Topic 1 Chem Edexcel

Question-and-answer flashcards covering sub-atomic particles, isotopes, mass spectrometry, ionisation energies, electronic structure, and periodic trends.

Which sub-atomic particle is located in the nucleus, has a relative mass of 1 and a charge of +1?

The proton.

Which sub-atomic particle is found in the nucleus, has a relative mass of 1 and zero charge?

The neutron.

Which sub-atomic particle occupies orbitals, has a relative mass of 1⁄1840 and a charge of –1?

The electron.

How do you calculate the number of neutrons in an atom?

Number of neutrons = mass number (A) – atomic number (Z).

Define an isotope.

Atoms of the same element with the same number of protons but different numbers of neutrons.

Why do isotopes of the same element have virtually identical chemical properties?

Because they have the same electronic structure.

State the definition of relative isotopic mass.

The mass of one atom of an isotope compared with one-twelfth of the mass of one atom of carbon-12.

State the definition of relative atomic mass (Ar).

The average mass of one atom of an element compared with one-twelfth of the mass of one atom of carbon-12.

State the definition of relative molecular mass (Mr).

The average mass of a molecule compared with one-twelfth of the mass of one atom of carbon-12.

What analytical instrument can be used to determine all the isotopes in a sample?

A mass spectrometer.

When quoting the species responsible for a peak in a mass spectrum, what must be included?

Both the mass number and the charge, e.g. 24Mg⁺.

Write the formula used to calculate relative atomic mass from percentage abundances.

Ar = (isotopic mass × % abundance) ÷ 100.

Explain why a 24Mg²⁺ ion gives an m⁄z value of 12 in the mass spectrometer.

Its mass is 24 and its charge is +2, so m⁄z = 24⁄2 = 12.

List two practical uses of mass spectrometry outside the laboratory.

Examples: drug testing in sport; identifying elements on planetary space probes; quality control in pharmaceuticals; radioactive dating of fossils.

In a molecular mass spectrum, what is the peak with the largest m⁄z normally called, and what does it represent?

The parent (or molecular) ion peak; it represents the Mr of the whole molecule.

Give the definition of first ionisation energy.

The energy required when one mole of gaseous atoms forms one mole of gaseous ions each with a single positive charge.

Write the general equation that represents a first ionisation energy.

X(g) → X⁺(g) + e⁻

Give the definition of second ionisation energy.

The energy required when one mole of gaseous ions with a single positive charge forms one mole of gaseous ions each with a double positive charge.

State three factors that determine the magnitude of ionisation energy.

Nuclear charge (number of protons), distance of the outer electrons from the nucleus, and shielding by inner-shell electrons.

Why is the second ionisation energy of an element always higher than the first?

After the first electron is removed, the remaining electrons experience a greater attraction to the now more positive ion, so more energy is needed to remove the next electron.

How can large jumps in successive ionisation energies be used to deduce an element’s group number?

A big jump occurs when an electron is removed from a new inner shell, so the number of electrons removed before the jump equals the group number.

Why does helium have the highest first ionisation energy of all elements?

Its electrons are in the first shell (closest to the nucleus) with no inner-shell shielding, and it has a high nuclear charge for that shell.

Explain why first ionisation energy decreases down a group.

Outer electrons are in higher shells further from the nucleus and are more shielded, reducing nuclear attraction.

Explain the general increase in first ionisation energy across a period.

Proton number increases while electrons are added to the same shell with similar shielding, so nuclear attraction increases.

Why is the first ionisation energy of Na much lower than that of Ne?

Na’s outer electron is in the 3s subshell, which is higher in energy and further from the nucleus than Ne’s valence electrons, and is more shielded.

Explain the small drop in first ionisation energy from Mg to Al.

Al’s outer electron enters the 3p subshell, which is slightly higher in energy and more shielded than the 3s subshell of Mg, making it easier to remove.

Explain the small drop in first ionisation energy from P to S.

The paired electrons in one 3p orbital in sulfur repel each other slightly, so the paired electron is easier to remove than an unpaired 3p electron in phosphorus.

In the A-level model, what is the maximum number of electrons an s-subshell can hold?

Two electrons.

How many electrons can a p-subshell hold in total?

Six electrons.

State the order in which the first eight subshells are filled according to increasing energy.

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p.

Describe the general shape of an s orbital.

Spherical.

Describe the general shape of a p orbital.

Dumb-bell shaped (two lobes on opposite sides of the nucleus).

According to Hund’s rule, how are electrons arranged in degenerate (equal-energy) orbitals?

Each orbital is filled singly before any pairing occurs, with electrons having parallel spins.

Write the full electron configuration of an O²⁻ ion.

1s² 2s² 2p⁶.

When forming positive ions in the 3d block, which subshell loses electrons first: 4s or 3d?

The 4s subshell loses electrons before 3d because it is higher in energy once filled.

What distinguishes an s-block element from a p-block element?

An s-block element has its highest-energy electron in an s subshell; a p-block element’s highest-energy electron is in a p subshell.

Define periodicity.

The repeating pattern of physical or chemical properties across different periods of the periodic table.

Across a period, why do atomic radii decrease?

Increasing nuclear charge pulls electrons in the same shell closer to the nucleus; shielding remains similar.

Across period 3, why do melting points rise from Na to Al?

Metallic bonding becomes stronger as charge density of the metal ions increases and more delocalised electrons are available.

Why does silicon have a very high melting point?

It forms a giant covalent (macromolecular) lattice with many strong covalent bonds that require much energy to break.

Why do Cl₂, S₈ and P₄ have much lower melting points than Si?

They are simple molecular substances held together by weak London forces, which require little energy to overcome.

Why does S₈ have a higher melting point than P₄?

S₈ has more electrons and a larger surface area, giving it stronger London dispersion forces.

Give two reasons metallic bonding is stronger in Al than in Na.

Al³⁺ ions have a higher charge and smaller radius than Na⁺ ions, and release more delocalised electrons to the metallic lattice.

What is the relationship between m⁄z and the number of electrons removed from an ion?

m⁄z equals the ion’s mass divided by its charge; removing more electrons (higher charge) lowers the m⁄z value.