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EL Flashcards

Elements of Life

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1
Q

Define an Isotope

A

Atoms of the same element with different numbers of neutrons but the same number of protons

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2
Q

What is the equation for relative atomic mass?

A

(Relative abundance x isotopic mass)+(relative abundance x isotopic mass) etc.
———————————————————
100

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3
Q

What is avagadros constant?

A

6.02 x1023

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4
Q

What is the equation linking moles, mass and Mr?

A

Moles(mol)= mass (g) / Mr

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5
Q

Define nuclear fusion

A

Lighter nuclei are fused together to form heavier nuclei, which releases and enormous amount of energy

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6
Q

In stars where are the heaviest elements found?

A

Centre

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7
Q

Why does a heavyweight star become unstable?

A

When the centre is iron, iron absorbs energy during fusion instead of releasing it

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8
Q

What happens to a heavyweight star?

A

Becomes unstable and explodes into a supernova. Disperses elements as gas and dust, restarting the cycle

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9
Q

What happens to a lightweight star?

A

Once they run out of hydrogen, the expand into a red giant. Eventually resulting in outer gases drifting away to leave a white dwarf

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10
Q

What are the differences between a lightweight and heavyweight star?

A
  • Lightweight stars can only do nuclear fusion of hydrogen, they are not as hot, they will last longer
  • Heavyweight stars can do nuclear fusion of elements up to iron, they are at higher temperatures and pressures but won’t last as long
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11
Q

What is mass spectrometry?

A

A measure of the atomic or molecular mass of different particles to find relative isotope abundances

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12
Q

How does mass spectrometry work?

A

Atoms or molecules with be ionised to positively charged cation. This ions are separated according to their mass (m) to charge(z) ratio, m/z

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13
Q

What is a mass spectrum?

A

A graph where the relative abundance of each ion can be calculated from the height of each peak

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14
Q

How are stars formed?

A

By the culmination of dust and gas originating from the big band

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15
Q

What is spectroscopy?

A

The study of how light and matter interact

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16
Q

How can we recognise elements in space?

A

Under certain conditions, a substance can absorb or emit electromagnetic radiation in a recognisable way. By analysing this radiation we can recognise and find out information about the substance

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17
Q

In the electromagnetic spectrum where is frequency to highest?

A

In the gamma ray area

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18
Q

State the electromagnetic spectrum starting and the largest wavelength

A

Radio frequency, microwaves, infrared, visible light, ultraviolet, x-rays, gamma rays

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19
Q

How does absorption spectra work?

A

Ions, atoms and molecules will absorb certain frequencies of the emitted radiation. These missing frequencies will show up as black lines and an absorption spectra

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20
Q

How is an absorption spectra carried out?

A

White light is passes through a cooler flame, then analysed

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21
Q

How does an emission spectra work?

A

When particle absorb radiation they will be raised from their ground state to an excited state. This energy then gets emitted and will appear as a coloured line on a black background

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22
Q

Define ground state

A

The lowest energy state

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23
Q

Define excited state

A

A higher state of energy

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24
Q

Lines ______ at higher frequencies

A

Converge

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25
Q

Define continuous spectrum

A

A seamless transition of colours or wavelengths, white light has a continuous spectrum

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26
Q

Light in absorption and emission spectra have a _________

A

Line spectrum

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27
Q

What term refers the the absorption and emission spectra of elements?

A

Atomic spectra

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28
Q

What is the balmer series?

A

Hydrogen emission spectrum in UV light

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29
Q

What is the lyman series?

A

Hydrogen emission spectrum in visible light

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30
Q

Explain wave theory

A

Light is a form of electromagnetic radiation and behaves like a wave which a characteristic wavelength and frequency

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31
Q

What is the speed of light travelling in a vaccuum?

A

3.00 x108 ms-1

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32
Q

What is the equation connecting speed of light, wavelength and frequency?

A

Speed of light,c,ms-1= wavelength, lambda, m x frequency,mew(v), s-1

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33
Q

Explain particle theory

A

Light can be explained as a stream of photons the energy of the photons depends on the light’s position in the EM spectrum
- proposed by Einstein in 1905

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34
Q

What are photons?

