1: Atomic Structure and the Periodic Table

Atomic Structure and Periodic Table

Sub-atomic Particles

Atomic Representation

An atom of Lithium (Li) can be represented as follows: ^7_3Li

• Atomic Number (Z): The number of protons in the nucleus.

• Mass Number (A): The total number of protons and neutrons in the atom.

• Number of neutrons = A - Z

Isotopes

Isotopes are atoms with the same number of protons but different numbers of neutrons.

• DEFINITION: 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.

• DEFINITION: Relative atomic mass is the average mass of one atom compared to one-twelfth of the mass of one atom of carbon-12.

• DEFINITION: Relative molecular mass is the average mass of a molecule compared to one-twelfth of the mass of one atom of carbon-12.

Isotopes have similar chemical properties due to the same electronic structure but may have slightly varying physical properties because they have different masses.

Mass Spectrometer

The mass spectrometer is used to determine isotopes present in a sample and identify elements. The relative atomic mass on the periodic table is a weighted average of all the isotopes.

m/z (mass/charge ratio) and abundance are measured for each isotope.

If asked to give the species for a peak in a mass spectrum then give charge and mass number e.g. ^{24}Mg^+

Calculating Relative Atomic Mass (R.A.M)

R.A.M = { \text{Σ} (isotopic \text{ mass} \times \% \text{ abundance}) } / 100

Example for Mg:

R.A.M = [ (78.7 \times 24) + (10.13 \times 25) + (11.17 \times 26) ] / 100 = 24.3

Mass Spectra for Diatomic Molecules

Cl has two isotopes: ^{35}Cl (75%) and ^{37}Cl (25%)
Br has two isotopes: ^{79}Br (50%) and ^{81}Br (50%)

These lead to spectra caused by diatomic molecules like Cl2 and Br2.

Sometimes two electrons may be removed from a particle forming a 2+ ion. \text{e.g. } ^{24}Mg^{2+} with a 2+ charge would have a m/z of 12

If relative abundance is used instead of percentage abundance use this equation:

R.A.M = { \text{Σ} (isotopic \text{ mass} \times \text{relative abundance}) } / \text{total relative abundance}

Uses of Mass Spectrometers

• Included in planetary space probes to identify elements on other planets, which may have different isotopic compositions.

• Drug testing in sports to identify chemicals and their breakdown products in the blood.

• Quality control in the pharmaceutical industry and to identify molecules with potential biological activity.

• Radioactive dating to determine the age of fossils or human remains.

Measuring Mr of a Molecule

When a molecule goes through a mass spectrometer, it breaks into fragments. The peak with the largest m/z is due to the complete molecule and equals the Mr of the molecule. This peak is called the parent ion or molecular ion.

Example: Molecular ion C4H{10}^+

Ionisation Energies

Definition: First Ionisation Energy

The first ionisation energy is the energy required when one mole of gaseous atoms forms one mole of gaseous ions with a single positive charge.

H(g) \rightarrow H^+(g) + e^-

Definition: Second Ionisation Energy

The second ionisation energy is the energy required when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions with a double positive charge.

Ti^+(g) \rightarrow Ti^{2+}(g) + e^-

Factors Affecting Ionisation Energy

1. The attraction of the nucleus: More protons mean greater attraction.

2. The distance of the electrons from the nucleus: Bigger atoms have weaker attraction.

3. Shielding of the attraction of the nucleus: Inner electrons repel outer electrons, weakening attraction.

Successive Ionisation Energies

The patterns in successive ionisation energies provide information about electronic structure.

Example: Notice the big jump between the 4th and 5th ionisation energies. The fifth electron is in an inner shell closer to the nucleus and experiences greater attraction, with less shielding.

Why are successive ionisation energies always larger?

Each successive ionisation energy is larger because removing an electron creates a positive ion, which increases the attraction on the remaining electrons.

How are ionisation energies linked to electronic structure?

