cie A level chemistry atomic structure

What are the relative charge and relative mass (a.m.u) for a Proton (P)?

Relative Charge: +1
Relative Mass/a.m.u: 1

What are the relative charge and relative mass (a.m.u) for a Neutron (n)?

Relative Charge: 0
Relative Mass/a.m.u: 1

What are the relative charge and relative mass (a.m.u) for an Electron (e-)?

Relative Charge: -1
Relative Mass/a.m.u: 1/1840

Where is the mass of an atom concentrated?

Mass is concentrated within the centre, known as the nucleus.

What are the relative charge and relative mass (a.m.u) for a Proton (P)?

Relative Charge: +1
Relative Mass/a.m.u: 1

What are the relative charge and relative mass (a.m.u) for a Neutron (n)?

Relative Charge: 0
Relative Mass/a.m.u: 1

What are the relative charge and relative mass (a.m.u) for an Electron (e-)?

Relative Charge: -1
Relative Mass/a.m.u: 1/1840

Where is the mass of an atom concentrated?

Mass is concentrated within the centre, known as the nucleus.

Why is an atom considered electrically neutral?

An atom is electrically neutral because the number of protons (P+) is equal to the number of electrons (e-).

Define Atomic number (Z) or proton number.

The Atomic number (Z) or proton number is the number of protons in an atom.

Define Atomic mass (A) or nucleon number.

The Atomic mass (A) or nucleon number is the total number of protons (P) and neutrons (N) in an atom.

What are Isoelectronic Ions?

Isoelectronic Ions are ions having the same number of electrons (e-s).

Define Isotopes.

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

Why do Isotopes have similar chemical properties?

Isotopes have similar chemical properties because they have the same number of protons and electrons, leading to similar chemical interactions.

Why do Isotopes have different physical properties?

Isotopes have different physical properties because they have different numbers of neutrons, causing them to have different masses and, therefore, different physical interactions.

Describe the behaviour of a beam of Protons (P+) in an electric field.

Protons are positively charged, so they are deflected towards the negative -ve pole.

Describe the behaviour of a beam of Neutrons (n) in an electric field.

Neutrons have no charge, so they are not deflected by an electric field.

Describe the behaviour of a beam of Electrons (e-) in an electric field.

Electrons are negatively charged, so they are deflected towards the positive +ve pole.

Why are electrons deflected at a greater angle than protons in an electric field?

Electrons are lighter than protons, so they are deflected at a greater angle.

1. What are energy levels for electrons called?
2. What describes each shell?

1. Electrons are arranged in energy levels called shells.
2. Each shell is described by a principle quantum number (P.Q.).

How does the energy of a shell change as the Principal Quantum Number (P.Q.) increases?

As the P.Q. increases, the energy of the shell increases.

Name the subshells found inside the electron shells.

The subshells are s, p, d, and f.

What is an Orbital?

An orbital is a region in space where there is a maximum probability of finding an electron.

How many electrons can each orbital hold and in what manner?

Each orbital can hold 2 electrons (2e-s) in opposite directions (spins).

Explain Hund's Rule for placing electrons in orbitals of equal energy.

When electrons (e-s) are placed in a set of orbitals of equal energy, they occupy them singly first, and then pairing takes place.

Why are electrons placed in opposite directions within an orbital?

Electrons are placed in opposite directions, creating a spin, to reduce repulsion, as both have negative charges; if placed in the same direction, they would repel each other.

What kind of electron configurations are considered more stable?

Completely filled or half-filled electron configurations (i.e., one electron in each orbital) are more stable due to reduced repulsion.

Explain the electron filling order between the 4s and 3d orbitals.

Electrons would prefer the 4s orbital over 3d while filling up (e.g., 2 electrons in Titanium fill 4s before 3d) because the 4s orbital is a more stable (lower) energy level than the 3d orbital.

Explain the electron losing order between the 4s and 3d orbitals.

While losing electrons, electrons from the 4s orbital would be lost first, and then those from the 3d orbital, because the 4s orbital is outer than the 3d orbital.

For each subshell (s, p, d, f), state the number of orbitals and maximum electrons.

1. s:
   • Orbitals: 1
   • Max e-s: 2
2. p:
   • Orbitals: 3
   • Max e-s: 6
3. d:
   • Orbitals: 5
   • Max e-s: 10
4. f:
   • Orbitals: 7
   • Max e-s: 14

What is Aufbau's Principle?