A

Tiny packet of energy

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35
Q

State the equation to find the energy of a photon

A

Energy of photon,E,J= Planck constant,h,JHz x Frequency,mew, s-1

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36
Q

What is the value of Planck constant?

A

6.63 x10-34 JHz or Js-1

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37
Q

Explain Bohrs theory

A

Bohr’s theory explains why the hydrogen atom only emits a specific number of certain frequencies. The light that a substance absorbs or emits is specific to itself. When an atom is excited atoms will become excited and when the fall back to ground state, they emit the extra energy or photo as EM radiation. Bohr’s theory explained both emission and absorption spectra

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38
Q

The energy of a photon is…..

A

The difference between two energy levels

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39
Q

Why was Bohr’s theory considered controversial?

A

It relied on the pry of the quantisation of energy

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40
Q

An electron can only….

A

Possess definite quantities of energy

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41
Q

Define quanta

A

Discreta packets of energy

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42
Q

If an electron is far away from the nucleus, does it have a higher or lower energy level?

A

Higher

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43
Q

A large gap between energy levels is an energy level diagram mean what?

A

The frequency is high

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44
Q

Why do lines on a spectra converge?

A

Lines on a spectra converge at higher frequencies as the higher energy levels are closer together

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45
Q

What are the similarities of an absorption and emission spectra?

A
  • Each element with have a characteristic atomic spectra
  • Both use light to identify elements/atoms
  • Lines at the same place and same thickness
  • lines converge at higher frequencies
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46
Q

What does an emission spectra look like?

A

Coloured lines on a black background

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47
Q

What does an absorption spectra look like?

A

Black lines on a coloured background

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48
Q

How do you choose the appropriate amount of significant figures?

A

Recognise the least amount of s.f. used in the equation

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49
Q

Why do flame tests work?

A

As energy raises the electrons to an excited state they will drop back down again. The colours are seen as the electrons fall back down to ground state. The colour depends on the wavelength which will be different for every element as the energy difference between energy levels are unique

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50
Q

Why are atomic spectra helpful?

A
  • Can be used to identify elements
  • Intensity of the lines show abundance
  • Provides information about the element’s structure
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51
Q

What is each shell labelled by?

A

A principle quantum number, with the higher numbers referring to the outer shells

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52
Q

What is the maximum number of electrons in the 2nd shell?

A

8

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53
Q

What is the maximum number of electrons in the 1st shell?

A

2

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54
Q

What is the maximum number of electrons in the 3rd shell?

A

18

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55
Q

What is the maximum number of electrons in the 4th shell?

A

32

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56
Q

What is the Aufbau principle?

A

Electrons fill the lowest energy shells first

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57
Q

What subshells are there, and how many electrons can they hold?

A
  • s-sub-shell - 2 electrons
  • p-sub-shell - 6 electrons
  • d-sub-shell - 10 electrons
  • f-sub-shell - 14 electrons
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58
Q

Is the energy of sub-shells fixed?

A

No, there are different energies per type of sub-shell and the energy of each sub-shell depends on the rise and fall off the charge on the nucleus

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59
Q

How many orbitals does a s-sub-shell have?

A

1 s-orbital

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60
Q

How many orbitals does a p-sub-shell have?

A

3 p-orbitals

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61
Q

How many orbitals does a d-sub-shell have?

A

5 d-orbitals

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62
Q

How many orbitals does a f-sub-shell have?

A

7 f-orbitals

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63
Q

What is Heisenburg’s Uncertainty Principle?

A

The idea that the position of an electron can’t be mapped out exactly so an atomic orbital refers the the 95% probability of finding an electron that region

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64
Q

What are atomic orbitals?

A

A reference to a particular region of space around the nucleus

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65
Q

What is the max. number of electrons able to fit in an orbital?

A

2

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66
Q

Electrons can only occupy the same orbital is they have ___________? This is know as the ______________?

A

Opposite spins (clockwise and anti-clockwise)
Pauli Exclusion Principle

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67
Q

What shape in a s-orbital?

A

Spherical

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68
Q

What shape is a p-orbital?

