Module 2 Foundations in Chemistry Flashcards
1
Q
State the location, charge and mass of a proton
A
Located in nucleus. Charge is +1, mass is 1
2
Q
State the location, charge and mass of a neutron
A
Located in nucleus. Charge is 0, mass is 1
3
Q
State the location, charge and mass of an electron
A
Located in electron shells. Charge is -1, mass is 1/1840
4
Q
What does the atomic number show?
A
Number of protons in the nucleus of an atom
5
Q
What does the mass number show?
A
Number of protons and neutrons in the nucleus of an atom
6
Q
Define the term ‘isotope’
A
Atoms of the same element with the same number of protons (and electrons) but differing numbers of neutrons in the nucleus. Therefore isotopes of the same element have the same atomic number but different mass numbers.
7
Q
Why do isotopes of the same element have the same chemical behaviour as one another?
A
They have the same electron configuration so will react the same as one another
8
Q
Define the term ‘relative atomic mass’
A
The weighted mean mass of one atom of an element compared to 1/12 of the mass of an atom of carbon-12
9
Q
Define the term ‘relative molecular mass’ and state when it should be used.
A
The weighted mean mass of one molecule compared to 1/12 of the mass of an atom of carbon-12. Used when referring to a covalent compound.
10
Q
Define the term ‘relative formula mass’ and state when it should be used.
A
The weighted mean mass of one formula unit compared to 1/12 of the mass of an atom of carbon-12. Used when referring to an ionic compound.
11
Q
Define the term ‘relative isotopic mass’
A
The mass of one atom of an isotope compared to 1/12 of the mass of an atom of carbon-12.
12
Q
Define the term ‘strong acid’
A
A strong acid is a proton donor that fully dissociates into solution
13
Q
Define the term ‘weak acid’
A
A weak acid is a proton donor that only partially dissociates into solution. Some hydrogen ions will remain attached to the acid molecule
14
Q
Define the term ‘base’
A
A base is a proton accepter
15
Q
Define the term ‘alkali’
A
A water soluble base
16
Q
What is formed in the reaction between an acid and an alkali?
A
Salt and water
17
Q
What is a salt?
A
A salt is the product of a reaction in which the hydrogen ions of the acid are replaced by metal or ammonium ions
18
Q
Why is the mole used?
A
To measure the amount of substance without having to record a very large number of significant figures.
19
Q
What is Avogadro’s number?
A
The number of atoms in 12.000g of carbon-12
20
Q
What is the molar mass?
A
The mass, in grams, of one mole of a substance
21
Q
What are the units of molar mass?
A
grams per mole (g mol-1)
22
Q
How do you find a substance’s molar mass?
A
Look on the periodic table for its relative atomic/ molecular mass, then add correct units of grams per mol.
23
Q
What is the calculation linking moles, mass and molar mass?
A
number of moles (mol)= mass (g) / molar mass (g/mol)
24
Q
What is the calculation linking moles, number of particles and Avogadro’s constant?
A
number of moles (mol)= number of particles / Avogadro’s constant
25
One mole of any gaseous substance will occupy ______dm³ at room temperature and pressure
24
26
How do you convert between dm³ and cm³ ?
1000cm³ =1dm³
27
What is the equation linking gas volume and number of moles?
gas volume (dm³ )= moles (mol) x 24
28
Define 'concentration'
Concentration is a measure of how many moles there are of a substance in a particular volume of an aqueous solution
29
What are the units of concentration?
mol dm⁻³
30
What is a monobasic acid?
An acid that donates 1 hydrogen ion per molecule of acid
31
How do you calculate the pH of a strong monobasic acid?
-log (concentration)= pH
32
What is a standard solution?
A solution of known concentration
33
Why are standard solutions used in chemistry?
By reacting other solutions of unknown concentration with the standard solution, we can work out the unknown concentration e.g. in titration
34
Describe how to weigh out an exact mass of solid during the preparation of a standard solution (weighing by difference).
