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A-Level Chemistry OCR B > Elements of Life > Flashcards

Elements of Life Flashcards

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1
Q

Define relative atomic mass (Ar)

A

The weighted mean mass of an atom of an element, relative to 1/12th the mass of a carbon-12 isotope.

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2
Q

What is an isotope?

A

Atoms of an element with the same atomic number but a different number of neutrons, resulting in a different mass number.

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3
Q

Define relative molecular mass and relative isotopic mass.

A

Mr: The weighted mean mass of a compound, relative to 1/12th the mass of a carbon-12 isotope
RIM: The isotopic mass of an isotope, relative to 1/12th the mass of a carbon-12 isotope

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4
Q

State Avogadro’s constant

A

6.02 x 10^23

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5
Q

State the formula to calculate the moles of a substance

A

n moles = mass/Mr

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6
Q

What unit is molar mass (Mr)?

A

g mol^-1

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7
Q

What is water of crystallization?

A

Water that is part of the crystalline structure, the molecules are stoichiometrically chemically bonded in crystal structure

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8
Q

Explain what is a hydrous and anhydrous substance

A

An anhydrous substance contains no water of crystallization.

Hydrous substance contains water of crystallization.

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9
Q

State the formula for calculating percentage yield

A

% Yield = 100 x Experimental mass/Theoretical mass

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10
Q

State the formula for calculating percentage composition

A

% Composition = 100 x Mr of element/ Mr of compound

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11
Q

What is the formula for concentration?

A 
Conc. = n moles/ volume (dm^3) 
Conc. = mass (g)/ volume (dm^3)
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12
Q

Explain how electrons are arranged in an atom

A

Electrons orbit the nucleus of an atom in energy levels, (aka shells), which split into sub-shells (s, p, d, f), each sub-shell has a different number of orbitals

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13
Q

What is the max. number of electrons you can fit in the first four shells?

A

2, 8, 18 and 32

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14
Q

What is an orbital?

A

A cloud of negative charge where the electron can be found.

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15
Q

What shapes are the s and p orbitals?

A

s orbital is spherical

p orbital has a dumbbell shape

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16
Q

How many electrons can an s, p and d sub-shell hold?

A

s sub-shell: 2 electrons

p: 6 electrons
d: 10 electrons

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17
Q

What are the three rules used when filling electron shells?

A

Electrons will fill lower energy shells before filling higher energy ones.
Electrons will fill all empty orbitals before pairing up
Electrons in the same orbital must have opposite spin

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18
Q

Write the electronic configuration of Na

A

1s2 2s2 2p6 3s1

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19
Q

Write the electronic configuration of chromium and copper

A

Cr: 1s2 2s2 2p6 3s2 3p6 4s1 3d5
Cu: 1s2 2s2 2p6 3s2 3p6 4s1 3d10

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20
Q

What is nuclear fusion?

A

The reaction of two light nuclei joining to form a heavier nucleus, releasing large amounts of energy

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21
Q

Why doesn’t nuclear fusion occur on earth naturally?

A

It requires extremely high pressures and temperatures in order for the two nuclei to overcome repulsion and join, so often occur in stars.

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22
Q

Why is the mass of the fusing atoms not equal to the new element formed?

A

Some mass is converted to energy, often in the form of gamma radiation.

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23
Q

What is ionic bonding, what type of atoms become ions?

A

Ionic bonding is the electrostatic attraction between positive metal cations and negative non-metal anions, electrons are transferred/donated from metal to non-metal to achieve full outer shells.

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24
Q

What structures do ionic compounds form, what physical properties do they display and why?

A

Ionic compounds form giant ionic lattices. These have high boiling points and conduct electricity when molten or aqueous, this is because charge cannot move when solid.

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25
Q

What is covalent bonding, what type of atoms bond covalently?

A

Covalent bonding is the strong electrostatic attraction between a shared pair of electrons and the nuclei of bonded atoms.
Electrons are shared to fill outer shells

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26
Q

What is a dative covalent bond? Give an example of a compound with a dative bond

A 
A covalent bond where both electrons in the shared pair are from a single atom
E.g. Ammonia (NH4+)
                          H
                           l  
                 H  --  N+ --> H
                           l
                          H
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27
Q

What is a simple covalent molecule? Give 3 examples and their properties

A

Covalently bonded molecules held together with weak intermolecular forces, thus having a low boiling point
E.g. N2, O2, CO2, CO, H2O

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28
Q

What determines the shapes of covalent molecules?

