3.1.1 Atomic structure

What are the three subatomic particles

protons, neutrons, electrons

What is the relative mass and relative charge of protons

1, +1

What is the relative mass and relative charge of Neutron

1, 0

What is the relative mass and relative charge of Electron

0.0005, -1

Mass number (A)

The total number of nucleons (protons and neutrons) in the nucleus.

Atomic number (Z)

The number of protons in the nucleus, which uniquely identifies the element.

Which is where?

A at the top and Z at the bottom

Isotopes, explain how properties is affected

Atoms of the same element that have the same number of protons but different numbers of neutrons. Isotopes behave similarly in chemical reactions but have different physical properties, such as mass.

Why do they have similar chemical properties but different physical properties

Chemical properties:
-Isotopes have the same electron configuration.

-This means they have the same chemical properties (e.g. reactivity).

Physical properties:
-Isotopes have slightly different physical properties (e.g. mass and density).

-This is because physical properties depend on atomic mass.

Relative Atomic Mass

The average mass of an atom of an element, taking into account all its isotopes and their relative abundances, compared to 1/12 of the mass of a carbon-12 atom.

Relative Isotopic Mass

The mass of a specific isotope of an element relative to 1/12 of the mass of carbon-12.

How to determine the number of protons, electrons, and neutrons in a neutral atom

1.) The number of protons is equal to the atomic number (Z).

2.) In a neutral atom, the number of electrons is equal to the number of protons.

3.) The number of neutrons is calculated by subtracting the atomic number (Z) from the mass number (A).

How to determine the number of protons, electrons, and neutrons in an ion:

1.) The number of protons is equal to the atomic number.

2.) The number of electrons is calculated by subtracting the ion's charge from the atomic number (for cations) or adding the ion's charge to the atomic number (for anions).

3.) The number of neutrons is calculated by subtracting the atomic number from the mass number.

Stages of TOF Mass Spectrometry

1. Ionisation
2. Acceleration
3. Ion Drift
4. Detection

Two types of ionisation

1.) Electrospray Ionisation

2.) Electron Impact Ionisation

Electrospray Ionisation

-The sample is dissolved in a volatile solvent and injected through a fine needle, producing a mist.

-A high voltage is applied to the needle, causing the particles to gain a proton H+, forming positive ions.

-The solvent evaporates, leaving behind gaseous ions.

What is the formula for electrospray ionisation

X(g) + H+ -> XH+ (g)

Electron Impact Ionisation

-The sample is vaporised, and high-energy electrons are fired at it using an electron gun.

-This knocks off an electron from each particle, producing +1 ions.

Give two reasons why it is necessary to ionise the isotopes of chromium before they can be analysed in a TOF mass spectrometer.

-So Ions will interact with and be accelerated by an electric field
-So Ions create a current when hitting the detector

Explain why it is necessary to ionise molecules when measuring their mass in a TOF mass spectrometer.

-Ions, not molecules, will interact with and be accelerated by an electric field.
-Only ions will create a current when hitting the detector

Acceleration

-The positively charged ions are accelerated by an electric field.

-All ions are given the same kinetic energy KE= 1/2 mv^2 but their velocities will differ depending on their mass.

-Lighter ions move faster than heavier ions.

3. Ion Drift

-The ions enter a flight tube where there is no electric field.

-Ions drift through the tube, with the lighter ions (which move faster) arriving at the detector before the heavier ones.

Detection

-ion hits the detector and gains an electron
-relative abundance is proportional to the current

Time-of-flight calculations

t = time of flight (s)

d = length of flight tube (m)

Ke = kinetic energy (J)

m = mass of ion (kg)

Time-of-flight calculations for distance

Same as time just replace time with distance

Time-of-flight calculations for velocity

If question asks for one mole of that ion, how do u calculate it

Mass of that mole x Avogadro constant. This is in grams then x 1000 to find in kg.

What is the key features of mass spectra interpretation

-Isotopes produce different peaks on the spectrum because each isotope has a different mass.

-The peak height indicates the relative abundance of each isotope.

-The base peak is the tallest and represents the most abundant ion.

