A-Level Chemistry

A scanning tunnelling microscope can image individual atoms.

Everything is made of atoms 原子. An atom is mostly empty space. At its centre is a tiny, heavy nucleus 原子核. The nucleus holds two kinds of particle: protons 质子 and neutrons 中子. Around the nucleus, in the empty space, move the electrons 电子. The electrons stay in shells 壳层 — layers at set distances from the nucleus.

The nucleus is very small but holds almost all the mass. The electrons take up almost all the space but have almost no mass.

An atom is mostly empty space: protons and neutrons sit in the tiny central nucleus, while electrons move in shells around it

Relative charge and relative mass

We compare the three particles using relative charge 相对电荷 and relative mass 相对质量. These are simple numbers, not real units.

A proton and a neutron have almost the same mass. An electron is about 1836 times lighter. The proton is positive, the electron is negative, and the neutron has no charge — it is neutral 中性.

Proton number and nucleon number

Two numbers describe the nucleus:

• the proton number 质子数 (also called the atomic number 原子序数), symbol $Z$ — the number of protons.
• the nucleon number 核子数 (also called the mass number 质量数), symbol $A$ — the total number of protons and neutrons. Protons and neutrons are both nucleons 核子.

So the number of neutrons is $A - Z$.

Counting particles in an atom or ion

For a neutral atom, the number of electrons equals the number of protons, which equals $Z$.

An ion 离子 is an atom that has lost or gained electrons, so it has a charge:

• a positive ion has fewer electrons than protons.
• a negative ion has more electrons than protons.

Example: $^{27}_{13}\text{Al}^{3+}$ has $13$ protons, $27 - 13 = 14$ neutrons, and $13 - 3 = 10$ electrons (it lost 3 electrons to become $3+$).

How mass and charge are spread out

Almost all the mass sits in the nucleus, because protons and neutrons are heavy and electrons are very light. All the positive charge is in the nucleus (the protons). The negative charge is spread out in the shells (the electrons).

Beams of particles in an electric field

Imagine beams of protons, neutrons and electrons moving at the same speed into an electric field 电场 between two charged plates:

• the proton beam bends towards the negative plate (protons are positive).
• the electron beam bends the other way, towards the positive plate. It bends much more, because the electron is far lighter — the same force gives a bigger deflection 偏转 to a smaller mass.
• the neutron beam goes straight through. It has no charge, so the field gives it no force.

In an electric field the proton bends towards the $-$ plate and the electron bends the opposite way and far more (it is much lighter); the neutron passes straight through

Atomic radius and ionic radius

The atomic radius 原子半径 is the size of an atom. The ionic radius 离子半径 is the size of an ion.

Across a period 周期 (left to right), the atomic radius gets smaller. The nuclear charge 核电荷 (the pull from the protons) rises, but the electrons go into the same outer shell, so the shielding 屏蔽 by inner shells stays about the same. The stronger pull draws the outer shell inwards.

Down a group 族 (top to bottom), the atomic radius gets larger. Each step down adds a new shell, so the outer electrons are further out and feel more shielding from the nucleus.

For ions:

• a positive ion (cation 阳离子) is smaller than its atom. It has lost its outer shell, and the electrons that remain feel a stronger pull each.
• a negative ion (anion 阴离子) is larger than its atom. It has gained electrons, so there is more repulsion 排斥 between the electrons.
• among ions that have the same number of electrons, the one with more protons is smaller.

Isotopes 同位素 are atoms of the same element with the same number of protons but a different number of neutrons. So isotopes have the same proton number $Z$ but a different nucleon number $A$.

We write an isotope as $^{A}_{Z}\text{X}$: the nucleon number $A$ on top, the proton number $Z$ below. For example, chlorine has two main isotopes, $^{35}_{17}\text{Cl}$ and $^{37}_{17}\text{Cl}$.

Same chemical properties

Chemical properties 化学性质 depend on the electrons, especially the outer electrons. Isotopes of one element have the same number of electrons arranged in the same way. So they react in exactly the same way — they have the same chemical properties.

Different physical properties

Some physical properties 物理性质 depend on mass, so they differ between isotopes. A heavier isotope has more neutrons, so more mass, and therefore a higher density 密度. (The syllabus limits this difference to mass and density.)

Each element emits light at characteristic wavelengths — its emission spectrum.

Electrons are arranged in shells, sub-shells and orbitals.

Shells and the principal quantum number

Each shell is labelled by the principal quantum number 主量子数 $n = 1, 2, 3, \dots$ A larger $n$ means a shell that is further from the nucleus and higher in energy.

Sub-shells and orbitals

Each shell is split into sub-shells 亚层, named s, p and d. Each sub-shell is built from orbitals 轨道. An orbital is a small region that can hold up to two electrons.

So an s sub-shell holds 2 electrons, a p sub-shell holds 6, and a d sub-shell holds 10.

Order of increasing energy

Electrons fill the lowest-energy sub-shell first. For the first three shells, plus 4s and 4p, the order of rising energy is:

Notice the surprise: 4s is slightly lower in energy than 3d, so 4s fills first.

The sub-shells in order of increasing energy. 4s lies just below 3d, so 4s fills first

Electronic configuration

The electronic configuration 电子排布 lists how many electrons are in each sub-shell. The lowest-energy arrangement is the ground state 基态.

For iron (Fe, $Z = 26$):

You can write a shorthand using the nearest noble gas 稀有气体 in square brackets:

Here $[\text{Ar}]$ stands for the full configuration of argon.

For ions, you add or remove electrons. One key rule: when a transition metal forms a positive ion, it loses its 4s electrons before its 3d electrons. So $\text{Fe}^{3+}$ is $[\text{Ar}]\,3\text{d}^5$.

Electrons in boxes

The electrons in boxes notation draws each orbital as a box and each electron as an arrow. Two electrons in the same orbital must point opposite ways, because each electron has a property called spin 自旋, and a shared orbital needs opposite spins.

Within a sub-shell, electrons fill empty orbitals one at a time, with parallel arrows, before any orbital gets a second electron. Spreading out like this keeps the electrons apart and lowers the repulsion between them.

Electrons in boxes for nitrogen ($1\text{s}^2\,2\text{s}^2\,2\text{p}^3$): paired electrons point opposite ways, and the 2p orbitals fill singly with parallel spins

Why the configuration takes this shape

Electrons fill from low energy to high energy because that gives the most stable (lowest-energy) atom. Within a sub-shell they spread out singly first to reduce the repulsion between the negative electrons.

Shapes of s and p orbitals

• an s orbital is a sphere 球形 centred on the nucleus.
• a p orbital has two lobes, like a dumbbell, pointing along one axis. The three p orbitals point along three directions at right angles (the $x$, $y$ and $z$ axes).

An s orbital is a sphere; each p orbital is a dumbbell, and the three p orbitals point along the $x$, $y$ and $z$ axes

Free radicals

A free radical 自由基 is a species with one or more unpaired electrons 未成对电子. Free radicals are very reactive.

First ionisation energy

The first ionisation energy 第一电离能 (IE) is the energy needed to remove one electron from each atom in one mole of gaseous atoms, forming one mole of gaseous $+1$ ions.

