2.1 Atoms and reactions Flashcards
(119 cards)
1
Q
What are the symbols for atomic number and mass number?
A
Z and A respectively
2
Q
What is the Ar of a proton, neutron, and electron?
A
1, 1, and 1/1836 respectively
3
Q
Define isotopes
A
Atoms of the same element with the same number of protons but different numbers of neutrons
4
Q
What do chemical properties depend on?
A
Electronic structure
5
Q
How do isotopes of the same element vary?
A
In physical properties (e.g., density)
6
Q
Define cation
A
Positively charged ion
7
Q
Define anion
A
Negatively charged ion
8
Q
Define relative isotopic mass
A
The mass of an isotope compared to 1/12 of the mass of one carbon-12 atom
9
Q
Define relative atomic mass
A
The weighted mean mass of one atom compared to 1/12 of the mass of one carbon-12 atom
10
Q
What does a mass spectrometer measure?
A
Mass to charge ratio (m/z)
11
Q
What is the formula of a zinc ion?
A
Zn2+
12
Q
What is the formula of an ammonium ion?
A
NH4+
13
Q
What is the formula of a silver ion?
A
Ag+
14
Q
What is the formula of a nitrate ion?
A
NO3-
15
Q
What is the formula of a nitride ion?
A
N3-
16
Q
What is the formula of a carbonate ion?
A
CO32-
17
Q
What is the formula of a sulfate ion?
A
SO42-
18
Q
What is the formula of a phosphate ion?
A
PO43-
19
Q
What does a binary compound contain?
A
ONLY two different elements
20
Q
What is the mnemonic for remembering the diatomic elements?
A
Have No Fear Of Ice Cold Beer
21
Q
What are the steps for writing ionic equations?
A
- Balance the equation
- Break each (aq) substance into ions
- Cancel any ions on both sides (these being the spectator ions)
22
Q
What are spectator ions?
A
Ions that aren’t changing state or oxidation number
23
Q
Define one mole
A
The amount of substance which contains the same amount of particles as there are atoms in 12 grams of carbon-12
24
Q
Define molar mass
A
The mass in grams of 1 mol of a substance given in gmol-1
25
What is the equation linking mol (n), mass (m), and molar mass (Mr)?
n=m / Mr
26
What is the equation linking mol (n), volume of gas (dm3), and molar gas volume (24 dm3)?
moles =volume of gas / molar gas volume (24)
27
Define molecular formula
The ACTUAL number of atoms of EACH ELEMENT in an element/compound
28
Define empirical formula
The SIMPLEST ratio of atoms of EACH ELEMENT in an element/compound
29
In what contexts is 'relative molecular mass' and 'relative formula mass' used?
* Relative molecular mass - used for easily defined structures in which you know the exact no. of atoms (e.g., O2)
* Relative formula mass - used for giant structures in which you don’t know the exact no. of atoms (e.g., lattices such as NaCl)
30
What is a hydrated/hydrous salt?
A salt molecule loosely attached to water (called waters of crystallization)
31
What is a solute?
What is being dissolved into the solvent
32
What is a solvent?
The liquid in which a solute is dissolved in to form a solution
33
What is 1 L equal to (in volume)?
1dm3
34
What is the ideal gas equation and what is each unit measured in?
pV = nRT
Where p is pressure (Pa), V is volume (m3), n is moles (mol), R is the gas constant, T is temperature (K)
35
What is the conversion factor between Kelvin and Celsius?
Add 273 to C to get K
36
What is a limiting reagent?
A reactant that runs out first and thus limits how much product can be formed
37
Give 2 reasons for the percentage yield being less than 100%
* Reaction may have not had enough time to go to completion
* A side reaction has taken place (e.g., in the glassware)
* Product was lost under purification
38
What is a salt?
The product of a neutralisation reaction where the H+ IONS from the acid is replaced by metal or ammonium IONS
39
What are standard solutions?
A solution for which its concentration is known accurately
40
What 3 things should you ensure when making standard solutions?
* Ensure you have clean dry equipment
* Make washings to remove any salt BEFORE MAKING up the solution to the specified amount
* Invert the volumetric flask to ensure everything is dissolved
41
How does ammonia react with acid?
ammonia + acid → ammonium salt
NH3(aq)+HNO3 to (aq)NH4 + NO3(aq)
42
Give 3 strong acids and 1 weak acid
* Strong acids include: HCl, H2SO4, and HNO3
* Weak acids include carboxylic acids such as CH3COOH
43
Give 1 weak base
NH3
44
What are diprotic and triprotic acids (with examples) and what does this mean for neutralisation?
