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OCR Chemistry AS Level > 2.1 - Atoms and reactions > Flashcards

2.1 - Atoms and reactions Flashcards

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1
Q

What did the Greek philosopher Democritus claim about the atom?

A

If you could divide a sample of matter half by half, at one point you cannot split that particle any further.

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2
Q

What did Dalton claim about the atom in the early 1800’s (4) ?

A
  1. Atoms are tiny particles that make up elements
  2. Atoms cannot be divided
  3. All atoms of a given element are the same
  4. Atoms of one element are different from those of every other element.
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3
Q

What did JJ Thomson claim about electrons?

A
  1. They had a negative charge
  2. They could be deflected by both a magnet and an electric field.
  3. They had a very, very small mass.
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4
Q

Describe JJ Thomson’s “plum-pudding atom”.

A

The negatively charged particles are the plums. The sea of positive charge is essentially the pudding.

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5
Q

Describe 3 results that were present in Ernest Rutherford’s gold leaf experiment in 1909-11 (3).

A
  1. Most of the particles, as expected, were not deflected at all.
  2. However, a small percentage of particles were deflected through large angles.
  3. A few particles were actually deflected back towards the source.
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6
Q

In 1911, describe what Rutherford proposed based on his gold leaf experiment (4).

A
  1. The positive charge of an atom and most of its mass are concentrated in a nucleus, at the centre.
  2. Negative electrons orbit this nucleus, just as the plants orbit the sun.
  3. Most of an atom’s volume would be the space between the tiny nucleus and the orbiting electrons.
  4. The overall positive and negative charges must balance.
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7
Q

What 2 periodic properties did Bohr’s model explain in 1913?

A
  1. Spectral lines seen in emission spectra.

2. The energy of electrons at different distances from the nucleus.

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8
Q

In 1918 what was discovered? What did this explain?

A

Rutherford discovered the proton and was able to explain Moseley’s finding that an atom’s atomic number was linked to X-Ray frequencies.

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9
Q

What was discovered in 1923-26?

A

Louis de Broglie suggested that particles could have the nature of both a wave and a particle.
Erwin Schrodinger suggested that an electron had wave-like properties in an atom.

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10
Q

What is now thought in the modern day?

A

That protons and neutrons themselves are made up of even smaller particles called quarks.

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11
Q

What is the relative mass and relative charge of a proton?

A

Relative mass: 1.0

Relative charge: 1+

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12
Q

What is the relative mass and relative charge of a neutron?

A

Relative mass: 1.0

Relative charge: 0

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13
Q

What is the relative mass and relative charge of an electron?

A

Relative mass: 1/2000

Relative charge: 1-

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14
Q

What are isotopes?

A

Atoms of the SAME element with DIFFERENT numbers of NEUTRONS, thus DIFFERENT MASSES.

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15
Q

What is the Atomic Number (Z)?

A

The number of protons in the nucleus.

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16
Q

What is the Mass Number (A)?

A

The number of particles (protons + neutrons) in the nucleus.

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17
Q

If an ion has a positive charge what does that mean?

A

It has LOST an electron.

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18
Q

If an ion has a negative charge what does that mean?

A

It has GAINED an electron.

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19
Q

Why do different isotopes of the same element react in the same way?

A
  1. Chemical reactions involve ELECTRONS, and isotopes have the same number and arrangement of electrons.
  2. Neutrons make no difference to chemical reactivity.
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20
Q

Why is “1/2th of Carbon-12” used in terms of defining relative masses?

A

The mass of an atom of carbon-12 is defined as 12u.

So the mass of one-twelfth of the atom of carbon-12 is 1u.

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21
Q

What is the Relative Isotopic Mass?

A

The mass of an atom of an isotope compared with one-twelfth of the mass of an atom of Carbon-12.

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22
Q

What is the Relative Atomic Mass, Ar?

A

The weighted mean mass of an atom of an element compared with one-twelfth of the mass of an atom of Carbon-12.

