility of ionic lattice depend on

A

The relative strength of the attraction within the giant ionic lattice and the attraction between ions and water molecules

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117
Q

Where do we see stronger electrostatic forces of attraction between (ionic lattice)

A

ions with greater charge densities. Charge density is the charge an ion carries comapred to its size,

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118
Q

What determines the boiling point of a lattice

A

The attraction between the positive and negative ion

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119
Q

What holds a ionic lattice together

A

The electrostatic attraction between the positive ions and negative ions and the negative ions produce a strong giant ionic lattice

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120
Q

What are ionic compounds generally

A

Crystals with straight edges.
Suggesting that the ions within the crystal line up alternatively in repeating straight lines

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121
Q

What is standard solution

A

A solution of a known conc.

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122
Q

What is avogadros law

A

At the same temp. And pressure, equal volumes of different gases contain the same number of molecules

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123
Q

What is molar volume

A

The molar gas volume Vm is the volume per mole of gas molecules at a stated temp and pressure

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124
Q

What is Dative covalent bonding

A

Both bonding electrons originate from the same atom

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125
Q

What is molar gas volume

A

The space that one mole of a gas occupies at a certain temp and pressure

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126
Q

Units of molar gas volume

A

Dm3 mol-1

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127
Q

At room temp and pressure molar gas volume is 24
Add more info (numbers)

A

298K (25°c) and 101.3KPa

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128
Q

Number of moles =
(Molar gas volume)

A

Volume in dm3
———————
Molar gas volume

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129
Q

What is rhe ideal gas equation

A

pV = nRT

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130
Q

What does the ideal gas equation let you find

A

The number of moles in a certain volume at any pressure and temperature

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131
Q

What do the letters stand for in pV = nRT
+ units

A

P = pressure (Pa)
v = volume (m ^3)
n = number of moles
R = 8.314 J K-1 mol -1
T = temperature (K) (K = C° + 273)

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132
Q

What do balanced equations have

A

The same number of atoms on both sides

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133
Q

How do you know if an ionic equation is wrong

A

if the charges arnt balanced as well as the atoms
If the charges don’t balance the equation isn’t right

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134
Q

What does the reaction stoichiometry tell you

A

The ratios reactants to products

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135
Q

There are lots of ions that are made up of a group of atoms with an overall change
What are these called

A

Molecular ions

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136
Q

Charges in ionic compounds are always …

A

Balanced

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137
Q

What is formed when acids and bases react

A

Water and salt

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138
Q

What are salts

A

Ionic compounds

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139
Q

All solid salts consist of..

In some salts..

A

a lattice of positive and negative ions.

Water molecules are incorporated in the lattice too

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140
Q

What is the water in a lattice called

A

Water of crystallisation

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141
Q

As solid slat contains water of crystallisation is …

A

Hydrated

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142
Q

When is a salt anhydrous

A

When it doesn’t contain water of crystallisation

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143
Q

T or f
Ome mole of a particular hydrated salt always has the same number of moles of water of crystallisation

A

T

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144
Q

What do many hydrated slats loose when heated

A

Their water of crystallisation

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145
Q

The electrons in energy levels are all given numbers known as what?

A

Primcipal quantum numbers

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146
Q

Shells further away from the nucleus have … energy ( and a … principle quantum number)

A

Higher
Larger

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147
Q

What are energy levels further divided into

A

Sub-shells

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148
Q

What are the subshells called

A

S
P
D
F

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149
Q

How many orbitals in s subshell

A

1

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150
Q

How many orbitals in the p subshell

A

3

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151
Q

How many orbitals in the d subshell

A

5

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152
Q

How many orbitals in the d subshell

A

5

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153
Q

How many orbitals in the f subshell

A

7

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154
Q

Max electroms in s subshell

A

2

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155
Q

Max elevtrons in p subshell

A

6

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156
Q

Max electrons in d subshell

A

10

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157
Q

Max electrons in f subshell

A

14

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158
Q

Max electrons in first energy level

A

2

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159
Q

Max electrons in second energy level

A

8

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160
Q

Max electrons in 3rd energy level

A

18

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161
Q

Max electrons in 4th energy level

A

32

162
Q

T or f
Orbitals with in the same subshell have the same energy

A

T

163
Q

If there are two electrons in an orbital they have to - —-
+ what is this called

A

Spin in opposite directions
Spin- pairing

164
Q

What shape are s orbitals

A

Spherical

165
Q

What shape are p orbitals

A

Dumbbell shape

166
Q

There are 3 p orbitals and they are at ……. To each other

A

Right angles

167
Q

Rules for filling up energy levels

A

1) electrons fill up the lowest energy sub shells first

2) electrons fill orbitals with the same energy singly before they start pairing up

