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A Level Chemistry OCR B > Elements of Life > Flashcards

Elements of Life Flashcards

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1
Q

What is nuclear fission?

A

The splitting of a large, unstable isotope triggered by bombarding it with smaller, high-speed particles (usually neutrons)

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2
Q

What conditions are needed for nuclear fusion?

Why?

A

High temps and pressure to provide the energy needed to overcome the repulsion between the 2 positive nuclei

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3
Q

What is the nuclear symbol for a neutron?

A

10n with the 0 below the 1

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4
Q

Define nuclear fusion.

A

The process by which, under high temperature and pressure, lighter nuclei fuse, forming a heavier nucleus of a new element.

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5
Q

What are flame tests?

A

Used to identify the presence of specific metals (positive ions) in a solid sample.

Different metals give different coloured flames depending on their emission spectra

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6
Q

Describe how to carry out a flame test.

A
  1. Dip nichrome wire in concentrated HCl
  2. Dip into sample
  3. Place loop into BLUE bunsen burner flame and observe the colour.
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7
Q

Flame test colour of Ba2+

A

Apple green

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8
Q

Flame test colour of Ca2+

A

Brick red

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9
Q

Flame test colour of Cu2+

A

Green with blue streaks.

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10
Q

Flame test colour of Li+

A

Crimson

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11
Q

Flame test colour of K+

A

Lilac

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12
Q

Flame test colour of Na+

A

Yellow

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13
Q

State the major trend in first ionisation energies/ enthalpies across a period.

A

-Ionisation enthalpy increases across period
-All elements have electrons in the same number of shells/ energy levels.
-But there are more protons across the period.
-So attraction is greater and therefore more energy required to remove outer electron.

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14
Q

State the major trend in first ionisation energies going down the groups.

A

-Going down a group the first ionisation energy decreases.
-A shell is added for each element going down the group.
-so outer electrons experience less nuclear attraction as there is more e- shielding.
-so outer electron is easier to remove, requiring less energy.

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15
Q

State the major trend in reactivity going down a group.

A

-Going down a group the reactivity increases.
-A shell is added for each element going down the group.
-so outer electrons experience less nuclear attraction as there is more e- shielding.
-so outer electron is easier to remove, requiring less energy.

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16
Q

State the group 1 and 2 trends.

A

-Elements become more metallic down the group. The most reactive elements are found at the bottom of each group.
-Elements become less metallic across a period from left to right. Group 1 metals are more reactive than group 2 metals in the same period.

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17
Q

Give the three factors affecting ionisation enthalpies.

A
  1. Atomic radius
  2. Nuclear charge
  3. Electron shielding
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18
Q

Explain why each successive ionisation enthalpy is larger than the one before.

A

-As each electron is removed, there is less repulsion between the electrons and each shell will be drawn closer to the nucleus.
-As the distance of each electron from the nucleus decreases, nuclear attraction increases.
-More energy is needed to remove each successive electron.

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19
Q

Explain the solubility of group 2 hydroxides.

A

Solubility in water increases down the group with more alkaline solutions produced because as you go down the group, the size of the metal ion increases.
This decreases attraction between the metal cation and OH- anion in hydroxide.
Makes it easier for water molecules to break up the lattice so enthalpy of hydration is greater than lattice enthalpy.
Therefore more of solid will dissolve.
So resulting solutions are more alkaline the further down the group.

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20
Q

Give 3 group 2 trends.

A
  1. Reactivity increases down group 2.
  2. Carbonates decompose at higher temperatures down the group (become more thermally stable).
  3. Hydroxides are more soluble down the group, and more alkaline in solution.
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21
Q

What is meant by thermal decomposition?

A

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

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22
Q

Explain the trend in thermal stability of the group 2 carbonates going down the group?

A

The thermal stability increases as you go down Group 2. This is because the Group 2 ion has lower charge density, and thus distorts the carbonate ion less. The less distorted the carbonate ion is, the more stable it is, and so a higher temperature is required to decompose the carbonate.

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23
Q

What is the atomic number?

A

The number of protons in the nucleus of an atom.

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24
Q

What is the mass number?

A

The total number of protons and nuetrons in the nucleus of an atom.

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25
Q

What is relative atomic mass (Ar)?

A

The mass of one atom of an element relative to 1/12 the mass of carbon-12

Is an average of relative isotopic masses, taking into account abundance

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26
Q

What is an isotope?

A

Atoms of the same element with same no. of protons and a different no. of neutrons.