A

Dumbbell-shaped

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69
Q

What is electronic configuration?

A

The arrangement of electrons in shells and orbitals

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70
Q

What is Hund’s principle of maximum multiplicity? And explain

A

Orbital are first all filled singly and only fill with two electrons once every orbital has been singly filled. This happens because this arrangement is one which keeps the electrons as far apart as possible.

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71
Q

How are orbitals filled?

A

To produce the lowest energy arrangement possible

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72
Q

Electrons in singly occupied shells have __________.

A

Parallel spins

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72
Q

Write the electronic configuration of oxygen

A

1s1 2s2 2p2
This can be shown with boxes for each orbital and arrows within the boxes to signify electrons. Usually drawn going up the page

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73
Q

What elements do not follow the electronic configuration pattern? And why?

A

Elements in Period 3 up to nickel (besides chromium and copper). This is because The 4s sub-shell has a lower energy level that the 3d sub-shell when empty, but higher when filled. So 4s will fill first and empty first

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73
Q

How can electronic configurations be abbreviated?

A

Put the noble gas symbol and the the extra sub-shells e.g. [Ar] 4s2 3d1

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73
Q

How did Mendeleev arrange the elements?

A
  • In order of increasing atomic mass
  • Elements with the same properties were in the same vertical group
  • Mendeleev left gaps for elements that he predicted and made predictions about these elements properties
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74
Q

Who are the main chemists who contributed to the development of the periodic table?

A
  • Johann Dobereiner
  • Lothar Meyer
  • John Newlands
  • Dmitri Mendeleev
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75
Q

How are some elements made synthetically?

A

Made by bombarding uranium atoms with nuetrons

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75
Q

Why was Mendeleev’s grouping inaccurate?

A

Due to the existence of unknown isotopes

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76
Q

How is the modern periodic table arranged?

A
  • In order of increasing atomic number
  • Based of Mendeleev’s model from 1869
  • Split into four blocks: s,p,d,f
  • Vertical columns known as groups that are based on the electron number in the outer shell, and similar physical properties
  • Horizontal rows called periods that are based on the number of the shell being filled
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76
Q

Define periodicity

A

the occurrence of periodic patterns

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77
Q

What happens when elements are melted/boiled?

A

the intermolecular forces between the atoms must be broken. The energy needed to break these bonds depends on the strength of the bonds

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78
Q

How are the chemical properties of an element decided?

A

Decided by the (number of) electrons in the outer shell

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78
Q

What is the pattern across periods for m.p. and b.p.?

A

An initial increase but the fall dramatically

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78
Q

Define closed shell arrangements

A

A particularly stable arrangement where sub-shells have been fully occupied by electrons

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79
Q

How are elements sorted into s/p/d/f block?

A

Based off of the current sub-shell being filled by electrons

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80
Q

How does atomic radii change the reactivity of an atom?

A

As the number of shells increase, shielding increases and this leads to the first ionisation enthalpy decreasing (reactivity increases)

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81
Q

How does atomic radii change across a period?

A

Across a period the number of protons in the nucleus increases but the number of shells stays the same. The number of electrons in the outer shell also increase, this increases attraction between the outer shell and the nucleus leading the shell to be held closer to the nucleus and therefore decreasing atomic radii

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82
Q

What is shielding?

A

Shielding refers to the core electrons repelling the outer electrons, which lowers the effective charge of the nucleus on the outer electrons

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83
Q

To become stable an atom must what?

A

Accept or donate an electron/s in order to make a full outer shell and therefore form an ion

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84
Q

Define organic species

A

Molecules that contain not only carbon, but also at least one other element

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85
Q

Describe the Miller-Urey experiment

A

The scientist put methane, ammonia, carbon dioxide and water into a flask. They heated the mixture and subjected it to electrical discharge that would have mimicked the earths early atmosphere. They found that amino acids had been produced meaning that proteins could’ve been made in earths atmosphere or dense clouds in space. This was used to explain the origin of life on earth.

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86
Q

How many electrons are linked to the stability of noble gases?

