Chapter 3: Electronic Structure Flashcards
What is electromagnetic spectrum
All the wavelengths of radiation in different forms of different energy energy
All waves travel at the same speed. How can they be distinguished
Wavelength
What is frequency
No of waves passing a point per second
Relate frequency and wavelength and energy
Shorter the wavelength the more the frequency and hence the more the energy
State wave equation
c=√f
velocity=wavelength×frequency
How do w, f and e vary with the spectrum
Decreasing from radio waves to gamma waves
Energy/freqeuncy increasing from radio to gamma
How is white light split into an atomic emmision spectrum
Wavelengths/frequencies are divided into colors by diffraction through a prism
What is line spectrum
White light split into different wavelengths/frequencies consisting of discrete lines
What happens when atoms are supplied with energy
Electrons gain energy move from lower(ground) to higher(excited state) energy levels
What happens as electrons gradually lose energy
Move from higher to lower energy levels and emit energy of particular wavelength
What prevents electrons from leaving the atom
Electrostatic force of attraction between oppositepy charged protons in nucleus and electrons in orbits
Why are neutrally charged neutrons important
Stability of elements with more than 1 proton, which would repel and nucleus falls apart
State the essentials of Bohr’s model
Electrons move in shells
For any atom there’s a fixed set of allowable orbits/stationary states
As long as electrons remain in their energy levels their energy is constant
An electron can only pass from one stationary state to another giving out/absorbing discrete/definite quantities of energy (quanta)
An electron moves up a level when energy is absorbed. Why does it eventually fall back and what is released
The excited electron is unstable
Electromagnetic radiation/packet of energy/photon
Photons of UV have more energy than of IR
Energy of photon emitted is =
Energy change in atom
Explain in depth the Planck equation (evidence of bohrs model)
Atoms emit photons of certain energies which give likes of certain frequencies as they can only occupy certain orbits
They can either be at this orbit or the other not in between meaning produce continuous energy changes.
Energy is changed only by discrete amounts aka quantized and line spectrum is quantized and not continuous
What happens when an electron is at it’s highest energy level
Atom is ionized
What is ionization energy
Energy needed to remove an electron from each atom in a mole of gaseous atoms
Why do lines with higher energies converge
Energy levels in the atoms are closer at higher energy
What foes quantum theory suggest
Electrons/particled having wave properties
What was wrong with Bohrs model
It applied well to hydrogen. Not other complex atoms. It assumes electrons trajectory can be precisely described. An electrons position cannot known by means of radiation as this will give it a kick and disturb the motion and send hurtling off in a random direction
What’s heiseinbergs uncertainty principle
We cannot know the position of an electron at a given moment of time
We can only predict where it’s likely to be ie spread in an electron cloud
What is an atomic orbital
Volume if space around nucleus in which there is maximum probability of finding a specified electron
The higher the energy of an electron the - it is from the nucleus. And - varies with energy
Further
Shape
Energy levels have
Principle quantum numbers
How many electrons in s sublevel
2
How many electrons in p sublevel
6
How many electrons in d sublevel
10
How many electrons in f sublevel
14
Expression for sub level maximum electrons?
2n^2
State the order of orbitals
1s, 2s, 2p, 3s, 3p, 4s, 3p, 4p, 5s, 4d, 5p, 6s so on
How is it that p sublevel onwards there are more than 2 electrons
There are 3 or more orbitals as in axis orientations
What is paulis exclusion principle
Each orbital can hold 2 electrons so long as they have opposite spin which reduces effect of repulsion . This cases a small change in overall energy
What is the Aufbau principle
ORBITALS ARE FILLED IN THE INCREASING ORDER OF THEIR ENERGIES
What is hunds rule
When orbitals of identical energy are available electrons occupy these singly rather than in pairs as far as possible to reduce mutual repulsion
Total energy of electron is determined by 4 factors:
Energy level
Sublevel
Slightly different energies of orbitals from Same principle levels due to magnetic field generated by moving electrons
Opposite spin
Why are the outer most electrons lost first?
