Electrons, Bonding and Structure

Section 1 of 19Shells and subshells

Electron configuration

Electron configuration the arrangement of electrons in the shells, subshells and orbitals of an atom

Principal quantum shells

• Principal quantum shells numbered by the principal quantum number n
• A lower n means a shell closer to the nucleus and lower in energy; n = 1 is closest.
• The shells hold up to 2 (n = 1), 8 (n = 2), 18 (n = 3) and 32 (n = 4) electrons.

Subshells

• Subshells each shell is split into subshells labelled s, p and d
  • (and f for elements with more than 57 electrons)
• Energy increases in the order s < p < d.
• The subshell order overlaps for higher shells, so it does not follow a regular pattern at n = 3 and above.

Section 2 of 19Atomic orbitals

• Orbital = a region of probability for finding an electron, existing at a specific energy
  • Each orbital holds a maximum of two electrons with opposite spins.
  • This replaces the earlier Bohr-like model.

Subshell	Orbitals	Maximum electrons
s	1	2
p	3	6
d	5	10
f	7	14

Section 3 of 19Orbital shapes

s orbital

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

p orbital

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

d orbitals

• d orbitals their shape is not required

Section 4 of 19Filling orbitals

1Electron spin

• Electron spin an electron behaves as a tiny spinning charge, clockwise or anticlockwise, creating a tiny magnetic field
• Spin is shown by an arrow pointing up or down.

2Hund's rule

• Hund's rule electrons occupy separate orbitals of the same subshell first, with parallel spins, and only then pair up
• This minimises repulsion between electrons.
• Three electrons in a p subshell go one each into p_x, p_y and p_z.

3Pauli exclusion principle

• Pauli exclusion principle an orbital holds at most two electrons and they must have opposite spins

4Energy of electrons

• Energy of electrons it is set by the principal quantum number
• A 2p electron is in the second shell and has energy corresponding to n = 2.
• Electrons pair up in lower-energy orbitals because promoting one to a higher empty orbital costs more energy than the repulsion saved.

Section 5 of 19Ground state and orbital box diagrams

Ground state

Ground state the most stable, lowest-energy electron configuration of an atom

• It is reached by filling the lowest-energy subshells first, starting with 1s.

Orbital box diagram

• Orbital box diagram each box is an orbital, arranged from lowest energy upwards
• Electrons are drawn as arrows, with opposite directions in the same box.
• Boxes must be labelled and carry the correct number of electrons.

Section 6 of 19Writing configurations of atoms and ions

Full configuration

• Full configuration lists every subshell from 1s upwards
• Potassium (19 electrons) is 1s² 2s² 2p⁶ 3s² 3p⁶ 4s¹.

Shorthand configuration

• Shorthand configuration uses the symbol of the nearest preceding noble gas in brackets, then the remaining electrons
• Potassium is [Ar] 4s¹, calcium is [Ar] 4s² and gallium is [Ar] 3d¹⁰ 4s² 4p¹.

Ions

• Ions negative ions gain electrons in the outer subshell and positive ions lose them from the outer subshell
• Ca²⁺ is 1s² 2s² 2p⁶ 3s² 3p⁶, identical to argon.
• Transition metals fill 4s before 3d but also lose electrons from 4s before 3d.

Blocks

• Blocks the Periodic Table is divided by the subshell holding the outer electrons
• s block, p block, d block and f block.

Section 7 of 19Exceptions: chromium and copper

Chromium

• [Ar] 3d⁵ 4s¹, not [Ar] 3d⁴ 4s²

Copper

• [Ar] 3d¹⁰ 4s¹, not [Ar] 3d⁹ 4s²

Reason

• Reason promoting one electron from 4s to 3d gives a more stable half-full (3d⁵) or full (3d¹⁰) d subshell

Section 8 of 19Forming ions

Ionic bonding

• Ionic bonding electrons are transferred from a metal to a non-metal
• The ions formed have full outer shells and noble gas configurations.

Cations

• Cations metals lose outer electrons to form positive ions
• Li⁺ has the configuration of He, Ca²⁺ that of Ar and Al³⁺ that of Ne.

Anions

• Anions non-metals gain electrons into the outer shell to form negative ions
• O²⁻ and N³⁻ have the configuration of Ne and Cl⁻ that of Ar.

Section 9 of 19The ionic bond and giant ionic lattice

Ionic bond

Ionic bond the electrostatic attraction between oppositely charged ions

• It acts in all directions, so it is non-directional.
• It is very strong and needs a lot of energy to overcome, giving ionic compounds high melting points.

Giant ionic lattice

• Giant ionic lattice a 3D network of oppositely charged ions
• It is held together by strong electrostatic attractions between ions in all directions.
• Sodium chloride forms when electrons transfer from Na to Cl, giving Na⁺ and Cl⁻.

Section 10 of 19Structure of the ionic lattice

• Lattice
  • a regular repeating pattern of alternating cations and anions
  • The arrangement depends on the relative sizes of the ions.
  • Many common lattices, such as NaCl and MgO, are cubic.

Section 11 of 19Melting and boiling points of ionic compounds

State

• State most ionic compounds are solids at room temperature
• The strong electrostatic attractions between oppositely charged ions need large amounts of energy to overcome.