A

8 electrons besides Helium

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87
Q

Define covalent bonds

A

A bond formed by a shared pair of electrons, usually between two none metals

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88
Q

In terms of charges, describe covalent bonding.

A

The negatively charged electrons are attracted to the positive charges of both nuclei. The attraction overcomes the repulsion.

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89
Q

Define bonding pairs

A

The electron pairs that form the bonds

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90
Q

Define lone pairs

A

Electron pairs that aren’t involved in the bonds / form bonds

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91
Q

How do instantaneous induced dipole bonds form?

A

Electrons are always moving so at any instant there can be a concentration of electrons on one side. This causes a difference in charge, leading to an attraction/force between the two atoms.
Can also be known as intermolecular bonds

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92
Q

Define a dative covalent bond

A

A shared pair of electrons that have been supplied by only one of the atoms
Can be shown in dot and cross models as an arrow instead of a line
Also known as coordinate bonds

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93
Q

What are covalent intramolecular bonds?

A

The bonds within a molecule that are usually very strong

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94
Q

What bonds are weaker? Intramolecular or electrons static attractions(intermolecular)
And what does this mean?

A

Intermolecular bonds are weaker meaning that little energy is needed to overcome these bonds. This results in low melting and boiling points.

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95
Q

True or false: simple molecules can conduct electricity?

A

False, there are no free electrons/charged particles

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96
Q

What is the electron repulsion theory?

A

As similar charges repel, the electron pairs repel each other as far apart as possible

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97
Q

How do lone pairs affect the shape of a molecule?

A

Lone pairs repel more strongly, so will reduce the bond angle by 2.5º

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98
Q

In drawings of molecules, what does the line represent?

A

The bonds that line on the plane of the paper

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99
Q

In drawings of molecules what do solid triangles represent?

A

The bonds that come towards you

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100
Q

What do solid triangles represent in molecules drawings?

A

Bonds that come towards you

101
Q

What do dotted lines/triangles represent in molecules drawings?

A

Bonds that go away from you

102
Q

What is the octet rule

A

Atoms should have a complete outer shell of electrons

103
Q

Covalent bonds form to pair up electrons to obey the ________?

A

Octet rule

104
Q

What is expanding the octet?

A

Where bonding atoms may have more then 8 electrons in the outer shell
Groups 5-7

105
Q

In group 5 how many covalent bonds can be made and which elements?

A

3 or 5 bonds, P and As

106
Q

In group 6 how many covalent bonds can be made and which elements?

A

2,4 and 6 bonds, S, Se and Te

107
Q

In group 7 how many covalent bonds can be made and which elements?

A

1,3,5 or 7, Cl,Br,I and At

108
Q

What is a better rule that the octet rule?

A
  • Unpaired electrons pair up
  • The maximum number of electrons that can pair it is the equivalent to the number of electrons in the outer shell
109
Q

Describe a linear molecule

A

Regions of charge: 2
Bonding pairs: 2
Lone pairs: 0
Bond angle: 180

110
Q

Describe a planar triangular molecule

A

Regions of charge: 3
Bonding pairs: 3
Lone pairs: 0
Bond angle: 120

111
Q

Describe a tetrahedral molecule

A

Regions of charge: 4
Bonding pairs: 4
Lone pairs: 0
Bond angle: 109.5

112
Q

Describe a pyramidal molecule

A

Regions of charge: 4
Bonding pairs: 3
Lone pairs: 1
Bond angle: 109.5-2.5= 107

113
Q

Describe a non linear molecule

A

Regions of charge: 4
Bonding pairs: 2
Lone pairs: 2
Bond angle: 109.5-5= 104.5
OR
Regions of charge: 3
Bonding pairs: 2
Lone pairs: 1
Bond angle: 120-2.5= 117.5

114
Q

Describe a bipyramidal molecule

A

Regions of charge: 5
Bonding pairs: 5
Lone pairs: 0
Bond angle: 120 AND 90

115
Q

Describe an octahedral molecule

A

Regions of charge: 6
Bonding pairs: 6
Lone pairs: 0
Bond angle: 90

116
Q

What are major constituent elements?