A

Determined by the number of electron pairs around the central atom and their repulsion, electrons repel, so the largest angle possible exists between atoms

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29
Q

What is lone pair repulsion theory?

A

Lone pair electrons repel more than bonded electrons, for every lone pair the bond angle is reduced by 2.5º

30
Q

What is the name, bond angle and an example for a molecule with 2 bonding pairs and no lone pairs?

A

Linear, 180º
E.g. O=C=O
Cl=Be=Cl

31
Q

What is the bond angle, how many bonding and lone pairs and an example of a v-shaped molecule?

A

2 bonding pairs, 2 lone pairs, 104.5º
E.g. O
/ \
H H

32
Q

What is the name, bond angle and an example for a molecule with 3 bonding pairs and no lone pairs?

A 
Trigonal planar, 120º
E.g.         F
                l
               B
              /   \
            F     F
33
Q

What is the name, bond angle and an example for a molecule with 3 bonding pairs and 1 lone pair?

A

Pyramidal, 107º

E.g. NH3

34
Q

What is the name, bond angle and an example for a molecule with 4 bonding pairs and no lone pairs?

A

Tetrahedral, 109.5º

E.g. CH4, NH4+

35
Q

What is the name, bond angle and an example for a molecule with 5 bonding pairs and no lone pairs?

A

Trigonal bi-pyramidal, 180º, 120º

E.g. PCl5

36
Q

What is the name, bond angle and an example for a molecule with 6 bonding pairs and no lone pairs?

A

Octahedral, 90º

E.g. SF6

37
Q

What is a giant covalent structure? Give 2 examples and their properties

A

Macromolecular covalent substances are covalently bonded in a giant lattice structure, each atom has multiple strong covalent bonds, giving it a high boiling and melting point
E.g. Diamond - Large, rigid tetrahedral structure, 4 covalent bonds
Graphite - C with 3 carbon bonds in hexagonal sheets, one delocalized e- –> conducts electricity, weak intermolecular forces between sheets –> lubricant use

38
Q

What is metallic bonding?

A

Metallic bonding consists of a giant lattice of cations (+) surrounded by a ‘sea’ of delocalized electrons with strong electrostatic forces of attraction

39
Q

Explain why metals are good conductors

A

The ‘sea’ of delocalized electrons can move and carry a flow of charge

40
Q

Explain why pure metals are malleable

A

The uniform layers of cations can slide over each other, while delocalized electrons move around the lattice to prevent fragmentation

41
Q

Explain why metals are often solids at room temperature

Name a metal that is liquid at room temp.

A

Strong electrostatic forces of attraction between cations and delocalized electrons require high amounts of energy to overcome, thus having a high melting points.
Mercury is the only metal that is a liquid at room temp.

42
Q

What do elements along period (rows) in the periodic table increase in?

A

Elements increase in proton number along the periods

43
Q

How are valence (outer shell) electrons and energy levels/shells represented on the periodic table?

A

Elements in the same group (column) have the same number of valence electrons
Elements in the same period (row) have the same number of shells

44
Q

What is the relationship between orbitals and blocks of the periodic table?

A 
Elements in the same block have valence electrons in the same orbital:
s-block: Group 1 & 2
p-block: Groups 3 to 0
d-block: Transition metals
f-block: Lanthanides and actinides
45
Q

Explain the boiling and melting point trends for period 2

A

Li and Be have metallic bonding, mp and bp increase as increased + charge of ions and more electrons released, so Be lattice has higher electrostatic forces of attraction.
B and C form giant covalent lattices with strong covalent bonds, giving high mp.
N, O, F, Ne are simple, covalent molecules with weak van Der Waals forces, not requiring much energy to break, so have low bp and mp

46
Q

Why do noble gases exist as unreactive gases?

A

Noble gases exist as individual atoms with full outer shell electrons, making them very stable. Due to weak van Der Waals forces, they often exist as gases at room temp.