We use gaseous atoms so there are no forces between the particles. The unit is $\text{kJ mol}^{-1}$. As an equation, for an element X:

The $(\text{g})$ shows the species is a gas.

Successive ionisation energies

After removing one electron, you can remove another. The second ionisation energy removes one electron from each $+1$ ion:

You can keep going. These are the successive ionisation energies 逐级电离能. Each is larger than the one before, because every electron is pulled away from a more positive ion.

What ionisation energy depends on

Ionisation energy comes from the attraction between the positive nucleus and the outer electron. Three main factors set how strong that attraction is:

• nuclear charge: more protons pull the electrons more strongly, so the ionisation energy is higher.
• atomic radius: the further the outer electron sits from the nucleus, the weaker the pull, so the ionisation energy is lower.
• shielding: inner shells block some of the pull on the outer electron. More inner shells mean more shielding and a lower ionisation energy.

There is a smaller effect too — spin-pair repulsion 自旋成对排斥. When two electrons share one orbital, they push each other a little, so one is easier to remove.

Trends in first ionisation energy

Across a period, the first ionisation energy generally rises. The nuclear charge grows while shielding stays about the same, so the outer electrons are held more tightly.

Down a group, the first ionisation energy falls. Lower elements have more shells, so more shielding and a larger radius, and the outer electron is easier to remove.

The dips are evidence for sub-shells

The rise across a period is not smooth. Two small dips appear, and you should be able to explain both:

• Group 2 to Group 13 (for example Mg to Al): the electron removed from Al comes from a 3p sub-shell, which is higher in energy than the full 3s sub-shell in Mg. A 3p electron is easier to remove, so the value dips.
• Group 15 to Group 16 (for example P to S): in S, one 3p orbital now holds a pair of electrons. Spin-pair repulsion makes one of them easier to remove, so the value dips.

These dips are evidence that sub-shells exist.

First ionisation energy rises across Period 3 but dips at Al and at S — evidence that sub-shells exist

Successive ionisation energies are evidence for shells

If you plot the successive ionisation energies of one element, the values rise, with big jumps at certain points. A big jump happens when the next electron must come from a shell closer to the nucleus.

Count how many electrons come off easily before the first big jump — that is the number of electrons in the outer shell, which tells you the group the element is in. You can also use the pattern to work out the electronic configuration and the position of the element in the Periodic Table.

Successive ionisation energies of sodium (log scale): the big jumps reveal the $2,8,1$ shell structure

Atoms 原子 are far too light to weigh in grams, so we compare every mass to one standard. The standard is the unified atomic mass unit 统一原子质量单位 (symbol u), defined as exactly one twelfth of the mass of one carbon-12 atom.

Using this unit, we state masses as simple numbers:

• the relative atomic mass 相对原子质量 $A_r$ of an element is the average mass of its atoms compared with $\tfrac{1}{12}$ of a carbon-12 atom. It is an average over all the isotopes 同位素, weighted by how common each one is.
• the relative isotopic mass 相对同位素质量 is the mass of one atom of a single isotope, compared with $\tfrac{1}{12}$ of a carbon-12 atom.
• the relative molecular mass 相对分子质量 $M_r$ of a molecule 分子 is the sum of the relative atomic masses of all its atoms.
• the relative formula mass 相对式量 is the same idea for a substance that is not made of molecules (such as an ionic compound). Add up the relative atomic masses shown in the formula.

A laboratory balance measures mass — the basis of mole calculations.

Chemists count particles in groups called moles, just as we count eggs in dozens.

One mole 摩尔 (symbol mol) is the amount of substance that contains the same number of particles as there are atoms in exactly 12 g of carbon-12. That number is the Avogadro constant 阿伏伽德罗常量:

So one mole of anything contains $6.02 \times 10^{23}$ particles. The particles may be atoms, molecules or ions 离子 — always say which.

The mass of one mole in grams equals the relative mass ($A_r$ or $M_r$). This is the molar mass 摩尔质量, with units $\text{g mol}^{-1}$. The key equation is:

where $n$ is the amount in moles, $m$ is the mass in grams, and $M$ is the molar mass.

The mole is the hub of every amount calculation: convert to mass ($n=m/M$), particles ($\times N_A$), gas volume ($n=V/24$) or solution ($n=cV$)

A compound 化合物 is a substance made of two or more elements chemically joined.

Charges and formulas of ionic compounds

In an ionic compound 离子化合物 the total positive charge balances the total negative charge, so the compound is neutral overall.

You can predict the charge of many ions from the element's position in the Periodic Table:

Hydrogen forms $\text{H}^+$, and the Group 18 noble gases do not normally form ions.

Some ions you must know by name and formula:

For a metal that can have more than one charge, a Roman numeral shows the oxidation number 氧化数. For example, iron(II) is $\text{Fe}^{2+}$ and iron(III) is $\text{Fe}^{3+}$. To write a formula, balance the charges: iron(III) oxide is $\text{Fe}_2\text{O}_3$, because two $\text{Fe}^{3+}$ balance three $\text{O}^{2-}$.

Equations and state symbols

A chemical equation must be balanced 配平 — the same number of each kind of atom on both sides. Add state symbols 状态符号 to show the state of each species: (s) solid, (l) liquid, (g) gas, and (aq) aqueous 水溶液 (dissolved in water).

An ionic equation 离子方程式 shows only the ions and molecules that actually change. The ions that do not change are spectator ions 旁观离子, and you leave them out. For example, the reaction that forms silver chloride is:

Empirical and molecular formulas

The empirical formula 实验式 is the simplest whole-number ratio of the atoms of each element in a compound. The molecular formula 分子式 shows the actual number of atoms of each element in one molecule.

For example, ethane has empirical formula $\text{CH}_3$ but molecular formula $\text{C}_2\text{H}_6$.

Hydrated and anhydrous solids

Some solids hold water inside their crystals. This water is the water of crystallisation 结晶水. A solid that contains it is hydrated 水合的; the same solid with the water removed is anhydrous 无水的.

For example, hydrated copper(II) sulfate is $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$. Heating it drives off the water to leave anhydrous $\text{CuSO}_4$.

Calculating empirical and molecular formulas

To find the empirical formula from masses (or percentages by mass):

1. divide each element's mass by its $A_r$ to get the moles.
2. divide all the mole values by the smallest one.
3. round to the nearest whole numbers — that ratio is the empirical formula.

To get the molecular formula, you also need $M_r$. Find how many times the empirical formula mass fits into $M_r$, then multiply the formula by that number.

A titration finds reacting volumes precisely.

Reacting masses and percentage yield

The numbers in front of each species in a balanced equation give the mole ratio — this is the stoichiometry 化学计量. To find a reacting mass: change the known mass to moles, use the mole ratio to find the moles you want, then change back to mass.

In real reactions you usually get less product than the maximum. The percentage yield 产率 compares the amount you actually made with the most you could make:

Limiting and excess reagent

When two reactants are mixed, one usually runs out first. The limiting reagent 限量试剂 is the one that runs out — it decides how much product forms. The other is the excess reagent 过量试剂, because there is more than enough of it. Always base the calculation on the limiting reagent.