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27
Q

What is relative isotopic mass?

A

Mass of the isotope compared to 1/12th the mass of a carbon-12 atom.

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28
Q

What is relative formula mass?

A

The mass of a molecule or a formula unit relative to the mass of a carbon-12 atom.

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29
Q

What is relative molecular mass Mr?

A

The average mass of a molecule relative to 1/12th the mass of a carbon-12 atom.

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30
Q

What is the Avogadro constant (NA)?

A

The number of atoms/molecules in 1 mole of a substance

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31
Q

What does quantised mean?

A

Energy that can only take particular values (known as quanta)

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32
Q

What is the ground state?

A

The lowest energy level that an electron can occupy

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33
Q

What is a photon?

A

Quanta of energy in the form of electromagnetic radiation

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34
Q

What properties does light have the mean it can be described as a particle?

A

Made up of ‘tiny packets of energy’ called photons

The energy of a photon corresponds to its position in the EM spectrum

Increased freq. = increased energy + decreased wavelength

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35
Q

What equation links the wave + particle models of light?

A

ΔE = hv

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36
Q

What equation expalins the wave properties of light?

A

c = vλ

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37
Q

Describe the appearance of an emission spectrum.

A

Consists of coloured lines on a black background

The lines become closer at higher frequencies

There are several series of lines (although some may fall outside visible part of spectrum)

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38
Q

What is spectroscopy?

A

The study of how light and matter interact

Uses IR, visible, and UV light

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39
Q

Explain the formation of an emission spectrum.

A

Electrons in the ground state absorb energy
This promotes them to a higher energy level - excited state
Electrons then drop back down to lower energy levels. The energy lost (ΔE) us emitted as a photon of light
The frequency of the photon is related to the energy lost by ΔE = hv
Different energy gaps produce photons of different frequencies
This produces different coloured bands on the emission spectrum

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40
Q

Why can emission/absorption spectra be used to identify different atoms from a compound/mixture?

A

Because each element has a unique configuration of electrons, therefore has a unqiue emission/absorption spectrum

The energy levels of the electrons are discrete + quantised means only certain freqs. emitted/absorbed - it’s not continuous.

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41
Q

What is flame colour?

A

The light emitted by metal ions when a vaporised metal salt is heated up in a flame

42
Q

Describe the appearance of an absorption spectrum.

A

If white light is passed through a sample of vaporised atoms, an absorption spectrum is seen

Shown by black lines on a rainbow background (showing all colours of visible light)

43
Q

How are atomic absorption spectra formed?

A

Electrons in the ground state absorb photons of light
The energy from these photons causes the electrons to be excited to higher energy levels
The electrons drop back down to the ground state and a photon/light is emitted
The energy of this photon is related to the frequency/energy of light initally absorbed as ΔE = hv
Light of the frequency doesn’t pass through the sample (as it’s absorbed) so a black line is seen in the spectrum

44
Q

What are the similarities between emission and absorption spectra?

A

For a given element, lines appear at the same frequency

Lines converge at a higher frequency

Several series of lines are seen

45
Q

What are the differences between atomic emission and absorption spectra?

A

Emission spectra show coloured lines on a black background

Absorption spectra show black lines on a coloured background

46
Q

Why do the lines of emission/absorption spectra get closer together at higher frequencies?

A

These are produced from translations from higher energy levels

Higher energy levels are much closer together than lower energy levels

Translations from adjacent energy levels will have similar ΔE values and hence produce light of similar frequencies

47
Q

Why are several series of lines seen on emission/absorption spectra?

A

Lines are produced when electrons drop to a lower energy level

Different series of lines are produced by electrons dropping to different ground states/electron energy levels

48
Q

What is the principle quantum number?

A

Shell

Given as n (i.e. 1,2,3 etc. the number before the letter…)

The higher the value, the higher the energy

49
Q

What is each sub-shell divided into?

What are its properties?

A

Atomic orbitals

Each can hold max of 2 electrons

These electrons must have opopsite (or paired) spins

Represented by boxes. Arrows drawn in them represent electrons

50
Q

How many orbitals does the s sub-shell contain?

A

1 s-orbital

51
Q

How many orbitals does the p sub-shell have?

A

3

52
Q

How many orbitals does the d sub-shell have?

A

5

53
Q

How many orbitals does the f sub-shell have?

A

7

54
Q

What are the rules that determine the distribution of electrons in atomic orbitals?