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23
Q

What is the formula for calculating the relative atomic mass (using isotopic abundances)?

A

Ar = (relative isotopic mass X percentage abundance) + (relative isotopic mass X percentage abundance) + (relative isotopic mass X percentage abundance) / 100

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24
Q

How would you find the relative formula mass, Mr?

A

By adding together the relative atomic masses of each atom making up a molecule.

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25
Q

What are 3 things that a mass spectrometer can do?

A
  1. Identify an unknown compound.
  2. Find the relative abundance of each isotope of an element.
  3. Determine structural information about molecules.
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26
Q

How does a mass spectrometer determine the mass of a molecule or isotope?

A

By measuring the mass-to-charge ratio of ions.

It does this by causing substances to become positive ions and are then passes through the apparatus and separated according to their mass and charge.

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27
Q

How would you calculate the relative abundance from a mass spectrum?

A
  1. Work out the total heights of all peaks on spectra.
  2. Use the formula to work out the height of each peak:
    % abundance = (height of peak / total height of peaks) x 100
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28
Q

How would you calculate the relative atomic mass from a mass spectrum?

A

Ar = (relative isotopic mass X percentage abundance) + (relative isotopic mass X percentage abundance) + (relative isotopic mass X percentage abundance) / 100

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29
Q

Atoms of metals in groups 1-13 ….. electrons and form …. ions with the electron configuration of the … noble gas in the periodic table.

A

Atoms of metals in groups 1-13 LOSE electrons and form POSITIVE ions with the electron configuration of the PREVIOUS noble gas in the periodic table.

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30
Q

Atoms of metals in groups 15-17 …. electrons and form …. ions with the electron configuration of the …. noble gas in the periodic table.

A

Atoms of metals in groups 15-17 GAIN electrons and form NEGATIVE ions with the electron configuration of the NEXT noble gas in the periodic table.

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31
Q

Name the formulae of the 1+ molecular ions you need to know.

A

Ammonium, NH4+

Silver, Ag+

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32
Q

Name the formulae of the 1- molecular ions you need to know.

A

Hydroxide, OH-

Nitrate, NO3-

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33
Q

Name the formulae of the 2- molecular ions you need to know.

A

Carbonate, CO3 2-

Sulphate, SO4 2-

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34
Q

Name the formulae of a 2+ molecular ion you need to know.

A

Zinc, Zn2+

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35
Q

Define “amount of substance”.

A

The quantity that has moles as its unit. Chemists use ‘amount of substance’ as a way of counting atoms.

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36
Q

What is Amount of substance based on?

A

A standard count of atoms called the Avogadro’s constant, NA.

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37
Q

What is the Avogadro’s constant?

A

The number of atoms per mole of the carbon-12 isotope (6.02 x 10^23 mol-1)

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38
Q

What is a mole?

A

The amount of any substance containing as many particles as there are carbon atoms in exactly 12g of the carbon-12 isotope.

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39
Q

What is Molar Mass?

A

The mass per mole of a substance. The units of molar mass are g mol-1

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40
Q

What’s the formula to finding the number of moles?

A

no. of moles (mol) = mass (g) / Mr (g mol-1)

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41
Q

Define “empirical formula”.

A

The simplest whole number ratio of atoms of each element present in a compound.

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42
Q

What are the steps to calculate the empirical formulae?

A
  1. Divide the amount of each element present by its molar mass - this will give you the molar mass.
  2. Divide the answer for each element by the smallest number - this ensures your ratio is in the format 1:x.
  3. If necessary, multiply the answer by a suitable value to make sure the ratio is in whole number only.
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43
Q

Define “molecular formulae”.

A

Tells you the number of each type of atom that make up a molecule.

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44
Q

What are the steps to calculate the molecular formulae?

A
  1. Work out the empirical formula mass of the compound.
  2. Find the number of units of that compound in a molecule (Mr / empirical formula mass)
  3. Work out the molecular formulae by multiplying the number of units by the compound. (e.g. 4 x CH2 = C4H8)
45
Q

Define “molar gas volume”.

