The covalent model

A **shared pair of electrons** between two (usually non-metal) atoms.

A **non-bonding** pair of electrons that stays on one atom (drawn as two dots).

A **bonding pair** (one shared pair of electrons).

Atoms tend to gain a full outer shell of **8 electrons** by sharing (or transferring) electrons.

Number of **shared pairs**: 1, 2, 3 — bond order 1, 2, 3. Higher order → shorter, stronger.

O=C=O — **two double bonds**, two lone pairs on each oxygen, none on carbon.

N≡N — a **triple bond** with **one lone pair on each** nitrogen.

**BF_{3}** (boron has 6 electrons) and **BeCl_{2}** (beryllium has 4) — electron-deficient.

Count valence electrons → least electronegative atom central → single bonds → complete outer octets → multiple bonds if the centre is short.

**One** (three bonding pairs to H, one lone pair).

**V**alence **S**hell **E**lectron **P**air **R**epulsion.

Any group of electrons around the central atom — a single/double/triple **bond (each = 1 domain)** or a **lone pair**.

**Linear**, 180° (e.g. CO_{2}, HCN).

**Trigonal planar**, 120° (e.g. BF_{3}).

**Tetrahedral**, 109.5° (e.g. CH_{4}).

**Trigonal pyramidal**, ~107° (e.g. NH_{3}).

**Bent**, ~104.5° (e.g. H_{2}O).

Lone pairs repel **more** than bonding pairs, so they **reduce** the bond angle.

Each double bond is **one** electron domain; 2 domains, 0 lone pairs → linear, 180°.

CH_{4} (109.5°) > NH_{3} (107°) > H_{2}O (104.5°) — angle falls as lone pairs increase.

A measure of how strongly an atom **attracts the shared (bonding) electrons** in a covalent bond.

A **difference in electronegativity** between the two atoms — the electrons are pulled towards the more electronegative atom.

The **more electronegative** atom (it gets a bigger share of the electrons); the less electronegative atom is **δ+**.

Δχ = 0 → **pure covalent**; small Δχ → **polar covalent** (δ+/δ−); large Δχ → **ionic**.

The small separation of charge (δ+ → δ−) along a polar bond; drawn as an **arrow** pointing to the δ− atom.

When the molecule is **symmetrical**, so the bond dipoles **cancel** (e.g. CO_{2}, CCl_{4}, BF_{3}).

It is **linear** — the two equal C=O dipoles point in opposite directions and **cancel**.

It is **bent** (lone pairs on O), so the two O–H dipoles **do not cancel** and give a net dipole.

Yes — it is **trigonal pyramidal** (a lone pair on N), so the N–H dipoles do not cancel; NH_{3} is polar.

(1) the bonds are **polar** (electronegativity difference) and (2) the **shape** means the dipoles **do not cancel**.

No — both atoms are identical, so Δχ = 0; the bond is **non-polar** and there is no dipole.

A continuous lattice of atoms joined by **covalent bonds** in every direction — there are **no separate small molecules**.

Melting requires breaking **many strong covalent bonds**, which needs a large amount of energy.

Different structural forms of the **same element** — e.g. diamond and graphite are both pure carbon.

To **four** other carbons in a rigid **3-D tetrahedral** network.

Its **rigid 3-D framework** of strong covalent bonds cannot be pushed out of shape.

All **four** outer electrons of each carbon are used in bonds, so there are **no delocalised electrons** to carry charge.

To **three** others in flat **layers**; the **fourth** electron is **delocalised**.

The **delocalised electrons** between the layers are free to move and carry charge.

**Weak forces** between the layers let the **layers slide** over each other (the covalent bonds within a layer stay strong).

**Diamond**, **graphite** (carbon allotropes), **silicon (Si)** and **silicon dioxide (SiO_{2})**.

Giant covalent → break **strong covalent bonds**; molecular → only overcome **weak intermolecular forces**.

Diamond uses all 4 electrons in bonds (**no** delocalised e⁻ → no conduction); graphite has **1 delocalised** e⁻ per carbon (conducts).

