TOPIC 1 - ATOMIC STRUCTURE & PERIODIC TABLE Flashcards
(34 cards)
Define the terms ‘atomic (proton) number’ and ‘mass number’
The atomic number is the number of protons in the nucleus of one atom.
The mass number is the number of protons and neutrons in the nucleus of one atom.
Define the term ‘isotopes’
Isotopes are atoms with the same number of protons, but different numbers of neutrons.
Define the terms ‘relative isotopic mass’ and ‘relative atomic mass’
Relative atomic mass is the average mass of one atom compared to one twelfth of the mass of one atom of carbon-12.
Relative isotopic mass is the mass of one atom of an isotope compared to one twelfth of the mass of one atom of carbon-12
What is ‘relative molecular mass’ and ‘relative formula mass’?
Relative molecular mass is the average mass of a molecule compared to one twelfth of the mass of one atom of carbon-12. Relative molecular mass refers to compounds containing molecules.
Define the terms ‘first ionisation energy’ and ‘successive ionisation energies’
‘first ionisation energy’ - Energy required to remove one mole of electrons from one mole of gaseous atoms
‘successive ionisation energies’ - successive ionisation energies are always larger because the second ionisation energy of an element is always bigger than the first ionisation energy. When the first electron is removed a positive ion is formed. The ion increases the attraction on the remaining electrons and so the energy required to remove the next electron is larger.
Be able to analyse and interpret data from mass spectrometry to calculate relative atomic mass from relative abundance of isotopes and vice versa
A mass spectrometer produces a mass spectrum which shows lines where ions of that mass are present.- The relative heights of the peaks on the mass spectrum show the relative abundance of the different ions present.- The Ar of an element can be calculated by measuring how high each peak is and therefore working out the relative abundance of each isotope.
How are ionisation energies are influenced by the number of protons, the electron shielding and the electron sub-shell from which the electron is removed?
The attraction of the nucleus (The more protons in the nucleus the greater the attraction) 2. The distance of the electrons from the nucleus (The bigger the atom the further the outer electrons are from the nucleus and thus the weaker the attraction to the nucleus) 3. Electron shielding (An electron in an outer shell is repelled by electrons in complete inner shells, weakening the attraction of the nucleus)
What are the reasons for the general increase in first ionisation energy across a period?
Across a period , the number of protons increases making the effective attraction of the nucleus greater. The electrons are being added to the same shell which has the same shielding effect and the electrons are pulled in closer to the nucleus.
What are the reasons for the decrease in first ionisation energy down a group?
Increase in electron shielding down a group, so a weaker attraction of the nucleus due to the increased inner repulsion of the electron shells.
Number of electron shells increases down a group, so the atomic radius increases, so the outer electrons are further away, so they are more easily removed.
How did electron configuration develop?
- emission spectra - evidence of quantum shells
- successive ionisation energies - evidence for quantum shells within atoms and suggest the group of which the element belongs to
- first ionisation energy of successive elements - evidence of electron sub shells
What is an orbital?
An orbital is a region within an atom that can hold up to two electrons with opposite spins
What is the number of electrons in each sub shell?
s sub shell - 2
p-sub shell - 6
d- sub shell - 10
How do electrons fill subshells?
Electrons fill subshells singly, before pairing up, and that two electrons in the same orbital must have opposite spins
What is periodicity?
A repeating pattern across a period is called periodicity
Period 3 trend in melting points
For Na, Mg, Al- Metallic bonding : strong bonding - gets stronger the more electrons there are in the outer shell that are released to the sea of electrons. A smaller sized ion with a greater positive charge also makes the bonding stronger. Higher energy is needed to break metallic bonds. Si is Macromolecular: many strong covalent bonds between atoms high energy needed to break covalent bonds- very high mp +bp Cl2 (g), S8 (s), P4 (S)- simple molecular : weak London forces between molecules, so little energy is needed to break them - low mp+ bp S8 has a higher mp than P4 because it has more electrons (S8 =128)(P4=60) so has stronger London forces between molecules Ar is monoatomic weak London Forces between atoms.
Period 3 trend in first ionisation energies
The general trend across is to increase. This is due to increasing number of protons as the electrons are being added to the same shell There is a small drop between Mg + Al. Mg has its outer electrons in the 3s sub shell, whereas Al is starting to fill the 3p subshell. Al’s electron is slightly easier to remove because the 3p electrons are higher in energy. There is a small drop between phosphorous and sulfur. Sulfur’s outer electron is being paired up with an another electron in the same 3p orbital. When the second electron is added to an orbital there is a slight repulsion between the two negatively charged electrons which makes the second electron easier to remove.
Electron configuration of Cr and Cu
Only one electron in 4s sub shell
Cr: (Ar)4s^1 3d^5
Cu: (Ar)4s^1 3d^10
What is the relative mass of protons, neutrons and electrons?
Protons: 1
Neutrons: 1
Electrons: 1/1840
What affects the physical properties of an isotope?
Mass of atom affects physical properties, so isotopes will have slightly different physical properties e.g. different densities, rates of diffusion.
What affects the chemical properties of an isotope?
Number and arrangement of electrons determine the chemical properties, so isotopes have the same chemical properties.
How to use mass spectrometry to work out relative atomic masses?
- Multiply the relative isotopic mass by its relative isotopic abundance. (not % abundance)
- Add up the results.
- Divide by the sum of the isotopic abundances.
How to predict the mass spectra for diatomic molecules?
- Express the abundances as decimals.
- Create a table showing all the possible molecules that could be formed.
- For each molecule, multiply the decimal abundances of the isotopes to get the relative abundance of each one.
- Look for molecules in the table that are the same, if they are, add up their abundances.
- Divide all the relative abundances by the smallest relative abundance to get the smallest whole number ratio.
- Use the RMM of the molecules to predict the mass spectrum.
What are the steps in Mass Spectrometry?
- Vaporisation
- Ionisation
- Atomisation
- Deflection
- Detection
What are electrons grouped together in?
quantum shells (energy levels)