168
Q

Definition of ionincbonding

A

An ionic bond is an electrostatic attraction between two oppositely charged ions

169
Q

What do dot and cross diagrams show

A

The arrangement of electrons in an atom or ion.
Each electron is represented by a dot or cross.
They can show which atom the electrons in a bind originally came from

170
Q

Why is a giant ionic lattice called giant

A

It’s made up of the same basic repeating unit over and over again

171
Q

Why does giant ionic bonding form

A

Ion is electrostically attracted in all directions to ions of the opposite charge

172
Q

Why do atoms form covalent bonds

A

To gain a full outer shell

173
Q

What is a covalent bond

A

A covalent bond is the strong electrostatic attraction. Between a shared pair of electron and the nuclei of the bonded atoms

174
Q

Do all atoms in covalent bonds have 8 electrons in their outer shells

A

No
Some have less
A few use d orbitals to ‘expand the octet’ - and have more than 8

175
Q

What is average bond enthalpy

A

Measures energy required to break a specific bond

176
Q

The stronger the covalent bond the …. Energy required to break it so the …. Value of bond enthalpy

A

More
Greater

177
Q

In covalent bonds
Is there ever more than one pair of electrons

A

Yes
Some have double or triple

178
Q

What is the ideal gas constant

A

8.314

179
Q

How to turn c° in K

A

C° + 273

180
Q

How to turn cm3 into m 3

A

Divide 1 000 000

181
Q

What can the ideal gas equation be drivied from

A

Coming bottles law, Charles law + avogadros law

Also can be from kinetic theory of gases

182
Q

What two key assumptions is the ideal gas equation based on

A
  • no intermolecular forces between gas particles
  • the volume occupies by the molecules themselves is negligible compared to he volume of the container
183
Q

What other assumptions are made in the ideal gas equation

A
  • gases are made of molecules which are in constant random motion in straight lines
  • the molecules behave as rigid spheres
  • pressure is sure to collisions between molecules are the walls of the container
  • all collisions are perfectly elastic (no loss of kinetic energy )
  • the temperature of the gas is proportional to he average kinetic energy of the molecules l
184
Q

What does VSEPR stand for

A

Valence shell electron pair repulsion theory

185
Q

What is electron pair repulsion theory

A

Electrons will take up positions as far away from each other as possible to minimise repulsion

186
Q

What can co2 be shape be described as

A

Linear

187
Q

Explain a linear molecule

A

2 atoms around a central atom
2 regions of bonded pairs on opposite sides
180°

188
Q

What is the molecule shape when there are 3 atoms around a central atom and with what angle

A

Trigonal planar shape

189
Q

What are the angles in a trigonal planar shape

A

120°

190
Q

What does a dotted line when what draw rings molecule shapes

A

Away

191
Q

What does the wedge mean when drawring the 3D structures of molecules

A

Towards

192
Q

What is rhe shape when 4 atoms are bonded around the central atom

A

TetrHedral shape

193
Q

What are the angles in a tetrahedral shape

A

109.5

194
Q

What is the name of the shape for 5 atoms around the centre atom

A

Trigonal bipyramidal

195
Q

What are the angles in the Trigonal bipyramidal shape

A

120°
And
90°

196
Q

What is the name of the shape where there are 6 atoms around the central atom

A

Octrahedral

197
Q

What are the angles in an octrahedral

A

All 90°

198
Q

What are the order of biggest to smallest angles in shapes of molecules

A

Lone pair / lone pair angles are the biggest

Lone pair / bonding pair angles are the second biggest

Binding pair / bonding pair bond angles are the smallest

199
Q

What is the way of predicting molecular shape known as

A

Electron pair repulsion theory

200
Q

What are the bond angles in methane and why

A

No lone pairs so all angles 109.5°

201
Q

What are the angles of ammonia and why

A

1 lone pair
All 3 angles are 107°

202
Q

Water angles in molecular shape and why

A

2 lone pairs reduce the bond angle
104.5°

203
Q

How do you predict the shape of a molecule

A

Find the central atom work out the number of electrons in the outer shell of the central atom.

