Atomic Structure Flashcards
How will protons, neutrons and electrons behave when moving at the same velocity in an electric field?
- The electrons will deflect very readily towards the positively charged region as they have a low mass
- The protons will deflect towards the negatively charged region but not as readily as electrons
- The neutrons will continue in a straight line
What is atomic radius?
Half the distance between two nuclei of covalently bonded atoms of the same type
In the case of metals, it is half the distance between two adjacent nuclei within the metallic lattice
What is the trend in atomic radii down a group?
They increase
What is the trend in atomic radii across a period?
They decrease
What is the explanation behind the trend in atomic radii down a group?
- As you move down a group, the electrons surrounding the atoms occupy a greater number of shells
- The electrons in inner shells shield the electrons in the outer shells from the nucleus and repel them, weakening the pull of the nucleus on those electrons
- They therefore are at a higher energy level and are held less tightly to the nucleus, moving more
- This is enough to counteract the increasing number of protons
What is the explanation behind the trend in atomic radii across a period?
- The number of protons increase, increasing the positive charge of the nucleus
- This leads to the outer electrons experiencing a larger effective nuclear charge, causing them to be held more tightly to the nucleus
- While the number of electrons increases, they do not occupy new shells and shielding/repulsion does not increase significantly
What is the trend in cationic radii with increasing positive charge and what is the explanation behind it?
- Cationic radii decrease with increasing positive charge
- With increasing positive charge, the nuclear charge stays the same but there are fewer electrons
- This means that the remaining electrons experience less shielding and spin-pair repulsion and are at a lower energy level; this causes them to be held more tightly to the nucleus
- If enough electrons are lost, the outer electrons may begin to occupy a shell closer to the nucleus, significantly decreasing atomic radius
What is the trend in anionic radii with increasing negative charge and what is the explanation behind it?
- Anionic radii increase with increasing negative charge
- With increasing negative charge, the nuclear charge stays the same but there are more electrons
- This means that the electrons experience more shielding and spin-pair repulsion and are at a higher energy level; this causes them to be held less tightly to the nucleus
- If enough electrons are added, the outer electrons may begin to occupy a shell further from the nucleus, significantly increasing atomic radius
- Ionic radii will always increase down a group regardless of the type of ion as long as the number of shells increases
- Assuming they are in the same period, anions will always have a greater atomic radius
What is an isotope?
Atoms of the same element that contain the same number of protons and electrons but a different number of neutrons
Why do isotopes display similar chemical properties?
- They have the same number of electrons
- This means they have the same number of valence electrons
- They therefore participate in chemical reactions in the same way
Why do isotopes display similar physical properties?
- Isotopes have different masses and densities as they have different numbers of neutrons but are the same size
- An atom of one isotope may have a higher mass than another; as a result, a material consisting of this heavier isotope would have a higher density
What is the ground state of an atom
- It is the lowest energy state of an atom
- It will involve the electron configuration that the atom is most stable at when in isolation
- It usually involves having the same number of protons and electrons
What are principal quantum shells?
- Commonly referred to as shells, they are the highest level of organisation of electrons
- Each shell is assigned a principal quantum number, denoted by n: for example, the shell closest to the nucleus has n = 1
They are also called principal energy levels, as the energy of the shells increases as distance from the nucleus increases
What is the maximum number of electrons each principal quantum level can hold?
- n = 1: 2 electrons
- n = 2: 8 electrons
- n = 3: 18 electrons
- n = 4: 32 electrons
What are subshells?
- Subshells are the second highest level of organisation of electrons
- They are found within shells
- The four different types of subshells are s, p, d and f
It is more complicated than this, but for A-Level, this level of detail is sufficient
What is the maximum number of electrons each type of subshell can hold?
- s - 2 electrons
- p - 6 electrons
- d - 10 electrons
- f - 14 electrons
Which subshells does each principal quantum shell contain?
- n = 1: s
- n = 2: s, p
- n = 3: s, p, d
- n = 4: s, p, d and f
What are orbitals?
- The lowest level of organisation of electrons
- They are specific energy levels that electrons are found at; electrons cannot be found between these energy levels
- They contain up to two electrons
Within the same subshell, orbitals are known as degenerate as they have the same levels of energy
How many orbitals does each subshell contain?
s - 1
p - 3
d - 5
f - 7
What is the order of energy levels for each subshell in each shell from 1s to 4p?
