Textbook Ch.6: Electronic Structure and the Periodic Table Flashcards
1
Q
Wavelength (λ)
A
the distance between two consecutive crests or troughs, most often measured in meters or nanometres
1nm = 10⁻⁹ m
2
Q
Frequency (v)
A
the number of wave cycles (successive crests or troughs) that pass a given point in unit time.
3
Q
Hertz (Hz)
A
the frequency unit that represents one cycle per second
v = 10⁸m/s = 10⁸ Hz
4
Q
Speed of light in a vacuum (c)
A
2.998*10⁸ m/s
5
Q
How can you find the speed at which a wave moves through space?
A
λv = c
- λ should be expressed in meters
- v should be expressed in reciprocal seconds (hertz)
6
Q
Photons
A
a stream of particles that we consider light and has the energy E
7
Q
What equation gives the energy of photons?
A
E.= hv = hc/λ
8
Q
Joule (J)
A
an SI unit for energy
9
Q
Kilojoules
A
1kJ = 10³ J
10
Q
Planck’s equation / Planck’s constant
A
h = 6.62610⁻³⁴ Js
11
Q
Energy is ____ related to wavelength
A
inversely
12
Q
Bohr’s atomic model
A
assumed that a hydrogen atom consists of a central proton about which an electron moves in a circular orbit
13
Q
Bohr’s equation for the energy of the hydrogen electron:
A
E = -R / n^2
E = energy of the electron R = Rydberg constant (2.180 * 10⁻¹⁸ J) n = principal quantum number (depends on the state of the electron)
14
Q
Rydberg constant
A
2.180 * 10⁻¹⁸ J
15
Q
In Bohr’s model, where did he designate the zero energy point?
A
the point at which the proton and electron are completely separated
- Energy has to be absorbed to reach that point
- The electron must have an energy below zero
16
Q
Ground state / ground level
A
the hydrogen electron being in its lowest energy state
n=1
17
Q
Excited state
A
when the hydrogen electron absorbs enough energy and moves to a higher state
1st excited state: n = 2
2nd excited state: n = 3
and so on…
18
Q
When will an electron drop back to a lower energy state?
A
when if gives off energy as a photon of light
- It can return to ground state (from n=2 to n=1)
- It can go to a lower excited state (from n=3 to n=2)
19
Q
The energy of the photon (hv) evolved is equal to:
A
the difference in energy between the two states
∆E = hv = Eₕᵢ - Eₗₒ
20
Q
What equations can be used to relate the frequency of light emitted to the quantum numbers nₕᵢ and nₗₒ of the two states:
A
hv = -Rₕ - ( 1/(nₕᵢ)² - 1/(nₗₒ)² )
v = Rₕ / h( 1/(nₕᵢ)² - 1/(nₗₒ)² )
21
Q
Quantum mechanics
A
the science dealing with the behavior of matter and light on the atomic and subatomic scale
22
Q
How does the quantum mechanical atom differ from the Bohr model?
A
- The kinetic energy of an electron is inversely related to the volume of the region to which it is confined
- It is impossible to specify the precise position of an electron in an atom at a given instant
23
Q
𝚿
A
- the symbol for wave function
- for the hydrogen electron, 𝚿² is directly proportional to the probability of finding the electron at a particular point
24
Q
Electron cloud diagram
A
shows how 𝚿² for the hydrogen in its ground state (n=1) varies moving out from the nucleus
25
Orbital
a more common way of showing electron distribution in the hydrogen atom
26
Quantum number
```
a value that is used when describing the levels available to atoms and molecules. Associated with the solutions to the wave
function 𝚿
```
27
What are the quantum numbers?
n, l , mₗ, mₛ
28
First quantum number (n)
- of primary importance when determining the energy of an electron
- as n increases, energy increases, and the farther the electron is found from the nucleus
- Can only take integral values starting with 1 (1, 2, 3, 4...)
29
First principal level
n=1
30
Second principal level
n=2
31
Second Quantum Number (l)
- the sublevels that each principal energy level has is denoted by l
- the general shape of the electron cloud is associated with l --> larger l values produce more complex shapes
- l and n are related because l can be any integral value starting with 0 --> l=n-1
32
Sublevels (s, p, d, f)
- In the nth principal level, there are n different sublevels
- letters can be assigned (s,p,d, or f) can be used to indicate l=0, 1, 2, or 3
33
If l = 0
| What is the sublevel?
sublevel s
34
If l=1
| What is the sublevel?
sublevel p
35
If l=2
| What is the sublevel?
sublevel d
36
If l=3
| What is the sublevel?
sublevel f
37
If n=1
l = ?
Possible sublevels?
If n=1
l can only have one possible value, 0
1s sublevel
38
If n=2
l = ?
Possible sublevels?
If n=2
l has two possible values, 0 and 1
2s and 2p sublevels
39
If n=3
l = ?
Possible sublevels?
If n=3
l has three possible values 0, 1, and 2
3s, 3p, 3d sublevels
40
If n=4
l = ?
Possible sublevels?
l has 4 possible values, 0, 1, 2, and 3
| 4s, 4p, 4d, and 4f sublevels
41
Third Quantum Number, mₗ; Orbitals
- m relates to l
- It determines the direction in space of the electron cloud surrounding the nucleus
- m can have any integral value, including 0, between l and -l
42
In the s sublevel (l=0):
What are the possible values of mₗ?
How many orbitals are there?
mₗ = 0
| 1 orbital
43
In the p sublevel (l=1):
What are the possible values of mₗ?
How many orbitals are there?
mₗ = 1, 0, -1
| 3 orbitals
44
In the d sublevel (l=2):
What are the possible values of mₗ?
