1.2. The nuclear atom Flashcards
Dalton’s model of the atom
John Dalton noticed that Hydrogen and Oxygen always combine in fixed proportions
He explained this by…
- All matter is composed by indivisible atoms
- Atoms cannot be created nor destroyed
- Atoms of the same element are alike in every way
- Atoms of different elements are different
- Atoms can combine together in small numbers to form molecules
This explains reacting ratios.
Plum-pudding model
J.J. Thomson believed negatively charged electrons were placed in a positively charged sponge-like substance since the atom had no net charge.
FALSE
Rutherfords experiment - Conclusions about the atom
Fired alpha particles at a piece of gold foil.
Most particles passed straight through. A very small number was repelled.
- Large no. of undeflected atoms = Atom is mostly empty space.
- Only a few bounce back = Postive nucleus is very small
- Electrons have a low mass = Deflected atoms are just deflected slightly
Subatomic particles
Why do they have relative charges / masses.
Mass is so small it is given relative to the mass of 1 proton / 1 neutron.
Masses are found in the data booklet.
What are protons and neutrons made up of?
Quarks
What is the name of the opposing particle which all particles have?
Anti-particles
What is the electrons antiparticle?
Positron.
It has the same mass but an equal and oppisite (positive) charge.
What happens when particles and anti-particles collide?
They destroy each other and release energy in the form of high.energy photons called gamma rays.
Bohr model
Electrons in orbit around a positive nucleus.
Most volume in the atom is empty space.
Changes from dalton, to thomson, to rutherford, to bohr are called paradigm shifts. Scientific theory evolves.
Atomic symbol
AZX
A - Mass number
Z - Atomic number
X - Symbol of element
Cation
Positive ion
Anion
Negative ion
Isotopes
Atoms of the same element with different numbers of neutrons / different mass numbers.
They have the same chemical properties (same no. of electrons) but different physical properties –> eg. boiling + melting point.
Relative atomic mass
77.5% Cl-35, 22.5% Cl-37
Average mass of an element dependent on the abundance of isotopes found in nature compared to one atom of C-12 (12.000).
(0.77535)+(0.22537) = 35.45
Radioisotopes
Stability of a nuclues depends on balance between the no. of protons + neutrons.
A nucleus containing too many / too few neutrons to be stable = radioactive.
It changes to be stable by giving out radiation.
How does a mass spectrometer work?
- The tested element is vaporised (allows individual atoms to be analysed –> no bonds.)
- The atoms are ionised by high-energy electrons which knock out an electron to create a cation.
- The cations are attracted to a negatviely charged plate + deflected by the magnetic field at right angles to its path.
- The amount the cation deflects is then measured.
Amount of deflection is inversely proportional to their mass/charge ratio.
If the ions have a singular charge, deflection is equal to mass.
Fragmentation pattern
I don’t really understand the mass spectrometer but can’t be arsed to research it now.
Mass spectrum
Graph representing results of mass spectrometer.
y-axis = % abundance
x-axis = mass/charge
Electromagnetic radiation
Decreasing wavelength
Radio waves
Microwaves
Infrared
Visible light
Ultraviolet
X-rays
Gamma rays
Continous spectrum.
Spectrum of visible light - Colours merge into each other.
Absorption spectra
Continuous spectra with discrete black lines.
Show wavelengths of light which have been absorbed by an electron during excitation.
Emission line spectrum
Black spectrum with discrete lines showing wavelengths of photons emitted by an electron that relaxes.
Energy change of electron = Energy of photon
What is the energy of a photon proportional to?
The frequency
E = hf
E = hc/wavelength
E is inversely proportional to wavelength + proportional to frequency.
Calculating the energy between quantum levels
Energy change of electron = Energy of photon
Energy of photon = hf
h - plancks constant
f - frequency
What does it mean that the energy of an atom is quantised?
Electrons can only change their energy by discrete amounts/ quanta of energy.
Ground state
The first energy level of an atom. It has quantum number 1.
n = 1
First ionisation energy
Second ionisation energy
The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions under standard conditions.
Energy needed to excite an electron past the convergence limit (it becomes a free electron). Once removed the electron is in the ‘n = ∞’ energy level-.