List in terms of increasing energy
1s
2s
2p
3s
3p
4s
3d
4p
- The order of energy determines the order at which they will be filled; the subshells with the lowest energy are filled first
- After the 3d and 4s subshells are filled, the 4s subshell will have more energy than the 3d subshell. This causes electrons to be removed from 4s more readily as the 3d subshell is held more tightly to the nucleus due to it experiencing less shielding than the 4s subshell
- This order is determined using the Aufbau principle
How do orbitals within a subshell fill with electrons and why?
- Each orbital will gain a single electron, then each will gain a second
- This is to minimise electron spin-pair repulsion; electrons are at a higher energy level when sharing an orbital with another electron
What is the electronic configuration for iron in ground state?
Write both the longhand configuration and the shorthand configuration using the noble gas convention
- 1s², 2s², 2p⁶, 3s², 3p⁶, 3d⁶, 4s²
- [Ar] 3d⁶, 4s²
What would the electronic configuration for iron look like with the noble gas shorthand convention and the electrons in boxes notation?
What is the shape of an s orbital?
- Spherical and symmetric
- The size of the s orbitals increases with increasing principal quantum number
What is the shape of a p orbital?
- Dumbbell shaped, as shown in the image
- They become larger with increasing principal quantum number
- The three orbitals within a p subshell occupy the three different axes, so are perpendicular to one another (Pₓ is shown below)
What is a free radical?
A species with one or more unpaired electrons
This means it has an electron or electrons that do not share their orbital with another electron
What is the definition of first ionisation energy?
The amount of energy required to remove one mole of electrons from one mole of gaseous atoms of an element to form one mole of gaseous ions under standard conditions
For example: Ca (g) → Ca⁺ (g) + e⁻
What is the definition of second ionisation energy?
The amount of energy required to remove one mole of electrons from one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions under standard conditions
What is the trend in first ionisation energies down a group and what is the explanation for this trend?
- They decrease
- The number of shells increases as one descends the group, meaning atomic radius increases and the valence electron to be removed becomes further from the nucleus
- This means it is shielded more by inner electrons, held less tightly to the nucleus and at a higher energy
- Thus, it requires less energy to be removed
What is the trend in first ionisation energies across a period and what is the explanation for this trend?
- They increase
- This is because the elements increase in nuclear charge and decrease in atomic radius, so there is a stronger force of attraction between the valence electrons and the nucleus
- Thus, more energy is required to remove an electron further across a period
- This is enough to offset the increase in shielding and spin-pair repulsion engendered by the increased number of valence electrons as no new shells are formed
There are some exceptions to the last point
Why is the first ionisation energy of magnesium higher than the first ionisation energy of aluminium?
- Despite of the increased nucelar charge, aluminium has a lower ionisation energy
- This is because the valence electron to be removed is in the 3p subshell, whereas the valence electron to be removed in magnesium is in the 3s subshell
- The 3p subshell is higher in energy and further from the nucleus than the 3s subshell, so less energy is required to remove an electron from it
Why is the first ionisation energy of phosphorus higher than the first ionisation energy of sulfur?
- Phosphorus has three electrons in the 3p subshell, whereas sulfur has four
- This means that two electrons in sulfur’s 3p subshell are experiencing spin-pair repulsion as they share an orbital
- The electrons experiencing spin-pair repulsion are at a higher energy and move more than electrons that do not share an oribtal, so require less energy to be removed from the subshell
- This is enough to counteract the increased nuclear charge
Why are successive ionisation energies higher?
- It is much more difficult to remove an electron from a positive ion
- This is because there is a stronger electrostatic force of attraction between the nucleus and the remaining electrons
- In some cases, the electron to be removed will be occupying a subshell closer to the nucleus or even a principal quantum shell closer to the nucleus, so it will experience less shielding by inner electrons
- The electron to be removed may also no longer be experiencing spin-pair repulsion after the electron is was sharing an orbital with is removed, meaning more energy will be required to remove it from its subshell
- The last two points are important; they are what causes there to be major jumps between consecutive ionisation energies of an element
- These jumps can be used to predict the electron configuration of an element