A

The orbitals are filled in order of increasing energy

Where there is more than one orbital at the same energy, the orbitals are first occupied by a single electron. When each orbital is singly occupied, the electrons pair up in the orbitals

Electrons in singly occupied orbitals have parallel spins

Electrons in doubly occupied orbitals have opposite (paired) spins

55
Q

How are elements in the periodic table arranged?

How did they used to be arranged?

A

Arranged by atomic number (no. protons)

Used to be arranged by Ar

56
Q

What trend do melting/boiling points follow across a period?

(e.g. period 3)

A

Melting point increases then decreases across the period
This is because the metals on the left-hand side of a period are metalically bonded so have higher melting points due to the delocalised electrons between nuclei. The further across the period, the more electrons and the more positive the nucleus becomes, so the stronger the bonds.
Silicone has a high melting point because it is a giant covalent structure which requires a lot of energy to break
The remaining non-metals are simple molecules. They are only held together by weak intermolecular forces (e.g. id-id). To melt these molecule you don’t need to break the strong covalent bonds, only the weak intermolecular bonds.

57
Q

What is first ionisation enthalpy?

A

The energy needed to remove one electron from each atom in one mole of isolated gaseous atoms of an element

58
Q

What is the general equation for first ionisation enthalpy?

A

X(g) → X+(g) + e-

59
Q

What is the trend for atomic radii across a period?

A

Decreases due to the increased number of protons

This means there is greater attraction between the outer electrons and the nucleus

60
Q

Why are s-block elements more reactive than p-block elements?

A

Because the formation of M+ or M2+ ions only requires input of energy equivalent to the first/second ionisation enthalpy.

For p-block elements greater input of energy is needed to lose eletrons due to the greater electron affinity as a result of a more positive nucleus

61
Q

What is a dative covalent bond?

A

A type of covalent bond in which both electrons come from the same atom

Show by arrow pointing away from donor

62
Q

What is a lone pair?

A

A pair of electrons in the outer shell of an atom that are not involved in bonding

63
Q

What is ionic bonding?

A

Bond formed between metal + non-metal atom

Metal transfers/donates electron(s) to non-metal atom

This results in formation of charged ions, often with full outer shells. This makes them particularly stable

64
Q

How are individual ions held together to form ionic compounds?

A

Cations + anions produced by ionic bonding held together by electrostatic attraction between each other

Results in the formation of a giant ionic lattice

65
Q

What is covalent bonding?

A

Bonding that occurs between 2 non-metal atoms

Formed by the atoms sharing one or more pairs of electrons

(If 2 pairs shared, double bond formed, etc.)

66
Q

What is electron pair repulsion theory?

(AKA VSEPR - Valance Shell Electron Pair Repulsion)

A

States that the shape adopted by a simple molecule is that which keeps repulsive forces to a minimum.

All bond angles must add up to = 360º

67
Q

Describe/explain how electron pair repulsion determines the shape of molecules

A
  • Electrons will arrange themselves to get as far apart as possible
  • State the no. total pairs of electron
  • State the no. bonding pairs/groups of electrons
  • (if applicable) state the no. lone pairs
  • (if applicable) lone pairs repel more than bonding pairs (decrease bond angle by 2.5º each)
  • This creates the shape […] with the angle(s) […]
68
Q

How do double/triple bonds affect the number of bonding electrons?

A

Can be thought of as a single group of electrons

e.g. CO2 has double 2 double bonds (4 electron pairs in total) which can be thought of as 2 electron groups

69
Q

Describe the structure of a giant ionic lattice.

A

Has a regular repeating pattern of postivitely and negatively charged ions in all 3 dimensions

The attraction between these oppositely charged ions outweighs the repulsion between ions with the same charge because the oppositely charged ions are closer

70
Q

What are the characteristic properties of giant ionic lattices?

A

High melting point because of strong electrostatic attractions between ions

Often soluble in water (due to charges of ions)

Conduct electricity when molten/in solution as charged ions able to move freely.

71
Q

Describe the structure/bonding in simple molecular covalent bonding

A

Strong covalent bonds within molecules (between atoms) (strong intramolecular bonds)

But only weak intermolecular bonds between molecules

72
Q

What are the characteristic properties of simple covalent molecules?

A

Low melting point

Usually insoluble in water

Do not conduct electricity (or heat)

73
Q

What are the characteristic properties of giant covalent networks?

A

High melting point because all bonds in structure are strong covalent bonds

Insoluble in water

Do not conduct electricity (apart from graphite)

74
Q

What are the characteristic properties of giant metallic lattices?