Same definition but energy required to remove two electrons…
Process of excitation and relaxation of an electron.
- An electron in the ground-state becomes excited (eg. by passing a current through the atom → Electrical discharge)
- This causes the electron to move to a higher energy level
- It stays in this excited state for a fraction of second. It then falls down to a lower level as it relaxes
- When doing this it emits a photon (discrete energy)
- The photon has a specific wavelength, depending on the energy difference between the two energy levels.
Do we need to know lyman balmer and paschen.
What is the convergence limit and why does it exist?
Energy levels get closer together (converge) at higher energies)
At the limit of this convergence, the lines merge, forming a continuum.
Beyond this continuum the electron can have any energy; it is no longer under the influence of the nucleus and is therefore outside the atom.
–> Delocalised.
Wave vs particle model of light
How does this relate to electrons?
Light is a wave since it can be described by its frequency.
Light is made of particles called photons.
Both models are needed to explain the characteristics of light.
Relates to electrons as they also behave like waves sometimes. (Wave-particle duality)
Why can the electrons trajectory not be precisely described?
- Any attempt to locate an electron will disturb its motion.
- eg. focusing radiation on an electron to locate it causes it to divert from its path.
Heisenbergs uncertainty principle - We can’t know where an electron exists at any given moment in time, but we can make a probability picture for where it is likely to be.
Schrödinger model of the Hydrogen atom
Atomic orbital
A region around an atomic nucleus with a 90% probability of finding an electron.
The shape of orbitals depends on the energy of electrons.
Electron in higher energy orbital = higher probability of being foudn further from the nucleus
Sublevels in energy levels. Each atomic orbital can contain 2 electrons of opposite spin.
What is the spin of electrons?
Electrons spin on their axis, either in a clockwise (upwards arrow), or anti-clockwise direction (downwards arrow).
Electrons with opposite spin can exist in the same area despite their repulsion.
Pauli exclusion principle
An orbital can only hold two electrons of opposite spin.
–> Opposing spin means the electrons don’t repel each other.
1st energy level - Sublevels
1 atomic orbital (1s2)
2 electrons
s orbital is shaped like spere.
Second energy level - Atomic orbitals
(2 sublevels, s and p)
1 * 2s atomic orbital (2s2)
+
3 * 2p atomic orbitals (2p6)
8 electrons
S orbital is shaped like sphere (2s is larger than 1s + electron is likely to be further out).
P orbital is shaped like dumbells (along y, x and z axis)
Third energy level - Atomic orbitals
(3 sublevels - s, p and d)
1 * 3s aotmic orbital (3s2)
+
3 * 3p atomic orbitals (3p6)
+
5 * 3d atomic orbitals (3d10)
18 electrons
Four of the 3d atomic orbitals are made up of 4 lobes centred on the y-axis, x-axis, z-axis.
The fifth 3d atomic orbital looks like a pool floatie, with two nodes sandwiching it on the y-axis.
Fourth energy level - Atomic orbitals
Four sublevels
1 * 4s atomic orbital (4s2)
+
3 * 4p atomic orbitals (4p6)
+
5 * 4d atomic orbitals (4d10)
+
7 * 4f atomic orbitals (4f14)
32 electrons
1st energy level - Elements
Hydrogen, Helium
1s²
1st period
2nd energy level - Elements
Lithium, Beryllium, Boron, Carbon, Nitrogen, Oxygen, Flourine, Neon
2s² 2p⁶
2nd period
3rd energy level - Elements
4th energy level - Elements
Number of electrons within an energy level
2n2
Determining the level of energy that orbitals have. ???
- ## Depends on atomic number
Degenerate orbitals
Electron orbitals having the same energy levels.
Eg. 2s, 2p
Energy level vs sublevel vs orbital
3 p orbitals form a p sublevel.
A p sublevel and an s sublevel form the second energy level.
Aufbau principle
Electrons fill atomic orbitals of the lowest available energy level before occupying higher-energy levels.
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p
This is essentially the order of the periodic table.
Ground state
The lowest possible energy level that the atom can occupy.
This is the energy state that would be considered normal for the atom.
