Topic 1 Atomic structure and periodic table Flashcards
What does an atom contain?
It contains protons and neutrons in the nucleus and electrons in orbits around.
The nucleus is where most of the mass of an atom is and it is very small. It contains protons and neutrons
Electrons orbit the nucleus in shells and take most of the space of an atom
What is the mass number and what is the atomic number?
Mass number - Tells us the number of protons and neutrons in the nucleus. (bigger number)
Atomic number - Tells us the number of protons in the nucleus (smaller number)
since all atoms are neutral - Number of protons = number of electrons
What are ions?
They have different number of electrons and protons
Negative ions have gained electrons to gain a full shell of electrons. O2-
Positive ions have lost electrons to gain a full shell of electrons. Na+
What are isotopes?
They are elements with the same number of protons but a different number of neutrons
What is relative atomic mass and the relative isotopic mass
Relative atomic mass - It is the weighted mean mass of an atom of an element, compared to 1/12th of the mass of an atom of carbon-12
Relative isotopic mass - It is the mass of an atom of an isotope, compared to 1/12th of the mass of an atom of carbon-12
What is m/z?
It is just the mass of an isotope divided by the charge
How to calc relative atomic mass?
(Abundance of A x mass/charge ratio) + (Abundance of B x mass/charge ratio) / 100
How to predict mass spectra?
- Write the %’s as decimals
- Create a table showing the isotope combinations in a molecule. Multiply the decimal form of abundance of each isotope to get the relative abundance of each molecule
- Any molecules which are the same add the abundances up
- Divide all the relative abundances worked out before by the smallest value. This will give you a whole number ratio which can be used to predict your spectra
What do the peaks show on a mass spectra molecule version?
Peaks show fragments of the original molecule. The last peak is the M+1 peak or the molecular ion peak. This is the same as the relative molecular mass of the molecule
What subshells are electron shells split into? and how many orbitals do they have and how many electrons can they hold?
S P D F
S shell - Contains 1 orbital which holds 2 electrons
P shell - Contains 3 orbitals and can hold 6 electrons
D shell - Contains 5 orbitals and can hold 10 electrons
F shell - Contains 7 orbitals and can hold 14 electrons
What is the shape of the s orbital?
The s orbital is spherical and the 2 electrons can move anywhere within this sphere
What is the shape of the P orbital?
There are 3 p orbitals in the shape of dumbbells and can hold up to 2 electrons. They are 90 degrees to each other.
What is the order of electronic configuration?
1s2 2s2 2p6 3s2 3p6 4s2 3d10
You fill orbitals singly and then we pair up. This is due to electron repulsion.
Electron configuration of ions
For ions, you just add or remove from the highest energy level first
Ca2+ = would lose 2 electrons from the 4s
Electron configuration of transition metals
Chromium and copper behave differently
An electron from the 4s orbital moves into the 3d orbital to create a more stable half-full or full 3d sub shell
Same thing with iron. Fe3+ loses 3 electrons, 2 from the 4s and 1 from the 3d
The electron blocks
S block is on the left
D block is in the middle
P block is on the right
F block is at the bottom
What does the EM spectrum do?
The EM spectrum shows types of radiation at different frequencies`
“Raging Martians Invaded Venus Using X-ray Guns,”
Radio waves
Micro waves
Infra red
Visible light
Ultra Violet
X-rays
Gamma rays
As you go down, energy increases and so does the frequency of the radiation
As you go up, the wavelength of radiation increases
Atoms can release energy named in the EM spectrum
What is line spectra?
It shows the frequency of light given out when an electron moves down energy levels. We see them as coloured bands.
Every element has a different electron configuration and so will absorb and emit different frequencies of radiation. The emission spectra is unique to different elements
Explaining quantum shells
n=1 is known as the ground state and is the shell closest to the nucleus
The atom has shells known as quantum shells or energy levels
When an electron absorbs energy, it moves up to a higher quantum shell. The electron is excited.
Eventually, the electron will move back down and release energy.
How line spectra work?