* Diprotic acids donate two protons (e.g., H2SO4)
* Triprotic acids donate three protons (e.g., H3PO4)
* They need 2x or 3x the number of moles of base to be neutralised
45
Give 4 bases/types of bases
* Metal oxides
* Metal hydroxides
* Metal carbonates
* Ammonia
46
What are bases and alkalis?
* Bases are H+ ion acceptors
* Alkalis are soluble bases that release HO- ions into solution
47
What is atom economy and how is it calculated?
atom economy=Mr of desired products / Mr of reactants
A measure of how efficient a reaction is
48
What are the 2 types of reactions under atom economy?
* Addition - the reactants form a single DESIRED product (100%)
* Substitution - the atoms from a reactant are substituted by others leading to more than one product
## Footnote
Addition reactions are usually where a reactant is added to an unsaturated molecule to make saturated molecule.
49
Give 3 reasons for why a high atom economy is good
* Environmental friendly (due to less waste)
* Economical (due to not separating materials)
* Sustainable (as it prolongs finite materials)
50
Give 2 ways of improving atom economy
* Use an alternate reaction pathway with a higher atom economy
* Find a use for the waste products
51
What is oxidation number?
The no. of electrons removed from an atom
52
What are the 4 main rules for oxidation numbers?
* Elements always have an oxidation no. of 0
* Sum of oxidation no.’s in a compound = 0
* Sum of oxidation no.’s of an ion = its charge
* The most electronegative element in a compound is always negative
53
What are the 3 exceptions to oxidation rules?
* Hydrogen is always +1 except when in a hydride where it’s -1
* Oxygen is always -2 except with fluorine and as peroxides where it’s -1
* Chlorine is always -1 except when bonded with oxygen and fluorine
54
What is a hydride?
A binary compound of hydrogen and a metal
55
What is a peroxide?
A compound containing two oxygen atoms bonded together (O-O)
56
What is an -ide?
A binary compound in which the nonmetal is given an -ide suffix (e.g., sodium oxide)
57
What are oxidising and reducing agents?
* Oxidising agents gain electrons leading to oxidation elsewhere
* Reducing agents lose electrons leading to reduction elsewhere
58
What is a reduction reaction?
The loss of oxygen atoms or gaining of electrons
59
What is an oxidation reaction?
The opposite of reduction
60
What do -ate compounds contain?
Oxygen and at least one other element
61
What is the end point in a titration?
The point where the indicator changes colour
62
What do most noble gases have an outermost electron structure of?
ns2np6
63
Which shells have which subshells?
* 1 has s
* 2 has s p
* 3 has s p d
* 4 has s p d f
64
What are orbitals?
Regions around the nucleus that can hold 2 electrons with opposite spins
65
What are degenerate orbitals?
Orbitals that have the same energy levels
66
How many orbitals does each subshell have?
* s has 1 orbital
* p has 3 orbitals
* d has 5 orbitals
* f has 7 orbitals
67
How do atoms fill up with electrons?
In subshells, in order of increasing energy
68
What is the shape of an s-orbital/s-subshell?
Sphere
69
What are the shapes of the p-orbitals?
Dumbbells
70
What are the 2 rules for all orbitals of the same energy?
* Electrons won't pair in the same subshell if they don't have to
* Electrons in the same orbitals must have opposite spin
71
Which key orbital fills first and empties first, when, and why?
The s-orbital, past and including the 4th shell; it has less energy than the d-orbital
72
What is a σ-bond?
The overlap of orbitals directly between atoms
73
Define ionic bond
The electrostatic force of attraction between oppositely charged ions formed by electron transfer
74
Describe and explain the conductivity of ionic compounds under different states
* Doesn't conduct when solid because the ions are fixed in a lattice
* Conducts when molten or dissolved because the ions can move
75
Why are ionic compounds soluble?
As water molecules are polar solvents so will surround each ion
76
Why are some ionic compounds less soluble than others?
Greater charge difference meaning stronger attraction which is less easily overcome
77
Define covalent bond
The strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
78
Give an example of a covalent molecule whose central atom has fewer than 8 electrons in its outermost shell
BF3
79
What must an atom have access to for expansion of the octet to occur?
Its d subshell meaning the atom must have 3 or more shells
80
Give an example of a molecule which expands the octet
PF5
81
How is PF5 formed (electron-wise)?
By freeing up 5 orbitals by moving electrons into unoccupied d subshell orbitals
82
What is a coordinate/dative bond?
A shared pair of electrons where both electrons are from the same atom
83
What is used to show a coordinate bond?
An arrow from the atom providing the coordinate bond
84
Give the 3 bond angles largest to smallest
* The bond angle between lone pairs
* The bond angle between a lone pair and a bonding pair
* The bond angle between bonding pairs
85
What are the 3 wedges for shape of molecules?