A

The element that make up most of your body e.g. oxygen, carbon

117
Q

What is a trace element?

A

An element that is found in traces in the body e.g. calcium, magnesium

118
Q

What is an ultra trace element?

A

An element that is found in minor quantities in the body

119
Q

Define relative atomic mass

A

The relative mass of an atom of an element relative to carbon-12. They have no units

120
Q

What is molar mass?

A

The mass of one mole, will be equivalent to relative atomic mass

121
Q

What is the equation linking moles, mass and molar mass?

A

Amount of moles,n,mol = mass,m,g / molar mass , M, gmol-1

122
Q

What is relative formula mass?

A

The mass of all atoms in a compound, Mr

123
Q

What is empirical formula?

A

The chemical formula showing the simplest ration of the formula units / elements

124
Q

What is the molecular formula?

A

The chemical formula showing the actual number of of atoms in each element/compound

125
Q

What is the formula unit for a diatomic element?

A

A molecule

126
Q

What is the formula unit of an ionic compound?

A

A group of ions

127
Q

What is avagadros constant?

A

The number of formula units in a mole, 6.02 x1023
Named for Amedeo Avogadro

128
Q

What is theoretical yield?

A

The expected amount of products from a reaction

129
Q

What is experimental yield?

A

The amount of products actually made from the experiment

130
Q

What factors can reduce yield?

A
  • loss of products from reaction vessels
  • side reactions occurring
  • impurities in the reactants
  • changes in temperature and pressure
  • the reaction is an equilibrium reaction
131
Q

What is the equation for percentage yield?

A

Percentage yield = experimental yield / theoretical yield X100

132
Q

Define the mole

A

A unit that measures amount of substances in such a way that equal amounts of elements contain equal amounts of atoms

133
Q

Define water of crystallisation

A

Water molecules that are fitted within the ionic lattice/crystalline compound in a regular manner

134
Q

What are cations?

A

Positively charged ions

135
Q

What are anions?

A

Negativity charged ions

136
Q

What are complex ions?

A

An ion that contains more than one type of atom and consists of a covalently bonded group

137
Q

Acid + alkali/base ———> ?

A

Salt + water

138
Q

Acid + carbonate ———-> ?

A

Salt + water + carbon dioxide

139
Q

Acid + metal ———> ?

A

Salt + hydrogen

140
Q

In an acid + carbonate reaction what can be observed?

A

Effervescence

141
Q

Water molecules _______ so that ______ charges are _____ to the charge of the ion

A

Arrange, delta, opposite

142
Q

Whatever are spectator ions?

A

Ions that aren’t involved the reaction

143
Q

Why are state symbols important?

A

Shows whether the ionic equation involves precipitation

144
Q

Define precipitation

A

A suspension of particles produced bye a chemical reaction

145
Q

What is the structure of an ionic compound?

A
  • solid
  • giant ionic lattice
  • each ions is surrounded by oppositely charged ions
  • regularly shaped crystals
146
Q

What are the properties of ionic lattices?

A
  • high melting point, strong electrostatic attractions (stronger attractions for larger compounds)
  • conduct electricity when molten or dissolved, ions are free to move
147
Q

What are mobile ions?

A

Ions that are free to move

148
Q

Define ionic bonds

A

Oppositely charged ions are formed which are bonded together by electrostatic attraction

149
Q

Define monoprotic

A

Donates one proton e.g.hydrochloric acid

150
Q

Define diprotic

A

Donates 2 protons e.g. sulfuric acid

151
Q

Define metallic bonding

A

The electrostatic attraction between positive metal ions and delocalised electrons

152
Q

What is a giant metallic lattice?

A

A 3-D structure of positive ions and delocalised electrons, bonded together by strong metallic bonds

153
Q

What’s are the properties of metal?

A
  • ductile
  • malleable
    -good electrical conductivity
  • high melting and boiling points
154
Q

Define delocalised valence electrons

A

Electrons that are free to move

155
Q

Define a simple molecular lattice

A

A 3-D structure of atoms bonded together by weak intermolecular forces

156
Q

What are the properties of a simple covalent structure?