To find it: work out the moles of each reactant, divide each by its number in the equation, and the smallest result is the limiting reagent.

The limiting reagent runs out first and decides how much product forms; the leftover reactant is in excess

Volumes of gases

At the same temperature and pressure, equal volumes of any gases contain equal numbers of molecules. At room temperature and pressure (r.t.p.), one mole of any gas takes up $24.0\ \text{dm}^3$, so:

Equal volumes of gases at the same temperature and pressure hold equal numbers of molecules, whatever the gas

This is used when burning hydrocarbons 碳氢化合物 (compounds of only carbon and hydrogen), where you compare gas volumes.

Volumes and concentrations of solutions

The concentration 浓度 of a solution is the amount of solute 溶质 in each cubic decimetre of solution 溶液, measured in $\text{mol dm}^{-3}$:

where $c$ is the concentration and $V$ is the volume in $\text{dm}^3$. Remember that $1000\ \text{cm}^3 = 1\ \text{dm}^3$.

This is the basis of a titration 滴定, where you find an unknown concentration by reacting it with a solution whose concentration you already know.

In a titration a burette adds a solution of known concentration to the unknown in the conical flask, until the indicator changes

Significant figures

Give your answer to a sensible number of significant figures 有效数字 — usually match the data in the question. Do not write more digits than the data supports, and do not round so early that you lose accuracy.

Electronegativity 电负性 is the power of an atom to attract the electrons 电子 in a bond towards itself.

Three factors decide how electronegative an atom is:

• nuclear charge 核电荷: more protons pull the bonding electrons more strongly.
• atomic radius 原子半径: the closer the bond is to the nucleus, the stronger the pull.
• shielding 屏蔽 by inner shells and sub-shells: more inner electrons weaken the pull on the bonding electrons.

So electronegativity rises across a period (more nuclear charge, smaller radius) and falls down a group (larger radius, more shielding). Fluorine is the most electronegative element.

Electronegativity rises across a period and falls down a group, so fluorine is the most electronegative element

You can use the difference in Pauling electronegativity 鲍林电负性 values to predict the bond type. A large difference gives an ionic bond; a small difference gives a covalent bond.

Rock salt (halite): an ionic compound held in a giant lattice.

Ionic bonding 离子键 is the electrostatic attraction 静电引力 between oppositely charged ions 离子 — positive cations 阳离子 and negative anions 阴离子.

It forms when a metal gives electrons to a non-metal. Good examples are sodium chloride ($\text{NaCl}$), magnesium oxide ($\text{MgO}$) and calcium fluoride ($\text{CaF}_2$). The ions pack into a regular giant lattice 晶格, held together by the attraction in every direction.

Ionic bonding in NaCl: sodium transfers its single outer electron to chlorine, giving Na$^+$ and a full-octet Cl$^-$

A real crystal of rock salt (halite, NaCl); the cubic shapes mirror the giant ionic lattice inside

Metallic bonding 金属键 is the electrostatic attraction between positive metal ions and a "sea" of delocalised electrons 离域电子.

The outer electrons are free to move through the whole metal. This explains why metals conduct electricity and are strong.

Metallic bonding: positive metal ions sit in a sea of delocalised electrons that are free to move

Covalent bonding 共价键 is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.

Simple molecules with covalent bonds include $\text{H}_2$, $\text{O}_2$, $\text{N}_2$, $\text{Cl}_2$, $\text{HCl}$, $\text{CO}_2$, $\text{NH}_3$, $\text{CH}_4$, $\text{C}_2\text{H}_6$ and $\text{C}_2\text{H}_4$. A double bond shares two pairs; a triple bond (as in $\text{N}_2$) shares three pairs.

Atoms in Period 3 and below can expand the octet 扩展八隅体 — hold more than eight electrons in their outer shell. Examples are $\text{SO}_2$, $\text{PCl}_5$ and $\text{SF}_6$.

A coordinate bond 配位键 (also called a dative covalent bond) is a covalent bond where both shared electrons come from the same atom. For example, when ammonia and hydrogen chloride gases meet, the lone pair on the nitrogen forms a coordinate bond to $\text{H}^+$, making the ammonium ion $\text{NH}_4^+$. Coordinate bonds also join the two halves of the $\text{Al}_2\text{Cl}_6$ molecule.

A coordinate (dative) bond: nitrogen's lone pair forms the fourth N–H bond, both electrons coming from N

Sigma and pi bonds

Covalent bonds form when orbitals 轨道 overlap:

• a sigma bond σ键 forms by the direct, head-on overlap 重叠 of orbitals between the two atoms.
• a pi bond π键 forms by the sideways overlap of two p orbitals, above and below the sigma bond.

A single bond is one sigma bond. A double bond (as in $\text{C}_2\text{H}_4$) is one sigma plus one pi bond. A triple bond (as in $\text{N}_2$ and $\text{HCN}$) is one sigma plus two pi bonds.

A $\sigma$ bond forms by direct head-on overlap; a $\pi$ bond forms by the sideways overlap of two p orbitals, above and below

Hybridisation

Hybridisation 杂化 mixes orbitals in the same shell to make new, equal orbitals for bonding:

• $\text{sp}$: two equal orbitals, used in a linear molecule.
• $\text{sp}^2$: three equal orbitals, used in a flat molecule like $\text{C}_2\text{H}_4$.
• $\text{sp}^3$: four equal orbitals, used in $\text{CH}_4$.

Bond energy and bond length

• bond energy 键能 is the energy needed to break one mole of a particular covalent bond in the gas state.
• bond length 键长 is the distance between the centres of the two bonded atoms.

A shorter bond is usually stronger (higher bond energy). Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds. Stronger bonds make a molecule harder to react.

A molecular model shows the three-dimensional shape of a covalent molecule.

To work out a shape, use VSEPR theory 价层电子对互斥理论: the pairs of electrons around the central atom push apart as far as possible, because like charges repel.

A lone pair 孤对电子 (not in a bond) pushes more strongly than a bonding pair 成键电子对. Each lone pair squeezes the bond angle 键角 by about $2.5°$.

$\text{NH}_3$ has one lone pair and $\text{H}_2\text{O}$ has two, which is why their angles drop below the $109.5°$ of $\text{CH}_4$. You can predict the shapes of similar molecules and ions in the same way.

The seven shapes from VSEPR theory; the lone pairs on NH$_3$ and H$_2$O push harder, squeezing the bond angle below $109.5°$

Intermolecular forces 分子间作用力 are the forces between molecules 分子. They are much weaker than the ionic, covalent and metallic bonding inside substances.

Bond polarity and dipoles

When two atoms with different electronegativity share a bond, the electrons sit closer to the more electronegative atom. The bond then has a polarity 极性: one end is slightly negative ($\delta-$) and the other slightly positive ($\delta+$). This separation of charge is a dipole 偶极.

If the dipoles in a molecule do not cancel, the whole molecule has a dipole moment 偶极矩 and is polar. If they cancel by symmetry (as in $\text{CO}_2$), the molecule is non-polar.

Van der Waals' forces

Van der Waals' forces 范德华力 is the general name for all intermolecular forces. There are two main types.

