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Foundations in Chemistry Flashcards

1
Q

What is the relative atomic mass of an electron?

A

1/1836

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2
Q

What is atomic number?

A

The number of protons in 1 atom of an element.

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3
Q

Give an equation for the charge on an ion.

A

number of protons - number of electrons

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4
Q

Define isotopes.

A

Atoms of the same element with different numbers of neutrons, so different masses.

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5
Q

What is relative isotopic mass?

A

The mass of one atom of an isotope compared to 1/12 the mass of one atom of Carbon 12.

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6
Q

What is relative atomic mass?

A

The weighted mean of the atoms of an element compared to 1/12 the mass of one atom of Carbon 12.

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7
Q

How is relative atomic mass calculated?

A

(Mass of isotope) x (relative abundance of isotope)…
________________________________________
Sum of relative abundances

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8
Q

What graph does a mass spectrometer produce?

A

Relative abundance against mass-to-charge ratio

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9
Q

What does the mass to charge ration represent?

A

The mass of an ion with +1 charge.

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10
Q

How can you interpret a mass spectrometer graph to get relative atomic mass?

A

Divide the length of each line by the sum of the lengths of all the lines (x100 for %)

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11
Q

Define ‘relative molecular mass’

A

The mean mass of one molecule of an element or compound compared to 1/12 the mass of Carbon 12

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12
Q

Define ‘relative formula mass’

A

The mean mass of one formula unit of an element or compound compared to 1/12 the mass of Carbon 12.

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13
Q

When is it better to use relative molecular mass, over Mr?

A

When dealing with ionic compounds.

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14
Q

Define ‘first ionisation energy’.

A

The energy required to remove the first electron from every atom in 1 mole of gaseous atoms.

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15
Q

What 3 things effect ionisation energy?

A

Atomic radius, number of protons in nucleus, and number of shielding electrons.

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16
Q

What is the general trend of IE across a period and why?

A

IE increases across a period, as the number of shielding electrons is constant, but the number of protons increases, and this draws electrons closer, slightly decreasing atomic radius. This means the attraction of electrons to the nucleus increases.

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17
Q

What are the 4 types of orbital?

A

s,p,d,f

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18
Q

What shape is an s orbital?

A

spherical

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19
Q

What shape is a p orbital?

A

dumbbell shaped

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20
Q

Describe the number of electrons and orbitals in the first 4 periods of the periodic table.

A

1: 1 s orbital, 2 electrons
2: 1 s orbital, 3 p orbitals, 8 electrons
3: 1 s orbital, 3 p orbitals, 8 electrons
4: 1 s orbital, 3 p orbitals, 5 d orbitals, 18 electrons.

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21
Q

What is a subshell?

A

The set of all the orbital of a particular type in the same valence shell.

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22
Q

What order do subshells fill with electrons.

A

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d

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23
Q

How do electrons fill orbitals of the same energy?

A

Empty orbitals are filled first.

