Atoms, Molecules and the Periodic Table Flashcards
Final Exam Content
1
Q
Atom definition
A
The smallest unique species of an element.
2
Q
Atomic scale name
A
Angstrom
3
Q
Angstrom value
A
1 Å = 10^-10m
4
Q
Prefixes from small things
A
m = 10^1
mm = 10^-3
μm = 10^-6
nm = 10^-9
pm = 10^-12
fm = 10^-15
5
Q
Prefixes for large things
A
m = 10^0
km = 10^3
Mm =10^6
Gm = 10^9
Tm = 10^12
Pm =10^15
6
Q
Dimensional analysis process
A
desired unit = current unit x conversion rate.
rearrange the conversion rate so you have a fraction equal to 1.
Substitute in conversion rate.
7
Q
Why is angstrom used in chemistry
A
The radius of an atom ranges on the Angstrom scale, which means that chemists get their own unique measuring scale.
8
Q
Why is dimensional analysis important
A
It makes transferring between different units easy.
9
Q
number of atoms in a substance formula
A
number of atoms = (mass of substance) / (mass of one atom)
10
Q
Atomic Mass Units (amu) to kg
A
1amu = 1.661 x 10^-27kg
11
Q
kg to Atomic Mass Units (amu)
A
1 kg =6.02 x 10^26 amu
12
Q
The Avogadro Number
A
L= 6.02 x 10^23amu
13
Q
The Avogadro Number definition
A
The number of species within a mole.
14
Q
Number of atoms in a substance whole calculation
A
mass of one atom in kg = RAM amu x 1.661 x 10^-27kg/amu.
# of atoms = mass of substance / mass of one atom.
15
Q
mass of a proton
A
1.0073 amu
16
Q
mass of a neutron
A
1.0087 amu
17
Q
charge of a proton
A
+1.6 x 10^-19C
relative charge = +1
18
Q
charge of an electron
A
-1.6x 10 ^-19C
relative charge = -1
19
Q
nuclide notation
A
A over Z then chemical symbol
A = mass number
Z = atomic number
20
Q
mass number
A
The mass of the protons + the mass of the neutrons.
mass number = 1.0073(#protons) + 1.0087(#neutrons)
21
Q
atomic number
A
The number of protons in the nucleus.
22
Q
RAM of an element calculation
A
RAM = (%abundance y) x (isotope y RAM) + (%abundance z) x (isotope z RAM)
23
Q
Isotope definition
A
An atom of the same element with the same atomic number but a different mass number.
24
Q
Nuclear binding energy definition
A
The energy required for the nucleus of one atom to be disassembled into its protons and neutrons.
25
Nuclear binding energy formula
ΔE = ΔMc^2
26
ΔM formula
theoretical mass - actual mass
27
nuclear binding energy whole calculation
Theoretical mass = protons mass + neutrons mass.
ΔM = theoretical mass - actual mass.
Dimensional analysis to convert ΔM from amu to kg.
ΔE =ΔMC^2
Multiply by Avogadro's Number for J per mol.
28
energy units
J - joules
1J = 1 (m^2) x kg x (s^-2)
29
electronic structure stability comparison
electronic structure is less stable than nuclear structure since electrons require less energy than protons or neutrons to cause structural changes.
30
Changes in electronic structure
Electron energy shell.
Location of electron within an atom.
Which atom an electron is bonded to.
31
Quantum particles examples
Electrons and Photons
32
Photons definition
Electromagnetic radiation which behaves with wave-particle duality.
33
Wave-particle duality
where quantum particles behave as both waves and particles at the same time because of how small they are.
34
Energy equation for a photon acting as a wave
E = hv
35
Planck's constant value
6.626 x 10^-34 Js
36
energy equation for a photon acting as a particle
E = Mc^2
37
speed of light
c = 3 x 10^8 ms^-1
38
ways of describing waves
wavelength
frequency
wavenumber
39
wavelength
The distance that one between to peaks on a wave.
40
frequency
The number of periods of a wave per second.
Measured in s^-1.
Symbol: v.
41
wavenumber
The number of periods which pass through one cm.
Measured in cm^-1
Symbol: a v with a dash over it.
42
wavenumber equation
V = 1/λ cm^-1
43
wave particle duality proof
E=Mc^2 but also E=hv
so Mc^2=hv
v= c/λ
so Mc^2 = hc/λ
Mc = h/λ
This is also true at other velocities.
Mv =h/λ
Mv = p
p =h/λ or λ=h/p
44
wavelength and frequency formula
c=vλ
v=c/λ
λ = c/v
45
momentum formula
p =h/λ
momentum = Planck's constant /wavelength
46
momentum definition
The mass times the velocity of a particle. (p = Mv)
47
Heisenberg's uncertainty principle
ΔxΔp≥h/4π
You don't need to know this just the rule that you can never know both the exact momentum and position of a quantum particle at the same time.
48
orbital definition
An orbital describes the area outside the nucleus where there is a high probability of finding an electron.
