1: Atomic Structure and the Periodic Table Flashcards
Explain the structure of an atom in terms of electrons, protons and neutrons.
Protons and neutrons contained in the nucleus.
Electrons in shells around the nucleus.
Give the relative mass of protons, neutrons and electrons
Proton: 1
Neutron: 1
Electron: 1/1840
Give the relative charge of protons, neutrons and electrons
Proton: +1
Neutron: None
Electron: -1
If an atom has an atomic number = Z and a mass number = M, determine the number of each type of sub-atomic particle.
Proton: Z
Neutron: M - Z
Electron: Z
Explain the term ‘isotopes’.
Atoms with the same atomic number, but different mass numbers.
Define atomic number
The number of protons in the nucleus of an atom
Define mass number
the total number of protons and neutrons in the nucleus of an atom
Define relative isotopic mass
The mass of an atom of a particular isotope of an element, relative to 1/12th the mass of a carbon-12 atom.
Define relative atomic mass
The weighted average mass of all isotopes of an element, relative to 1/12th the mass of a carbon-12 atom.
What is relative molecular mass?
The average mass of a molecule of a substance, relative to 1/12th the mass of a carbon-12 atom.
What is relative formula mass?
The sum of the atomic masses of all the atoms in a particular formula.
- Used for compounds with giant structures.
Mass spectrometry
What happens in the ionisation stage?
Electrons from the electron gun collide with the particles and form charged ions.
Mass spectrometry
What happens in the acceleration stage?
An electric field accelerates the particles to a uniform speed.
Mass spectrometry
What happens in the deflection stage?
A magnetic field deflect the particles.
Particles with less mass and more charge are deflected more.
An electromagnet varies the strength of the field.
Mass spectrometry
What happens in the detection stage?
The ions create an electric current which is detected.
Define 1st ionisation energy.
The energy required to
remove 1 electron from each atom
in 1 mole of gaseous atoms
to form 1 mole of gaseous 1+ ions
Define successive ionisation energies.
The energy required to
remove 1 electron from each ion
in 1 mole of gaseous ions
to form 1 mole of gaseous (2+, 3+, 4+, …) ions
How does the number of protons influence ionisation energy?
More protons -> stronger attraction -> higher IE
How does electron shielding influence ionisation energy?
More shielding -> weaker attraction -> lower IE
How does the electron subshell from which the electron is removed influence ionisation energy?
Greater distance -> weaker attraction -> lower IE
State and explain the general trend in first ionisation energy as you go across a period.
Nuclear charge increases,
Shielding increases slightly,
Atomic radius decreases,
Strength of attraction between nucleus and outer electron increases,
so first ionisation energy increases.
State and explain the trend in first ionisation energy as you go down a group.
Nuclear charge increases,
Shielding increases,
Atomic radius increases,
Strength of attraction between nucleus and outer electron decreases,
so first ionisation energy decreases.
What ideas provide evidence for the existence of quantum shells?
- Atomic emission spectra
- Successive ionisation energies
- First ionisation energies of successive elements
What do the atomic emission spectra provide evidence for?
The existence of quantum shells.
What do successive ionisation energies provide evidence for?
The existence of quantum shells and the group to which the element belongs.
What do the first ionisation energies of successive elements provide evidence for?
Electron sub-shells.
How many electrons can fill the first four quantum shells?
1: 2
2: 8
3: 18
4: 32
What is an orbital?
A region within an atom that can hold up to 2 electrons with opposite spins.
What is the shape of an s orbital?
2 electrons in an s subshell, therefore 1 orbital in the shape of a sphere.
What is the shape of a p orbital?
6 electrons in a p subshell, therefore 3 orbitals
What is the order that the subshells are filled and the maximum number of electrons in each subshell?
1s - 2 2s - 2 2p - 6 3s - 2 3p - 6 4s - 2 3d - 10 4p - 6
What rules describe how electrons fill subshells and orbitals?
Electrons enter the lowest energy orbital available.
Electrons fill subshells singly, before pairing up.
Two electrons in the same orbital must have opposite spins.
What is the electronic configuration of copper in 1s notation? (group 11 period 4)
1s2 2s2 2p6 3s2 3p6 3d10 4s1
What is the electronic configuration of chromium in 1s notation? (group 6 period 4)
1s2 2s2 2p6 3s2 3p6 3d5 4s1
What are the four blocks of the periodic table?
What determines the chemical properties of an element?
Electronic configuration
Define periodicity.
A repeating trend in across different periods.
State and explain the trend in the melting and boiling points from sodium to aluminium.
Ioinic radius decreases,
Charge of ions increases,
So number of delocalised electrons increases,
So attraction between ions and electrons increases,
So melting and boiling point increases.
Explain why silicon has a high melting and boiling point compared to other period 3 atoms.
Silicon forms a giant covalent structure,
Held together by strong covalent bonds.
Other non-metals form simple molecular structures held together by weak intermolecular forces.
