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A Level Chemistry > Module 2.1 - Atoms and Reactions > Flashcards

Module 2.1 - Atoms and Reactions Flashcards

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1
Q

What are isotopes?

A

Atoms of the same element with different masses, different number of neutrons but same number of protons and electrons.

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2
Q

Describe the current model of the atom (5)

A

> Protons and Neutrons found in nucleus (at centre)
Electrons orbit the nucleus in ‘shells’
The nucleus is tiny compared to total volume of atom
Nucleus is very dense and accounts for almost all the atoms mass
Most of the atom is empty space between nucleus and electron’s shells

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3
Q

What are the relative masses and relative charges of a proton, electron and neutron?

A

Proton - 1.0 - 1+
Electron - 1 ÷ 2000 - 1-
Neutron - 1.0 - 0

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4
Q

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

A

> Chemical reactions involve electrons and isotopes have the same number and arrangement of electrons
Neutrons don’t affect chemical reactivity

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5
Q

What is relative isotopic mass?

A

Mass of an atom of an isotope compared with 1/12 of the mass of an atom of carbon 12

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6
Q

What is relative atomic mass?

A

Weighted mean mass of an atom of an element compared with 1/12 of the mass of an atom of carbon-12.

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7
Q

Explain the unified atomic mass unit, u.

A

> 1u is a tiny mass; 1.660540210 x 10^-27
The mass of an atom of carbon-12 is defined as 12u
So the mass of 1/12th of an atom of carbon-1 is 1u

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8
Q

Work out the relative atomic mass: A sample of bromine contains 53.00% of bromine-79 and 47.00% of bromine-81.

A

(53 x 79) + (47x 81) = 7994

7994 ÷ 100 = 79.94

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9
Q

Work out the RFM: CaBr2

A

40.1 + (79.9 x 2) = 199.9

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10
Q

What is mass spectrometry used for? (2)

A

> The determination of relative isotopic masses and relative abundances of the isotopes.
Calculation of relative atomic mass from the relative abundance of its isotopes.

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11
Q

How does a mass spectrometer determine the mass of a molecule or isotope? (5)

A

> Measures the mass to charge ratio
Turns substances into positive ions
Ions passed through apparatus and separated according to mass and charge
A computer analyses data on the ions and makes a mass spectrum
Similar to bar graph and gives information about the abundance of ions

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12
Q

Describe a mass spectrum

A

Shows a relative (percentage) abundance on the y-axis and mass-to-charge ratios on the x-axis

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13
Q

What is m/z?

A

Mass/Charge

*Charge is usually 1 on the ion

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14
Q

How would you work out the relative abundances without values or scale on the y-axis?

A

Measure each line. Find a percentage of each peak out of the total height of peaks.

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15
Q

What do atoms of metals in group 1-13 form?

A

Lose electrons to form positive ions with the electron config of the previous noble gas in the periodic table

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16
Q

What do atoms of non-metals in groups 15-17 form?

A

Gain electrons to form negative ions with the electron config of the next noble gas in the periodic table.

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17
Q

What do atoms of Beryllium (Be), Boron (B), Carbon (C) and Silicon (Si) form?

A

> Do not normally form ions

> Require too much energy to transfer the outer shell electrons to form ions

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18
Q

Give two examples of elements that can form more than one ion

A

> Iron(II) Fe^2+ and Iron(III) for Fe^3+

> Copper(I) Cu^+ and Copper(II) Cu^2+

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19
Q

What is the charge and formula of Ammonium?

A

> 1+

> NH4 ^+

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20
Q

What is the charge and formula of Hydroxide?

A

> 1-

> OH^-

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21
Q

What is the charge and formula of Nitrate?

A

> 1-

>NO3 ^-

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22
Q

What is the charge and formula of Carbonate?

A

> 2-

> CO3 ^2-

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23
Q

What is the charge and formula of Sulfate?

A

> 2-

> SO4 ^2-

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24
Q

How do work out the ionic formula from ionic charges? Do it for Calcium Chloride (3)

A

> Ionic compound - overall charge is zero
Ca is 2+, Cl is 1-
CaCl2

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25
Q

Define Amount of Substance

A

A quantity for counting atoms

  • Symbol: n
  • Unit: mole (mol)
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26
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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27
Q

What is avogadro’s constant? (Na)

A

The number of particles per mole - 6.02 x 10^23

1 mol of atom = 6.02 x 10^23 atoms

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28
Q

Define Molar Mass, M

A

> Mass per mole
gmol^-1
For CO2: 12.0 + (16.0 x 2) = 44.0 gmol^-1

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29
Q

How do you work out n.o of moles in a substance?