The first type is the instantaneous dipole–induced dipole force, also called the London dispersion force 伦敦色散力. Moving electrons make a brief instantaneous dipole 瞬时偶极, which then creates a matching induced dipole 诱导偶极 in a nearby molecule. These forces act between all molecules and get stronger when there are more electrons.

The second type is the permanent dipole–permanent dipole force. It acts between molecules that are always polar, because each one has a permanent dipole 永久偶极.

Hydrogen bonding

Hydrogen bonding 氢键 is a strong, special case of permanent dipole forces. It forms when hydrogen is bonded to a very electronegative atom — nitrogen, oxygen or fluorine — and is attracted to a lone pair on an N, O or F atom in a neighbour. Look for N–H and O–H groups, as in ammonia and water.

Hydrogen bonding in water: a $\delta+$ hydrogen is attracted to a lone pair on the $\delta-$ oxygen of a neighbouring molecule

Hydrogen bonding explains the strange behaviour of water:

• its high melting and boiling point 沸点, because many hydrogen bonds must be broken.
• its high surface tension 表面张力.
• ice is less dense than liquid water, because hydrogen bonds hold the molecules in an open, spread-out structure, so ice floats.

A dot-and-cross diagram 点叉图 shows the outer electrons of each atom, using dots for one atom and crosses for the other. This makes it clear where each bonding electron came from. You can draw them for ionic, covalent and coordinate bonding, including molecules with an expanded octet or an odd number of electrons.

Covalent dot-and-cross: each shared pair is one electron from each atom. Water has two bonding pairs and two lone pairs; nitrogen shares three pairs (a triple bond)

Boiling turns liquid water into steam — a change between states of matter.

Where gas pressure comes from

Gas molecules move fast in all directions. They keep hitting — colliding 碰撞 with — the walls of their container. Each hit gives the wall a tiny push. The pressure 压强 of the gas is the overall result of these many collisions on the walls.

Gas pressure: fast molecules move in all directions and collide with the walls; the many tiny pushes add up to the pressure

Ideal gases

An ideal gas 理想气体 is a simple model. We assume two things:

• the particles themselves take up zero volume.
• there are no intermolecular forces 分子间作用力 of attraction between the particles.

A real gas 实际气体 follows this model closely at low pressure and high temperature. It behaves least like an ideal gas at high pressure and low temperature, when the particles are squeezed close together and the forces between them start to matter.

An ideal gas is a model: point particles with no forces between them. A real gas behaves least like this at high pressure and low temperature, when the particles are crowded and their real size and attractions start to matter

The ideal gas equation

The ideal gas equation 理想气体方程 links pressure, volume, amount and temperature:

where $p$ is the pressure in Pa, $V$ is the volume in $\text{m}^3$, $n$ is the amount in moles, $T$ is the temperature in kelvin 开尔文 (K), and $R$ is the gas constant 气体常量 ($8.31\ \text{J K}^{-1}\,\text{mol}^{-1}$).

Always change the units first: °C to K (add 273), and $\text{cm}^3$ or $\text{dm}^3$ to $\text{m}^3$.

You can also use the equation to find a molar mass 摩尔质量. Since $n = m/M$:

This lets you work out $M_r$ from the mass (or the density) of a gas.

Quartz is a giant covalent structure of silicon and oxygen.

How a substance behaves depends on how its particles are joined. There are four main structures of a crystalline solid 晶体.

The four structures of a crystalline solid — the structure decides the melting point, conductivity and solubility

Giant ionic

A giant ionic 离子晶体 structure is a huge regular lattice 晶格 of positive and negative ions 离子, held together by strong attraction in every direction. Examples are sodium chloride and magnesium oxide.

Simple molecular

A simple molecular 分子晶体 structure is made of small molecules 分子. The bonds inside each molecule are strong, but the intermolecular forces between the molecules are weak. Examples are iodine ($\text{I}_2$), fullerene 富勒烯 ($\text{C}_{60}$) and ice.

Giant molecular

A giant molecular 原子晶体 structure (also called giant covalent) is a huge network of atoms joined by strong covalent bonds 共价键. Examples are silicon(IV) oxide 二氧化硅, graphite 石墨 and diamond 金刚石.

Two giant covalent forms of carbon: diamond is a rigid 3D network (very hard); graphite has sliding layers and spare electrons that conduct

Giant metallic

A giant metallic 金属晶体 structure is a lattice of positive metal ions in a "sea" of delocalised electrons 离域电子. An example is copper.

Physical properties

The structure decides the physical properties:

• melting point 熔点 and boiling point 沸点 are high when strong forces (ionic, covalent or metallic) must be broken, and low when only weak intermolecular forces break.
• electrical conductivity 导电性 needs charged particles that can move — ions that are free (when molten or dissolved) or delocalised electrons. Graphite conducts because some of its electrons are delocalised.
• solubility 溶解度 in water is usually high for ionic solids and low for molecular and giant covalent solids.

You can work backwards too: from the melting point, conductivity and solubility of an unknown substance, deduce the type of structure and bonding it has.

Burning fuel is exothermic, releasing energy to the surroundings.

An instant cold pack uses an endothermic reaction that takes in heat.

Every chemical reaction takes in or gives out energy. This energy change, measured at constant pressure, is the enthalpy change 焓变, with symbol $\Delta H$.

• in an exothermic 放热 reaction the system gives out heat, so the products have less energy than the reactants and $\Delta H$ is negative.
• in an endothermic 吸热 reaction the system takes in heat, so the products have more energy than the reactants and $\Delta H$ is positive.

Reaction pathway diagrams

A reaction pathway diagram 反应路径图 shows the energy of the reactants and products, and the energy "hill" between them. The height of the hill is the activation energy 活化能 — the least energy the particles need before they can react.

• exothermic: products sit lower than reactants ($\Delta H < 0$).
• endothermic: products sit higher than reactants ($\Delta H > 0$).

Exothermic reactions end lower than they start ($\Delta H<0$); endothermic reactions end higher ($\Delta H>0$)

Standard conditions and types of enthalpy change

Energy values are compared under standard conditions 标准条件: $298\ \text{K}$ and $101\ \text{kPa}$, shown by the symbol $^{\ominus}$. Each substance is in its normal physical state at those conditions.

Energy from breaking and making bonds

During a reaction, old bonds break and new bonds form. Breaking a bond needs energy (endothermic); making a bond releases energy (exothermic). The enthalpy change of the reaction is the difference between the two:

The bond energy 键能 is the energy needed to break one mole of a particular bond in the gas state, so it is always positive. Some bond energies are exact (for one specific molecule); others are averages taken over many different molecules, so calculations using them are only approximate.

Breaking bonds takes energy in; making bonds gives energy out. $\Delta H$ is the difference between the two

Measuring enthalpy change in the lab

When a reaction heats up (or cools down) a known mass of water or solution, the heat transferred is:

where $m$ is the mass, $c$ is the specific heat capacity 比热容 (how much energy raises 1 g by 1 K), and $\Delta T$ is the temperature change. The enthalpy change per mole is then:

The minus sign makes $\Delta H$ negative when the temperature rises (an exothermic reaction).