-For single-charged ions +1+1 ions), the m/z ratio corresponds directly to the ion's mass.

As energy levels increase…

The orbitals begin to overlap. For example, the 4s orbital is of lower energy than the 3d orbital, which is why electrons fill the 4s orbital before the 3d orbital.

What is the feature of electron configurations

-Electron configurations are arranged to minimise the overall energy of the atom or ion.

-This lowest energy arrangement corresponds to the most stable electronic structure.

Explain how transition metals have unusual electron configurations

Chromium (Cr) and copper (Cu) exhibit unusual electron configurations:
Cr: 1s2 2s2 2p6 3s2 3p6 3d5 4s1 (instead of 3d4 4s2)

Cu: 1s2 2s2 2p6 3s2 3p6 3d10 4s1 (rather than 3d9 4s2)

-These exceptions occur because configurations with a half-filled (d5) or fully filled (d10) d sub-shell are energetically more favourable.

-For transition metals forming ions, the loss of electrons happens from the 4s orbital before the 3d orbital.

What is first ionisation energy

The first ionisation energy is defined as the amount of energy required to remove one mole of electrons from one mole of gaseous atoms, forming one mole of gaseous ions with a charge of +1. It is measured in kilojoules per mole (kJ/mol).

What is the equation for first ionisation energy

X(g) → X+(g) + e−

Successive Ionisation Energies

Successive ionisation energies refer to the energy needed to remove more than one electron from an atom.

As more electrons are removed, the ionisation energy

increases

Explain why this occurs

As more electrons are removed, the ionisation energy increases because:

-The electron being removed is closer to the nucleus.

-There are fewer electrons to shield the attraction from the nucleus.

Atomic Radius

The greater the atomic radius, the further the outer electron is from the nucleus, requiring less energy to ionise.

Atomic radius across a period and down a group

Atomic radius decreases across a period and increases down a group.

Nuclear Charge

The more protons in the nucleus, the stronger the attraction between the nucleus and the outer electron, making it harder to remove an electron.

Nuclear charge across a period

The nuclear charge increases, but the number of electron shells remain the same

Shielding Effect

Inner electrons shield the outer electrons from the full attraction of the nucleus. The more inner electrons, the greater the shielding, reducing the energy required to remove outer electrons.

Shielding down a group

Shielding increases down a group, leading to lower ionisation energies.

Ionisation Energy Down a Group

As you move down a group, the first ionisation energy decreases. This is because:

-Atomic radius increases, meaning outer electrons are farther from the nucleus.

-Shielding by inner electrons increases, reducing the attraction from the nucleus.

Ionisation Energy Across a Period

As you move across a period from left to right, the first ionisation energy increases. This is due to:

-Decreasing atomic radius, with electrons being closer to the nucleus.

-Increasing nuclear charge, which pulls outer electrons closer.

-Shielding remains constant because electrons are added to the same energy level.

There are some zig zag patterns, explain why

-Group 3 elements (e.g., B, Al) show a drop in ionisation energy. This is because the added electron enters a higher-energy sub-shell (p-orbital), which is farther from the nucleus and more shielded.

-Group 6 elements (e.g., O, S) also show a drop due to electron pairing in the p-orbital, which causes repulsion, making it easier to remove the electron.

Explain how ionisation energy trend in Period 3 supports the idea of sub-shells:

The ionisation energy trend in Period 3 supports the idea of sub-shells:

-Sodium (Na) starts with a low first ionisation energy as it has a single electron in the 3s orbital.

-As you move across the period, ionisation energy gradually increases as more electrons are added to the 3p sub-shell, with increasing nuclear charge.

Explain how ionisation energy trend in group 2 supports the idea of sub-shells

-In Group 2 elements, successive ionisation energies give evidence of electron configurations in different shells:

-For example, in Beryllium (Be), removing the second electron requires significantly more energy than the first because it is being removed from a full 1s shell, closer to the nucleus.

As you move down Group 2, the first ionisation energy decreases due to:

-Increasing atomic radius.

-Greater shielding from inner electrons.