A

Mass (g) ÷ Molar Mass (gmol^-1)

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30
Q

Define Empirical Formula

A

The SIMPLEST WHOLE NUMBER ratio of atoms of each element present in COMPOUND.

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31
Q

Work out the empirical formula:

0.0675g magnesium combines with 3.995g bromine to form a compound [Ar : Mg - 24.3 , Br - 79.9]

A

Find the molar ratio - 0.0675 ÷ 24.3 : 3.995 ÷ 79.9
- 0.025 : 0.050
Divide by smallest number ( 0.025) - 1 : 2

Empirical formula is therefore - MgBr2

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32
Q

Define Molecular Formula

A

The NUMBER and TYPE of atoms of each element in a molecule

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33
Q

Calculate the molecular formula of a compound with an empirical formula of CH2 and a relative molecular mass of 56..

A

Empirical formula mass = 12.0 + (1.0 x 2) = 14

56 ÷ 14 = 4

4 x CH2 = C4H8

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34
Q

Define Molar gas volume

A

Volume per mole of gas - at RTP, it is 24.0 dm^3 mol^-1

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35
Q

Why do different gases of 1 mole have the same volume but different densities?

A

Have different densities because there are same number of particles but different mass per particle.

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36
Q

How do you work the n.o of moles of a volume, V, of a gas at RTP?

A

> (dm^3) - n = V ÷ 24.0

> (cm^3) - n = V ÷ 24,000

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37
Q

What is the ideal gas equation?

A

pV = nRT

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38
Q

What is p in pV = nRT?

A

p - pressure - Pa

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39
Q

What is V in pV = nRT?

A

V - Volume - m^3

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40
Q

What is n in pV = nRT?

A

n - mol

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41
Q

What is R in pV = nRT?

A

R - Gas Constant - 8.314 J mol^-1 K^-1

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42
Q

What is T in pV = nRT?

A

T - temperature - K

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43
Q

How do you work out mass concentration?

A

Mass ÷ Volume

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44
Q

How do you work out mole concentration?

A

Mole ÷ Volume

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45
Q

How do you work out mole concentration using mass concentration?

A

Mass concentration ÷ Mr

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46
Q

Balance:

C3H8 + O2 —> CO2 + H2O

A

C3H8 + 5O2 —> 3CO2 + 4H2O

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47
Q

What is a Standard solution?

A

A solution of KNOWN concentration. Normally used in titration to determine unknown information about another substance.

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48
Q

Define Concentrated

A

LARGE amount of SOLUTE, mol, per dm^3

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49
Q

Define Dilute

A

SMALL amount of SOLUTE, mol, per dm^3

50
Q

How do you work out percentage yield?

A

[Actual amount (mol) ÷ Theoretical amount (mol)] x 100

51
Q

4 reasons why percentage yields less than 100%?

A

> Side Reaction may occur
Reactants/Products may be left behind
Reaction maybe at equilibrium so will not be fully completed

52
Q

How do you work out atom economy?

A

(Mr of desired product ÷ Sum of Mr of all products) x 100

53
Q

Which type of reaction have an atom economy of 100%?

A

Addition reactions

54
Q

Which type of reactions have an atom economy less than 100%?

A

Substitution and Elimination reactions

55
Q

How can atom economy benefit society?

A

> Using processes with a higher atom economy, chemical companies can reduce the amount of waste produced. This is beneficial for companies as it can be expensive to treat the waste produced.
Makes processes sustainable as they can be maintained at productive level without depleting resources.

56
Q

What is the difference between atom economy and percentage yield?

A

Percentage yield tells you the EFFIENCY of reactants turning into products whereas atom economy tells you the PROPORTION of desired products to all the products formed.

57
Q

Name four common acids and their formulae

A

> Sulfuric acid - H2SO4
Hydrochloric acid - HCl
Nitric Acid - HNO3
Acetic (Ethanoic) acid - CH3COOH

58
Q

What do acids do in aqueous solution? include eq

A

Acids release H+ ions in water. HCl (g) —> H+ (aq) + Cl- (aq)

59
Q

Is acid a proton donor or acceptor?

A

Releases H+ ions in water so proton DONOR

60
Q

Explain Strong acids in terms of dissociation

A

Very good releasing H+ ions - FULLY DISSOCIATE

61
Q

Explain Weak acids in terms of dissociation

A

Not very good at releasing H+ ions as they are quickly taken back by the acid - PARTIALLY DISSOCIATE

62
Q

Are bases proton donors or acceptors?

A

Proton ACCEPTORS

63
Q

What do bases neutralise? include eq

A

Acids - H+ (aq) + OH- (aq) —> H2O (l)

64
Q

What are some common bases?(e.g.)

A

> Metal hydroxides - Mg(OH)2, NaOH

> Metal oxides - MgO, CuO

65
Q

What are alkalis?