The line in an emission spectrum shows the electrons moving to different energy levels. We call a group of lines a series. Series of lines are created when electrons move to the same energy level from different ones
The arrows from the shells going down to the ground state show a potential path an excited electron could follow.
Electrons that fall to ground state n=1 produce a series of lines in the UV part of the spectrum
Electrons that fall to the second energy level n=2 produces a series of lines in the visible part of the spectrum
Electrons that fall to the third energy level n=3 produces a series of lines in the infra-red part of the spectrum
We call a group of lines a series
Evidence for quantum shells?
The radiation emitted will have a fixed frequency as the energy of shells is fixed.
Electrons can only exist in quantum shells. They can’t be between shells
Each shell has a fixed energy
EM radiation is absorbed to move electrons to higher energy shells and emitted when they drop to lower ones
The defined lines proves electrons exist in shells only. They can never exist between them
What is ionisation energy?
It is the minimum amount of energy required to remove 1 mole of electrons from 1 mole of atoms in the gaseous state.
Requires energy so always endothermic with a + value
Na —> Na+ + e-
What is sheilding?
The more electron shells between the positive nucleus and negative electron that is being removed, the less energy is required. There is a weaker attraction.
What is atomic size?
The bigger the atom, the further away the outer electrons are from the nucleus. The attractive force between the nucleus and the outer electron reduces which makes it easier to remove electrons
What is nuclear charge?
The more protons in the nucleus, the bigger the attraction between the nucleus and the outer electrons. This means more energy is required to remove the electron
What are the 1st IE trends in groups
IE decreases as you go down a group.
This is as the atomic radius increases as you go down a group. Outer electrons are further away from the nucleus. Attractive forces are weaker which means less energy is required to remove an electron
Shielding increases as we go down the group. More shells between nucleus and outer shell which means the attractive force is weaker. The energy required to remove an electron decreases
What is successive IE?
This is the removal of more than 1 electron from the same atom
There is a general increase in energy as removing an electron from an increasingly more + ion needs more energy. So removing an electron from Mg+ will require more energy than removing something from Mg
You see big jumps in energy when you change the shell you are removing electrons from.
Removing from shells that are closer to the nucleus is much more difficult due to the attractive forces
Trend in atomic radius across a period
As you go across a period, the atomic radius decreases.
There is an increased nuclear charge as there is an increasing number of protons. This pulls the outer shell of electrons further in towards the nucleus
The extra electrons that elements gain across the period go into the same shell therefore shielding effect is similar.
Trend in atomic radius down a group?
It increases down a group due to extra electron shells added for each element down the group
What are the 1st IE trends across periods?
It increases as we go across a period.
As we go across, there is an increasing number of protons in the nucleus. This increases the nuclear attraction
Shielding is similar and distance from nucleus marginally decreases
More energy is required to remove an outer electron so IE increases
A decrease in sulfur?
It is evidence for electron repulsion in an orbital. Removing an electron from sulfur involves taking it from an orbital with 2 same charge electrons in it.
However, removing it from phosphorus is harder as it has no paired arrows
A decrease in aluminium?
It is evidence for atoms having sub-shells.
The outermost electron in aluminium sits in a higher energy sub-shell slightly further from the nucleus rather than the outer electron in magnesium. This makes it easier to remove
MP across a period?
Na Mg, and Al are metals so they have metallic bonding
Mg is higher than Na as it has 2+ ions reacting which means more electrons are being delocalised. Bigger electrostatic attraction so the MP increases
A general increase in MP as metal ions have an increasing + charge, increasing number of delocalised electrons and a smaller ionic radius. This means a stronger metallic bond
Silicon has the highest MP in period 3 as it has a giant covalent structure
Many strong covalent bonds hold the silicon atoms together. A large amount of energy is needed to overcome these strong covalent bonds
Phosphorus has a lower MP than Silicon due to a weaker simple molecular structure. The melting point is determined by weaker London forces
Sulfur has a higher MP than phosphorous due to a larger simple molecular structure (s8). It has larger London forces and a higher MP
Chlorine has a lower MP than phosphorous and sulfur due to a smaller simple molecular structure (cl2). It has smaller London forces and so a lower MP
