Atomic Structure Flashcards
Electron relative mass
1/1840
Nucleus
- carries virtually all of the mass of the atom
* positively charged
The chemical nature of an atom is determined by…
The number of electron it has, and their arrangement
Relative atomic mass
The average mass of one atom compared to 1/12 that of the mass of one atom of carbon-12
Relative isotopic mass
The mass of one atom of an isotope compared to 1/12 of the mass of one atom of carbon-12
Relative molecular mass
The average mass of a molecule compared to 1/12 of the mass of one atom of carbon-12
Relative formula mass
- the same as relative molecular mass, but for ionic compounds (as these are not molecules and giant covalent substances - the size of the molecule varies - it cannot have a precise molecular mass)
- used for compounds with giant structures
First ionisation energy
The energy required to remove one electron from each atom in one mole of gaseous atoms to produce one mole of gaseous ions with a 1+ charge
Equation for first ionisation energy
X(g) -> X+(g) -> e-
The first ionisation energy size provides information about…
… the force of attraction between the nucleus and the outer electrons
The attraction of an electron to the nucleus depends upon
- the nuclear charge
- the distance from the nucleus (greater = less attraction)
- shielding (more inner electrons = less attraction)
Nuclear charge
- the amount of protons relative to the amount of electrons
* the higher the charge, the greater the attraction
Trend of first ionisation energies across a period
- increase
- electrons are added to the same electron shell (same distance, shielding)
- nuclear charge increases; attraction to nucleus increases
- therefore more energy is needed to remove the electron
Trend across the periodic table for first ionisation energy
- group 1 has the lowest
* noble gases have the highest
Trend for first ionisation energy across a group
- decrease
* distance and shielding increase
Second ionisation energy
The energy required to remove one electron from each 1+ ion in one mole of gaseous ions to produce one mole of 2+ gaseous ions
Ionisation energy
The energy required to remove the readily available electron from an atom, by overcoming the attraction to the nucleus
General principles of first ionisation energy
- as nuclear charge increases, there is a stronger attraction, so ionisation energy increases
- as distance increases, attraction falls very rapidly - ionisation energy decreases
- shielding - inner shells repel outer electrons away from the nucleus
Spin-pairing
Causes repulsion; offsets attraction to the nucleus; paired e-s are removed more easily
Successive ionisions énergies
- give us important information about the electronic structure of an element
- determined which group an element is in
Trend for successive ionisation energies
- the second ionisation energy of an element is always greater than the first
- after each ionisation, nuclear charge increases
- the next electron must be removed from a greater attraction
E.g. successive ionisation energies for potassium
- first electron is removed from the outer-shell
- second electron has to be removed from an inner-shell; requires more energy
- the tenth electron has to be removed from the next inner shell
If there is a large jump in successive ionisation energy
You have moved inwards a cell- all previous electrons were in the outer shell
Electronic structure
Determines an atom’s chemical and physical properties
Aufbau principle
The lowest energy sub-levels are occupied first
Shell 1
- s subshell
* 2 electrons
Shell 2
- s and p subshells
* 8 electrons
Shell 3
- s, p and d sub-shells
- 18 electrons
- can expand the octet
Shell 4
- s, p, d and f subshells
* 32 electrons
The higher the energy of the shell…
… the more subshells it contains
2n^2
Gives the number of electrons in any shell
Where n is the shell number
Orbitals
An area occupied by an electron wave
S orbitals are
Spherical
Electrons in the first shell have the …
… same energy value, but opposite spin values
Electrons in the p subshell
- are of a higher energy level than those at the s
* found in 3 different regions of the atom: x, y and z
The 4s subshell fills up…
… before the 3d subshell
Because an exactly half-full/full d sub-shell is particularly stable
Hund’s rule
Electron orbitals are filled such that once an electron has entered an orbital of a p or d subshell, the orbital will not take on another electron until all other orbitals of the same energy level also have one electron
s-block element
An element in which the last electron is in the s orbital
Similar properties are portrayed by…
… elements in the same group; elements with the same number of electrons and electrons in the same type of sub-shell
Copper and chromium
- exceptions to the Aufbau principle
- an electron from the 4s subshell goes to the 3d subshell to give chromium a half-full d subshell and copper a full subshell
Cr: 1s^22s^22p^63s^23p^64s^13d^5
Cu: 1s^22s^22p^63s^23p^64s^13d^10
Noble gases
Have shells which are completely full
[Ar]
- has the same electron configuration as argon
* used because it is a noble gas
Group 4 elements
- tend to form giant covalent structures
- elements towards the bottom of the group are metallic
- Pb and Sn are metallic because they are large; electrons are further from the nucleus and there is more shielding; e-s are lost more easily
Groups 5,6 and 7
- generally have simple covalent structures
* bottom of the groups become more metallic
Group 8
- noble gases
- have complete electron shells
- monoatomic
Moving horizontally…
… electron control increases as non-metallic character increases
Moving downwards…
… electron control decreases due to larger atoms; electrons are further from the nucleus with greater shielding; metallic character increases
Transition metals
A d-block element that forms one or more stable ions with incompletely filled d orbitals
Melting temperatures across a period
Rise until group 4, then decrease
Metals
- strong forces of attraction between cations and delocalised electrons
- increases across a period because charge increases (charge density increases)
- no of delocalised electrons increases; attraction of metallic bonding increases
Giant molecules
- strong forces of attraction between nuclei and shared pair of electrons
- melting temperature is at peak because vast amounts of strong covalent bonds must be broken
Simple molecules
- covalent bonds gold atoms strongly in the molecule, but there are weak intermolecular forces
- melting temperatures are low because intermolecular forces are weak and roughly relative to the site of the molecules
Noble gases have the lowest melting temperatures because…
… they are monatomic
Phosphate
(PO4)3-
Soluble:
- Na+, K+, NH4+, NO3-
- Cl- except Ag+ and Pb2+
- (SO4)2- except Ca2+, Ba2+ and Pb2+
Insoluble:
- carbonate
- oxide
- hydroxide except Ca2+