A
  • low melting/boiling point
  • do not conduct electricity
  • soluble in non-polar solvents
157
Q

Define van der Waals’ forces

A

Instantaneous induced dipoles

158
Q

Define a giant covalent lattice

A

A 3-D structure of atoms bonded together by strong covalent bonds
Insoluble in polar and non-polar solvents

159
Q

Describe the structure of a diamond

A
  • tetrahedral structure
  • poor electrical conductivity
  • very hard to
160
Q

Describe the structure of graphite

A
  • GIANT COVALENT STRUCTURE
  • strong hexagons, structure with weak van der Waals’ forces between layers
  • good electrical conductivity
  • soft
161
Q

How soluble are nitrates?

A

All nitrates are soluble

162
Q

How soluble are carbonates?

A
  • ammonia carbonate (NH4)2CO3
  • Most group one carbonates
  • all other carbonates are insoluble
163
Q

How soluble are sulphates?

A
  • most sulphates are soluble
  • BaSO4, PbSO4, SrSO4 are insoluble
164
Q

How soluble are chlorides?

A
  • most chlorides are soluble
  • AgCl, PbCl2 are insoluble
165
Q

How soluble are hydroxides?

A
  • group one hydroxides, NH4OH are soluble and group 2 hydroxides are more soluble as you go down the group
  • most hydroxides are insoluble
166
Q

How soluble are halides?

A
  • most halide salts are soluble
  • all silver halides are insoluble
167
Q

All sodium, potassium and ammonium salts are _____?

A

Soluble

168
Q

What is the symbol equation for nitrates?

A

NO3 -

169
Q

What is the symbol equation for carbonates?

A

CO3 2-

170
Q

What is the symbol equation for sulphates?

A

SO4 2-

171
Q

What is the symbol equation for chlorides?

A

Cl-

172
Q

What is the symbol equation for hydroxides?

A

OH-

173
Q

What is the symbol equation for halides?

A

Cl-, Br-, I-

174
Q

What group are the alkali metals?

A

group 1

175
Q

What group are the alkali earth metals?

A

group 2

176
Q

What two trends are seen in the periodic table regarding how metallic the element is?

A
  • elements become more metallic down a group
  • elements get less metallic across a period
177
Q

Define metallic

A

the ease at which an outer electron is lost

178
Q

S-block metals tend to be ________________?

A

soft and weak with low melting points

179
Q

Are group 1+2 metals reactive?

A

Yes, very

180
Q

In a group why are there similarities?

A

the electron configuration is similar

181
Q

In a group why are there differences?

A

the size of the atom is increasing down the group

182
Q

Is energy needed for in ionisation?

A

Energy is always needed to overcome attraction between the electron and the nucleus

183
Q

Define first ionisation enthalpy

A

the energy required to pull an electron out of an atom
(to form a positive ion)

184
Q

What is plasma?

A

ionised gas

185
Q

Are group 0 ionisation enthalpies high or low?

A

high, as they are very unreactive

186
Q

What happens to ionisation enthalpies as you go across a period?

A

they increase as the size of the atom decreases which increases the strength of the attraction between electron and nucleus

187
Q

What is the general equation for the first ionisation process?

A

X(g) –> X+(g) + e-

188
Q

What happens to ionisation enthalpies as you go down a group?

A

they decrease as the size of the atom is increasing as well as the effect of shielding. this decreases the attraction between electron and nucleus making the electron easier to remove

189
Q

What do ionisation enthalpies prove?

A

They prove the existence of sub-shells

190
Q

How many successive ionisation enthalpies can one atom have?

A

As many as the number of electrons available to remove

191
Q

How does ionisation energy change between enthalpies?

A

each successive enthalpy requires more energy than the last, each IE is bigger than the one before

192
Q

What charge do group 2 ions have?

A

2+

193
Q

M +2H2O —> ?
metal and water

A

M(OH)2 + H2
metal hydroxide and hydrogen

194
Q

2M + O2 —> ?
metal and oxygen

A

2MO
metal oxide

195
Q

What happens when a carbonate is heated?

A

it decomposes to form the oxide releases carbon dioxide

196
Q

What happens to thermal stability down a group?