An insulated cup and a thermometer measure the temperature change of a known mass of solution

Hess's law 盖斯定律 says that the total enthalpy change for a reaction is the same, no matter which route you take from reactants to products. This is because energy is conserved.

This lets you draw an energy cycle 能量循环: link the reactants and products by a direct step and by an indirect route, then add the steps so that both routes give the same total.

The direct route equals the indirect route, so $\Delta H_r = \Delta H_1 + \Delta H_2$

Hess's law is useful in two ways:

• it lets you find an enthalpy change that you cannot measure directly (for example, the formation of a compound that forms slowly or with side reactions).
• it lets you calculate $\Delta H_r$ from bond energy data, or from formation or combustion data given in the question.

The oxidation number 氧化数 (also called the oxidation state) shows how many electrons 电子 an atom has lost or gained compared with the free element. You work it out using simple rules:

• an uncombined element has an oxidation number of $0$.
• a simple ion has an oxidation number equal to its charge (so $\text{Mg}^{2+}$ is $+2$).
• Group 1 is always $+1$, Group 2 is always $+2$.
• hydrogen is $+1$ (but $-1$ in metal hydrides).
• oxygen is $-2$ (but $-1$ in peroxides).
• fluorine is always $-1$.
• the oxidation numbers in a neutral compound add up to $0$; in an ion they add up to the charge.

A Roman numeral shows the size of the oxidation number of an element, for example iron(II) means $+2$ and manganese(VII) in $\text{KMnO}_4$ means $+7$.

A battery uses redox reactions to push electrons round a circuit.

A redox 氧化还原 reaction is one where electrons move from one species to another.

• oxidation 氧化 is the loss of electrons. The oxidation number goes up.
• reduction 还原 is the gain of electrons. The oxidation number goes down.

A useful memory aid is OIL RIG: Oxidation Is Loss, Reduction Is Gain.

Oxidation and reduction always happen together, because the electrons lost by one species are gained by another. This electron transfer 电子转移 is why we call it a redox reaction.

Redox is electron transfer: the reducing agent loses electrons (oxidised, number up); the oxidising agent gains them (reduced, number down)

Changes in oxidation number help you balance a redox equation:

1. find the element whose oxidation number rises (it is oxidised) and the one whose number falls (it is reduced).
2. the total rise must equal the total fall, because every electron lost is gained somewhere.
3. choose the ratio of the two species so the rise and fall match, then balance the rest of the equation.

Balancing a redox equation: the total rise in oxidation number must equal the total fall, which fixes the ratio

Disproportionation 歧化 is a special redox reaction in which the same element is both oxidised and reduced at the same time. For example, when chlorine reacts with cold water, some chlorine atoms are reduced (to $\text{Cl}^-$) and others are oxidised (to $\text{ClO}^-$).

Disproportionation: chlorine ($0$) is both oxidised to ClO$^-$ ($+1$) and reduced to Cl$^-$ ($-1$) at once

Rusting is the oxidation of iron by oxygen.

• an oxidising agent 氧化剂 takes electrons away from another species. In doing so, it is itself reduced.
• a reducing agent 还原剂 gives electrons to another species. In doing so, it is itself oxidised.

So in any redox reaction, the oxidising agent causes the oxidation of the other species, while the reducing agent causes the reduction.

Cobalt chloride changes colour as its equilibrium shifts.

A reversible reaction 可逆反应 can go both ways. The forward reaction 正反应 makes products; the reverse reaction 逆反应 turns products back into reactants. We show this with the sign $\rightleftharpoons$.

In a closed system 封闭系统 (nothing enters or leaves), the reaction reaches a dynamic equilibrium 动态平衡 when:

• the rate of the forward reaction equals the rate of the reverse reaction.
• the concentrations of reactants and products stay constant.

It is called dynamic because both reactions are still happening — they just cancel out. A closed system is needed, or products would escape and equilibrium could never be reached.

At dynamic equilibrium the forward and reverse rates have become equal, so the concentrations stay constant

Le Chatelier's principle 勒夏特列原理 says: if you change a system at equilibrium, the position of equilibrium 平衡 moves to oppose (reduce) that change.

A catalyst lets equilibrium be reached faster, but it does not change the position of equilibrium.

Le Chatelier's principle: the equilibrium always shifts to oppose the change you make

For a reaction at equilibrium, the equilibrium constant 平衡常数 links the amounts of products and reactants. For the reaction $a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D}$:

where the square brackets mean concentration in $\text{mol dm}^{-3}$.

For reactions of gases, we use partial pressure 分压 instead of concentration. The partial pressure of a gas is the share of the total pressure that it provides. It is found from the mole fraction 摩尔分数 (the fraction of all the moles that are that gas):

The constant written with partial pressures is $K_p$.

Only temperature changes the value of $K_c$ or $K_p$. Changing concentration or pressure, or adding a catalyst, shifts the position of equilibrium but leaves the constant unchanged.

Ammonia for fertiliser is made by the Haber process — a key industrial equilibrium.

These two industrial processes are chosen by balancing yield, rate and cost using Le Chatelier's principle.

• the Haber process 哈伯法 makes ammonia: $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$ (exothermic). It uses about $450\,°\text{C}$, $200\ \text{atm}$ and an iron catalyst. A low temperature would give more ammonia but too slowly, so a moderate temperature is a compromise.
• the Contact process 接触法 makes sulfur trioxide: $2\text{SO}_2 + \text{O}_2 \rightleftharpoons 2\text{SO}_3$ (exothermic). It uses about $450\,°\text{C}$, near $1$–$2\ \text{atm}$ and a vanadium(V) oxide catalyst.

The Haber equilibrium gives more ammonia at lower temperature and higher pressure; about $450\,°$C and $200$ atm is the working compromise

A proton 质子 is simply an $\text{H}^+$ ion. The Brønsted–Lowry theory defines acids and bases by what they do with protons:

• an acid 酸 is a proton donor 质子供体 (it gives away $\text{H}^+$).
• a base 碱 is a proton acceptor 质子受体 (it takes $\text{H}^+$). A base that dissolves in water is called an alkali.

Brønsted–Lowry: the acid donates a proton ($\text{H}^+$) to the base, making two conjugate acid–base pairs

Common acids you must know: hydrochloric acid ($\text{HCl}$), sulfuric acid ($\text{H}_2\text{SO}_4$), nitric acid ($\text{HNO}_3$) and ethanoic acid ($\text{CH}_3\text{COOH}$). Common alkalis: sodium hydroxide ($\text{NaOH}$), potassium hydroxide ($\text{KOH}$) and ammonia ($\text{NH}_3$).

This is about how fully an acid or base splits up in water — not how concentrated it is.

• a strong acid 强酸 or strong base 强碱 is fully dissociated 解离 in water (almost every molecule splits into ions).
• a weak acid 弱酸 or weak base 弱碱 is only partly dissociated (most molecules stay whole).

Same concentration, different ionisation: a strong acid is fully dissociated into ions; a weak acid stays mostly as whole molecules

The pH scale measures how acidic a solution is: pure water is pH 7, acids are below 7, and alkalis are above 7. A strong acid has a lower pH than a weak acid of the same concentration.