1. Using a spatula, weigh out approximately 2.5g of solid into a weighing boat.
2. Record the exact mass of solid and weighing boat using a balance.
3. Transfer the solid to a beaker, and wash the weighing boat with distilled water into the beaker to ensure that all the solid is dissolved
4. Re-weigh and record the mass of the used weighing boat.
5. You can now work out the exact mass of solid used by taking the mass of the boat away from the total mass of boat and solid.
35
Describe how to make a standard solution after finding the exact mass of solid used (weighed by difference)
1. Transfer the solid into an 100cm³ beaker. Wash the weighing boat with distilled water into the beaker to ensure all of the solid is transferred
2. Dissolve all of the solid in distilled water from a wash bottle. Stir with a glass stirring rod until all of the solid is dissolved.
3. Transfer this solution into a 250cm³ volumetric flask through a funnel. Rinse the beaker and stirring rod with more distilled water and transfer the washings to the volumetric flask to ensure all of the solid is transferred.
4. Add distilled water up until you have almost reached the 250cm³ graduation line. Then remove the funnel and add distilled water with a pipette until the bottom of the meniscus is touching the graduation line.
5. Stopper the flask and gently invert a few times so that the concentration is the same throughout.
6. Calculate the concentration of your solution using mass/Mr to find moles, then use moles/ volume to find concentration.
36
Define the term 'ionic bond'
The electrostatic attraction between oppositely charged ions
37
How are ionic bonds formed?
When a metal transfers its valence electrons to a non-metal.
38
Define the term 'covalent bond'.
The STRONG electrostatic attraction between a shared pair of electrons and the nuclei of the two bonded atoms
39
How are ionic compounds drawn?
Dot and cross diagrams with brackets showing the charge drawn around ions.
40
How are covalent compounds drawn?
Dot and cross diagrams with dots and crosses showing where electrons are shared.
41
Covalent molecules form when two ___-___ ____ electrons in covalent bonds.
non-metals, share
42
What are oxidation and reduction?
oxidation- loss of electrons
reduction- gain of electrons
43
What is an oxidation state?
a + or - number that tells us the number of electrons an element has donated or accepted during a chemical reaction.
44
When is the oxidation state of hydrogen NOT +1?
When hydrogen is bonded to a metal
45
When is the oxidation state of oxygen NOT 2-?
When oxygen is bonded to fluorine or is part of a peroxide
46
When do halogens have an oxidation state other than 1-?
When chlorine, bromine and iodine are bonded to oxygen
47
Describe how you calculate the relative atomic mass shown in the periodic table?
Ar= (mass x percentage)+ (..)... all divided by 100
48
Describe how a mass spectrometer is used
- a sample of an element is placed into the mass spectrometer
- the mass spectrometer determines what % of the sample is a certain isotope
- an average is then calculated.
49
What does m/z represent?
The mass charge ration. For mass spectrometry, it is the same as the mass number.
50
The ____ model of the atom suggested that electrons are distributed into ___. The shells are numbered, with shell 1 being _____ to the nucleus. The closer to the nucleus a shell is, the ____ its energy.
Bohr, shells, closest, lower
51
What is the maximum number of electrons in each shell 1-4?
1- 2 electrons
2- 8 electrons
3- 18 electrons
4- 32 electrons
52
Each shell is divided into ______. The shells closest to the nucleus have _____ subshells than those further away. The subshells are labelled _,_,_ and _.
subshells, fewer, s,p,d and f
53
How many electrons can each subshell hold?
s- 2 electrons
p- 6 electrons
d- 10 electrons
f- 14 electrons
54
Finally, the subshells are split into ____. No matter where these are they can only hold a maximum of __ electrons.
orbitals, 2
55
How many orbitals are in each subshell?
s- 1 orbital
p- 3 orbitals
d- 5 orbitals
f- 7 orbitals
56
Electron configurations can be shown on an _________. The diagram fills up with electrons from ___ to ___.
energy level diagram, low, high
57
What does Hund's Rule state regarding an energy level diagram?
