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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
Define Amount of Substance
A quantity for counting atoms - Symbol: n - Unit: mole (mol)
26
What is a mole?
The amount of any substance containing as many particles as there are carbon atoms in exactly 12g of the carbon-12 isotope.
27
What is avogadro’s constant? (Na)
The number of particles per mole - 6.02 x 10^23 1 mol of atom = 6.02 x 10^23 atoms
28
Define Molar Mass, M
> Mass per mole > gmol^-1 > For CO2: 12.0 + (16.0 x 2) = 44.0 gmol^-1
29
How do you work out n.o of moles in a substance?
Mass (g) ÷ Molar Mass (gmol^-1)
30
Define Empirical Formula
The SIMPLEST WHOLE NUMBER ratio of atoms of each element present in COMPOUND.
31
Work out the empirical formula: | 0.0675g magnesium combines with 3.995g bromine to form a compound [Ar : Mg - 24.3 , Br - 79.9]
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
32
Define Molecular Formula
The NUMBER and TYPE of atoms of each element in a molecule
33
Calculate the molecular formula of a compound with an empirical formula of CH2 and a relative molecular mass of 56..
Empirical formula mass = 12.0 + (1.0 x 2) = 14 56 ÷ 14 = 4 4 x CH2 = C4H8
34
Define Molar gas volume
Volume per mole of gas - at RTP, it is 24.0 dm^3 mol^-1
35
Why do different gases of 1 mole have the same volume but different densities?
Have different densities because there are same number of particles but different mass per particle.
36
How do you work the n.o of moles of a volume, V, of a gas at RTP?
> (dm^3) - n = V ÷ 24.0 | > (cm^3) - n = V ÷ 24,000
37
What is the ideal gas equation?
pV = nRT
38
What is p in pV = nRT?
p - pressure - Pa
39
What is V in pV = nRT?
V - Volume - m^3
40
What is n in pV = nRT?
n - mol
41
What is R in pV = nRT?
R - Gas Constant - 8.314 J mol^-1 K^-1
42
What is T in pV = nRT?
T - temperature - K
43
How do you work out mass concentration?
Mass ÷ Volume
44
How do you work out mole concentration?
Mole ÷ Volume
45
How do you work out mole concentration using mass concentration?
Mass concentration ÷ Mr
46
Balance: | C3H8 + O2 —> CO2 + H2O
C3H8 + 5O2 —> 3CO2 + 4H2O
47
What is a Standard solution?
A solution of KNOWN concentration. Normally used in titration to determine unknown information about another substance.
48
Define Concentrated
LARGE amount of SOLUTE, mol, per dm^3
49
Define Dilute
SMALL amount of SOLUTE, mol, per dm^3
50
How do you work out percentage yield?
[Actual amount (mol) ÷ Theoretical amount (mol)] x 100
51
4 reasons why percentage yields less than 100%?
> Side Reaction may occur > Reactants/Products may be left behind > Reaction maybe at equilibrium so will not be fully completed
52
How do you work out atom economy?
(Mr of desired product ÷ Sum of Mr of all products) x 100
53
Which type of reaction have an atom economy of 100%?
Addition reactions
54
Which type of reactions have an atom economy less than 100%?
Substitution and Elimination reactions
55
How can atom economy benefit society?
> 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
What is the difference between atom economy and percentage yield?
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
Name four common acids and their formulae
> Sulfuric acid - H2SO4 > Hydrochloric acid - HCl > Nitric Acid - HNO3 > Acetic (Ethanoic) acid - CH3COOH
58
What do acids do in aqueous solution? include eq
Acids release H+ ions in water. HCl (g) —> H+ (aq) + Cl- (aq)
59
Is acid a proton donor or acceptor?
Releases H+ ions in water so proton DONOR
60
Explain Strong acids in terms of dissociation
Very good releasing H+ ions - FULLY DISSOCIATE
61
Explain Weak acids in terms of dissociation
Not very good at releasing H+ ions as they are quickly taken back by the acid - PARTIALLY DISSOCIATE
62
Are bases proton donors or acceptors?
Proton ACCEPTORS
63
What do bases neutralise? include eq
Acids - H+ (aq) + OH- (aq) —> H2O (l)
64
What are some common bases?(e.g.)
> Metal hydroxides - Mg(OH)2, NaOH | > Metal oxides - MgO, CuO
65
What are alkalis?
Type of BASE
66
Name three common alkalis
> Sodium hydroxide - NaOH > Potassium hydroxide - KOH > Ammonia - NH3
67
What do alkalis do in aqueous solution? include eq
- Alkalis release OH- ions in water | - NaOH (s) + aq —> Na+ (aq) + OH- (aq)
68
Describe ammonia as a base. include eq
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
What type of compound is salt?
IONIC
70
What makes the cation of a salt?
METAL or AMMONIUM ion
71
What makes the anion of a salt?