They are least attracted by the effective nuclear charge of protons
What is the first ionization energy of an element
Energy needed to remove one electron from each atom in a mole of gaseous atoms to form one mole of gaseous singly charged ions (+)
What is the second ionization energy of an element
Energy needed to remove one electron from each ion in a mole of gaseous singly charged ions to form one mole of gaseous double charged ions (2+)
The successive ionization energy value _
Increases
Why is it that successive ionization energy increases?
As each electron is removed the remaining ion becomes more positively charged. Fewer outer electrons are being attracted by the same no of protons. Successive electrons are closer to the nucleus and of lower energy. A strong force of attraction has to be overcome to remove these electrons
What is the attraction between protons and electrons called
Electric field
Electrons furthest from the nucleus are of
Highest energy
Explain the 3 (+1) actors affecting ionization energy
•Size of nuclear charge
As no of protons increases as electrons are removed, +ve nuclear charge increases, and attractive force between nucleus and electrons increases so more (ionization) energy is needed to over come it.
•Atomic radius (distance from electrons to nucleus/half tge distance between nuclei of 2 closest atoms)
Size of atom depends on how strongly protons attract outermost electrons, and increases as shells increase.
Fα1/d^2 . Force of attraction between protons and electrons decreases as quantum shells/distance increase and the lower ionization energy
•Shielding effect
Electrons in filled inner shells repel those in outermost shells and reduce the effect of +ve nuclear charge. The net +ve charge outer electrons feel is called effective nuclear charge. The greater the shielding effect, the lower the ionization energy
(•Stability of electron configuration
More energy required to remove electrons from inert gases/full stable shell)
Explain the ionization energy trend down a group
Decreases.
Down a group (period no increases) and the outer shell moves further away from nucleus and shielding effect increases due to increase in inner shells. So attraction for outer shell decreases and so does energy
In the periodic table, group number denotes
Outer Shell electrons
In the periodic table, period number denotes
Number of shells
Explain the ionization energy trend across a period
Increases in general.
Effective nuclear charge increases as protons increase, but atomic radius stays the same as shell no remains constant. Attraction for outermost electron increases and energy also generally increases
Name some elements which are an exception to the general trend across a period that energy increases
For boron, oxygen, aluminium, silicon, sulfur it decreases. Because of sublevels (group increases but sublevels are further than orbits)
Give 2 exceptions to electron configuration that you know and why
Chromium 24
Should be: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d4 but is 3d5, 4s1
Copper 29
Should be: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2, 3d9 but is 3d10, 4s1
This is because eelectrons prefer to be single than paired so the available spaces in the 3d orbital are occupied by an electron from the 4s orbital to reduce repulsion and make it more stable while half filled
The easier it is to remove an outer electron the _ it’s energy level
Higher
And vice versa
Why is a logarithmic scale used
As electrons are lost the same protons hold other electrons more tightly hence increasing ionization energy
In a graph, what 4 things can be defined?
Successive energies show the number of electrons in each level
Levels contain different numbers of electrons before they are full
Energies for each electron/level
By seeing the first large jump the no of valence electrons can be determined and hence the group (first electron is easier to remove and is thus in energy level)
What are s, p d, f blocks?
Valence electrons of elements in these blocks are in these sublevels
Which elements are in s, p and d blocks
S- groups 1 and 2 (eg ns2)
P- groups 3,4,5,6,7,0 (except He which is in S) (eg ns2np6)
D- transition (eg ns2np6nd8)
How do graphs provide evidence for the existence of sublevels?
Graph is divided into subsections between noble gases, this corresponds to filling of one she’ll with electrons. Successive energies decrease even as groups increase (ie proton increases, same shell but energy is still less) this is due to order of sublevels that is, p is further than s in same energy level n
Define atomic, neutron, nucleon/mass number
Atomic : no of protons in the nucleus
Neutron : no of neutrons
Nucleon/mass : no of protons and neutrons in a nucleus
Write the 1st and 2nd ionization energy equations for an element X including state symbols
X(g)->X+(g) + e ΔHi1 ^(top) -O-
X+(g)->X++(g) + e ΔHi2 ^(top) -O-
-O- is standard conditions