A

Volume per mole of a gas.

The units of molar volume are dm^3 mol-1.

46
Q

How much does one mole of gas approximately occupies?

A

24.0 dm3 (24 000 cm3)

47
Q

What is the volume per mole of gas molecules?

A

24.0 dm^3 mol^-1

48
Q

State the ideal gas equation and state each variable (with their units).

A 
pV = nRT
p = Pressure (Pascals)
V = Volume (m^3) 
n = number of moles (mol)
R = Gas constant, 8. 314 (J mol-1 K-1)
T = Temperature (K)
49
Q

How would you convert from m^3 to dm^3 to cm^3?

A

m^3 → dm^3: x 1000

m^3 → cm^3: x 1 000 000

50
Q

State the formula to work out the number of moles of a gas if RTP is present.

A

if V is in dm^3: n = V / 24.0

if V is in cm^3: n = V / 24000

51
Q

What is the formula to calculate the number of moles of a solution?

A

n (mol) = c (mol dm^-3) x V (dm^3)

n (mol) = c (mol dm^-3) x V (cm^3) / 1000

52
Q

Define “concentration”.

A

The concentration of a solution is the amount of solute, in mol, dissolved per 1 dm^3 (1000 cm^3) of solution.

53
Q

What is a standard solution?

A

A solution of known concentration. Normally used in titrations to determine unknown information about another substance.

54
Q

What are the 5 steps to make up a standard solution?

A
  1. Weigh out the solute.
  2. Completely dissolve the solute in solvent in a beaker. Transfer the solution to the flask and rinse the beaker repeatedly, using more solvent, adding the rinsings to the flask.
  3. Add solvent to the flask, but do not fill all the way up to the graduation line.
  4. Carefully add solvent drop by drop up to the line on the flask, until the bottom of the meniscus sits exactly on the graduation mark on the flask. If the solution goes over the meniscus line, you must throw it away and start again.
55
Q

Define “mass concentration”.

A

The mass dissolved in 1dm^3 of solution (g dm^-3)

56
Q

Define the terms “concentrated” and “dilute”.

A

Concentrated - a large amount of solute per dm^3.

Dilute - a small amount of solute per dm^3.

57
Q

State the 4 state symbols.

A

(s) = solid
(l) = liquid
(g) = gaseous
(aq) = aqueous

58
Q

What is stoichiometry about?

A

The study of amounts of substances that are involved in a chemical reaction.

59
Q

What is the formula for percentage yield?

A

% yield = actual amount, in mol, of product / Theoretical amount, in mol, of product x100

60
Q

What are the steps to calculate percentage yield?

A
  1. Find the MOLES of REACTANT.
  2. Use this (and balanced equation) to find THEORETICAL MOLES of PRODUCT.
  3. Find the THEORETICAL MASS of product.
  4. DIVIDE ACTUAL MASS by THEORETICAL MASS and MULTIPLY BY 100.
61
Q

Define % yield.

A

The mass/moles of priduct obtained in a reaction expressed as as percentage of what you should’ve expected.

62
Q

Give 5 reasons of low % yields.

A
  1. Reaction may be at EQUILIBRIUM and may not go to completion.
  2. SIDE REACTIONS may occur, leading to by-products.
  3. The reactants may NOT be PURE.
  4. Some of the reactants or products may be left behind in the apparatus used in the experiment.
  5. Separation and purification may result in the loss of some of the product.
63
Q

Define Atom Economy.

A

The amount of starting material that ends up as USEFUL PRODUCTS.

64
Q

What is the formula of Atom Economy?

A

Mr of Desired Product / Mr of ALL Products x 100

65
Q

What are the 3 steps to calculate Atom Economy?

A
  1. Identify the useful product and find its Mr (use balanced equations)
  2. Work out the TOTAL Mr of all products.
  3. Use the formula to calculate Atom Economy.
66
Q

State the 4 common acids.