You can tell strong and weak acids apart by:

• reaction with a reactive metal: a strong acid fizzes faster.
• pH: measured with a pH meter or universal indicator 通用指示剂.
• electrical conductivity: a strong acid conducts better, because it has more ions.

Neutralisation 中和 happens when the hydrogen ions from an acid react with the hydroxide ions from an alkali:

A salt 盐 is also formed, from the rest of the acid and base.

In a titration 滴定 you add one solution to another and follow the pH. The titration curve 滴定曲线 has a steep, almost vertical jump near the end point. The exact shape depends on whether each reactant is strong or weak.

To pick an indicator 指示剂, choose one whose colour change falls inside that steep jump. For a strong acid with a strong base most indicators work; for a weak acid with a strong base you need one that changes in the higher pH range.

Titration curves each have a steep pH jump at the end point. The weak-acid jump sits higher, so the indicator must change colour in that range

The rate of reaction 反应速率 is how fast reactants turn into products. We measure it as the change in concentration (or amount) in each unit of time.

To react, particles must collide 碰撞. The collision frequency 碰撞频率 is how often the particles hit each other. But not every collision leads to a reaction:

• an effective collision 有效碰撞 has enough energy and the correct direction, so a reaction happens.
• a non-effective collision 无效碰撞 does not have enough energy, or the particles hit at the wrong angle, so nothing happens.

So the rate depends on the frequency of effective collisions — how many useful collisions happen each second.

A collision only reacts with the right orientation and enough energy (coloured ends = the reactive part)

Concentration and pressure

If you increase the concentration of a solution (or the pressure of a gas), the particles are packed closer together. They collide more often, so there are more effective collisions each second, and the rate goes up.

More particles in the same volume collide more often, so the rate rises

You can calculate a rate from experimental data — for example, the volume of gas made divided by the time taken.

The activation energy 活化能 ($E_A$) is the minimum energy a collision needs in order to be effective.

The Boltzmann distribution 玻尔兹曼分布 is a graph showing how the energies of the molecules are spread out at one temperature. The curve starts at the origin, rises to a peak, then falls away in a long tail. The total area under the curve is the total number of molecules. Only the molecules to the right of $E_A$ have enough energy to react.

When you raise the temperature:

• the curve flattens and spreads to the right, so a much larger fraction of molecules now have energy greater than $E_A$.
• the molecules also move faster and collide more often.

The first effect is the bigger one. This is why a small rise in temperature gives a large rise in rate.

At higher temperature the curve spreads to the right, so a larger fraction of molecules can react

A catalytic converter speeds up the reactions that clean a car's exhaust gases.

A catalyst 催化剂 speeds up a reaction but is not used up itself. Catalysis 催化作用 is the name for this action.

A catalyst works by giving the reaction a different reaction mechanism 反应机理 — a new route with a lower activation energy. On the Boltzmann distribution, lowering $E_A$ moves the line to the left, so more molecules now have enough energy. This means more effective collisions each second, and a faster rate.

On a reaction pathway diagram 反应路径图, the catalysed route has a lower energy "hill". The enthalpy change of the reaction, $\Delta H$, is not changed by the catalyst.

A catalyst gives a route with lower activation energy; the enthalpy change is unchanged

There are two types:

• a homogeneous catalyst 均相催化剂 is in the same physical state as the reactants — for example, an acid catalyst dissolved in a solution of liquids.
• a heterogeneous catalyst 多相催化剂 is in a different state from the reactants — for example, solid iron speeding up the reaction of gases in the Haber process.

A car's catalytic converter is a heterogeneous catalyst; its honeycomb gives a huge surface area

Properties repeat in a regular pattern across each period of the table.

Periodicity 周期性 means that properties repeat in a regular pattern as you go across each period 周期 of the Periodic Table. Period 3 (Na to Ar) is the standard example.

Atomic radius decreases across Period 3: the rising nuclear charge pulls the same outer shell inwards

The melting point and conductivity follow from the structure and bonding:

• Na, Mg, Al are giant metallic 金属晶体. Melting points rise (Na → Al) because each atom gives more delocalised electrons and the ions get smaller, so the bonding is stronger. They conduct well.
• Si is giant molecular 原子晶体 (giant covalent). It has the highest melting point, because strong covalent bonds must be broken. It barely conducts.
• P, S, Cl, Ar are simple molecular 分子晶体 (or single atoms). Their melting points are low, because only weak intermolecular forces break. They do not conduct.

Melting point across Period 3 peaks at silicon (giant covalent); it is high for the metals and low for the simple molecular elements

Reactions with oxygen, chlorine and water

With oxygen:

With chlorine:

With water (only Na and Mg react):

Sodium reacts fast; magnesium reacts only very slowly with cold water.

Oxidation number of the oxides and chlorides

The oxidation number 氧化数 of the Period 3 element in its oxide or chloride rises across the period, because it equals the number of outer-shell (valence shell 价层) electrons 电子 the atom uses in bonding:

The chlorides $\text{NaCl}$, $\text{MgCl}_2$, $\text{AlCl}_3$, $\text{SiCl}_4$, $\text{PCl}_5$ show oxidation numbers $+1$ to $+5$ in the same way.

Oxides with water, and acid–base behaviour

Across the period the oxides 氧化物 change from basic to acidic:

The Period 3 oxides change from basic (the metals) through amphoteric (Al$_2$O$_3$) to acidic (the non-metals)

Metal oxides (left) are basic; non-metal oxides (right) are acidic. $\text{Al}_2\text{O}_3$ and its hydroxide 氢氧化物 $\text{Al(OH)}_3$ are amphoteric — they react with both acids and bases:

Aluminium oxide is amphoteric — it reacts with acids (behaving as a base) and with bases (behaving as an acid)

Chlorides with water

• $\text{NaCl}$ and $\text{MgCl}_2$ are ionic. They simply dissolve, giving a near-neutral solution.
• $\text{SiCl}_4$ and $\text{PCl}_5$ are covalent. They undergo hydrolysis 水解 (react with water) to make acidic solutions and fumes of $\text{HCl}$:

Explaining the trends

These trends follow from the change in bonding and electronegativity 电负性. On the left, the elements are metals with low electronegativity, so their oxides and chlorides are ionic and basic (or neutral). On the right, the elements are non-metals with high electronegativity, so their oxides and chlorides are covalent and acidic. You can use a chloride's or oxide's properties (melting point, conductivity, effect on water) to suggest whether its bonding is ionic or covalent.

The same idea works for any group. If you know the pattern down a group and across a period, you can:

• predict the properties of an element from its position (for example, a Group 1 element will be a reactive metal forming a $+1$ ion).
• deduce the likely position and identity of an unknown element from its physical and chemical properties.

Magnesium, a Group 2 metal, burns with a brilliant white flame.

Limestone is calcium carbonate — a compound of the Group 2 metal calcium.

Group 族 2 holds the metals magnesium, calcium, strontium and barium. They all have two outer electrons, which they lose to form $2+$ ions. Going down the group, the atoms get larger and the outer electrons are easier to lose, so the metals get more reactive — their reactivity 反应活性 increases down the group.