Then work out how many electrons are shared with the central atom

Add up the electrons and divide by 2 to find the number of electron pairs on the central atom and if you have an ion remember to count for its charge

compare the number of electron pairs with a number of bonds to find the number of lone pairs.

You can then use the number of electron pairs and the number of lone pairs and bonding centres around the central atom to work at the shape.

204
Q

What is rhe shape when the era re no lone pairs but 4 electron pairs around the centrato,

A

Tetrahedral

205
Q

What is the shape when there is one lone pair and 4 electron pairs around th e central atom

A

Trigonal pyramidal

206
Q

What is the shape when there are 2 lone pairs and 4 electron pairs around the central atom

A

Nonlinear or bent

207
Q

What is theoretical yield

A

Mass of products that should be made in a reaction if no chemicals are ‘lost’ in the process.

208
Q

What is percentage yield

A

The actually amount of product you collect

209
Q

What is the percentage yield equation

A

Actual yield
——————. X 100
Theoretical yield

210
Q

What is atom economy

A

Measure of the proportion of reactant to atoms that become part of the desired product in the balenced chemical equation

211
Q

Atom economy equation

A

Molecular mass of desired products
————————————————— X 100
Sum of molecular masses of all products

212
Q

When is the atom economy always 100%

A

In an addition reaction
The reactants combine to form a single product

213
Q

Why is low atom economy bad

A

Lots of waste
Costs money to separate desired prices form the waste product and more money to dispose of the water products safely so they don’t harm the environment

Reactant chemicals are usually costly so it’s a waste of money if a high proportion end up as useless products

214
Q

Why are reaction with low atom economy less sustainable

A

Many raw materials are in limited supply, so it makes sense to use them efficiently so they last as long as possible.

215
Q

Apart form atom economy and % yield what else is considered in making a reaction sustainable

A

Temp and pressure
Lower temps and pressure are cheaper and better for the environment

Raw materials that come from renewable sources are better for the environment,ny that materials from non renewable sources

216
Q

Why might a reaction not give a 100% yield

A

1) some reactions may not go to completion (not finished / reversible)
2) some products may be lost when it is separated for the reaction mixture
3) unexpected reaction might occur

217
Q

What is a standard solution

A

A solution of a known concentration made by dissolving a primary standard in a suitable solvent

218
Q

What is a concentrated solution

A

A solution that has a high concentration of solute

219
Q

What is a primary standard

A

a soluble solid compound that is very pure, with a consistent formula that does not change on exposure to the atmosphere, and has a relatively high molar mass.

220
Q

In order for a solid to be suitable for a primary standard a compound must…?

A

· be readily available (and inexpensive)
· be very pure (analytical reagent grade, 99.9% pure)
· have a known formula (e.g. degree of hydration must not change)
· be unaltered in air during weighing, that is, the compound must
NOT(a) absorb moisture from the air
(b) absorb carbon dioxide from the air
(c) be oxidised by air
· have a relatively large molar mass (to minimise errors during
weighing)
· be soluble (in the solvent used, and under the conditions used, to
make the solution)

221
Q

Why are most alkali hydroxides not suitable for use as a primary standard

A

· readily absorb moisture, H2O, from the atmosphere
· readily absorb carbon dioxide, CO2, from the atmosphere

222
Q

Whya re HCL and sulfuric acid not suitable for use as a primary standard

A

although they are both commercially available as concentrated solutions that are easily diluted, the concentration of the “concentrated” solution is NOT accurately known.

223
Q

Why is nitric acid not suitable for a primary standard

A

Always contains a little nitrous acid, which has destructive action on many acid-base indicators

224
Q

What are some standards that are suitable as primary standards for acid-base tritiations

A

· anhydrous sodium hydrogen carbonate, NaHCO3, a weak base (we will use this one)
· potassium hydrogen phthalate, KH(C8H4O4), a weak acid

225
Q

What is the method to make a standard solution

A
  1. Weigh by difference 2.65g of Na2CO3 into a beaker
  2. Add a small amount of distilled water to the beaker containing the Na2CO3 to dissolve it, stirring with a glass rod.
  3. Rinse weighting boat with distilled water and add the washing to the boat
  4. Use a funnel boat with distilled water and add the washing to the Beaker
  5. Rinse beaker + glass rod with distilled water + add the washing to the flask
  6. Make flask up to the graduation, made with distilled water (bottom of miscues on the line)
  7. Add stopped and shake the flask by inverting 15 times
226
Q