A

Type of BASE

66
Q

Name three common alkalis

A

> Sodium hydroxide - NaOH
Potassium hydroxide - KOH
Ammonia - NH3

67
Q

What do alkalis do in aqueous solution? include eq

A
  • Alkalis release OH- ions in water

- NaOH (s) + aq —> Na+ (aq) + OH- (aq)

68
Q

Describe ammonia as a base. include eq

A

A gas that dissolves in water to form a WEAK alkaline solution. Dissolved ammonia reacts with water: NH3 (aq) + H20 (l) —> NH4+ (aq) + OH - (aq)

69
Q

What type of compound is salt?

A

IONIC

70
Q

What makes the cation of a salt?

A

METAL or AMMONIUM ion

71
Q

What makes the anion of a salt?

A

Derived from an acid

72
Q

Name 4 common salts

A

> Sodium sulfate - Na2SO4
Sodium chloride - NaCl
Calcium nitrate - Ca(NO3)2
Potassium sulfate - K2SO4

73
Q

How are salts formed from carbonates? (General eq, e.g., ionic of e.g., state symbols)

A
  • acid + carbonate —> salt + carbon dioxide + water
  • 2HCl + CaCO3 —> CaCl2 + H2O + CO2
  • 2H+ + CaCO3 —> Ca^2+ + H2O + CO2
  • (aq) + (s) —> (aq) + (l) + (g)
74
Q

How are salts formed from metal oxides? (General eq, e.g., ionic of e.g., state symbols)

A
  • acid + metal oxide —> salt + water
  • 2HCl + CaO —> CaCl2 + H2O
  • 2H+ + CaO —> Ca^2+ + H2O
  • (aq) + (s) —> (aq) + (l)
75
Q

How are salts formed from alkalis? (General eq, e.g., ionic of e.g., state symbols)

A
  • acid + alkali —> salt + water
  • HCl + NaOH —> NaCl + H2O
  • H+ + OH- —> H2O
  • (aq) + (aq) —> (l)
76
Q

How are salts formed from metal? (General eq, e.g., state symbols)

A
  • acid + metal —> salt + hydrogen
  • 2HCl + 2Li —> 2LiCl + H2
  • (aq) + (s) —> (aq) + (g)
77
Q

What type of reaction is the formation of a salt from a metal?

A

REDOX

78
Q

Define Hydrated crystals

A

a crystalline compound containing WATER MOLECULES

79
Q

Define Anhydrous crystals

A

A substance that contains NO water molecules

80
Q

What is water of crystallisation?

A

Water molecules that form an essential part of a the CRYSTALLINE structure of a compound

81
Q

What is the formula of anhydrous copper sulfate?

A

CuSO4

82
Q

What is the formula of hydrated copper sulfate?

A

CuSO4.5H2O

83
Q

What is the dot formulae?

A

Gives the RATIO between the number of COMPOUND molecules and WATER molecules

84
Q

Find the dot formula:

The empirical formula for hydrated magnesium chloride is MgCl2H10O5

A

10 H atoms so there must be 5 water molecules, .5H2O

MgCl2.5H2O

85
Q

Find the dot formula:
From an experiment to determine the formula of hydrated magnesium sulfate:
> Mass of MgSO4.xH2O - 4.312g
> Mass of MgSO4 - 2.107g

A

Mass of H2O = 4.312 - 2.107 = 2.205g

Mol of MgSO4 = 2.107 ÷ 120.4 = 0.0175
Mol of H2O = 2.205 ÷ 18.0 = 0.01225

0.01225:0.0175 = 1:7 - x = 7 - MgSO4.7H2O

86
Q

What are the 4 steps for completing a titration?

A

1) Using a pipette, add a measured volume of acid into a conical flask. Add a suitable indicator
2) Place the base into the burette
3) Add the base into the conical flask slowly until the reaction is complete and record the volume of base used
4) Repeat until concordant results are achieved

87
Q

What is the colour of methyl orange in acids, alkali and the end of a titration?

A

> Acid - red
Alkali - yellow
End - orange

88
Q

What is the colour of bromothyl blue in acids, alkali and the end of a titration?

A

> Acid - yellow
Alkali - blue
End - green

89
Q

What is the colour of phenolphthalein in acids, alkali and the end of a titration?

A

> Acid - colourless
Alkali - pink
End - pale pink

90
Q

Calculate the concentration of H2SO4:
In a titration, 25.0cm^3 of 0.15mol dm^-3 NaOH reacted exactly with 23.40cm^3 of H2SO4.