Derived from an acid
72
Name 4 common salts
> Sodium sulfate - Na2SO4 > Sodium chloride - NaCl > Calcium nitrate - Ca(NO3)2 > Potassium sulfate - K2SO4
73
How are salts formed from carbonates? (General eq, e.g., ionic of e.g., state symbols)
* acid + carbonate —> salt + carbon dioxide + water * 2HCl + CaCO3 —> CaCl2 + H2O + CO2 * 2H+ + CaCO3 —> Ca^2+ + H2O + CO2 * (aq) + (s) —> (aq) + (l) + (g)
74
How are salts formed from metal oxides? (General eq, e.g., ionic of e.g., state symbols)
* acid + metal oxide —> salt + water * 2HCl + CaO —> CaCl2 + H2O * 2H+ + CaO —> Ca^2+ + H2O * (aq) + (s) —> (aq) + (l)
75
How are salts formed from alkalis? (General eq, e.g., ionic of e.g., state symbols)
* acid + alkali —> salt + water * HCl + NaOH —> NaCl + H2O * H+ + OH- —> H2O * (aq) + (aq) —> (l)
76
How are salts formed from metal? (General eq, e.g., state symbols)
* acid + metal —> salt + hydrogen * 2HCl + 2Li —> 2LiCl + H2 * (aq) + (s) —> (aq) + (g)
77
What type of reaction is the formation of a salt from a metal?
REDOX
78
Define Hydrated crystals
a crystalline compound containing WATER MOLECULES
79
Define Anhydrous crystals
A substance that contains NO water molecules
80
What is water of crystallisation?
Water molecules that form an essential part of a the CRYSTALLINE structure of a compound
81
What is the formula of anhydrous copper sulfate?
CuSO4
82
What is the formula of hydrated copper sulfate?
CuSO4.5H2O
83
What is the dot formulae?
Gives the RATIO between the number of COMPOUND molecules and WATER molecules
84
Find the dot formula: | The empirical formula for hydrated magnesium chloride is MgCl2H10O5
10 H atoms so there must be 5 water molecules, .5H2O MgCl2.5H2O
85
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
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
What are the 4 steps for completing a titration?
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
What is the colour of methyl orange in acids, alkali and the end of a titration?
> Acid - red > Alkali - yellow > End - orange
88
What is the colour of bromothyl blue in acids, alkali and the end of a titration?
> Acid - yellow > Alkali - blue > End - green
89
What is the colour of phenolphthalein in acids, alkali and the end of a titration?
> Acid - colourless > Alkali - pink > End - pale pink
90
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
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
What is the oxidation number of a uncombined element? (e.g.)
> 0 | > C, Na, O2, P4
92
What is the oxidation number of a combined oxygen? (e.g.)
> -2 | > H2O, CaO
93
What is the oxidation number of a combined oxygen in peroxides? (e.g.)
> -1 | > H2O2
94
What is the oxidation number of a combined hydrogen? (e.g.)
> +1 | > NH3, H2S
95
What is the oxidation number of combined hydrogen in metal hydrides? (e.g.)
> -1 | > LiH
96
What is the oxidation number of a simple ion? (e.g.)
> Charge on ion | > Na+: -1, Mg2+: -2
97
What is the oxidation number of combined fluorine? (e.g.)
> -1 | > CaF2, NaF
98
What are the oxidation number for all the element in sulfur dioxide?
* S: +4 * O: -2 * O: -2
99
What are the oxidation numbers for the elements in a carbonate ion?
* C: +4 * O: -2 * O: -2 * O: -2
100
What are the oxidation numbers for Fe in (II)FeCl2 and (III)FeCl2?
* (II) FeCl2: +2 | * (III) FeCl2: +3
101
What is the oxidation number for nitrogen in the nitrate ion?
• NO2^- : + 3
102
What is the oxidation number of sulfur in the sulfate ion?
• SO4^2-: +6
103
What is oxidation in terms of electrons?
LOSS of electron
104
What is Oxidation in terms of oxygen?
GAIN of oxygen
105
What is Oxidation in terms of oxidation numbers?
Increases in oxidation number
106
How can oxidation be shown in an equation?
Mg —> Mg^2+ + 2e^-
107
What is reduction in terms of electrons?
GAIN of electrons
108
What is reduction in terms of oxygen?
LOSS of oxygen
109
What is reduction in terms of oxidation number?
DECREASE in oxidation number
110
How can reduction be shown in an equation?
Cl2 + 2e^- —> 2Cl^-
111
What is a redox reaction?
A reaction where reduction and oxidation both takes place
112
What are the oxidation numbers? Which undergo reduction and oxidation? Mg + Cl2 —> MgCl2
> Mg: 0 —> +2: oxidation | > Cl: 0 —> -1: reduction
113
What are the oxidation numbers? Which undergo reduction and oxidation? MnO2 + 4HCl —> MnCl2 + 2H2O + Cl2
> Mn: +4 —> +2: reduction | > Cl: -1 —> 0: oxidation
114
Describe the development of the atom (9).
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
What is relative molecular mass, Mr?
> The addition of the relative atomic masses of each atom making up a molecule. > Used for simple molecules.
116
What is relative formula mass?
> The addition of the relative atomic masses of each atom making up a formula unit . > Used for compounds.
117
What is the charge and formula of a Zinc ion?
> 2+ | > Zn^2+
118
What is the charge and formula of a Silver ion?
> 3+ | > Ag^3+
119
What are the ionic charges of group 1-13 elements?
> Lose electrons | > Positive ions
120
What are the ionic charges of group 15-17 elements?
> Gain electrons | > Negative ions
121
What atoms do not normally form ions? Why?
> Be, B, C and Si | > Requires too much energy to transfer outer shell electrons

A Level Chemistry (7 decks)

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