Aufbau's Principle is a method of showing how atomic orbitals are filled in a definite order to give the lowest energy arrangement possible.

What happens due to the small energy difference between 4s and 3d orbitals?

The energy difference between 4s and 3d is very small, so an electron from 4s can be promoted to half-fill or full-fill a 3d orbital to make the atom more stable.

Describe the shapes of s and p orbitals.

s orbitals are spherical, with the nucleus at the centre. p orbitals are dumbbell-shaped.

What is a free radical?

A free radical is a species with one or more unpaired electrons.

Define 1st Ionisation Energy (I.E).

The 1st I.E. is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous unipositive ions.

Why is each successive Ionisation Energy higher than the previous one?

Each successive I.E. is higher because as electrons are removed, the number of protons becomes greater than the number of remaining electrons, which increases the attraction between the protons and the remaining electrons.

What indicates a large jump in successive Ionisation Energies?

Successive I.E.s have a large jump in their value when electrons are removed from a lower energy shell.

How can the group number of an element be deduced from Ionisation Energies?

The group number can be deduced by checking when the first big jump in successive Ionisation Energies occurs.

How does Nuclear Charge affect Ionisation Energy?

Greater nuclear charge (positive charge due to protons in the nucleus) means greater Ionisation Energy.

How does the Shielding Effect affect Ionisation Energy?

Inner shells of electrons repel outermost electrons, shielding them from the positive nucleus. The more electron shells, the greater the shielding effect, which leads to a lower I.E. because there is less attractive force between the nucleus and outer electrons.

How does Atomic Radius affect Ionisation Energy?

As the atomic radius increases, the Ionisation Energy decreases. This is because the distance of the outermost electron to the nucleus is large, so less energy is needed to remove that electron.

How does a Stable Configuration affect Ionisation Energy?

A high I.E. is needed to remove electrons from completely filled or half-filled orbitals because these configurations are more stable.

Describe the trend of 1st Ionisation Energy down a group.

Down a group, the 1st I.E. generally decreases because:

• New shells are added.
• The attraction of the nucleus to valence electrons decreases.
• The shielding effect increases.

Describe the trend of 1st Ionisation Energy across a period.

Across a period, the 1st I.E. generally increases because:

• The shell number remains the same.
• The proton number increases.
• The effective nuclear charge increases.
• The atomic radius decreases.

Why is the 1st I.E. of Aluminium (Al) lower than Magnesium (Mg)?

The 1st I.E. of Al is lower than Mg because the electron removed in Al is from a higher energy 3p orbital, which is further away from the nucleus than the 3s electron being removed from Mg. Nuclear attraction is less for 3p than 3s, so the I.E. of Al is lower than Mg.

Why is the 1st I.E. of Sulfur (S) lower than Phosphorus (P)?

The 1st I.E. of S is lower than P because the electron being removed in P is in a half-filled, more stable 3p orbital. In S, the pairing of electrons in the 3p orbital results in increased repulsion, so less energy is needed to remove an electron.

What is Ionic Radius?

Ionic Radius describes the size of an ion.

Why is a positive ion (cation) smaller than its original neutral atom?

A positive ion has a smaller radius than its original neutral atom because:

• The shell number decreases.
• The screening effect decreases.
• The attraction of the nucleus to the remaining electrons increases.

Why is a negative ion (anion) larger than its original neutral atom?

A negative ion has a larger ionic radius than its neutral atom because electrons are added, while the nuclear charge remains the same, leading to increased electron-electron repulsion and a larger electron cloud.

Describe the trend of ionic radius across a period.

Across a period, the proton number and effective nuclear charge increase, which causes the ionic radius to decrease.

Compare the size of negative ions to positive ions in the same period.

Negative ions are always larger than positive ions in the same period, as they typically have one more shell or a greater number of electrons for the same nuclear charge.

Describe the trend of ionic radius down a group.

Ionic radius increases down the group since the number of electron shells increases.

How does the negative charge on an anion affect its ionic radius?

As the negative charge on an anion increases, the ionic radius increases because the number of electrons gained increases such that the number of electrons exceeds the number of protons, leading to greater electron-electron repulsion.

How does the positive charge on a cation affect its ionic radius?

As the positive charge on the cation increases, the number of electrons lost increases, so the electrostatic attraction between