What are acids

A

Proton donors.
When mixed with water, all acids release hydrogen ions - H+ (these are just protons, but you never get them by themselves in water they always combine with h20 to form, h30+)

227
Q

What are alkalis

A

Bases that are soluble in water
They release OH- ions in solution

228
Q

What do alkalis produce in an aqueous solution

A

OH-

229
Q

How would u describe the reaction between acid and water and bases and water

A

Reversible

230
Q

Explain a strong acid

A

Nearly all the acid will dissociate / ionise in water, and nearly all the H+ ions released

231
Q

Are strong acids reversible,

A

Yes, the equilibrium lies far on the right

232
Q

Explain strong bases

A

The forward reaction is favoured so nearly all the bases dissociate in water and lots of OH- ions are released

233
Q

Explain weak acids

A

Backwards reaction favoured
So only small amounts of the acid will dissociate in water and only a few H+ ions are released

234
Q

Explain weak bases

A

Ionise only slightly in water
Backwards reaction is favoured so only a s,all amount of base dissociated and only a few OH- Ions are released

235
Q

Give an example of a strong acid

A

HCL

236
Q

Give an example of a strong base

A

NaOH

237
Q

Give an example of a weak acid

A

CH3COOH
Ethanoic acid

238
Q

Give an example of a weak base

A

NH3

239
Q

Acid + alkalis —->

A

salt + water

240
Q

Explain how acid + alkali = salt + water

A

1) it’s the hydrogen ions released by the acid and the hydroxide ions released by the alkali that combine to form water
2) you get a salt when the hydrogen ions in the acid are replaced by metal ions or ammonium ions form the alkali

241
Q

What are salts called from sulfuric acid

A

Sulfates

242
Q

What are salts called from hcl

A

Chlorides

243
Q

What are salts called from nitric acid

A

Nitrates

244
Q

Why is ammonia an exception for neutralisation equation

A

It doesn’t directly produce ions, but aqueous ammonia is still an alkali. This is because the reaction between ammonia and water produces hydroxide ions, ammonia accepts a hydrogen ion form water molecules forming an ammonium ion and a hydroxide ion. So Ammonia can neutralise acids

245
Q

Metal + acid =

A

Metal salt + hydrogen

246
Q

Metal oxide + acid =

A

Salt + water

247
Q

Metal hydroxide + acid =

A

Slat + water

248
Q

Metal carbonate + acid =

A

Metal salt + carbon dioxide + water

249
Q

What does the neutralise ionic equation show + what is it
( metal hydroxide + acid)

A

OH- + H20 -> H20
Shows that a proton is transferred from the acid to the hydroxide ion .

250
Q

Ammonia + acid =

A

Ammonium salt

251
Q

What do all acids contain

A

Hydrogen

252
Q

When acids dissolve in water and release 1 H+ ion what is it called

A

Monoprotic

253
Q

When acids dissolve in water and release 2 H+ ion what is it called

A

Diprotic

254
Q

Examples of bases

A

Metal oxides
Hydroxides
Carbonates
Ammonia

255
Q

Base definition

A

a compound that neutralises an acid to form a salt / a proton acceptor

256
Q

Salt definition

A

The products of a reaction in which the H+ ions form the acid are replaced by metal or ammonium ion

257
Q

What is electronegativity

A

An atoms ability to attract the electron pair in a covalent bond

258
Q

What is the most electronegative element

A

Fluorine

259
Q

What is electronegativity measured on

A

Pauling scale

260
Q

Describe the trends ofelectronegitivity on the Pauling scale

A

Increases across periods and decreases down groups ignoring noble gases)

261
Q

What makes a covalent bond polar

A

When the two atoms have different electro negativities, the bonding electrons are pulled towards the more electronegative atom

262
Q

In apolarbond thedifference inelectronegativity between the two atoms causes what?

A

A permanent dipole

263
Q

What is a dipole

A

A difference in charge between the two atoms caused by a shift in electron density in the bond, the greater the difference in electronegativity I the more polar the bond

264
Q

Are covalent bonds in diatomic gases polar? And why?