Don't pair electrons in an orbital unless you have to.
58
What do the arrows show in an energy level diagram?
Shows that electrons must have opposite spin in order to go in the same orbital
59
What is the maximum pattern of the shells? (strontium)
1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶ 5s²
60
Why is the 4s subshell written before the 3d subshell?
4s comes first as it has a lower energy than the 3d subshell. Therefore electrons are put into and removed from this shell before the 3d.
61
What is the shape of an s orbital?
sphere shaped
62
What is the shape of a p orbital?
dumbell shaped- lying on the x, y or z axis
63
Define 'first ionisation energy'.
The minimum amount of energy required to remove one mole of electrons from one mole of gaseous atoms.
64
Write an equation showing the first ionisation energy of Magnesium
Mg (g) → Mg ⁺(g) + e
65
Describe the trend in ionisation energy down a group.
First ionisation energy decreases
66
Explain the trend in ionisation energy down a group.
As you move down a group, atomic radius increases. This results in shielding increasing- due to more inner shells. The shielding means that the nuclear attraction on the outer valence electrons decreases. This means an electron can more easily be removed- which decreases first ionisation energy.
67
Describe the trend in ionisation energy across a period.
First ionisation energy increases
68
Explain the trend in ionisation energy across a period.
Across a period, atomic radius decreases. Shielding remains constant across a period, but nuclear charge increases. This means that nuclear attraction on the outer electrons also increases. This means that it is harder to remove a valence electron- increasing first ionisation energy.
69
As you travel across a period, why does atomic radius decrease?
As you go across a period, the number of shells remains the same. However the number of protons increases, making the nucleus more positively charged. Therefore, across a period, the same number of electron shells are being attracted by a greater number of protons. This means that electrons orbit closer to the nucleus, decreasing atomic radius.
70
Why does aluminium have a lower first ionisation energy than magnesium, despite being further along the same period?
In magnesium, the valence electrons are both in the 3s subshell. However in aluminium, the valence electron is in the 3p subshell. This means that there is less nuclear attraction on the the 3p subshell than the 3s subshell. Therefore, the 3p electron in aluminium requires less energy to be removed than the 3s electron in magnesium.
71
Why does sulfur have a lower first ionisation energy than phosphorus, despite phosphorus being in group 5 and sulfur being in group 6?
In both phosphorus and sulfur, the highest energy valence electrons are found in a 3p subshell. However, in sulfur, there is an electron paired in an orbital that is not found in phosphorus. The fact that two electrons are in the same orbital causes extra repulsion. This extra repulsion means that the valence electron of sulfur can be removed more easily than the valence electron of phosphorus.
72
Why is there a general increase in energy required with each successive ionisation energy?
With each electron removed, there are now fewer electrons being attracted by the original number of protons. Therefore, the nuclear attraction on the remaining electrons increases - making it more difficult to remove a valence electron. This means that with each successive ionisation energy, more energy is required to remove an electron- increasing ionisation energy.
73
Explain the 'big jumps' in ionisation energy shown on a successive ionisation energy graph.
The big jumps indicate large increases in ionisation energy. This occurs when the next electron is being removed from a more inner shell, closer to the nucleus. This means that the next electron is more attracted to the nucleus, and requires a large increase in ionisation energy to be removed.
74
Define the term 'molar mass'.
the mass of one mole of a substance.
75
What are the units of molar mass?
g mol⁻¹
76
How many grams are in one tonne?
10⁶ grams
77
How many milligrams are in one gram?
1000mg
78
How many grams are in one kilogram?
1000g
79
How do you convert between mol dm⁻³ and g dm⁻³
Multiply by Mr. Mol dm⁻³ x Mr = g dm ⁻³
80
What is the equation linking concentration, moles and volume, including units?
concentration (mol dm⁻³)= moles (mol)/ volume (dm³)
81
What is the equation linking moles to gas volume? Include units.
number of moles (mol)= volume (dm³) / 24
82
What is the molar gas volume?