Effect of charge

• Effect of charge bond strength increases with ionic charge, so lattices with more highly charged ions have higher melting and boiling points
• Na₂O melts at 1405 K but CaO melts at 2845 K.
• Size is not the factor here because Na⁺ and Ca²⁺ are similar in size.

Section 12 of 19Solubility of ionic compounds

Dissolving

• Dissolving many ionic compounds dissolve in polar solvents such as water
• The polar molecules break down the lattice and surround each ion in solution.
• The δ+ end of the molecule surrounds anions and the δ− end surrounds cations.

Balance

• Balance solubility depends on ionic bonding in the lattice against attraction between ions and the polar solvent

General trend

• General trend the greater the ionic charge, the less soluble the compound, though there are many exceptions
• 356.9 g of NaCl dissolve in 1 dm³ of water but only 74.4 g of calcium chloride.

Section 13 of 19Electrical conductivity of ionic compounds

Solid

• Solid does not conduct because the ions are in fixed positions, so there are no mobile charge carriers

Molten or aqueous

• Molten or aqueous conducts because the lattice has broken down
• The ions are free to move and carry charge.

Section 14 of 19Covalent bonds and bond strength

• Covalent bond = the electrostatic attraction between the nuclei of bonding atoms and a shared pair of outer-shell electrons
  • It occurs between two non-metals and gives each atom a noble gas configuration without electron transfer.

Bond	Representation	Shared electrons
Single	C–C	2
Double	C=C	4
Triple	C≡C	6

Bond energy

Bond energy the energy required to break one mole of covalent bonds in the gaseous state, in kJ mol⁻¹

• The larger the bond energy, the stronger the bond.

Average bond enthalpy

Average bond enthalpy the average energy to break a bond type, measured across many compounds

• It may differ from the value for a specific compound, and sources vary slightly.

Multiple bonds

• Multiple bonds double and triple bonds have higher bond enthalpies than single bonds because of greater electron density between the nuclei
• Stronger bonds are shorter, so triple bonds are the shortest and strongest.

Section 15 of 19Dot-and-cross diagrams

Content

• Content they show only the outer-shell electrons
• Dots represent the electrons of one atom and crosses those of the other.
• Electrons are shown in pairs.

Single bonds

• Single bonds one shared pair, e.g. Cl₂, HCl and NH₃

Double bonds

• Double bonds two shared pairs, e.g. O₂, CO₂ and C₂H₄

Triple bonds

• Triple bonds three shared pairs, e.g. N₂

Outer electrons

• Outer electrons use the Periodic Table group to count the outer electrons of each non-metal atom

Section 16 of 19Exceptions to the octet rule

Expanded octet

Expanded octet a central atom holding more than 8 electrons in its outer shell

• It occurs with period 3 elements, e.g. SO₂ and PCl₅.

Electron deficient

Electron deficient a central atom holding fewer than 8 electrons in its outer shell

• BCl₃ has only 6 electrons around boron.

Section 17 of 19Dative covalent (coordinate) bonding

Dative covalent bond

Dative covalent bond a covalent bond in which both shared electrons come from the same atom

• It forms when one atom has a lone pair and the other has an incomplete outer shell.
• It is drawn as an arrow from the lone pair donor to the acceptor.

Ammonium ion

• Ammonium ion the lone pair on nitrogen in NH₃ is donated to an H⁺ ion, which has no electrons, forming NH₄⁺

Aluminium chloride dimer

• AlCl₃ is electron deficient and exists as such at high temperature
• At lower temperature two AlCl₃ molecules join to form Al₂Cl₆.
• Two chlorine atoms each donate a lone pair to an aluminium atom, forming two dative bonds.

Section 18 of 19Key terms

Term	Meaning
Principal quantum number	The number n that labels each shell, with higher n further from the nucleus
Orbital	A region of probability holding up to two electrons of opposite spin
Hund rule	Electrons occupy separate orbitals of a subshell with parallel spin before pairing
Pauli exclusion principle	An orbital holds a maximum of two electrons with opposite spins
Ground state	The lowest-energy electron configuration of an atom
Ionic bond	Electrostatic attraction between oppositely charged ions

Term	Meaning
Covalent bond	Electrostatic attraction between nuclei and a shared pair of outer electrons
Bond energy	Energy required to break one mole of covalent bonds in the gaseous state
Dative covalent bond	A covalent bond in which both electrons come from one atom
Expanded octet	More than eight electrons in the outer shell of a central atom
Electron deficient	Fewer than eight electrons in the outer shell of a central atom
Giant ionic lattice	A regular 3D network of alternating cations and anions

Section 19 of 19Exam technique

1

Writing Cr as [Ar] 3d⁴ 4s² — it is [Ar] 3d⁵ 4s¹ for a half-full d subshell

2

Removing 3d electrons first when forming transition metal ions — 4s electrons are lost first

3

Pairing electrons in a p subshell before each orbital has one — apply Hund rule

4

Drawing both arrows in one box in the same direction — they must be opposite

1

Saying ionic compounds conduct when solid — ions are fixed in the lattice, so they only conduct when molten or aqueous

2

Stating that ionic compounds melt high because of "ionic molecules" — they have a giant lattice with strong electrostatic attractions

3

Showing the dative bond arrow pointing from acceptor to donor — it points from the lone pair donor

4

Using bond energy for a single compound as the average bond enthalpy — averages are taken across many compounds