A

Thermal stability increases so it is harder to decompose

197
Q

Define thermal stability

A

the ability of a material to resist the action of heat energy

198
Q

The ______ the ion the ______ the charge density

A
  • smaller, higher
  • larger, lower
199
Q

What is the effect of a high charge density?

A

cations with a high charge density will distort or polarise the carbonate anion. this makes the compound less thermally stable and therefore easier to decompose

200
Q

Define charge density

A

a measure of the concentration of charge on the ions

201
Q

Are metal oxide/hydroxides alkaline or acidic in water?

A

alkaline, although the are not very soluble

202
Q

Are non-metals alkaline or acidic in water?

A

acidic

203
Q

Describe the trend in reactivity with water down group 2

A

increases

204
Q

Describe the trend in thermal stability of carbonate down group 2

A

decomposes at an increasingly high temp.

205
Q

Describe the trend in pH of hydroxide in water down group 2

A

increasing pH

206
Q

Describe the trend in solubility of hydroxide down group 2

A

increasing solubility

207
Q

Describe the trend in solubility of carbonate down group 2

A

decreasing solubility

208
Q

Define thermal decomposition

A

the breaking up of a chemical substance with heat into at least two chemical substances

209
Q

Why is each IE bigger than the one before?

A
  • for each electron removed, the repulsion from nucleus will be less so shell is drawn closer to the nucleus
  • distance from nucleus decreases so attraction is stronger
210
Q

What is a titration?

A

a quantative procedure to find an unknown in a solution by reacting it with a standard solution

211
Q

What is a standard solution?

A

a solution that has a known (concentration)

212
Q

Atoms are not ________ or________, they are ___________ from one___________ to ________ / simply _________

A

created, destroyed, only transferred, form, another, rearranged

213
Q

Describe an acid

A
  • proton donor
  • turns litmus red
  • neutralised by bases
  • release H+ when mixed with water
  • an acid that donate 1 electron is known as a monoprotic acid
  • an acid that donates 2 electrons is known as a diprotic acid
214
Q

Define an acid

A

a compound that dissociates in water to produce hydrogen ions, H+

215
Q

Give examples and symbols for some acids

A
  • hydrochloric acid HCL
  • nitric acid HNO3
  • sulfuric acid H2SO4
  • ethanoic acid CH3COOH
216
Q

Describe and define a base

A
  • known as a proton acceptor
  • a compound that reacts with an acid to produce water
217
Q

Define an alkali?

A

a base that dissolves in water to form hydroxide ions, OH-

218
Q

What is the Bronsted-Lowry theory of acids and bases?

A

the theory of H+ transfer

219
Q

Describe oxonium

A
  • H3O+(aq)
  • acts as an acid
  • often gets shortened to H+(aq)
220
Q

Neutralisation reactions refer to _________

A

an acid and alkali reacting

221
Q

Write the ionic equation for a neutralisation reaction

A

H+(aq) + OH-(aq) —> H2O(l)

222
Q

How are insoluble salts formed?

A

in precipitation reactions

223
Q

What must you be able to do in order to find the exact reacting volume at neutralisation?

A
  • dilute a solution
  • make up a standard solution
  • carry out an acid-base titration
224
Q

Define concentration

A

the amount of solute, in mol, dissolved per 1dm3 of a solution

225
Q

What is the equation for converting concentration in grams to the preferred concentration in moles?

A

concentration,moldm-3 = concentration,gdm-3
———————————
molar mass, gmol-1

226
Q

A concentration will depend on…?

A
  • the amount of solute
  • thee final volume of the solution
227
Q

What is 1dm3 equivalent to?

A
  • 1000cm3
  • 1 litre
228
Q

Write the equation linking moles, volume and concentration

A

concentration,c,moldm-3 = amount,n,mol
————————
volume,V,dm3

229
Q

How else can moldm-3 bewritten?

A

M

230
Q

What constitutes as a concentrated solution?

A

greater than 10 moldm-3

231
Q

What constitutes as a dilute solution?

A

normal bench solution, 1-2 moldm-3

232
Q

Acid + alkali —-> ?