2NaOH + H2SO4 —> Na2SO4 + 2H2O

A

mol of NaOH = conc x vol
= 0.15 x (25 ÷ 1000)
= 3.75 x 10^-3 mol
mol of H2SO4 = 3.75 x 10^-3 ÷ 2
= 1.875x10^-3

Conc of H2SO4 = moles ÷ volume
= 1.875 x 10^-3 ÷ (23.4 ÷ 1000)
= 0.08 mol dm^-3 q

91
Q

What is the oxidation number of a uncombined element? (e.g.)

A

> 0

> C, Na, O2, P4

92
Q

What is the oxidation number of a combined oxygen? (e.g.)

A

> -2

> H2O, CaO

93
Q

What is the oxidation number of a combined oxygen in peroxides? (e.g.)

A

> -1

> H2O2

94
Q

What is the oxidation number of a combined hydrogen? (e.g.)

A

> +1

> NH3, H2S

95
Q

What is the oxidation number of combined hydrogen in metal hydrides? (e.g.)

A

> -1

> LiH

96
Q

What is the oxidation number of a simple ion? (e.g.)

A

> Charge on ion

> Na+: -1, Mg2+: -2

97
Q

What is the oxidation number of combined fluorine? (e.g.)

A

> -1

> CaF2, NaF

98
Q

What are the oxidation number for all the element in sulfur dioxide?

A
  • S: +4
  • O: -2
  • O: -2
99
Q

What are the oxidation numbers for the elements in a carbonate ion?

A
  • C: +4
  • O: -2
  • O: -2
  • O: -2
100
Q

What are the oxidation numbers for Fe in (II)FeCl2 and (III)FeCl2?

A
  • (II) FeCl2: +2

* (III) FeCl2: +3

101
Q

What is the oxidation number for nitrogen in the nitrate ion?

A

• NO2^- : + 3

102
Q

What is the oxidation number of sulfur in the sulfate ion?

A

• SO4^2-: +6

103
Q

What is oxidation in terms of electrons?

A

LOSS of electron

104
Q

What is Oxidation in terms of oxygen?

A

GAIN of oxygen

105
Q

What is Oxidation in terms of oxidation numbers?

A

Increases in oxidation number

106
Q

How can oxidation be shown in an equation?

A

Mg —> Mg^2+ + 2e^-

107
Q

What is reduction in terms of electrons?

A

GAIN of electrons

108
Q

What is reduction in terms of oxygen?

A

LOSS of oxygen

109
Q

What is reduction in terms of oxidation number?

A

DECREASE in oxidation number

110
Q

How can reduction be shown in an equation?

A

Cl2 + 2e^- —> 2Cl^-

111
Q

What is a redox reaction?

A

A reaction where reduction and oxidation both takes place

112
Q

What are the oxidation numbers? Which undergo reduction and oxidation?
Mg + Cl2 —> MgCl2

A

> Mg: 0 —> +2: oxidation

> Cl: 0 —> -1: reduction

113
Q

What are the oxidation numbers? Which undergo reduction and oxidation?
MnO2 + 4HCl —> MnCl2 + 2H2O + Cl2

A

> Mn: +4 —> +2: reduction

> Cl: -1 —> 0: oxidation

114
Q

Describe the development of the atom (9).

A
  1. Democritus developed the first idea of the atom. Believed that there was a limited n.o of times that matter could be spilt up (particle was the smallest).
  2. Dalton developed this - atoms cannot be divided, atoms are tiny particles that make up element, atoms of a given element are the same but different for different elements. Developed first table.
  3. Thomson discovered electrons - plum pudding model - negative electrons moving in a sea of positive charge.
  4. Rutherford disproved pp model. Proposed the nuclear atom.
  5. Bohr believed electrons followed certain paths. Helped explain some periodic properties.
  6. Rutherford discovered the proton
  7. de Broglie suggested particles have a wave and particles like behaviour.
  8. Chadwick discovered the neutron.
  9. Modern - protons and neutrons made up of smaller particles - quarks.
115
Q

What is relative molecular mass, Mr?

A

> The addition of the relative atomic masses of each atom making up a molecule.
Used for simple molecules.

116
Q

What is relative formula mass?

A

> The addition of the relative atomic masses of each atom making up a formula unit .
Used for compounds.

117
Q

What is the charge and formula of a Zinc ion?

A

> 2+

> Zn^2+

118
Q

What is the charge and formula of a Silver ion?

A

> 3+

> Ag^3+

119
Q

What are the ionic charges of group 1-13 elements?

A

> Lose electrons

> Positive ions

120
Q

What are the ionic charges of group 15-17 elements?

A

> Gain electrons

> Negative ions

121
Q

What atoms do not normally form ions? Why?

A

> Be, B, C and Si

> Requires too much energy to transfer outer shell electrons

A Level Chemistry (7 decks)

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