A

Nonpolar because the atoms have equal electronegativities and so the electrons are equally attracted to both nuclei

265
Q

Are bonds polar when the two elements have similar electronegativities

A

No

266
Q

What do polarbonds have

A

Permanent dipoles

267
Q

What determines whether or not the molecule will have an overall dipole

A

The arrangement of polar bonds in a molecule

268
Q

If the polar bonds are arranged symmetrically _ what does that mean for the polarity

A

Dipoles cancel each other out so the molecule has No overall dipole and is non-polar

269
Q

What if polar bonds are arranged so they don’t cancel each other out (dipole)

A

The charge is arranged unevenly across the whole molecule, and it will have an overall dipole these are polar

270
Q

What are the only bonds that ave purely covalent

A

Bunds between atoms of a single element/ like diatomic gases

271
Q

Most compounds are somewhere in between the two extremes of ionic covalent what does this mean for the properties

A

Often got coin + covalent properties

272
Q

What can you use electro negativity to predict

A

What type y bonding would occur between two atoms the higher the difference in electro negativitys, the more ionic n character the bonding becomes

273
Q

What are intermolecular forces

A

Forces between molecules

274
Q

Are intermolecular forces strong or weak

A

Weak
Much weaker than ionic, covalent or metallic

275
Q

What are the 3 types of intermolecular forces

A

1) induced dipole-dipole or London (dispersion) forces
2) permanent dipole-dipole interactions
3) hydrogen bonding

276
Q

What is the strongest type of intermolecular forces

A

Hydrogen bonding

277
Q

What types of intermolecular forces does van den walls forces refer to

A

Induced dipole-dipole or London (dispersion) forces
Permanent dipole-dipole interactions

278
Q

What do include dipole-dipole forces cause

A

All atoms and molecules to be attracted to each other

279
Q

Why are induced dipole-dipole forces found between all atoms and molecules

A

1) elections in charge clouds are always moving reculy quickly, at any particular moment the electrons in an atom are likely to be more to one side than the other_ at this moment the atom would have A temporary dipole
2) this dipole can cause another temporary induced) dipole in opposite directions on a neighbouring atom the two dipoles are then attracted to each other.
3) the second dipole can cause yet another dipole in a third atom
4) because the electrons are constantly moving - the dipoles are being created and destroyed all the time. Even though the dipoles keep changing, the overall effect is for the atoms to be attracted to each other

280
Q

Are all dipole-dipole forces the same strange, and why?

A

No, larger molecules have larger electron clouds, meaning stronger include dipole-dipole forces, molecules with greater surface areas also have stronger induced dipole-dipole forces because they have A bigger exposed electron cloud.

281
Q

Why do liquids with stronger induced dipole-dipole forces will have higher boiling points

A

When you boil a liquid you need to overcome the intermolecular forces, so that the particles can escape from the liquid surface

282
Q

What physical properties does dipole-dipole forces affect

A

Boiling/ meaning points
Viscosity

283
Q

What holds iodine molecules together in a lattice

A

Induced dipole-dipole forces.
Iodine atoms are held together in pairs by strong covalent bonds
Then molecules are held together in a molecular lattice arrangement by weak induced dipole-dipole attractions

284
Q

What are permanent dipole-dipole interactions

A

The positive and negative changes on polar molecules cause weak electrostatic forces of attraction between molecules

285
Q

Do permanent dipole-dipole interactions happen in addition to or instead of induced dipole-dipole interactions

A

In addition to

286
Q

When can hydrogen bonding happen

A

Can only happen when hydrogen is covalently bonded to fluorine - nitrogen or oxygen

287
Q

Why can hydrogen bonding only happen when its bonded toflorine/ nitrogen or O2

A

Hydrogen has a high charge density because its so small and fluorine,hydrogen , oxygen are very electronegative the bond is so polarised that a weak bond forms between the hydrogen of ome molecule and a lone pair of electrons on the fluorine, nitrogen or oxygen in another molecule

288
Q

Molecules which have hydrogen bonding usually contain what groups

A

-OH or -NH

289
Q

How do hydrogen bonds effect the properties of substances

A

They are soluble In water + have higher boiling and freezing points

290
Q

Why does water, ammonia and hydrogen fluoride generally have higher boiling points if u can compare then with other hydrides in their group

A

There is extra energy needed to break the hydrogen bonds

291
Q

In ice how are h2o molecules held together

A

In a lattice by hydrogen bonds

292
Q

Does ice or water have more hydrogen bonds

A

Ice, as when ice melts hydrogen bonds are broken

293
Q

Is ice or water more dense

A

Ice is less dense as ice has more hydrogen bonds (which are relatively long)

294
Q

In general what is the main factor that determines the boiling points of a substance

A

The strength of the induced dipole-dipole forces (unless the molecule can form hydrogen bonds)

295
Q

Explain why the boiling points of the group 7 increase from HCI to HI

A

Although the permanent dipole-dipole interactions are decreasing, the nu. Of elections in the molecule increases so the strength of the induced dipole-dipole interactions also increases.