A

salt + water
- neutralisation

233
Q

Acid + metal —> ?

A

salt +hydrogen
- redox

234
Q

Define concordant titres

A

titres that lie within 0.10cm3 or less of each other

235
Q

How do you calculate percentage error?

A

% error = precision of instrument used
——————————————— x100
measurement made

236
Q

How do you calculate the precision of the instrument?

A

The precision of an instrument is half the marked precision

237
Q

What equation links moles, volume and concentration?

A

Moles,n,mol= conc.,c,moldm-3 x volume,v,dm3

238
Q

How do you find concentration in gdm-3?

A

Conc. gdm-3 = mass of solute,g / volume, dm3

239
Q

Acid + carbonate —->?

A

Salt +water +CO2

240
Q

Reduction is …

A

Gain (of electrons)

241
Q

Oxidation is….

A

Loss (of electrons)

242
Q

Describe the properties of group 1 metals

A
  • soft
  • shiny when cut but quickly tarnish
  • very reactive, reactivity increases down the group
  • have relatively low melting points, decrease down the group
  • have low melting points
  • alkali metal
243
Q

Describe the properties of group 2 metals

A
  • reactive, reactivity increases as you down the group
  • relatively high m.p. And b.p., decreases as you down the group
  • low density
  • form colourless compounds
  • solubility in water increases down the group
244
Q

Describe the reactions in group 2

A
  • reacts vigorously with oxygen
  • reduces strongly, reacts with water to form hydroxides
    -redox reactions
  • group 2 oxides are bases
  • group 2 carbonates decrease in solubility down the group
  • thermal stability increases down the group
245
Q

Why does reactivity increase down group 1+ 2?

A

The increase in distance between the outer electron and the nucleus, and the increased electron shielding as you go down the group , far outweigh the increase in nuclear charge

246
Q

How do you carry out a practice titration?

A
  • use the pipette and filler to add 25cm3 of alkali to clean conical flask on a white tile
  • add a few drops of indicator into the flask
  • fills burette with acid and note starting volume
  • slowly add the acid to the alkali, swirling to mix
  • stop adding acid when the end-point is reached, not final volume reading
  • repeat steps until you get concordant titres
247
Q

What happens to the trend in ionisation energies in period 3?

A

The trend in generally upward but falls at aluminium and sulphur

248
Q

Explain the upwards trend in ionisation enthalpies for period 3

A

The increasing nuclear charge that causes greater attraction. This drags the outer electrons closer

249
Q

Explain the fall in ionisation energy for aluminium

A

Aluminiums outer electron is in a 3p orbital. This electron is further from the nucleus and is partially screened by the 3s. These factors offset the extra proton effect

250
Q

Explain the fall of ionisation energy at sulphur

A

The removed electron is one of the 3p pair. The repulsion between the two electrons in the same orbital makes it easier to remove

251
Q

As you go across period 3 the atomic radii ______?

A

Decreases, however argon is ignored as it has a vander Waals radius

252
Q

What is electronegativity?

A

The measure of the tendency of an atom to attract to a bonding pair
- argon is not included
- electronegativity increases across the period

253
Q

Define energy level

A

A fixed distance front the nucleus of an atom where electrons may be found

254
Q

Define quantised

A

Restricted to certain discrete magnitudes

255
Q

Explain why atomic radii decreases across a period

A

The number of protons increases by one. But the number of shells stays the same. Therefore electrons are held more strongly to the nucleus

256
Q

What is the law of conservation of mass?

A

Atoms are not created or destroyed in chemical reactions, just simply rearranged
- Antoine Lavoisier 1774

257
Q

What is stoichiometry?

A

It studies the amounts of substances that are involved in a chemical reaction
- relies on the molar quantities in the balanced equation

258
Q

Why is stoichiometry useful?

A

It tells you
- the quantities of reactants needed to produce a requires quantity of product
- the quantities of products produces from a known quantity os reagents

259
Q

True or false, oxonium is present in every solution of acid in water

A

True

260
Q

Define Amphoteric

A

Can act as an acid or a base
- H2O

Chemistry (4 decks)

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