49
wave function rule for electrons
When electrons behave like a wave, there is an equal probability of finding an electron at any point with the same amplitude.
50
what is the use of quantum numbers
To identify and describe any electron within the same atom.
51
n quantum number
principal quantum number - refers to electron energy shell.
52
n values
1st shell =1
2nd shell =2 ect
53
L quantum number
Angular momentum quantum number - refers to the shape of orbital.
54
L values
L= n-1 maximum
s=0
p=1
d=2
f=3
55
mL quantum number
the magnetic quantum number - refers to orbital orientation/ the dimension that the orbital lies on.
56
mL values
mL= [-L,+L]
57
ms quantum number
spin magnetic quantum number - shows the spin direction of the electron.
58
ms values
-1/2 or +1/2
59
Rydberg equation
vn = -Rh/n^2
wavenumber = Rydberg's constant / principal (quantum number)
60
Rydberg's constant
Rh = 1.097 x 10^5 cm^-1
61
what does the Rydberg's equation show
The wavenumber, and hence the energy of an atomic orbital of hydrogen.
62
orbital stability rule relating to energy
as the energy becomes more negative / decreases the greater the stability of the orbital.
63
energy of hydrogen atomic orbitals rule
Orbitals with the same principal quantum number (in same energy shell) will have the same energy.
64
1s orbital structure
a sphere with no nodes
65
2s orbital structure
sphere - node sphere.
66
3s orbital structure
sphere - node - sphere node - sphere.
67
2p orbital structure
a dumbbell with a nodal plane passing through the origin.
68
3p orbital structure
a dumbbell with a nodal plane and nodal sphere
69
3d orbital structure
A 3D four leaf clover with 2 nodal planes passing through the origin.
70
nodes definition
An area around the nucleus where an electron cannot be found.
71
number of orbitals explanation
the number of orbitals is based on the range of the mL quantum number: [-L,+L].
72
s orbitals number of orbitals
1
since mL = 0
73
p orbitals number of orbitals
3
since mL = -1, 0, +1
pxy, pxz, pyz
74
d orbitals number of orbitals
5
since mL= -2,-1,0,+1,+2
dxy, dxz, dyz, dx^2-Y^2, dz^2.
75
f orbitals number of orbitals
7
since mL=-3,-2,-1,0,+1,+2,+3
76
Explain why other atomic orbitals energy differs from hydrogens atomic orbitals
other elements on the periodic table will have more than one electron in their atom, this will cause electron repulsion which will result in electrons with the same principle quantum number to no longer being degenerate, like in hydrogen.
77
why is the 2s orbital lower energy than the 2p orbital
The 2s orbital has a small sphere where an electron can be found which is extremely close to the nucleus, this means the electron will spend some time much closer to the nucleus in 2s than in 2p's dumbbell shape. This means the orbital penetration in 2s is greater than in 2p. Resulting in a lower energy in 2s than 2p.
78
order of energy of other elements atomic orbitals
1s, 2s,2p,3s,3p,4s,3d,4p,5s,4d.
if you need to go further than 4d just draw the diagram out.
79
orbital penetration definition
orbital penetration is how close an electron can get to the nucleus in a specific orbital.
80
orbital penetration and orbital energy rule
As the orbital penetration increases the orbital energy decreases, since the electron is closer to the nucleus.
81
Pauli exclusion principal
no two electrons within the same atom will have the same 4 quantum numbers since only to electrons can be in the same orbital and these electrons will have opposite spin.
82
Aufbau principal
Electrons fill orbitals in order of increasing energy, meaning the lowest energy orbital is filled first.
83
Hund's rule
For degenerate orbitals electrons will fill singularly first keeping their spins parallel.
84
exceptions to orbital filling rules
lanthanum will fill 5d before 4f, this is because these orbitals are so close in energy.
uranium does whatever the fuck it wants to.
85
what is the use of orbitals
orbitals can be used to calculate the probability of finding an electron at one exact position.
86
Periodic table trends
metallic character
ionisation energy
electronegativity
atomic radius
87
metallic character definition
The measure of how easily an atom of an element will lose an electron to become stable.
88
metallic character along a period
As you go along a period the number of protons in the nucleus will increase, this will increase the positive charge of the nucleus and therefore, increase the force of attraction between the valence electrons and the nucleus. This will mean that it is more difficult for an electron to be lost, and therefore the metallic character will be decreasing.
89
metallic character down a group
As you go down a group the number of outer electron shells will increase, this will mean the valence electrons will have more electron shielding, this will result in a lower force of attraction between the valence electrons and the nucleus. this will result in electrons being lost more easily, and therefore metallic character will be increasing.
90
Electron shielding explanation
when there is a high number of electron energy shells, the valence electrons will have the force of the attraction between themselves and the nucleus decreased.
91
Metals
Elements on the left hand side of the periodic table with high metallic character.
92
Non metals
Elements on the right hand side of the periodic table with low metallic character.
93
metalloids
elements in between the metals and non metals on the periodic table which act as semiconductors.