Hund’s rule
If more than one orbital in a sublevel is available, electrons occupy different orbitals with parallel spins.
So in carbon (1s22s22p2)…
There are three possible 2p orbitals, instead of having two electrons of opposite spin in the first 2p orbital, there is one electron in the first 2p orbital, and one electron with parallel spin in the second 2p orbital.
This is done because it minimises repulsion between electrons. –> less overlap between sublevels.
Parallel spins because it leads to lower energy.
Relative energy of orbitals
??
Valence electrons
Electrons in the outer energy level. Responsible for compound formation.
4s vs 3d
Total energy level for n=3 is lower than total energy level for n=4.
HOWEVER…..
Total energy level for the orbital 4s is lower than total energy level for 3d. Therefore 4s gets filled before 3d.
HOWEVER….
Filling all of 3d10 and one electron in 4s is more stable than the alternative.
Filling all of 3d5 and one electron in 4s is more stable than the alternative.
Chromium - Condensed electron configuration
1s22s22p63s23p63d54s1
[Ar]3d54s1
This is done because it is more stable and 3d and 4s are similar in energy levels.
Which two elements are exceptions to the aufbau principle?
Chromium and copper
Copper - Condensed electron configuration
1s22s22p63s23p63d104s1
[Ar]3d104s1
This is done because it is more stable and 3d and 4s are similar in energy levels.
How do atoms lose electron?
/
How are cations formed?
Atoms always lose electrons from the outer sub-level first (highest quantum number).
Eg. all transition elements can form 2+ ions since 2 electrons can be removed from 4s2
S-block
First two elements of a period eg. Li, Be
P-block
Last 6 elements in a period.
eg. Al, Si, P, S, Cl, Ar
D-block
Transition element block
10 elements across
F-block
Lanthanides and actinides
Factors impacting ionisation energy
- Nuclear charge (attraction of the nucleus to the electron, no. of + protons - Z)
- As nuclear charge increases across a period, the ionisation energy increases (greater nuclear attraction)
- Electron shielding (electron-electron repulsion between energy levels)
- Full energy levels ‘shield’ valence electrons form the nuclear attraction of the nucleus.
- This decreases IE
- Electron spin repulsion is when an electron is added to an atomic orbital in a half-filled sublevel.
- O = 4 electrons in 2p (hund’s rule) –> 2 electrons in one orbital
- Technically opposite electron spins in one orbital allow them to coexist, but they still repel each other slightly → This makes it unstable → Easier to lose an electron → Lower ionisation energy.)
- This makes it easier to lose electrons –> IE decreases.
- Atomic radius inc. = smaller attraction between nucleus + valency electrons –> IE dec.
Trend in ionisation energy going across a period. eg. period 2.
Increases.
Nuclear charge increases –> IE inc.
Energy shielding stays constant (1 energy shield)
Explain the decrease in ionization energy between Be (Z = 4) and B (Z = 5).
Be - 1s22s2
B - 1s22s22p1
- Easier to remove an electron from 2p sublevel –> Further from nucleus.
2p is higher energy level
Explain the decrease in ionisation energy required between Z = 7 and Z = 8.
N = 1s22s22p3
O = 1s22s22p4
Electron-Electron repulsion.
Nitrogen has 1 electron in each orbital.
In oxygen there are two electrons in the first orbital (electron-electron repulsion) –> Less stable than oxygen. –> Easier to lose it –> Dec. IE.
Explain the decrease in ionisation energy going down a group in the periodic table.
Going down a group, means electron shielding is added each period.
This decreases the nuclear attraction (attraction between valence electrons and the nucleus).
This decreases IE.
Successive ionisation energy
First ionisation energy = Na(g)</sub) –> Na+</sub>(g)</sub) + e(-)</sup)
Second ionisation energy = Na+</sub>(g)</sub) –> Na2+</sub>(g)</sub) + 2e(-)</sup)
etc…
Explaining the successive ionisation energy for Na
Na = 1s2</sub>2s2</sub>2p6</sub>3s1</sub>
First electron is removed at a 500 kJmol-1</sub>
Second electron is removed at 5000 (much higher bcecause the energy shield between 2p6</sub>3s1</sub> has been removed.