A
  1. HCl (hydrochloric acid)
  2. H2SO4 (sulphuric acid)
  3. HNO3 (nitric acid)
  4. CH3 COOH (ethanoic (acetic) acid)
67
Q

When acid is added to water, what does the acid do?

A

Releases H+ ions into solution. This therefore makes them PROTON DONORS.

68
Q

What’s the difference between weak and strong acids?

A

Weak acids are not very good at giving H+ ions away so they only partially dissociate.
Strong acids are very good at giving up H+ ions so they almost fully dissociate.

69
Q

What are examples of common bases?

A

Metal oxides: e.g. MgO, CuO

Metal hydroxides: NaOH, Mg(OH)2

70
Q

What is the definition of a base?

A

Proton ACCEPTOR - they neutralise acids.

71
Q

What alkali is dissolved in water, what does the alkali do?

A

Releases OH- (aq) ions.

72
Q

Give 3 examples of alkalis.

A
  1. NaOH (Sodium Hydroxide)
  2. KOH (Potassium Hydroxide)
  3. NH3 (Ammonia, or aqueous ammonia)
73
Q

Why is ammonia a weak base?

A

Only a small proportion of the dissolved NH3 reacts with water.

74
Q

What are amphoteric substances? Give an example.

A

Substances that can behave as acids AND bases.

e.g. AMINO ACID (H2N = base, COOH = acid)

75
Q

What are the benefits of sustainability of developing chemical processes with a high atom economy?

A
  • chemical companies can reduce the amount of waste produced.
  • processes can be maintained at a product level without completely depleting resources.
76
Q

State the neutralisation reaction of an acid and a base.

A

H+ (acid) + OH- (base) → H2O (water)

77
Q

Define a “salt”

A

A chemical compound where the H+ ion has been replaced by a positive ion.

78
Q

State the four general equations for how salts are formed by reactions of acids.

A

(1) acid + carbonate → salt + water + carbon dioxide
(2) acid + metal oxide → salt + water
(3) acid + alkali → salt + water
(4) acid + metal → salt + hydrogen

79
Q

What happens when you mix AQUEOUS ammonia and an acid?

A

Neutralisation reaction - acids are neutralised by the aqueous ammonia and ammonium salts are formed, contianing the ammonium ion NH4+.

NH3 (aq) + HNO3 (aq) → NH4 NO3 (aq)

80
Q

What is the difference between hydrated and anhydrous crystals?

A

Hydrated - the crystalline compound contains water molecules.
Anhydrous - substance contains no water.

81
Q

Define WATER OF CRYSTALLISATION.

A

When a compound crystallises within water, the water can become part of the resulting crystalline structure.

82
Q

What does “.xH20” (dot formulae) mean?

A

The WATERS OF CRYSTALLISATION - gives the ratio between the number of compound molecules and the number of water molecules within the crystalline structure.

83
Q

Give 4 potential unknown information you can find out from a titration?

A
  1. Concentration of solution
  2. Molar mass
  3. A formula
  4. The number of molecules of water of crystallisation
84
Q

Give the 4 steps of carrying out a titration.

A
  1. Using a pipette, add a measured volume of one solution (standard solution) to a conical flask. Add a suitable indicator.
  2. Place the other solution in a BURETTE.
  3. Add the solution in the burette to the solution in the conical flask until the reaction has JUST been completed (END POINT) of the titration. Measure the volume of the solution added from the burette.
  4. You now know the volume of one solution that EXACTLY reacts with the volume of the other solution.
85
Q

How do we identify an end point?

A

Using an indicator that must be a different colour in the acidic solution than in the basic solution.

86
Q

State 3 common acid-base indicators.
State their colour in acid.
State their colour in base.
State their end point colour.

A

Methyl orange - red - yellow - orange
Bromothymol blue - yellow - blue - green
Phenolphthalein - colourless - pink - pale pink

87
Q

Define OXIDATION NUMBER.

A

The number of electrons that an atom uses to bond with atoms of another element.