Molar gas volume is 24dm³. It is the volume that one mole of a gas occupies at room pressure and temperature.
83
What is room temperature and pressure?
100kPa and 298K
84
What is the ideal gas equation?
pressure x volume = moles x gas constant x temperature
85
What are the units of the ideal gas equation?
pressure- Pa
volume- m³
moles- mol
gas constant- 8.314J K⁻¹
temperature- K
86
How many cm³ are in a m³?
1,000,000
87
How many dm³ are in a m³?
1000
88
Groups 1 and 2 are in the __ block of the periodic table.
s
89
Groups 3-0 are in the __ block of the periodic table.
p
90
Transition metals are in the __ block of the periodic table.
d
91
Lanthanoids and actanoids are in the __ block of the periodic table.
f
92
What is the equation for percentage yield?
percentage yield= (actual mass/ theoretical mass) x 100
93
Define the term actual mass
The mass that the person performing the experiment has actually physically achieved.
94
Define the term theoretical mass.
The mass that the experiment could have theoretically achieved, had the experiment been 100% perfect.
95
What is the equation for atom economy?
Atom economy= (Mr of desired/ total Mr of products) x 100
96
Define 'empirical formula.'
The simplest whole number ratio of atoms of each element in one molecule or one formula unit of a compound.
97
Define 'molecular formula.'
The actual whole number ratio of atoms of each element in the molecule/ formula unit of a compound.
98
What can be unusual about the empirical formula?
It may not be a molecule that can really exist e.g. hydrogen peroxide HO
99
How do you calculate the empirical formula?
1. work out the moles of each element using either given masses or percentages
2. divide by smallest number of moles
3. work out the whole number ratio
100
How do you find a molecular formula from an empirical formula, given the molecular formula's mass?
1. Calculate the molar mass of the empirical formula
2. Divide the molecular molar mass by the empirical molar mass
3. Multiply empirical molar mass by the value found
101
Define 'anhydrous'.
A crystalline compound that does not contain any water molecules
102
Define 'hydrated'.
A crystalline compound that contains water molecules
103
Define 'waters of crystallisation'.
Water molecules that are bonded into the crystalline structure of a hydrated compound.
104
Explain how a water molecule can form a hydrated cation.
Water is a non-linear molecule. The oxygen end of the water molecule is more electronegative . This means that water can interact with the cation, so that the oxygen ends form an attraction with the cation. The hydrogen ends of the water molecule will point away
105
Explain how a water molecule forms a hydrated anion.
Water is a non-linear molecule. The hydrogen ends of the water molecule have a low electron density and are less electronegative. Water can therefore interact with an anion, so that the hydrogen ends form an attraction with the anion. The oxygen ends will point away.
106
Describe heating to a constant mass.
Continuously heating and recording mass, then reheating, until a constant mass is recorded.
107
Define the terms acid, base and alkali.
Acids are proton donors; bases are proton acceptors; alkalis are soluble bases that contain hydroxide ions
108
What is an inorganic acid?
An acid that does not contain carbon
109
What is produced in the reaction between a metal oxide and an acid?
aqueous salt and water
110
What can be observed in the reaction between a metal oxide and acid?
solid metal oxide disappears
111
What is produced in the reaction between an acid and a metal carbonate?
aqueous salt + water + carbon dioxide
112
What will be observed in the reaction between acids and metal carbonates?
bubbles/ effervescence and the solid carbonate disappearing
113
What is produced in the reaction between a acid and a hydroxide?
aqueous salt and water
114
What is observed in the reaction between an acid and a hydroxide and why?
No observations because hydroxides are alkalis and therefore soluble, and so dissolve into the water. There are no observations unless indicators are added.
115
What is the reaction for an acid and ammonia?
ammonia + acid = ammonium salt
116
What is the reaction between a metal and an acid?
metal + acid --> salt + hydrogen
