Topic 1 - Atomic Structure & the Periodic Table Flashcards
What are the relative masses and charges of electrons, protons & neutrons?
>Electrons = 1/1840 & -1 >Protons = 1 & +1 >Neutrons = 1 & 0
What is meant by the term mass number?
The total number of protons & neutrons.
What is meant by the term atomic number?
The number of protons.
What is meant by the term isotope?
Atoms with the same number of protons but different numbers of neutrons.
What is meant by the term relative isotopic mass?
The mass of 1 atom of an isotope compared to 1/12th of the mass of 1 atom of carbon-12.
What is meant by the term relative atomic mass?
The average mass of 1 atom compared to 1/12th of the mass of 1 atom of carbon-12.
How do you calculate the relative atomic mass?
∑ (isotopic mass x % abundance) / 100 OR ∑ (isotopic mass x abundance) / total abundance
What does a mass spectrometer measure?
m/z (mass/charge ratio) & abundance
How can you work out the Mᵣ from the mass spec?
The larges peak on the mass spec is due to the complete molecule (the M+ peak). The M+1 peak is cause by the carbon 13 isotope.
What is meant by the term 1st ionisation energy?
The energy required when 1 mole of gaseous atoms forms 1 mole of gaseous ions with a single positive charge. (energy to take 1 electron)
H → H⁺ + e⁻
What is meant by the term 2nd ionisation energy?
The energy required when 1 mole of gaseous atoms with a single positive charge forms 1 mole of gaseous ions with a double positive charge.
Ti⁺ → Ti²⁺ + e⁻
How does the number of protons affect the ionisation energy?
The more protons there are in the nucleus, the greater the attraction will be ∴ the higher the ionisation energy.
How does electron shielding affect ionisation energy?
The more electrons there are, the more they repel the outer electrons ∴ the lower the ionisation energy.
How does the distance of the electrons from the nucleus affect the ionisation energy?
The further the outer electrons are from the nucleus, the weaker the attraction will be ∴ the ionisation energy will be lower.
Why does the 1st ionisation energy increase across a period?
The number of protons increases, increasing the attraction to the nucleus, without extra shells for more shielding.
Why does the 1st ionisation energy decrease down a group?
The outer electrons are in shells further away from the nucleus and have more shells shielding them.
Why are succesive ionisation energies larger?
When the 1st electron is removed a positive ion is formed which increases the attraction of the remaining electrons, making them harder to remove.
Why has helium got the largest ionisation energy?
It has no shells before the outer electron to shield it, and is bigger that hydrogen as it has 1 more proton.
Why are there drops from Be → B & Mg → Al?
B & Al start to fill a new orbital so the outer electrons have more shielding & energy ∴ decreasing their ionisation energies.
Why are there drops from N → O & P → S?
The outer electrons start to double up in the orbitals, causing slight repulsion & making them easier to remove.
What is an orbital?
A region within an atom that can hold up to 2 electrons with opposite spins.
In what order do electrons fill subshells?
Singly, before pairing up with an electron with an opposite spin.
Which elements are in the s, p & d blocks?
s = 1,2 p = 3,4,5,6,7,0 d = transition metals
What is the electronic structure for ions?
When a positive ion is formed electrons are lost, when a negative ion is formed electrons are gained.
What do orbitals represent?
The mathematical probability of finding an electron at a certain point around the nucleus.
What is meant by the term periodicity?
The repeating pattern of physical or chemical properties going across the periods.
Why does the atomic radii decrease across a period?
The electrons have the same shielding as they are in the same shell, but have an increased attraction to the nucleus, as the number of protons increases.
Why does metallic bond strength increase across a period?
The more electrons there are in the outer shell, the more electrons can be released into the sea. (Na, Mg, Al)
Why does Si have a high melting point?
It is macromolecular (lots of strong covalent bonds) ∴ needs lots of energy to break them.
Why are the mp & bp of molecular gases low?
They have weak london forces ∴ little energy is needed to break them.
Why are the mp & bp of Ne & Ar so low?
They are monoatomic, so have very weak london forces.
Why does S₈ have a high mp than P₄?
They are both low as they are simple molecular (weak london forces) but S₈ has more electrons ∴ slightly stronger london forces.
In what order do transition metals loose electrons?
Loose electrons from 4s before 3d.