88
Q

What oxidation number does an UNCOMBINED ELEMENT (e.g. C, Na, O2, P4) have?

A

0 (ZERO)

89
Q

What oxidation number does a COMBINED OXYGEN (e.g. H20, CaO) have?

A

-2

90
Q

What oxidation number does a COMBINED OXYGEN in PEROXIDES (e.g. H2O2) have?

A

-1

91
Q

What oxidation number does a COMBINED HYDROGEN (e.g. NH3, H2S) have?

A

+1

92
Q

What oxidation number does a COMBINED HYDROGEN in METAL HYDRIDES (e.g LiH) have?

A

-1

93
Q

What oxidation numbers do SIMPLE IONS (e.g. Na+ or Mg2+) have?

A

The charge on that ion, e.g. if it was Na+ then it would have a +1 oxidation number. If it was Cl- then it would have a -1 oxidation number.

94
Q

What oxidation number does a COMBINED FLUORINE (e.g. NaF, CaF2, AlF3) have?

A

-1

95
Q

What’s an easy way to remember the oxidation numbers?

A

123 FHOC (it’s similar to pronouncing f**k)

1 - Any group 1 element has an Oxidation No. of +1
2 - Any group 2 element has an Oxidation No. of +2
3 - Any group 3 element has an Oxidation No. of +3
F - Fluorine = -1
H = NEARLY always +1
O = NEARLY always -2
C (Chlorine) = NEARLY always -1 EXCEPT WHEN BONDED TO OXYGEN)

96
Q

In compounds, what must the sum of the oxidation numbers equal to (the overall charge)?

A

0 (ZERO)

97
Q

In molecular ions, what must the sum of the oxidation numbers equal to (the overall charge)?

A

The charge on that ion, e.g. CO3 2-

So therefore the individual atoms must add to make a charge of -2.

98
Q

What do the Roman Numerals (e.g. I, II, III, IV, V, VI) represent in transition elements (and oxyanions)?

A

THE OXIDATION NUMBER/STATE.

e.g. Iron(II)Chloride, Fe has an oxidation number of +2.

99
Q

What are oxyanions?

A

Negative ions that contain an element along with oxygen, e.g. SO4 2- , CO3 2-, NO3 -
They usually end in -ate, to indicate oxygen, e.g. carbonate.

100
Q

In terms of gaining or losing electrons, define what oxidation and reduction are.

A

OXIDATION - THE LOSS OF ELECTRONS

REDUCTION - THE GAIN OF ELECTRONS.

101
Q

What’s a good mnemonic to memorise the definitions of oxidation and reductions?

A

OIL RIG
OIL = Oxidation is LOSS
RIG = Reduction is GAIN

102
Q

Why is reduction the GAIN of electrons?

A

The substance that is reduced TAKES electrons from the substance that is oxidised.

103
Q

What is a redox reaction?

A

A reaction in which both REDuction and OXidation take place.

104
Q

X (Group 2) → X 2+ + 2e-

What type of reaction is this half equation?

A

OXIDATION, because it’s GIVING AWAY (losing) electrons for REDUCTION to TAKE.

105
Q

X2 + 2e- → 2X-

What type of reaction is this half equation?

A

REDUCTION, because it’s TAKING IN (gaining) the electrons.

106
Q

Using the science behind oxidation states, why do metals form positive ions and why do non-metals form negative ions?

A

Metals tend to be OXIDISED, LOSING electrons to form POSITIVE IONS.
Nonmetals, tend to be reduced, GAINING electrons to form negative ions.

107
Q

How ELSE can oxidation and reduction can be described as?

A

In terms of oxidation number:

  • Reduction is a DECREASE in oxidation number.
  • Oxidation is an INCREASE in oxidation number.
108
Q

Describe redox reactions of metals with acids.

A

The metal is oxidised, forming positive metal ions.

The hydrogen in the acid is reduced, forming the element hydrogen as a gas (H2).

OCR Chemistry AS Level (6 decks)

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