296
Q

If you have two molecules with a similar number of electrons then the strength of their induced dipole-dipole interactions will be – - —-?

A

Similar

297
Q

If one of the substances las molecules that are more polar than the other. What will it meanfor the permanent dipole -dipole interactions

A

Stronger permanent dipole-dipole interactions and so a higher boiling point

298
Q

Why do simple covalent compounds have low melting and boiling points

A

The intermolecular forces that hold together the molecules in simple covalent compounds are weak so don’t need much energy to break. So the melting and boiling points are normally low - they are often liquids or gases at room temperature

299
Q

AS intermolecular forces get stronger, melting and boiling points —–

A

Increase

300
Q

As water is a polar molecule, what does it only tend to dissolve

A

Other polar substances

301
Q

Why are compounds with hydrogen bonds soluble

A

They can form hydrogen bonds with water molecules

302
Q

Molecules that only have induced dipole- dipole forcesare soluble or insoluble

A

Insoluble

303
Q

Why don’t simple covalent compounds conduct electricity

A

They are overall uncharged so they can’t conduct electricity

304
Q

To melt or boil a simple covalent compounds _ you have to overcome?

A

The intermolecular forces that hold the molecule together
You don’t need to break the much stronger covalent bonds that hold the atoms together in the molecules

305
Q

What 3 factors effectelectronegativity

A

The distance between the nucleus and the outermost electrons
Ie. Atomic radius

Sheilding

The nuclear charge

306
Q

Describe how the distance between the nucleus and the outmost electrons ( atomic radius) affects electro negativity

A

The smaller the atom the closer the nucleus is to the shared outer electrons, so the greater the electronegativity

307
Q

Explain how The distance between the nucleus and the outermost electrons
Ie. Atomic radius is shown as a trend on the periodic table

A

The electro negativity increases as you go up the group as the closer the electrons are to the nucleus (smaller atomic ractius)

308
Q

Describe now shielding effects electro negativity

A

The larger the atom the more energy levels between the nucleus and the outer energy level so the lessor the electronegativity

309
Q

Explain how shielding is shown as a trend on the periodic table as effecting Electronegativity

A

The electronegativity increases as you go up the group as the less shielding

310
Q

Describe how nuclear charge effects nuclear charge

A

The large the nuclear charge (for a given shielding effect) the greater the electro negativity

311
Q

Explain how the nuclear charge is shown as a trend on the periodic table as effecting Electronegativity

A

Increasing electronegativity as you go right along the period table as there are more protons in the nucleus so more nuclear charge( number of charges on the nucleus increases) This attracts the bonding pair of electrons more strongly

312
Q

W hat is a polar bond

A

Covalent bond when there is a separation of charge between one end and the other

313
Q

How do polar bonds arise

A

When one is more slightly electronegative than the other.
The more electronegative atom will attract the electron pair more than the other. So one end becomes slightly negative and the other becomes slightly positive.

314
Q

What is polarity

A

The unequal sharing of electrons in a bond- if the electron sharing is equal the bond is described as non -polar

315
Q

Larger molecules may contain polar bonds but why are they not necessarily polar

A

It depends on the shape
A symmetrical molecule ww not be polar even if the individual bonds are polar. The individual dipoles on the bonds cancel out’ due to the symmetrical shape of the molecule.
There is No met dipole its non polar

316
Q

What are intermolecular forces

A

Wear attractive forces between molecules

317
Q

What is the bond enthalpy for London forces

A

1-10

318
Q

What is the bond ethalpy for permanent dipole-permanent dipole interactions

A

3-25

319
Q

What is the bond enthalpy for single covalent bond

A

IS0-500

320
Q

What is the bond ethalpy for hydrogen bonding

A

10 - 40

321
Q

What intermolecular forces doall molecules have

A

London forces

322
Q

What increases the strength of the London forces

A

More electrons

323
Q

What leads to a permanent dipole

A

Great difference in electronegativity

324
Q

Where is hydrogen bonding found

A

Large electro negative atoms
Lone pairs
So mainly oxygen, nitrogen, fluorine

325
Q

What 3 elements do hydrogen bunds form between

A

Oxygen
Nitrogen

Fluorine

326
Q

Why is water a liquid at room temp.

