Atomic Structure Flashcards
Calculate the maximum number of electrons in a given shell
We can calculate the maximum number of electrons that each shell can hold by using the equation 2n² (where n is the principal quantum number of the shell).
Describe what is meant by an atomic orbital and a subshell
An atomic orbital is a region around the nucleus that can hold up to 2 electrons with opposite spins.
-An electron is considered to be a cloud of negative charge and the negative charge cloud has the shape of the orbital occupied by the electron.
-Scientists can never be certain of the exact location of an electron so the atomic orbital shows us a 95% probability of where an electron will exist.
A subshell is all of the all of the orbitals of the same type in the same shell (like the 1s subshell).
Describe the shapes of the S and P orbitals
S orbitals are shaped like a sphere and P orbitals are shaped like a dumbbell, every shell contains 1 S orbital and 3 P orbitals.
Describe the electron configuration of Chromium and Copper
Both chromium and copper only fill the 4s subshell with 1 electron and the other electron goes to the 3d subshell to make it more stable.
Explain what is meant by first ionisation energy and by successive ionisation energies
The 1st ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of atoms in their gaseous state to form 1 mole of + ions (also in gaseous state).
The 2nd ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of + ions in their gaseous state to form 1 mole of 2+ ions (also in gaseous state).
Successive ionisation energy is when we continue removing electrons and measuring the ionisation energy each time.
Describe the factors that affect ionisation energy
The electrons in an atom are attracted to the positive protons in the nucleus. The greater the attraction between the outer electrons and the nucleus, the greater the ionisation energy.
3 factors that affect ionisation energy:
-The distance between the nucleus and the outermost electrons (atomic radius). As the atomic radius increases, the force of attraction between the positive nucleus and the outer electrons decreases.
-The charge on the nucleus (electrons are attracted to the positively charged protons in the nucleus). The greater the number of protons, the greater the force of attraction between the outer electrons and the nucleus.
-Shielding (when electrons in the outer shell are repelled by electrons in inner shells). This shielding effect reduces the attraction between the outer electrons and the nucleus.
Explain what successive ionisation energies tell us about how electrons are arranged in atoms
Each time we remove an outer electron from the outer shell, the remaining electrons are pulled slightly closer to the nucleus which means there is a greater attraction between the outer electrons and the nucleus causing the ionisation energy to gradually increase.
When we remove electrons from a new inner shell there is a massive increase in ionisation energy because the inner shell is closer to the nucleus and electrons in the inner shell experience much less shielding so these electrons have a greater attraction to the nucleus.
Identify an element from ionisation energy data
Work out the number of electrons in the outer shell by counting all the electrons that gradually increase in IE and stop counting once there is a massive increase in IE.
Describe and explain how the first ionisation energy varies down a group
The 1st ionisation energy decreases as we go down a group. This is due to 2 factors:
-Firstly, moving down a group, the atomic radius of the elements increase. This means that the outer electron shell is further away from the nucleus.
-Secondly, going down the group the number of internal energy levels also increase which means there is more shielding between the nucleus and the outer electrons.
Describe and explain how the first ionisation energy varies across a period
The 1st IE tends to increase as we move across a period.
As we move across a period, the nuclear charge increases as the number of protons increases. This increases the attraction between the nucleus and the electrons so the atomic radius decreases across a period. Both the increased nuclear charge and atomic radius mean that the outer electrons are more attracted to the nucleus and this causes the 1st IE to increase across a period.
In all of the elements in a period, we are removing an electron from the same electron shell as the period number. This means that the shielding effect due to the inner electron shell is similar for each element which explains the overall increase in the 1st IE.
Exceptions to the pattern in 1st IE across a period
Along period 2, there are 2 cases where the 1st IE decreases. These are going from Beryllium-Boron and Nitrogen-Oxygen.
In the 1st 2 elements in period 2 (Lithium and Beryllium) we are removing an electron from the 2s subshell but in the case of boron, the outer electron is in the 2p subshell. The 2p subshell has a higher energy than the 2s subshell which means that it takes less energy to remove the outer electron of boron compared to the outer electron of beryllium. This is why boron has a lower 1st IE than beryllium.
Ionisation energy continues to increase across carbon and nitrogen but falls again at oxygen. In the case of nitrogen , each electron is in a separate 2p orbital but in the case of oxygen, one of the orbitals contains a pair of electrons which repel each other. This means that it takes less energy to remove one of these electrons than if the electrons were in separate orbitals. This means that the 1st IE of oxygen is less than nitrogen.
Define relative isotopic mass and relative atomic mass
R.I.M - the mass of an isotope relative to 1/12th the mass of an atom of carbon-12
R.A.M - (weighted average mass of all the isotopes) / 1/12 mass of one atom of 12C
Define atomic number, mass number and isotopes
Atomic number is the number of protons in the nucleus of an atom.
Mass number is the number of protons and neutrons in the nucleus of an atom.
Isotopes are atoms of the same element with a different number of neutrons.