94
Atomic radius definition
The distance between the valence electrons and the nucleus, which is a measure of the size of an atom.
95
atomic radius along a period
As you go along a period the number of protons in the nucleus will increase, this will increase the positive charge of the nucleus and therefore, increase the force of attraction between the valence electrons and the nucleus. this will mean the valence electrons are pulled closer, so the atomic radius will be decreasing.
96
atomic radius down a group
As you go down a group the number of outer electron shells will increase, this will mean the valence electrons will have more electron shielding, this will result in a lower force of attraction between the valence electrons and the nucleus. This will mean that the distance between the nucleus and valence electrons will increase, increasing the atomic radius.
97
electron shielding equation
Z(eff) = z - δ
nucleus pull effected by shielding = nucleus pull - shielding constant.
98
Electronegativity definition
The measure of the force of attraction that an atom has for its electrons, or electrons of a bond.
99
Electronegativity along a period
As you go along a period the number of protons in the nucleus will increase, this will increase the positive charge of the nucleus and therefore, increase the force of attraction between the valence electrons and the nucleus. And this will mean the electronegativity is increasing.
100
Electronegativity down a group
As you go down a group the number of outer electron shells will increase, this will mean the valence electrons will have more electron shielding, this will result in a lower force of attraction between the valence electrons and the nucleus. Which will result in decreasing electronegativity.
101
positive ion name
cation
102
negative ion name
anion
103
why do atoms form ionic and covalent bonds
To allow the atom to achieve a full set of orbitals - which is the electronic structure with the highest stability.
104
ionic bond electronegativity
1.8 - 2.2
105
polar covalent bond electronegativity
1.2 -1.8
106
pure covalent bond electronegativity
less than 1.2 or equal to 1.2
107
polar covalent bond difference to pure covalent bond
In a polar covalent bond there one atom will pull the electrons much closer to the nucleus causing a dipole, whereas in a pure covalent bond they will share the electrons equally.
108
stable electron arrangements
2, 8, 8, 18, 18, 32
109
1st ionisation energy definition
The energy required for one electron to be lost from all the atoms in one mole of free atoms.
110
ionisation energy as you go along a period
As you go along a period the number of protons in the nucleus increases, this means the force of attraction between the nucleus and the valence electrons will increase, which will result in more energy being required for electrons to be lost, which will result in an increasing ionisation energy.
111
ionisation energy going down a group
As you go down a group the number of electron shells increases which will increase the electron shielding of the valence electrons. this will mean less energy is needed for electrons to be lost from the atoms, so the ionisation energy will be decreasing.
112
rules for ionisation energies
atoms must be in a gaseous state and ion formed must always be positive
113
electron affinity definition
The measure of how easily an electron is accepted by an atom.
114
Polarity of a molecule rule
A molecule will be polar if it contains polar covalent bonds and these polar bonds are unevenly distributed.
115
VSEPR stands for
Valence shell electron pair repulsion
116
VSEPR rule
electron pairs will always arrange themselves in the structure which produces the minimum possible repulsion between them.
117
order of repulsion for electron pairs
non bonding > bonding
118
bond order rule
as the bond order increases the energy of the bond increases and the bond length decreases. This is due to there being more electrons attracted by the positive nucleus.
119
bond order explanation
1st order = single bond
2nd order =double bond
3rd order = triple bond
120
hypervalent atoms
atoms which have more than 8 valent electrons in a compound.
121
EA 2 electron pairs
linear
121
electron deficient atoms
atoms which contain less than 8 valent electrons in a compound.
122
EA 3 electron pairs
trigonal planar
123
EA 4 electron pairs
tetrahedral
124
EA 5 electron pairs
trigonal bipyramidal
125
EA 6 electron pairs
octahedral
126
MS for 2 bonding electron pairs
linear
127
MS for 2 bonding electron pairs and 1 non bonding pair
angular
128
MS for 3 bonding electron pairs
trigonal planar
129
MS for 4 bonding electron pairs
tetrahedral
130
MS for 3 bonding and 1 non bonding electron pair
trigonal pyramidal
131
MS for 2 bonding and 2 non bonding electron pairs
angular
132
MS for 5 bonding electron pairs
trigonal bipyramidal
133
MS for 4 bonding and 1 non bonding electron pairs
distorted tetrahedron or see saw shape
134
MS for 3 bonding and 2 non bonding electron pairs
t shaped
trigonal planar
135
MS for 6 bonding electron pairs
octahedral
136
MS for 5 bonding electron pairs and 1 non bonding
square pyramidal
137
MS for 4 bonding and 2 non bonding electron pairs
square planar
138
Lewis structure process
Find # of electrons on molecule/ion
Preliminary sketch
Formal charge calculation
Rearrange bonds to fit formal charges.
139
number of valence electrons formula
number of electrons = valence electrons of each atom added up - charge.
140
Formal charge calculation
valence electrons - non bonding electrons -1/2 bonding electrons
141
what is the purpose of Lewis structure
to find the bonding structure of a molecule or ion.
