Atomic Structure

Section 1 of 22The structure of the atom

The structure of the atom

Atoms

• Atoms the smallest parts of an element that can take part in chemical reactions, and mostly empty space around a tiny dense nucleus

Nucleus

• Nucleus it contains protons and neutrons, called nucleons, and has an overall positive charge

Electrons

• Electrons negatively charged electrons are found in orbitals in the space around the nucleus

Model

• Model understanding of atomic structure has evolved with evidence from experiments

Section 2 of 22Subatomic particles

• Relative charges
  • actual charges are too small for coulombs, so they are compared with each other

Particle	Relative charge	Relative mass
Proton	+1	1
Neutron	0	1
Electron	-1	1/1836

Subatomic particles

Actual charge

• Actual charge an electron has -1.602 × 10^-19 C and a proton has +1.602 × 10^-19 C

Neutral atoms

• Neutral atoms the number of protons equals the number of electrons

Ions

• Ions an atom that loses electrons forms a positive ion, and one that gains electrons forms a negative ion

Section 3 of 22Atomic number and mass number

1Atomic (proton) number, Z

• Atomic (proton) number, Z the number of protons in the nucleus, equal to the number of electrons in an atom, and unique to each element

2Chemical change

• Chemical change electrons can be lost, gained or shared, but the proton number never changes

3Mass (nucleon) number, A

• Mass (nucleon) number, A the total number of protons and neutrons

4Location

• Location both numbers are on the Periodic Table

Section 4 of 22Calculating particles in atoms and ions

Calculating particles in atoms and ions

Protons

• Protons equal to the atomic number

Neutrons

• Neutrons mass number minus atomic number

Electrons in an atom

• Electrons in an atom equal to the atomic number

Electrons in an ion

• Electrons in an ion atomic number adjusted by the charge, because a 2+ ion has lost two electrons and a 1- ion has gained one

Species	Protons	Neutrons	Electrons
Mg	12	12	12
O2-	8	8	10
Al3+	13	14	10

Section 5 of 22Isotopes

Isotopes = atoms of the same element with the same number of protons and electrons but different numbers of neutrons

Notation

• Notation the symbol or name then the mass number, such as carbon-12 and carbon-14 with 6 and 8 neutrons

Hydrogen

• Hydrogen its isotopes have the same protons but different neutrons

Section 6 of 22Time of flight mass spectrometry: purpose and stages

1Purpose

• Purpose to find the relative atomic mass of an element from the abundance and mass of its isotopes, to identify it,
  • and to find relative molecular mass

2Common principle

• Common principle the sample is ionised, accelerated, and the ions are separated by their charge-to-mass ratio

3TOF stages

• TOF stages ionisation, acceleration, ion drift and detection

Section 7 of 22Time of flight: ionisation and acceleration

Time of flight: ionisation and acceleration

Ionisation

• The sample is dissolved in a volatile solvent and injected through a fine hollow needle to make a mist.
• The needle is on a high-voltage supply, so particles lose electrons: X(g) → X⁺(g) + e^-.
• Most ions have a 1+ charge because removing further electrons is difficult.
• The solvent evaporates, leaving only positive ions.

Acceleration

• Positive ions are attracted to a negatively charged plate and accelerate in an electric field.
• All ions gain the same kinetic energy, so their speed depends on mass.
• Lighter ions move faster and heavier ions move slower.

Section 8 of 22Time of flight: ion drift and detection

Ion drift

• The ions pass through a hole in the plate into a flight tube.
• The time of flight depends on velocity.

Detection

• Ions hit a negatively charged detector plate and gain an electron, producing a current.
• The size of the current is proportional to the abundance of the ions.
• Lighter ions arrive first and heavier ions arrive later.
• A computer converts the signal into the mass spectrum.