A

Hydrogen bonding requires lots of energy to overcome
This increase the melting point and boiling point of water

327
Q

How do London forcesform between molecules

A

Instantaneous dipoles form due to constant movement of electrons - which causes an induced dipole in a neighbouring molecules. These form London forces between the molecules (weak attractiveforces form)

328
Q

Why are pd-pd forces stronger than London forces

A

Greater electronegativity difference between bonded atoms form stronger attractive forces

329
Q

Why is ice less dense than water

A

Ice has an open lattlee, h bonds hold water molecules apart
When ice melts, H20 molecules move closer together
So ice is less dense then water

330
Q

Why does water exist as a liquid at room temperature?

A

Water has hydrogen bonds in combination with DP-DP and London forces takes more energy to overcome and therfore is a liquid at room temp.

331
Q

What are tritations

A

Techniques used to accurately measure the volume of one solution that reacts exactly with another solution.
Used for finding the concentration of a solution,. Identifying unknown chemicals, finding the purity of a substance

332
Q

What is needed for valid results in tritations

A

Precise t accurate measurements needed.

333
Q

Is tritation quantitative or qualitative

A

Quantitative

334
Q

Describe the method for a tritation

A
  1. Use a pipette and pipette filler to add 25 cm3 of the unknown solution to a clean conical flask.
  2. Add a few drops of a suitable indicator and put the conical flask on a white tile.
  3. Fill the burette with your standard solution. Flush the tap through to remove any air bubbles. Ensure the burette is vertical. Record the initial volume in the burette to the nearest 0.05cm3.
  4. Slowly run the solution from the burette into the conical flask, swirling to mix. (The mixture may at first change colour, and then back again when swirled.)
  5. Stop adding the acid when the end-point is reached (when the colour first permanently changes). Note the final volume reading to the nearest 0.05cm3. This is your rough/trial titration to determine your approximate titre.
  6. Repeat steps 1 to 5 accurately until two results are repeatable (in close agreement). Ideally these should lie within 0.10 cm3 (concordant) of each other.
335
Q

In a tritation where is the known concentration placed

A

In the burette

336
Q

In a tritation where do you put the solution of an unknown concentration

A

In the conical flask

337
Q

What is the volume of solution added from the curette called in a tritation

A

Titre

338
Q

Name a suitable indicator for a tritation

A

Phenolphthalein
Methyl orange

339
Q

Describe the colour change in phenolphthalein

A

Clear in an acid
Pink un an alkali

340
Q

Describe the colour change for methyl orange

A

Red un acid
Yellow in alkali

341
Q

Why can universal indicator not be used un imitations

A

Needs to change colours at a specific ph - so can be used to precisely identify when the neutralisation reaction is complete.

342
Q

What are concordant results

A

Within 0.10 of each other

343
Q

What does a pipette measure

A

Fixed volumes

344
Q

What does A b urrette measure

A

Different volumes + lets you aced solutions drop by drop

345
Q

When taking a reacting from a bureette, where should you take the reading from

A

The bottom of the meniscus

346
Q

To make tritation readings more accurate, what should the readings be to the nearest

A

0.05 cm3

347
Q

What is the analyte in tritations

A

Concentration of the unknown solution

348
Q

What is concentration measured in

A

Mol dm -3

349
Q

How many Cm 3 in a litre

A

1000

350
Q

What errors could occur during tritation

A

The burette should be rinsed out with the substance that will be put in it
-If it is not rinsed out, the acid or alkaline added, may be diluted by residual water in the barrettes, or may react with substances left from a previous tritation
-This would lead to the concentration of substances being lower and a larger titre being delivered

-If the Jet space in the burette is not filled properly prior to commencing the tritation it will lead to errors if it then fill during tritation -leading to a larger than expected, titre reading