How many orbitals are there?
mₗ = 2, 1, 0, -1, -2
| 5 orbitals
45
In the f sublevel (l=3):
What are the possible values of mₗ?
How many orbitals are there?
mₗ = 3, 2, 1, 0, -1, -2, -3
| 7 orbitals
46
Fourth Quantum Number mₛ; Electron spin
- mₛ is associated with electron spin (clockwise or counterclockwise)
- NOT associated with n, l, of mₗ
- mₛ has two possible values: +1/2 or -1/2
47
Opposed spins
when electrons have different values of mₛ
48
Parallel spins
when electrons have the same values of mₛ
49
Pauli exclusion principle
no two electrons in an atom have the same set of four quantum numbers
50
What is the capacity of the s sublevel?
2 electrons
51
What is the capacity of the p sublevel?
6 electrons
52
What is the capacity of the d sublevel?
10 electrons
53
What is the capacity of the f sublevel?
14 electrons
54
What are some characteristics of the shape of the s orbital?
- All s sublevels are spherical, they differ from one another ONLY by size
- As n increases, the radius of the orbital increases
- The electron in a 2s orbital is more likely to be found far out from the nucleus than is a 1 s electron
55
What are some characteristics of the shape of the p orbital?
- Consists of two lobes along an axis (x, y, or z) --> there is a zero probability of finding an electron at the nucleus at the atom
- The three p orbitals in a given sublevel are oriented at right angles to one another along the x, y, and z axis --> are designated
pₓ, pᵧ, and pz
56
Electron configuration
the simplest way to describe the electron arrangement in atoms
57
What does the electron configuration of: 1s²2s²2p⁵ mean?
- 2 electrons in the 1s sublevel
- 2 electrons in the 2s sublevel
- 5 electrons in the 2p sublevel
58
How are electron configurations obtained?
- Electron configuration is obtained easily if the order of filling sublevels is known
- Electrons enter sublevels in order of increasing sublevel energy
- A sublevel is filled to capacity before another sublevel starts to fill
59
What is the order of filling sublevels?
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 4d, 4f
60
How do you find the electron configuration of a neutral atom?
1. Find the number of electrons (atomic number)
| 2. Fill the sublevels in order until they are full, fill the last sublevel with the leftover electrons
61
Abbreviated electron configuration
electron configurations that are shortened to save space and start with preceding noble gas
Ex: Sulfur = 1s²2s²2p⁶3s²3p⁴
- Neon can be substituted for the first 10 electrons
- -> [Ne]3s²3p
62
The atoms of elements in a group of the periodic table have:
the same distribution of electrons in the outermost principal energy level
63
Elements in Groups 1 and 2 are filling an ___ sublevel
s sublevel
64
Elements in Groups 13-18 (6 elements in each period) fill a ___ sublevel
p sublevel
65
The transition metals fill the ___ sublevel
d sublevel
66
The lanthanides (the two rows of 14 elements that are listed at the bottom) fill the ___ sublevel
f sublevel
67
Orbital diagrams
used to show further how electrons are distributed among orbitals
- Represented by parenthesis ( ) and ↑ ↓ arrows depending on the spin
68
How do you determine the orbital diagrams of atoms?
- A pair of electrons must have opposite spins (+1/2, -1/2 or ↑↓)
- Distribute the arrows by the orbitals
69
Hund's rule
when several orbitals of equal energy are available, as in a given sublevel, electrons enter singly with parallel spins
- Only after all orbitals in a sublevel are half-filled do electrons pair up in orbitals
- In all filled orbitals, the two electrons have opposite spins
- With a given sublevel there are as many half-filled orbitals as possible
70
When a monatomic ion is formed from an atom, how are electrons added or removed?
electrons added or removed from sublevels in the highest principal energy level
71
Isoelectronic
ions that have the same electron configuration due to them losing or gaining electrons
72
Halide
a chemical compound that contains a halogen (Group 17 elements)
73
Halides of alkali metals have the general formula:
MX
- M= Li, Na, K...
- X=F, Cl, Br
74
Halides of alkaline earth metals have the general formula:
MX₂
- M= Mg, Ca, Sr...
- X=F, Cl, Br
75
Oxides of alkaline earth metals have the general formula:
MO
| - M=Mg, Ca, Sr...
76
In forming cations, electrons are removed from the sublevel of highest ___
n
77
When transition metal atoms for positive ions, the outer __ electrons are lost first
s
78
"First in first out" rule
electrons come out in the same order of sublevels they come in (ex: 4s before 3d)
79
Periodic law
the chemical and physical properties of elements are a periodic function of atomic number
80
Atomic radius
can be defined and measured, assuming a spherical atom.
Taken to be 1/2 the distance of the closest approach between
atoms in an elemental substance
81
What is the periodic trend for atomic radius?
Atomic radii
- Decrease across a period (from left to right in the periodic table)
- Increase down a group in the periodic table
82
Positive ions are ____ than the metal atoms from which they are formed.
smaller
83
Negative ions are ____ than the nonmetal atoms from which they are formed.
larger
84
Ionization energy
a measure of how difficult it is to remove an electron from a gaseous atom
- Energy must always be absorbed to bring about ionization
- Ionization energies are always positive quantities
- The more difficult it is to remove electrons, the higher the ionization energy
85
What is the periodic trend for ionization energy?
- Increases moving across the periodic table from left to right
- Decreases moving down a group in the periodic table
86
Electronegativity
measure the ability of an atom to attract itself to the electron pair forming a covalent bond
- The greater the electronegativity of an atom, the greater its attraction for electrons
87
What is the periodic trend for electronegativity?
- Increases moving across the periodic table from left to right
- Decreases moving down a group in the periodic table