Reactivity increases down Group 2: larger atoms hold their two outer electrons less tightly, so they are lost more easily

Reactions of the elements

With oxygen, they burn to form an oxide:

Magnesium burns in air with a brilliant white flame, forming white magnesium oxide

With water, they form a hydroxide and hydrogen. The reaction gets faster down the group:

Magnesium is slow with cold water but reacts fast with steam, giving $\text{MgO}$ and $\text{H}_2$.

With dilute hydrochloric or sulfuric acid, they form a salt and hydrogen:

With sulfuric acid the reaction slows down the group, because the sulfates 硫酸盐 formed (such as $\text{BaSO}_4$) are insoluble and coat the metal.

Reactions of the compounds

The oxides 氧化物 and hydroxides 氢氧化物 are basic. They react with water and with dilute acids:

The carbonates 碳酸盐 react with dilute acids to give a salt, water and carbon dioxide:

Thermal decomposition

Thermal decomposition 热分解 means breaking a compound apart by heating it.

• the carbonates break into the oxide and carbon dioxide:

• the nitrates 硝酸盐 break into the oxide, brown nitrogen dioxide gas and oxygen:

The thermal stability 热稳定性 of both the carbonates and the nitrates increases down the group. A larger metal ion pulls less on the carbonate or nitrate ion, so the compound is harder to break apart — it needs a higher temperature. So magnesium carbonate decomposes most easily, and barium carbonate is the hardest.

Thermal stability rises down the group: a small cation polarises (distorts) the carbonate ion more, weakening it so it decomposes more easily

Trends in solubility

Down Group 2 the hydroxides become more soluble while the sulfates become less soluble (note the log scale)

From these trends you can predict the properties of the next element down, radium, and of its compounds.

Bromine is a Group 17 halogen — a dark liquid that gives off an orange vapour.

The halogens 卤素 are the Group 族 17 elements. They exist as diatomic molecules ($\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$). Going down the group:

The volatility 挥发性 (how easily a substance turns to vapour) decreases down the group: chlorine is a gas, but iodine is a solid. This is because the molecules get larger and have more electrons, so the instantaneous dipole 瞬时偶极 and induced dipole 诱导偶极 forces between them get stronger. Stronger forces are harder to break, so the boiling point rises and volatility falls.

Down Group 17 the halogens change from a pale green gas to a red-brown liquid to a grey-black solid as volatility falls

The bond energy 键能 (bond strength) of the $\text{X}\text{–}\text{X}$ molecules generally falls from $\text{Cl}_2$ to $\text{I}_2$, because the shared electrons are further from the nuclei in the larger atoms.

Halogens as oxidising agents

Each halogen reacts by gaining one electron to form a $1-$ ion, so it acts as an oxidising agent 氧化剂. This power decreases down the group, because the larger atoms attract an extra electron less strongly. A more reactive halogen can push out a less reactive one from its salt:

Reactions with hydrogen

Each halogen reacts with hydrogen to form a hydrogen halide 卤化氢:

The reaction gets less vigorous down the group: chlorine reacts explosively in light, bromine needs heat, and iodine reacts slowly and only partly.

Thermal stability of the hydrogen halides

The thermal stability 热稳定性 of the hydrogen halides decreases down the group. The H–X bond gets weaker as the halogen atom gets larger, so HI breaks apart on gentle heating while HCl is very stable.

Halide ions as reducing agents

A halide ion 卤离子 (such as $\text{Cl}^-$) can give away an electron, acting as a reducing agent 还原剂. This power increases down the group, because a larger ion holds its outer electron less tightly.

Two opposite trends: the oxidising power of the halogens falls down the group, while the reducing power of the halide ions rises

Reaction with aqueous silver ions

Add aqueous silver nitrate, then aqueous ammonia 氨, to identify the halide from the colour of the silver halide precipitate 沉淀:

Silver halide precipitates: AgCl is white, AgBr cream and AgI yellow — and their solubility in ammonia confirms which halide is present

The silver halide test: AgCl is white, AgBr cream and AgI yellow

Reaction with concentrated sulfuric acid

All the halides first give the hydrogen halide. The lower halides are then oxidised by the acid, because they are stronger reducing agents:

• chloride gives only $\text{HCl}$ (no redox).
• bromide also gives some brown $\text{Br}_2$ and $\text{SO}_2$.
• iodide gives $\text{I}_2$ and the smelly gases $\text{H}_2\text{S}$ and $\text{SO}_2$, because $\text{I}^-$ is the strongest reducing agent.

Chlorine is added to pool water to kill microbes.

With sodium hydroxide

With cold, dilute sodium hydroxide, chlorine reacts to form chloride and chlorate(I):

With hot, concentrated sodium hydroxide, it forms chloride and chlorate(V):

In both, the oxidation number 氧化数 of chlorine goes both up and down (from $0$), so both are disproportionation 歧化 reactions.

Chlorine in water purification

A little chlorine is added to water for water purification 水净化. It reacts with water:

The active species $\text{HOCl}$ and $\text{ClO}^-$ kill bacteria 细菌, making the water safe to drink.

Nitrogen gas, $\text{N}_2$, makes up most of the air but reacts with very little. There are two reasons:

• the two nitrogen atoms are joined by a triple bond 三键, which has a very high bond energy 键能. A lot of energy is needed to break it.
• the molecule has no polarity 极性 — it is perfectly symmetrical, so nothing pulls other molecules towards it.

Nitrogen is unreactive for two reasons: a very strong triple bond, and a symmetrical, non-polar molecule

Ammonia is converted into nitrogen fertilisers on a huge scale.

Basicity of ammonia

The basicity 碱性 of ammonia (its ability to act as a base) comes from the lone pair 孤对电子 of electrons on the nitrogen atom. Using the Brønsted–Lowry theory, ammonia is a base because this lone pair can accept a proton 质子 ($\text{H}^+$):

The ammonium ion

When the lone pair forms a bond to $\text{H}^+$, it makes the ammonium ion 铵离子, $\text{NH}_4^+$. Because both shared electrons came from the nitrogen, this new bond is a coordinate bond 配位键. The ion has four identical N–H bonds and a tetrahedral shape.

Displacement of ammonia from its salts

If you warm an ammonium salt with a base (such as sodium hydroxide), you push out ammonia gas. This is an acid–base displacement 置换:

The sharp smell of ammonia, and damp red litmus turning blue, is a test for an ammonium salt.

Oxides of nitrogen and sulfur contribute to smog and air pollution.

Where they come from

Oxides of nitrogen ($\text{NO}$ and $\text{NO}_2$, together called $\text{NO}_x$) come from two sources:

• natural: lightning gives enough energy for nitrogen and oxygen in the air to combine.
• man-made: the high temperature inside an internal combustion engine 内燃机 makes nitrogen and oxygen react:

Removing them from car exhaust

A catalytic converter 催化转化器 cleans the exhaust gases 尾气. It lets the harmful gases react together to form harmless ones:

In a catalytic converter the harmful gases NO and CO react over the catalyst to form harmless N$_2$ and CO$_2$

Photochemical smog

In sunlight, $\text{NO}$ and $\text{NO}_2$ react with unburned hydrocarbons to form peroxyacetyl nitrate (PAN). PAN is a harmful part of photochemical smog 光化学烟雾, the brown haze seen over busy cities.