* Solid line is a bond on the plane of the page
* Solid wedge is a bond that comes out of the plane of the page
* Dotted line is a bond going into the plane of the page
86
What do lone pairs repel by and why?
An extra 2.5° as they’re closer to the nucleus and thus take up more space
87
Describe the linear shape of a molecule with an example
2 bonding pairs; 180 degrees; an example is CO2
88
Describe the trigonal planar shape of a molecule with an example
3 bonding pairs and 0 lone pairs; 120 degrees; an example is BF3
89
Describe the tetrahedral shape of a molecule with an example
4 bonding pairs; 109.5 degrees; an example is NH4
90
Describe the pyramidal shape of a molecule
3 bonding pairs and 1 lone pair; 107 degrees
91
Describe the nonlinear shape of a molecule
2 bonding pairs and 2 lone pairs; 104.5 degrees between bonding pairs
92
Describe the octahedral shape of a molecule with an example
6 bonding pairs; 90 degrees; an example is SF6
93
How should you tackle an ‘explain the shape of a molecule’ question in 4 steps?
* State the no. of BONDING PAIRS and lone pairs ‘surrounding the central atom’
* State that ‘electrons repel and try to get as far apart as possible’
* If there are lone pairs, state lone pairs repel more than BONDING PAIRS
* State the shape of molecule with the bond angle
94
Why are chemists able to predict the shapes of molecules?
* As electron pairs REPEL
* The shape is determined by the no. of bonding pairs and the no. of lone pairs
95
Define electronegativity
The ability of an atom to attract electrons in a covalent bond towards itself
96
When can you tell a bond is polar?
Generally, if the two atoms are of different elements, it’s polar
97
How can a molecule be nonpolar yet contain polar bonds?
As it’s symmetrical so dipoles cancel each other out
98
How can you work out if a molecule is symmetrical?
Generally, if it has no lone pairs and the same bond pairs then it’s symmetrical
99
What are permanent dipole-dipole interactions?
Electrostatic forces of attraction between polar molecules
100
Which will have a higher boiling point, HF or HCl, and why?
HF due to stronger permanent dipole-dipole interactions
101
What are ‘London Forces’ also called?
Induced-dipole-dipole interactions
102
Does every structure have London Forces? If so, which ones don't?
No. Lattices like SiO2
103
Describe induced dipole-dipole interactions
Unequal distribution of electrons leads to temporary dipole which induces a dipole in a nearby molecule
104
Why do the strength of induced dipole-dipole interactions increase down groups?
The number of electrons increases leading to stronger dipoles and stronger attraction
105
Does every structure have London Forces? If so which ones don't?
No. Lattices like SiO2
## Footnote
SiO2 does not exhibit London Forces due to its strong covalent bonding in a lattice structure.
106
Describe induced dipole-dipole interactions.
Unequal distribution of electrons ⇒ temporary/instantaneous dipole ⇒ induces a dipole in a nearby molecule leading to attraction.
107
Why do the strength of induced dipole-dipole interactions increase down groups?
The no. of electrons increases ⇒ stronger dipoles ⇒ stronger attraction.
108
What is a hydrogen bond?
The electrostatic attraction between a hydrogen atom bonded to an electronegative atom and a lone pair on an electronegative atom of a different molecule.
109
How is a hydrogen bond represented?
By a dashed line parallel to the covalent bond from the hydrogen.
110
Which electronegative elements will hydrogen bonding happen with and why?
Oxygen, nitrogen, and fluorine ∵ they’re the most electronegative.
111
Give 3 reasons why water has a higher m.p. and b.p. than structurally similar compounds.
Hydrogen bonding provides stronger intermolecular forces.
* Forms up to four hydrogen bonds.
* Oxygen is the second most electronegative.
112
Describe and explain the 2 anomalous properties of ice.
Ice is less dense than water due to hydrogen bonds in an open lattice.
* Ice has a higher m.p. than expected due to strong intermolecular forces.
113
What causes high surface tension in water and why?
Hydrogen bonding leading to a strong and flexible lattice structure.
114
What can polar molecules and nonpolar molecules dissolve into?
Polar solvents and nonpolar solvents respectively.
115
Why doesn’t water and oil mix?
Water is polar and oil is nonpolar; they cannot form stable interactions.
116
What are the strongest to the weakest intermolecular forces?
Hydrogen bonds, permanent dipole-dipole interactions, induced dipole-dipole interactions.
117
Why is a p-block element considered a p-block element?
As its highest energy electron occupies a p-orbital.
118
Why won’t all CO2 be released from a reaction involving a solution?
As CO2 is slightly soluble.
119
How does the structure of elements change across period 3?
Na, Mg, Al - giant metallic lattices.
* S - giant covalent lattice.
* P4, S8, Cl2, Ar - simple molecular.
