Topic 1- Atomic Structure And The Periodic Table

Question 1

Know the structure of an atom in terms of protons neutrons amd electrons Know the relative mass and relative charges of protons neutrons and electrons

Answer 1

Proton- mass 1 charge +1 (in nucleus) Neutron- mass 1 charge 0 (in nucleus) Electron- mass 0 charge -1 (in orbitals)

Question 2

Know what is meant by atomic number and mass number

Answer 2

Mass number - tells you the total number of protons and neutrons Atomic number - tells you number of protons (all atoms of the same element have the same number of protons )

Question 3

be able to determine the number of each type of sub-atomic particle in an atom, molecule or ion from the atomic (proton) number and mass number

Answer 3

Neutrons - take the amount of protons away from mass number Electrons - same as protons

Question 4

understand the term ‘isotopes’

Answer 4

Atoms with the same number of protons but different number of neutrons (Have same chemical properties bcs same electron configuration but different physical properties ie densities bcs different masses )

Question 5

be able to define the terms ‘relative isotopic mass’ and ‘relative atomic mass’, based on the 12C scale

Answer 5

Relative atomic mass- the weighted mean mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12 Relative isotopic mass- the mass of an atom of an isotope compared with 1/12th of the mass of an atom of carbon 12

Question 6

understand the terms ‘relative molecular mass’ and ‘relative formula mass’, including calculating these values from relative atomic masses Definitions of these terms will not be expected. The term ‘relative formula mass’ should be used for compounds with giant structures.

Answer 6

Relative molecular mass - add up all the relative atomic mass values of all the atoms in the molecule Relative formula mass-add up the relative atomic mass vales of all the ions or atoms in the formula unit.

Question 7

be able to analyse and interpret data from mass spectrometry to calculate relative atomic mass from relative abundance of isotopes and vice versa

Answer 7

(mass x abundance) + ( mass x abundance ) divide by 100

Question 8

be able to predict the mass spectra, including relative peak heights, for diatomic molecules, including chlorine

Answer 8

Show each percent as a decimal

make a table showing all the different molecules

look for any molecules in the table that are the same and add up their abundance’s

divide all relative abundance’s by the smallest relative abundance to get a ratio

Question 9

understand how mass spectrometry can be used to determine the relative molecular mass of a molecule
Limited to the m/z value for the molecular ion, M+, giving the relative molecular mass of the molecule

Answer 9

Look at the peak with the highest m/z value (ignore any small m+1 peaks). The m/z value of the molecular ion peak is the molecular mass

Question 10

be able to define the terms ‘first ionisation energy’ and ‘successive ionisation energies’

Answer 10

1st ionisation energy-the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ Ions (O > O+ +e-)

successive ionisation energy- the second ionisation energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions. (O+ > O2+ +e-)

Question 11

understand how ionisation energies are influenced by the number of protons, the electron shielding and the electron sub-shell from which the electron is removed

Answer 11

Nucleur charge- more protons more + charged the nucleus is so stronger attraction for the electrons

electron shell- attraction decreases with distance, electrons in shells closer to the nucleus are much more strongly attracted than those further away

shielding- number of electrons between the outer electrons and nucleus increases, the outer electrons feel less attraction towards the nucleus charge.

(high inionsation energy means strong attraction more energy needed to remove electron)

Question 12

understand reasons for the general increase in first ionisation energy across a period

Answer 12

Number of protons increases so stronger nucleur attraction and all the extra electrons are in same energy level even if in different orbitals so little effect of shielding or extra distance.

Question 13

understand reasons for the decrease in first ionisation energy down a group

Answer 13

Elements further down the group have extra electron shells compared to ones above, means the atomic radius is larger so electrons are further away amd reduces their attraction from the nucleus, the extra inner shells shield the attraction of the nucleus

evidence to prove shells exists.

Question 14

understand how ideas about electronic configuration developed from:
i )the fact that atomic emission spectra provide evidence for the existence of quantum shells

Answer 14

Emission spectra has clear lines for different energy levels, this supports the idea that energy levels are discrete (not continuous) shows electrons don’t move from one energy level to another it “jumps” with no in between stage backs up the idea that electrons exist in quantum shells

Question 15

ii) the fact that successive ionisation energies provide evidence for the existence of quantum shells and the group to which the element belongs

Answer 15

Within the shell ionisation energies increase, as electrons are being removed from an increasingly positive ion so less repulsion amongst remaining electrons so they’re held more strongly by the nucleus.

big jumps in ionisation energy show when a new shell is being broken into an electron is being removed from the shell closer to the nucleus evidence to prove their are shells

Question 16

the fact that the first ionisation energy of successive elements provides evidence for electron sub-shells

Answer 16

2 and 3 = sub shell structure- Al has a 3p orbital rather than a 3s, 3p has slightly higher energy than 3s and is found to be further from the nucleus also has additional shielding from 3s2, overrides effect of nucleus charge shows a drop proves evidence for the theory of electron sub shells.