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24
Q

Nitrate ion

A

NO3 -

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25
Carbonate ion
CO3 2-
26
Sulphate ion
SO4 2-
27
Hydroxide ion
OH -
28
Ammonium ion
NH4 +
29
Zinc ion
Zn 2+
30
Silver ion
Ag +
31
What is meant by 'amount of substance'.
Amount in moles.
32
What is a mole?
The mass of a substance containing the same number of fundamental units as there are atoms in exactly 12 g of 12C.
33
What is Avogadro's constant?
The number of particles in 1 mole, or the number of C12 atoms in exactly 12g of C12.
34
Define 'Molar mass'.
The mass of one mole of a substance.
35
What is molar gas volume?
The volume in dm3 of one mole of a gaseous substance at a given pressure and temperature- always 20dm3 at rtp.
36
What is empirical formula?
The simplest whole number ratio of each element present in a compound.
37
What is molecular formula?
The actual number and type of each element found in a molecule.
38
What is meant by the term 'anhydrous'?
Without water of crystallization.
39
What is water of crystallization?
Water that is trapped within the structure of a substance, but is not chemically bonded to that substance.
40
Give the molar mass equation.
``` Moles = Mass / Formula mass n = m / Mr ```
41
Give the solution volume and concentration equation.
``` Concentration = Moles / Volume c = n / v ```
42
Give the equation for gas volume at rtp.
``` Moles = Volume / 24dm3 n = v / 24 ```
43
Give the ideal gas equation, with units.
p- Pa V- m3 n- mols T- K R-constant pV=nRT pressure x volume = moles x temperature x R
44
What is a stoichiometric coefficient?
A number used before a species to balance an equation.
45
What is a stoichiometric ratio?
The ratio between amounts of different species in a balanced equation.
46
Give an equation for percentage yield.
(Actual yield / Theoretical yield) x 100
47
Give an equation for atom economy.
(Mr of desired products / Mr of total products) x 100
48
Hydrochloric acid
HCl
49
Sulphuric acid
H2SO4
50
Nitric acid
NO3
51
Ethanoic acid
CH3COOH
52
Phosphoric acid
H3PO4
53
Sodium hydroxide
NaOH
54
Potassium hydroxide
KOH
55
Ammonia
NH3
56
Define an acid
A proton donor that releases H+ ions in solution.
57
Define a base
A proton acceptor that can neutralise an acid.
58
Define an alkalis.
An aqueous base that releases OH- ions in solution.
59
What is a strong acid?
An acid in which a large majority of the H+ ions dissociate in water, eg. HCl
60
What is a weak acid?
An acid in which many of the H+ ions do not dissociate in water, creating an equilibrium.
61
What is a neutralisation reaction.
One in which OH- ions react with H+ ions to give water.
62
Describe the reactions of bases with acids.
metal + acid -> salt + hydrogen metal oxide + acid -> salt + water metal hydroxide + acid -> salt + water metal carbonate + acid -> salt + water + carbon dioxide
63
Describe the equipment used to make up a standard solution.
Volumetric flask, volumetric pipette, funnel, distilled water.
64
Describe the equipment used in a titration.
Burette, conical flask, volumetric pipette, indicator.
65
What is the oxidation state of a pure element, eg H2?
0
66
Describe the oxidation state of oxygen atoms.
-2 in all compounds except in peroxides (-1) and when bonded to fluorine (+2)
67
Describe the oxidation state of hydrogen atoms.
+1 in all compounds except metal hydrides (-1)
68
What do roman numerals in a compound's name mean?
They show the oxidation number of the positive ion in a compound, in situations where it may be ambiguous.
69
What is oxidation?
The loss of electrons, or the increase in oxidation state.
70
What is reduction?
The gain of electrons, or the decrease in oxidation state.
71
What is an orbital?
A region around the nucleus that can contain 2 electrons with opposite spin.
72
What order do we write electronic configuration in?
Numerical order.
73
What is an ionic bond?
An electrostatic attraction between positive and negative ions.
74
What is a giant ionic structure?
A regular lattice structure of positively and negatively charged ions, electrostatically attracted in all directions.
75
Describe the physical and structural properties of giant ionic structures.
Giant structures give them high boiling points as strong ionic bonds must be broken in order for the substance to change state. Polarity gives them solubility in polar solvents such as water, and also electrical conductivity when molten or aqueous as charged ions can flow past each other.
76
What is a covalent bond?
A strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonding atoms.
77
What is a dative covalent bond?
A bond in which both bonding electrons originate from the same atom.
78
What is average bond enthalpy?
The relative strength of a covalent bond.
79
Give the shape and bond angle and an example of a molecule with 4 bonding pairs and no lone pairs,
Tetrahedral- 109.5 degrees. | eg. CH4
80
Give the shape and bond angle and an example of a molecule with 1 lone pair and 3 bonding pairs
Pyramidal- 107 degrees | eg NH3
81
Give the shape and bond angle and an example of a molecule with 2 lone pairs and 3 bonding pairs.
Non-linear- 104.5 degrees | eg H2O
82
Describe the bonding of a trigonal planar molecule, and give the bond angle.
A molecule with 3 covalent bonds and no lone pairs, eg. BF3 | 120 degrees.
83
Describe the bonding of a linear molecule, and give the bond angle.
A molecule with 2 covalent bonds and no lone pairs, eg BeCl2 | 180 degrees
84
Describe the bonding of an octahedral molecule, and give the bond angle.
A molecule with 6 covalent bonds and no lone pairs, eg. SF6 | 90 degrees.
85
Describe electron pair repulsion theory.
Due to the negative charge on electrons, pairs of electrons repel other pairs. This means that they are positioned as far apart around a molecule as possible. However lone pairs repel more than bonding pairs, meaning bonding pairs are pushed further away from them, changing the bond angle around a molecule.
86
Define electronegativity.
A measure of the attractiveness of an atom's nucleus to the bonding pair of electrons in a covalent bond.
87
What is the periodic trend of electronegativity?
It increases towards group 7 and period 1.
88
What is a polar bond?
When an atom bonds covalently with another atom of a different electronegativity, the more electronegative atom will be more attractive to the bonding electrons, giving it a δ- charge, whilst the less electronegative atom carries a δ+ charge.
89
How can a polar bond cause a permanent dipole?
When a polar bond occurs in an asymmetrical molecule, this can cause one end of a molecule a δ charge, create permanent polarity throughout the molecule, provided the δ charges do not cancel out.
90
What is an instantaneous dipole?
Due to the random movement of electrons, molecules can sometimes become slightly polar if the electrons instantaneously gather on one side of the molecule causing momentary δ charges.
91
What are London forces?
Intermolecular forces caused when instantaneous dipoles induce charges in surrounding molecules and then form weak electrostatic charges. These intermolecular forces can appear in all simple covalent molecules.
92
What are permanent dipole-dipole forces?
Intermolecular forces caused by the electrostatic attraction between permanent dipoles in neighbouring molecules.
93
What is a hydrogen bond?
A strong electrostatic attraction between the δ+ charge on a hydrogen atom and and a lone pair of electrons on either, nitrogen (N), oxygen (O), or fluorine (F).
94
Why does ice float on water?
Ice is less dense than water because, when water freezes, a hydrogen bond forms from every lone pair and hydrogen atom (4 hydrogen bonds per water molecule). These hydrogen bonds force water molecules apart, creating a loosely packed lattice structure that is less dense than water.

Chemistry (13 decks)

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