Key point

• Key point all ions have the same kinetic energy, so separation depends only on mass

Section 9 of 22Interpreting mass spectra

Relative abundance

• Relative abundance the proportion of one isotope in the natural mixture, such as Cl-35 at 75% and Cl-37 at 25%

Peaks

• Peaks their heights show abundance, and the horizontal axis is m/z, which equals the mass number for 1+ ions

Boron

• Boron two peaks show two isotopes, B-10 at 19.9% and B-11 at 80.1%

Interpreting mass spectra

Multiply each isotope’s abundance by its m/z.

Add the products.

Divide by 100.

Interpreting mass spectra

Result

• A_r of boron = ÷ 100 = 10.8
  • (19.9 × 10 + 80.1 × 11)

Types

• Types low resolution mass spectrometry gives relative atomic mass, and high resolution gives relative molecular mass of a compound

Section 10 of 22Electron energy levels

Electronic configuration

• Electronic configuration the arrangement of electrons in energy levels (shells) around the nucleus

Principal quantum number, n

• Principal quantum number, n it numbers the shells from 1 to 4, and the lower the n the closer to the nucleus and the lower the energy

Shell (n)	Maximum electrons
1	2
2	8
3	18
4	32

Spacing

• Spacing energy levels get closer together further from the nucleus, with an overlap between the third and fourth levels

Sub-shells

• Sub-shells shells divide into sub-shells s, p, d
  • (and f for elements with more than 57 electrons)

Energy order

• Energy order s < p < d < f

Orbital

• Orbital a volume of space with a high probability of finding an electron

Content by level

• Content by level the first level has one s orbital, and the second level has an s and p orbitals

Section 11 of 22Orbitals in each shell

• Capacity
  • each atomic orbital holds a maximum of two electrons

Shell (n)	Sub-shell	Orbitals	Electrons in sub-shell	Electrons in shell
1	s	1	2	2
2	s	1	2	8
2	p	3	6	8
3	s	1	2	18
3	p	3	6	18
3	d	5	10	18
4	s	1	2	32
4	p	3	6	32
4	d	5	10	32
4	f	7	14	32

• Exam requirement
  • know the number of s, p and d orbitals in each level, and the shapes of s and p orbitals only

Section 12 of 22Orbital shapes

s orbital

• s orbital spherical
• Its size increases with shell number, so 3s is bigger than 1s.

p orbital

• p orbital dumbbell-shaped
• Every shell except n = 1 has three p orbitals.
• They lie along the x, y and z axes at right angles to each other.
• The lobes become larger and longer with increasing shell number.

Section 13 of 22Rules for filling orbitals

1Lowest energy first

• Lowest energy first electrons fill orbitals of lower energy before higher ones

2Maximum two

• Maximum two each orbital holds at most two electrons, with opposite spins

3Singly first

• Singly first orbitals of equal energy, such as the three 2p orbitals, fill singly before pairing

44s before 3d

• 4s before 3d the 4s orbital is lower in energy than 3d, because of the overlap between the third and fourth levels, so it fills first

Section 14 of 22Box notation

Boxes

• Boxes each box represents an atomic orbital, arranged in order of increasing energy from bottom to top

Electrons

• Electrons shown as opposite arrows to represent opposite spins

Examples

• Examples nitrogen, sodium, chlorine and titanium can all be drawn this way

Section 15 of 22Writing electronic configurations

Notation

• Notation each sub-shell is written as number, letter and a superscript giving the electrons, such as hydrogen 1s¹

Potassium

• 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹

Shorthand

• Shorthand use the nearest preceding noble gas in square brackets, so potassium is [Ar] 4s¹

Calcium

• 1s² 2s² 2p⁶ 3s² 3p⁶ 4s², or [Ar] 4s²

Gallium

• 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p¹, or [Ar] 3d¹⁰ 4s² 4p¹
• Configurations are often written in numerical order, so 3d is written before 4s even though 4s fills first.

Ions

• Ions remove electrons from the outermost shell
• Mg²⁺ loses its two 3s electrons, so it is 1s² 2s² 2p⁶ or [Ne].