5 and 6 = electron repulsion - phosphorus is being removes from a singly ocupied orbital in sulfuric its being removed from one which contains 2 electrons, revulsion between the two electrons means the electrons are easier to removed from shared orbitals, more evidence for electron structure model.

Question 17

know the number of electrons that can fill the first four quantum shell

know the number of electrons that occupy s, p and d-subshells

Answer 17

Question 18

know that an orbital is a region within an atom that can hold up to two electrons with opposite spins

know the shape of an s-orbital and a p-orbital

Answer 18

An orbital is a space where electrons move, orbitals within the same sub shell have the same energy.
electrons in each orbital have to spin in opposite directions called spin- pairing (electrons can be represented as arrows in boxes)

s orbitals are spherical, p orbitals have dumbbell shape, there are 3p orbitals at right angles to each other

Question 19

knowthatelectronsfillsubshellssingly,beforepairingup,andthattwoelectrons in the same orbital must have opposite spins

Answer 19

For example arrows in boxed 2p would have three boxes , there would be an up arrow in all three boxes before there would be a down arrow.
this is due to repulsion .

4s fills up before 3D because it has a lower energy level and electrons fill up the lowest energy subshells first.

Question 20

be able to predict the electronic configurations, using 1s notation and electrons- in-boxes notation, of:
i atoms, given the atomic number, Z, up to Z = 36
ii ions, given the atomic number, Z, and the ionic charge, for s and p block ions only, up to Z = 36

know that elements can be classified as s, p and d-block elements

Answer 20

S blocks have an outer shell electronic configuration of s1 or s2, p block have an outer shell electonic configuration of s2p1-s2p6, the 4s orbitals fills up before 3d, Cr and cu are exceptions 4s1 then fills d as they’re more stable with a full or half subshell)

to work out ions wrote the electronic configuration of the atom and add or remove electrons to ot from the highest energy ocupied subshell

Question 21

understand that electronic configuration determines the chemical properties of an
element

Answer 21

S block - they have 1 or 2 outer electrons which are easily lost to form positive ions with an inert gas configuration

p block - can gain 1,2 or 3 electrons to form negative ions with an inert gas configuration

groups 4-7 can share electrons when they covalent bond

d block - lose s and d electrons to form a positive ion

Question 22

understand periodicity in terms of a repeating pattern across different periods

Answer 22

All elects within a group have the same number of electrons in their outer shell, this means they have similar chemical properties, so there’s a repeating trend ( known as periodicity)

Question 23

understand reasons for the trends in the following properties of the elements from periods 2 and 3 of the Periodic Table:
i the melting and boiling temperatures of the elements, based on given data, in terms of structure and bonding
ii ionisation energy based on given data or recall of the plots of ionisation energy versus atomic number

Question 24

be able to illustrate periodicity using data, including electronic configurations, atomic radii, melting and boiling temperatures and first ionisation energies

Answer 24

Atomic radius- the number of protons increases, this increases the positive nucleur charge so means electrons are pulled closer to the nucleus making the atom smaller, extra electrons are added to the outer shell so don’t provide a shielding effect

Question 25

Atomic emission spectra - elecrons release energy in fixed amounts info on page 12

Answer 25

In their ground state atoms have electrons in their lowest energy levels , if an electron takes in energy from its surroundings they move to a higher energy level , electrons release energy by dropping from a higher energy level to a lower energy level, these have certain fixed values , a line spectrum (emission spectrum) shows the frequencies of light emitted when electrons drop from a higher energy level to a lower one ,shown as coloured lines on a dark background. Spectrum for each element is unique.

Question 26

Notes on page 12 still emission spectra.

Answer 26

Question 27

What are the four basic principles for electrons.

Answer 27

• electrons can only exists in fixed orbital sans not anywhere in between
• each shell has a fixed energy
• when an electron moves between shells electromagnetic radiation is emitted or absorbed
• bcs the energy of shells is fixed the radiation will have a fixed frequency.