Example:

There is a big jump between the 2nd and 3rd ionisation energies, indicating that this element is in group 2 of the periodic table. The 3rd electron is removed from an electron shell closer to the nucleus, with less shielding, resulting in a larger ionisation energy.

First Ionisation Energy of the Elements

The shape of the graph for periods two and three is similar. A repeating pattern across a period is called periodicity. The pattern in the first ionisation energy gives us useful information about electronic structure.

Key trends and explanations:

• Helium (He) has the largest first ionisation energy: Its first electron is in the first shell, closest to the nucleus, with no shielding effects. It has one more proton than H.

• First ionisation energies decrease down a group: Outer electrons are in shells further from the nucleus and are more shielded, reducing the attraction of the nucleus.

• General increase in first ionisation energy across a period: The number of protons increases, increasing the effective attraction of the nucleus. Electrons are added to the same shell with similar shielding, pulling them closer to the nucleus.

• Na has a much lower first ionisation energy than Neon (Ne): Na's outer electron is in the 3s shell, further from the nucleus and more shielded, making it easier to remove.

• Small drop from Mg to Al: Al starts filling a 3p subshell, with electrons slightly easier to remove because they are higher in energy and slightly shielded by the 3s electrons.

• Small drop from P to S: Sulfur has 4 electrons in the 3p subshell, with the 4th electron starting to doubly fill the first 3p orbital, causing slight repulsion and making it easier to remove.

Electronic Structure

Models of the Atom

• Bohr Model (GCSE model): Electrons in spherical orbits (2 electrons in the first shell, 8 in the second, etc.).

• Electrons are arranged on Principle energy levels numbered 1,2,3,4.. with 1 closest to nucleus

• Sub energy levels labelled s , p, d and f.

  • s holds up to 2 electrons

  • p holds up to 6 electrons

  • d holds up to 10 electrons

  • f holds up to 14 electrons

Atoms and ions with noble gas electron arrangements should be stable.

Orbitals

Orbitals represent the mathematical probabilities of finding an electron at any point within certain spatial distributions around the nucleus. Each orbital has its own approximate, three-dimensional shape.

Energy order of subshells: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p

Writing Electronic Structure

For oxygen: 1s^2 2s^2 2p^4

Using Spin Diagrams

For fluorine:

Arrows represent electrons, with opposite directions representing different spins. Boxes represent one orbital.

• s sublevels are spherical

• p sublevels are shaped like dumbbells

When filling up sub levels with several orbitals, fill each orbital singly before starting to pair up the electrons

Electronic Structure for Ions

When a positive ion is formed electrons are lost

Mg \text{ is } 1s^2 2s^2 2p^6 3s^2 \text{ but } Mg^{2+} \text{ is } 1s^2 2s^2 2p^6

When a negative ion is formed electrons are gained

O \text{ is } 1s^2 2s^2 2p^4 \text{ but } O^{2-} \text{ is } 1s^2 2s^2 2p^6

PERIODICITY

Classification of elements in s, p, d blocks

Elements are classified as s, p or d block, according to which orbitals the highest energy electrons are in.

Atomic radius

Atomic radii decrease as you move from left to right across a period, because the increased number of protons create more positive charge attraction for electrons which are in the same shell with similar shielding.

1st ionisation energy

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 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.

Melting and boiling 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)(P_4=60) so has stronger London forces between molecules

• Ar is monoatomic weak London Forces between atoms

Exactly the same trend in period 2

Period 2 = Li, Be, B, C, N, O, F, Ne

Period 3 = Na, Mg, Al, Si, P, S, Cl, Ar

Exactly the same trend in period 2 with drops between Be & B and N to O for the same reasons- make sure change 3s and 3p to 2s and 2p in explanation!

Similar trend in period 2 Li,Be metallic bonding (high mp) B,C macromolecular (very high mp) N2,O2 molecular (gases! Low mp as small London Forces) Ne monoatomic gas (very low mp)

Periodicity is the repeating pattern of physical or chemical properties going across the periods