Atomic structure

Why is the mass number (top number) a decimal?

• Because of the presence of isotopes

Define isotopes

• Atoms of the same element with different numbers of neutrons and masses

Why do all isotopes of an element react in the same way?

• They have the same electron configuration

What does abundance tell you?

• How common each isotope is

• E.g. 69% of copper atoms are copper 63

• 31% of copper atoms are copper 65

How can we determine the mass number and abundance of isotopes ?

• Use a machine called a mass spectrometer

Outline the stages of the time of flight mass spectrometer

• Take a sample of the element your interested in. Place it in the sample chamber

• The sample contains all of the different isotopes of that element

• Atoms go through a process of ionisation. Converting all the atoms into positive ions

• These positive ions are now attracted to a negatively charged plate

• The negatively charged plate causes the ions to accelerate increasing kinetic energy of the ions

• All of the ions with the same charge will have the same kinetic energy

• Once the ions pass through the negative plate. They stop accelerating

• They then drift down the flight tube at different velocities .With lighter ions moving faster than heavier ions

• After reaching the detector, each positive ion gains an electron from the detector.( An ion with a single positive charge will gain a single electron)

• This transfer of electrons causes a current to flow

• Ions of the lighter isotope will have a greater velocity and reach the detector first

• Size of current produced when each isotope hits the detector is proportional to its abundance

Why is the interior of the mass spectrometer a vacuum?

• To prevent the ions from colliding with air molecules

• As this can cause the ions to be deflected and slowed down

Explain electron impact

• Its a method of ionisation

• A vapourised sample is injected at a low pressure

• Electron gun fires high energy electrons at the sample

• Knocking out outer electron

• E.g. Ti → Ti⁺ + e^-

Explain electrospray ionisation

• Sample is dissolved in a volatile solvent

• Injected through a needle connected to a high voltage

• Molecules gain H⁺

• E.g.: Mg + H⁺ →MgH⁺

Interpretating the mass spectrum

• 2 peaks show that Copper has 2 main isotopes with an abundance of 69.2% and 30.8%

• M/Z = Ratio of the mass of each ion to its charge

• All of the ions have a single positive charge so the M/Z ratio is basically the relative mass of the ion

What happens when mass spectrometry is carried out on a molecule?

• A range of peaks are formed due to fragmentation

What does fragmentation do?

• When fragmentation happens a bond breaks, forming a molecular ion and a radical.

What can the fragmentations be used for?

• They can help identify the molecule

Mass spectrum for carbon dioxide

• The peak at 44 shows you a molecule that has an mr of 44. It can be assumed that it is co2. This peak is called a molecular ion M⁺

• The rest of the peaks depict other parts of the molecule that have been ionised

• The other peaks give you the mr of the atoms in the molecule that help figure out what that the molecule is

• The peak at 12 (mr of carbon) tells you there is a carbon in the molecule

How can the maximum number of electrons an energy level can hold be calculated

• 2n²

• n being the energy level

• E.g. the first energy level can hold 2 electrons 2(1)²

Electrons in energy levels are found in regions called……

• Atomic orbitals

Define atomic orbital

• A region around the nucleus that can hold up to 2 electrons with opposite spins

Shape of the s orbital

• spherical

Shapes of p orbitals

What happens to the energy of the shells as you move further from the nucleus ?

• energy of the shell increases

• weaker electrostatic force of attraction between nucleus and electrons means electrons on shells further from the nucleus require more energy to stay in these positions

Order of orbital filling

• 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁶

Rules for filling atomic orbitals

• Orbitals with the lowest energy are filled first

• Can have 2 electrons in each orbital but they must have opposite spins

• Fill 4s before 3d

• Fill orbitals with the same energy singly, before pairing. Because electrons in the same orbital repel

Name 2 elements exempt from orbital filling rules and why

• Chromium and copper

• When filling the orbitals for these 2 elements put only 1 electron in 4s and the rest in 3d

• Because 3d subshell is most stable when its completely or half full

Electron configuration is always written in order of …..

• shells not filling

Whats the shorthand configuration for sodium

What is the shorthand configuration for sulphur?

When an atom forms an ion which subshell does it loose or gain electrons from or in

• The subshell with the highest energy

What is the electron configuration of an Mg atom and an Mg²⁺ ion?

• Mg atom=1s² 2s² 2p⁶ 3s²

• Mg²⁺ ion = 1s² 2s² 2p⁶

• Mg loses 2 electrons from the 3s subshell to form a positive ion

What happens when a D block element becomes an ion and why

• It loses electrons from 4s subshell before 3d

• The 4s subshell always filled before 3d. As 4s has lower energy than 3d

• Once the 4s subshell contains electrons it has a higher energy than 3d subshell

Electron configuration for Fe and Fe²⁺

• Fe = 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²

• Fe²⁺ = 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶

• Fe is a d block element

What do you need to remove an electron from an atom?

• Energy

Define first ionisation energy

• Energy needed to remove one mole of electrons from one mole of atoms in their gaseous state to form one mole of 1+ ions in their gaseous state

First ionisation energy equation for Mg

What is meant by successive ionisation energies?

• Once you have removed 1 electron you can continue to remove electrons and measure the ionisation energy each time

Define second ionisation energy

• Energy required to remove one mole of electrons from one mole of 1+ ions in their gaseous state to form one mole of 2+ ions in their gaseous state

Second ionisation equation for Mg

The stronger the attraction between the outer electrons and the positive protons in the nucleus…

• The greater the ionisation energy

Factors affecting ionisation energy

• Atomic radius (distance between outer electrons and nucleus)

• Charge on the nucleus (the greater the amount of protons, the greater the force of attraction)

• Shielding (More shells reduce the attraction)

Explaining the successive ionisation energies of oxygen

• For the first to six electrons there was a gradual increase in the ionisation energy

• This is because all 6 electrons were removed from oxygens outer shell (low attraction between nucleus and outer electrons so little ionisation energy needed)

• But the remaining 2 electrons were on the shell closest to the nucleus. Higher ionisation energy was needed to remove the 7th and 8th electron

Explain the trend for first ionisation energy down a group

• FIE decreases down a group

• Atomic radius increases-meaning the outer electron is further from the nucleus/weaker force of attraction

• Number of energy levels increase-more shielding

• Although charge on the nucleus increases its effect is cancelled by the other 2 factors

Explain the trend for FIE across a period

• FIE generally increases

• Nuclear charge increases- higher number of protons/stronger attraction→ more energy is needed

• Strong attraction also decreases atomic radius as outer electrons are pulled closer to the nucleus

FIE of elements in period 2

• In period 2 the FIE of oxygen and boron actually decreases

• This is due to their electron configuration

• Oxygen has a pair of electrons in its p subshell. Less energy required to remove these electrons as they repel each other

• Boron has one electron in its p subshell whereas Be doesn’t have a p subshell.

• It is easier to remove that one electron in boron’s p subshell as it is further away from the nucleus