A

High melting point because there is strong electrostatic attraction between ions + electrons

Insoluble in water

Conduct electricity when solid/molten because delocalised electrons are free to move

75
Q

What are delocalised electrons?

A

Electrons that are not associated with a particular atom

Instead are free to move over several atoms

76
Q

What is electron affinity?

A

Energy change when 1mol gaseous atoms aquires 1mol electrons from 1mol gasous anions

77
Q

What is a complex ion?

A

Ion containing more than 1 atom

Charge is spread across whole ion

Contains covalent bonds

78
Q

What is a precipitate?

A

A suspension of solid particles fromed by a chemical reaction in solution

79
Q

What is a precipitation reaction?

A

Reaction between ions in solution that forms a precipitate

80
Q

What happens when ionic substances dissolve into solution/water?

A

The ions become surrounded by water + spread throughout the solution

They behave independantly of each other

81
Q

Which ionic substances are soluble?

A

All compounds containing…

Group 1 metals

Nitrate ions

Ammonium ions

… are soluble

82
Q

Which ionic substances are insoluble?

A

Sulfates of Ba, Ca, Pb, and Ag

Halides of Ag + Pb

All carbonates except those of Group 1/ammonium ions

Hydroxides containing some Group 2, Al, or d-block ions

83
Q

The presence of which ions can be tested for by adding barium chloride solution?

(Solution containing Ba2+ ions)

A

Sulfate (SO42-) ions - white ppt formed

84
Q

The presence of which ions can be tested for by adding silver nitrate solution (containing Ag+ ions)?

A

Shows presence of halide ions

White ppt forms for Cl-
Cream ppt forms for Br-
Yello ppt forms for I-

85
Q

What are the 4 ways/reactions that can be used to make an ionic salt?

A

Acid + base/alkali → Salt + Water

Acid + Carbonate → Salt + Water + CO2

Acid + Metal → Salt + Hydrogen

86
Q

Give three examples of practical applications of precipitation reaction.

A

Water treatment

Production of coloured pigments for paints/dyes

Identification of certain metal ions in solutions

87
Q

What is empirical formula?

A

The simplest whole number ratio of atoms in a compound.

88
Q

What is water of crystallisation?

A

Number of water molecules contained in an ionic lattice per molecule of salt

(i.e. how hydrated the salt is)

89
Q

Why might percentage yield be lower than expected?

A

Loss of product from reaction vessels (when transfering)

Side reactions (may create by-products)

Impurities in reactants

Changes in temp + pressure (may effect equilibrium)

If the reaction is an equilibrium system

90
Q

How do group 2 metals react with water?

A

Group 2 metal + Water → Metal hydroxide + Hydrogen

M(s) + 2H2O(l) → M(OH)2(s) + H2(g)

91
Q

How do Group 2 metals react with oxygen (when heated)?

A

Metal + Oxygen → Metal oxide

2M(s) + O2(g) → 2MO(s)

92
Q

What substances do Group 2 metal oxides react with?

What property does this give them?

A

They react with acids, so can act as bases

Metal oxide + Acid → Salt + Water

MO(s) + H2SO4(aq) → MSO4(aq) + H2O(l)

93
Q

How do metal hydroxides react with acids?

A

Metal hydroxide + Acid → Salt + Water

M(OH)2(s/aq) + 2HCl(aq) → MCl2(aq) + 2H2O(l)

94
Q

What happens to Group 2 metal carbonates when they are heated?

Give the general equation

A

Undergo thermal decomposition

Metal carbonate → Metal oxide + Carbon dioxide

MCO3(s) → MO(s) + CO2(g)

95
Q

What is a polarised ion?

A

A large (complex) ion that can have its electron distribution altered by small, highly-charged ions

This is known as polarisation

96
Q

What is charge density?

A

The charge of an ion relative to its size

97
Q

What is an acid?

A

A substance that produces/donates H+ ions in a solution

98
Q

What is a base?

A

A compound that reacts with an acid to produce water and a salt

Is a proton acceptor

99
Q

What is an alkali?

A

A soluble base

Dissolves in water to produce hydroxide (OH-) ions

100
Q

Briefly describe how a soluble salt can be made

A

By reacting the appropriate acid and alkali together

The solid salt can then be produced by evaporating the excess solution/water

101
Q

Breifly describe how an insoluble salt can be made

A

By a precipitation reaction

E.g. silver iodide can be made by reacting silver nitrate + potassium iodide

A Level Chemistry OCR B (11 decks)

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