351
Q

What is concentration

A

The amount of solute in moles, dissolved in 1dm3

352
Q

What is deliquescent

A

Absorbed moisture from air

353
Q

What is efflorescent

A

Loses water of crystallisation

354
Q

What is the equation to find % uncertainty

A

Percentage uncertainty =
How many times equipment was used x uncertainty
————————————————————————- X 100
Meausred value

355
Q

What do all acids contain

A

Hydrogen

356
Q

What is a monoprotic acid

A

If 1+ ion released

357
Q

What is a diprotic acid

A

If 2+ ions released

358
Q

Give examples of some bases

A

Metal oxides
Hydroxides
Carbonates
Ammonia

359
Q

Defition of an acid

A

A species that releases H+ ions in aqueous solution / a proton donor

360
Q

Define an alkali

A

A type of base that dissolves in water to release OH- ions

361
Q

Define a base

A

A compound that neutralises an acid to from a salt / a proton acceptor

362
Q

Define a salt

A

The products of a reaction in which the H+ ions from the acid are replaced by metal or ammonium ions

363
Q

What is oxidising agent

A

substance which oxidises something else

364
Q

What is a reducing agent

A

Reduces something else

365
Q

What is oxidation number

A

Based on a set of rules that apply to atoms and can be thought of as the number of electrons involved in bonding to a differnt elements

366
Q

What is the oxidation number for elements

A

0

367
Q

How is oxidation number written

A

Sign placed before the number
Eg,
+2

368
Q

What is the oxidation. Number the same as
However…

A

The charge
However the sign comes be for the number

369
Q

What is hydrogens oxidation number normally

A

+1

370
Q

What is oxygens oxidation number usually

A

-2

371
Q

What do oxidation numbers add to (not on a charged ion )

A

0

372
Q

How can oxidation be described as in terms of hydrogen

A

Oxidation is loss of hydrogen

373
Q

How can reduction be described as in terms of hydrogen y

A

Reduction is gain of hydrogen

374
Q

Oxidising agents themselves are…

A

Reduced

375
Q

Reducing agents themselves are …

A

Oxidised

376
Q

What is the oxidation number for nitrite (lll)

A

+3

377
Q

What does the Roman numeral after a compound / element mean for oxidation number

A

That’s its oxidation number
Eg, iron (ll) is +2

378
Q

Oxidation number of oxygen us -2 except what

A

Peroxides

379
Q

The oxidation number of hydrogen is +1 except what

A

Hydrides

380
Q

Are most elements oxidation states the same

A

Yes

381
Q

What are group ones oxidation number

A

+1

382
Q

What is group 2s oxidation number

A

+2

383
Q

What are groups 7s oxidation number mainly

A

-1

384
Q

What element in group 7s oxidation number is always -1

A

Fluorine

385
Q

The oxidation of a simple, monatomic ions is….

A

:the same as its charge

386
Q

What is rhe oxidation charges in peroxides

A

-1

387
Q

If an element can have multiple oxidation numbers or isn’t in its normal oxidation state what does it have

A

Shows using Roman numerals

388
Q

Ions with names ending with -ate
Contain ….

A

Oxygen and another element

389
Q

Sometimes the other element except oxygen in an -ate ion can exist in differnt oxidation numbers and so forms…

A

Differnt -ate ions

390
Q

What is the reaction stoichiometry

A

How many moles of one reactant react with how many moles of another reactant

391
Q

What are polyprotic acids

A

Donate more than one proton

392
Q

An element is simultaneously oxidised and reduces Is called……

A

Disproportionation

393
Q

What are the rules to help complex half equations

A

1) balence atoms
Is you need more oxygen add h20
If you need for hydrogen add H+ ions
2) balence charge
- add electrons

394
Q

What is a redox reaction

A

Both oxidation and reduction take place to differnt elements

395
Q

An increase in oxidation number =

A

Oxidation

396
Q

A decrease in oxidation number =

A

Reduction

397
Q

When metal forms compounds is their oxidation number usually positive or negative

A

Positive

398
Q

What non metals form compounds do they generally have positive or negative oxidation number s

A

Negative

399
Q

Metals are,……. when they react with acids

A

Oxidesed

400
Q

Explain the redox with metal and acid

A

Metal atoms are oxidised, loosing electrons to from positive metal ions
The hydrogen ions in solution are red used gaining electrons and forming hydrogen molecules

A-level Chemistry <3 (9 decks)

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