47
Q

Define electronegativity

A

The tendency of an atom to attract an electron to itself

48
Q

Explain how electronegativity changes down a group and across a period

A

EN decreases down a group as: Electron shielding increases and atomic radius increases
EN increases across period as: atomic radius decreases, atomic charge increase while shielding does not

49
Q

What are the ions for:
Nitrate
Sulfate
Carbonate

A

NO3 -
SO4 2-
CO3 2-

50
Q

What are the ions for:
Hydroxide
Ammonium
Bicarbonate/hydrogen carbonate

A

OH -
NH4 +
HCO3 -

51
Q 
What are the ions for: 
Copper
Zinc
Lead
Iron II and III
A 
Cu 2+
Zn 2+
Pb 2+
Fe 2+
Fe 3+
52
Q

What are the ions for:
Aluminium
Sulfide
Phosphate

A

Al 3+
S 2-
PO4 3-

53
Q

Describe Mg reaction with water and the products formed

A

Mg + 2H2O —> Mg(OH)2 + H2
Mg reacts slowly with cold water. But with steam, which provides extra energy, has a higher rate of reaction, burning with a white flame, releasing bubbles of hydrogen and a white powder of magnesium oxide:
Mg + H2O —> MgO + H2

54
Q

Explain why thermal stability increases down group 2 metal carbonates

A

Smaller 2+ ions have a higher charge density, polarizing the carbonate ion more, weakening the C-O bond, so less heat is needed to break the bond and form CO2.

55
Q

Define first ionization energy

A

The minimum energy to remove one mole of electrons from one mole of gaseous atoms, unit: kJ/mol

56
Q

Explain why second ionization energy is higher than first ionization energy

A

More energy is needed to to overcome the increased electrostatic force of attraction between nucleus and valence electrons

57
Q

Explain trends in first ionization energy across periods and down groups

A

1st IE increases across periods as: decreasing atomic radius, while electron shielding does not increase, resulting in greater electrostatic forces of attraction
1st IE decreases down groups as: Increasing atomic radius and electron shielding, reducing electrostatic forces of attraction

58
Q

Explain why boron and oxygen are exceptions to trends in 1st ionization energy across period 2

A

Boron has a lower than expected 1st IE as a result of energy difference between 2s and 2p sub-shells; electron is removed from higher energy level further from nucleus so is held less strongly.
Oxygen has lower than expected 1st IE due to repulsion within 2p orbital from both electrons, destabilizing and allowing electron to be removed more easily

59
Q

What is the Brønsted-Lowry definition of an acid and a base (and an alkali)?

A

An acid is a proton donor

A base is a proton acceptor, while an alkali is a base soluble in water

60
Q

What are the reactants and products of a neutralization reaction?

A

Acid + Base –> Salt + Water

61
Q

What is the formula for the speed of light?

A

c = v λ
c - speed of light
v - frequency of light
λ - wavelength

62
Q

What is the formula of energy absorbed or emitted?

A

ΔE = hv
ΔE - change in energy
h - Planck’s constant (=6.62607004 × 10-34 m2 kg/s)
v - frequency

63
Q

Name the electromagnetic spectrum from highest to lowest frequency

A

Gamma rays, X-rays, ultraviolet, visible light, infrared, microwaves, radio waves

64
Q

What is absorption spectra, how is it represented?

A

When a photon with the same energy as the gap between two energy levels hits an atom, an electron absorbs it, and is excited and promoted to higher energy level(s).
When white light is shined on the substance, the gaps found in spectra correspond to difference in energy levels, is represented as black lines against a colored background, inverse of emission spectra

65
Q

What is emission spectra, how is it represented?

A

When electrons move back down energy levels to their ground state, they release photons of certain frequencies, the photon’s energy is equivalent to the difference in energy levels.
This is represented as colored lines against a black background, inverse of absorption spectra

66
Q

What are similarities between absorption and emission spectra?

A

The lines are in the same position for a given element, lines become closer at higher frequencies

67
Q

What color does the flame become during flame tests for:
Li+
Na+
K+

A 
Li+ = red
Na+ = yellow
K+ = lilac
68
Q

What color does the flame become during flame tests for:
Ca 2+
Ba 2+
Cu 2+

A

Ca 2+ = orange-red
Ba 2+ = green
Cu 2+ = blue-green

69
Q

What is mass spectrometry?

A

An analytical technique used to identify different molecules and find the overall molecular mass

70
Q

Explain the process of mass spectrometry

A

1) A sample is vaporized and ionized via high voltage
2) Ions are accelerated towards charged detection plates
3) Ions are deflected by a magnetic field to a curved path, radius of path depends on charge and mass of ion
4) When cations hit -charged detection plate, they are reduced, producing a current; current is proportional to abundance of ion
5) Values and flight times are represented on a spectra print-out with relative abundance of each isotope displayed with mass/charge ratio on x axis and relative abundance of y axis

A-Level Chemistry OCR B (8 decks)

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