A force of attraction **between** separate molecules — much weaker than the covalent bonds **inside** a molecule.

The strength of its **intermolecular forces** — stronger IMFs need more energy, so a **higher** boiling point.

**London (dispersion) < dipole–dipole < hydrogen bonding.**

Forces from **temporary, instantaneous dipoles**; present between **all** molecules and the **only** force in non-polar ones.

**More electrons** (a larger, more polarisable molecule) — so they increase **down a group** and with molecular size.

When it is **polar** — it has a **permanent dipole** (δ+ and δ− ends) from an electronegativity difference.

The **strongest** IMF: a very δ+ H bonded to **N, O or F** is attracted to a lone pair on the N, O or F of a neighbour.

Hydrogen bonded directly to **N, O or F** ('H bonds to NOF').

NH_{3} has **hydrogen bonding** (H on N); PH_{3} has only weaker dipole–dipole/London forces.

Larger molecules have **more electrons → stronger London forces → higher boiling point**.

**No** — boiling only **separates the molecules** by overcoming intermolecular forces; the covalent bonds stay intact.

N, O and F are very electronegative, so the H is very δ+ and the attraction to a lone pair is especially strong.

A bookkeeping label comparing the electrons an atom 'owns' in a Lewis structure (lone pairs + half of each bond) with its normal valence count. It is **not** a real charge.

$\text{FC} = V - N - \tfrac{1}{2}B$ — valence electrons − non-bonding electrons − ½(bonding electrons).

The atom's **valence electrons as a free atom** — equal to its **group number** (C → 4, N → 5, O → 6).

Total electrons in the atom's bonds: **2 per single, 4 per double, 6 per triple**, summed over all its bonds.

V=5, N=0, B=8 → FC = 5 − 0 − 4 = **+1** (matches the ion's +1 charge).

The one with formal charges **closest to zero**, with any **negative** formal charge on the **most electronegative** atom.

When ≥2 valid Lewis structures can be drawn; the real species is a **hybrid** with electrons **delocalized** over the positions (drawn with a ↔ arrow).

Carbonate CO_{3}^{2-}, nitrate NO_{3}^{-} (three forms each) and ozone O_{3} (two forms).

**Equal bond lengths** — e.g. all three C–O bonds in CO_{3}^{2-} are identical, between a single and a double bond.

A molecule whose central atom has **fewer than 8** outer electrons, e.g. BeCl_{2} (4 on Be) and BF_{3} (6 on B).

A central atom holding **more than 8** outer electrons — only **period-3-or-below** atoms can, e.g. PCl_{5} (10) and SF_{6} (12).

Carbon is **period 2** — it has no available extra valence space, so it is limited to 8 outer electrons.

By **head-on (end-to-end)** overlap of orbitals, with electron density **along the bond axis**. Every single bond is a σ bond.

By **sideways** overlap of two parallel **p orbitals**, with electron density **above and below** the bond axis.

Single = **1 σ**; double = **1 σ + 1 π**; triple = **1 σ + 2 π**. Every multiple bond has exactly **one** σ.

The **mixing** of an atom's atomic orbitals (1 s + some p) into a set of equal-energy **hybrid orbitals**, one per electron domain.

**4** domains (1 s + 3 p), **tetrahedral**, 109.5°, **0** π bonds (all single bonds).

**3** domains (1 s + 2 p), **trigonal planar**, 120°, leaves **1** p orbital → **1 π** bond.

**2** domains (1 s + 1 p), **linear**, 180°, leaves **2** p orbitals → **2 π** bonds.

**Count its electron domains** (σ bonds + lone pairs): 4 → sp³, 3 → sp², 2 → sp.

**σ = total number of bonds** between atoms; **π = (number of doubles) + 2 × (number of triples)**.

The two p orbitals must stay **parallel** to overlap; twisting breaks the π bond, so a C=C cannot rotate → causes **cis–trans isomerism**.

Both are **sp²** (3 domains each), trigonal planar; the C=C is 1 σ + 1 π.

**sp** — it has 2 electron domains (the triple bond + one single bond), so it is linear.