Section 16 of 22Blocks of the Periodic Table

s block

• s block outer electrons in an s orbital

p block

• p block outer electrons in a p orbital

d block

• d block outer electrons in a d orbital

f block

• f block outer electrons in an f orbital

Section 17 of 22Ionisation energy and successive ionisation energies

First ionisation energy = the energy needed to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions

Conditions and units

• Conditions and units measured at 298 K and 101 kPa in kJ mol^-1, and always positive because it is endothermic

Equations

• Equations first: Na(g) → Na⁺(g) + e^-, second: Na⁺(g) → Na²⁺(g) + e^-, third: Na²⁺(g) → Na³⁺(g) + e^-

Starting species

• Starting species the second ionisation energy of magnesium starts from Mg⁺(g)

Increase

• Increase each successive ionisation energy is larger, because electrons are removed from increasingly positive ions

Sodium data (kJ mol⁻¹)

• 496, 4563, 6913, 9544, 13352, 16611, 20115, 25491, 28934, 141367, 159079

Section 18 of 22Using ionisation energies to find shells and group

Energy levels in sodium

• Energy levels in sodium the first electron is easy to remove, the next 8 are harder, and the last 2 closest to the nucleus are hardest

Group

• Group the largest jump between successive ionisation energies shows the group, as a jump between 1st and 2nd means Group 1

Reason

• Reason the next electron comes from an inner shell closer to the nucleus

Section 19 of 22Ionisation energy across Period 3

Trend

• Trend the first ionisation energy generally increases across the period

Nuclear charge

• Nuclear charge it increases, so the outer shell is pulled closer and atomic radius decreases

Shielding

• Shielding it stays nearly constant because electrons are added to the same shell

Result

• Result it is harder to remove an electron, so more energy is needed

Aluminium lower than magnesium

• Magnesium loses a 3s electron.
• Aluminium loses a 3p electron, which is slightly higher in energy.
• This is evidence for s and p sub-shells.

Sulfur lower than phosphorus

• Phosphorus has one electron in each of three 3p orbitals.
• Sulfur has a pair of electrons in one 3p orbital.
• Repulsion between the paired electrons makes one easier to remove.

Section 20 of 22Ionisation energy down Group 2

Ionisation energy down Group 2

Trend

• Trend the first ionisation energy decreases down the group

Reason

• Reason atomic radius increases and shielding by inner shells increases, outweighing the increase in nuclear charge

Result

• Result the outer electron is easier to remove

Exam skills

• Exam skills write ionisation equations with state symbols, explain the trends, and explain deviations as evidence for shells and sub-shells

Section 21 of 22Key terms

Term	Meaning
Nucleon	A proton or a neutron in the nucleus.
Isotopes	Atoms of the same element with different numbers of neutrons.
Relative isotopic abundance	The proportion of an isotope in a natural mixture.
m/z	The mass-to-charge ratio of an ion.
Principal quantum number	The number n that labels an electron shell.
Orbital	A region of space where an electron is likely to be found.

Term	Meaning
Ionisation energy	The energy to remove one mole of electrons from one mole of gaseous atoms.
Sub-shell	A division of a shell into s, p, d or f orbitals.
Shielding	The repulsion of outer electrons by inner-shell electrons.
Successive ionisation energies	The energies to remove electrons one after another from an atom.
Shorthand configuration	A configuration written using the preceding noble gas symbol.

Section 22 of 22Exam technique

1

Saying the electron has a mass of 1 — its relative mass is 1/1836.

2

Changing the proton number to form an ion — ions form by changing electrons only.

3

Saying all ions have the same speed in TOF — they have the same kinetic energy but different speeds.

4

Forgetting to divide by 100 in relative atomic mass — divide the weighted sum by 100.

1

Writing the second ionisation equation from the atom — start from the 1+ ion.

2

Forgetting state symbols — ionisation equations need (g).

3

Saying aluminium has a lower IE than magnesium because of smaller nuclear charge — it loses a higher-energy 3p electron.

4

Filling 3d before 4s — 4s fills first.