Acid rain

The oxides of nitrogen also help make acid rain 酸雨 in two ways:

• directly: $\text{NO}_2$ dissolves in rain to form nitric acid.
• as a catalyst: $\text{NO}_2$ speeds up the oxidation of atmospheric sulfur dioxide 二氧化硫 ($\text{SO}_2$) into $\text{SO}_3$, which then forms sulfuric acid in the rain.

Nitrogen and sulfur oxides make acid rain: NO$_2$ forms nitric acid directly and also catalyses the oxidation of SO$_2$ to sulfuric acid

Crude oil is a complex mixture of hydrocarbons — the feedstock for organic chemistry.

A hydrocarbon 碳氢化合物 is a compound of only carbon and hydrogen. Alkanes 烷烃 are the simplest hydrocarbons and have no functional group.

A functional group 官能团 is the reactive part of a molecule. It decides the physical and chemical properties of the compound, so molecules with the same functional group behave alike (for example the $\text{–OH}$ group in alcohols).

Types of formula

The same molecule (butane) shown four ways — molecular, structural, displayed and skeletal — each hiding more detail than the last

You can read off the empirical formula 实验式 (simplest ratio) from any of these.

Naming

Use systematic nomenclature 命名法: a stem for the number of carbons (meth-, eth-, prop-, but-, pent-, hex- for 1 to 6), an ending for the functional group, and numbers to show where groups are.

Some key terms

• a homologous series 同系列 is a family of compounds with the same functional group and general formula, where each member differs by $\text{CH}_2$.
• a saturated 饱和 compound has only single C–C bonds; an unsaturated 不饱和 compound has a C=C double bond (or a triple bond).

Breaking bonds

A covalent bond can break in two ways:

• homolytic fission 均裂: the bond splits evenly, one electron to each atom. This makes two free radicals 自由基 (species with an unpaired electron).
• heterolytic fission 异裂: the bond splits unevenly, both electrons going to one atom. This makes two ions.

Homolytic fission gives one electron to each atom (two radicals); heterolytic fission gives both electrons to one atom (two ions)

Attacking species

• a nucleophile 亲核试剂 is a species with a lone pair that is attracted to a positive (electron-poor) centre.
• an electrophile 亲电试剂 is a species attracted to a negative (electron-rich) centre, such as a C=C double bond.

A nucleophile uses its lone pair to attack an electron-poor centre; an electrophile is drawn to an electron-rich one such as a C=C bond

Types of reaction

For organic redox, the symbol $[\text{O}]$ stands for one oxygen atom from an oxidising agent, and $[\text{H}]$ for one hydrogen atom from a reducing agent.

Types of mechanism

A free-radical reaction happens in three steps: initiation 引发 (radicals are made), propagation 增长 (radicals react and make new radicals), and termination 终止 (two radicals join and stop the chain).

The main mechanisms you meet are free-radical substitution 自由基取代 (alkanes), electrophilic addition 亲电加成 (alkenes), nucleophilic substitution 亲核取代 (halogenoalkanes) and nucleophilic addition 亲核加成 (carbonyls). In mechanisms, a curly arrow 弯箭头 shows a pair of electrons moving; it starts at a bond or a lone pair 孤对电子.

Organic molecules have definite three-dimensional shapes.

Organic molecules can be straight-chained, branched or cyclic 环状 (in a ring).

The shape around a carbon depends on its hybridisation 杂化:

Every single bond is a sigma bond σ键, made by direct overlap. A double bond is one sigma bond plus one pi bond π键, made by sideways overlap of p orbitals. Ethene is planar because of its $\text{sp}^2$ carbons.

The shape at a carbon follows from its hybridisation: sp$^3$ is tetrahedral ($109.5°$), sp$^2$ is planar ($120°$), sp is linear ($180°$)

Isomers are compounds with the same molecular formula but a different arrangement of atoms. This is called isomerism 异构.

Structural isomerism

In structural isomerism 结构异构 the atoms are joined in a different order. There are three kinds:

• chain isomerism 链异构: the carbon chain is branched in different ways.
• positional isomerism 位置异构: the functional group is on a different carbon.
• functional group isomerism 官能团异构: the atoms form a different functional group (for example an alcohol and an ether).

The three kinds of structural isomerism — chain, positional and functional-group — each pair sharing the same molecular formula

Stereoisomerism

In stereoisomerism 立体异构 the atoms are joined in the same order but point in different directions in space.

• geometrical isomerism 几何异构 (cis/trans) happens at a C=C double bond. The pi bond stops the two carbons rotating (restricted rotation 受限旋转), so groups are fixed on the same side (cis 顺式) or opposite sides (trans 反式).

Cis–trans isomerism at a C=C bond: the methyl groups are fixed on the same side (cis) or opposite sides (trans) because the bond cannot rotate

• optical isomerism 旋光异构 happens at a chiral 手性 carbon — a chiral centre 手性中心 is a carbon with four different groups attached. Such a carbon gives two mirror-image forms called enantiomers 对映体. A molecule may have more than one chiral centre.

Optical isomerism: a carbon with four different groups gives two mirror-image forms (enantiomers) that cannot be superimposed

From a molecular formula you can deduce the possible isomers by trying different chains, positions and functional groups.

Natural gas — mainly methane, the simplest alkane — burns on a stove.

Alkanes 烷烃 are saturated hydrocarbons (general formula $\text{C}_n\text{H}_{2n+2}$).

Making alkanes

• hydrogenation 氢化: add hydrogen to an alkene 烯烃, using a $\text{Pt}$ or $\text{Ni}$ catalyst and heat.
• cracking 裂化: break a long-chain alkane into shorter ones by heating with $\text{Al}_2\text{O}_3$.

Combustion

In combustion 燃烧 an alkane burns in oxygen:

• complete combustion 完全燃烧 (plenty of oxygen) gives carbon dioxide and water.
• incomplete combustion 不完全燃烧 (not enough oxygen) gives water plus toxic carbon monoxide 一氧化碳 ($\text{CO}$) and soot (carbon).

Free-radical substitution

Alkanes react with chlorine or bromine by free-radical substitution 自由基取代, in ultraviolet light 紫外线. For ethane and chlorine the mechanism has three steps:

• initiation 引发 — UV light splits the halogen into two free radicals 自由基: $\;\text{Cl}_2 \rightarrow 2\,\text{Cl}\cdot$
• propagation 增长 — radicals react and make new radicals:
  $$\text{Cl}\cdot + \text{C}_2\text{H}_6 \rightarrow \text{C}_2\text{H}_5\cdot + \text{HCl} \qquad \text{C}_2\text{H}_5\cdot + \text{Cl}_2 \rightarrow \text{C}_2\text{H}_5\text{Cl} + \text{Cl}\cdot$$
• termination 终止 — two radicals join, ending the chain: $\;\text{Cl}\cdot + \text{C}_2\text{H}_5\cdot \rightarrow \text{C}_2\text{H}_5\text{Cl}$

Free-radical substitution in three steps: initiation makes radicals, propagation carries the chain, and termination ends it

Why cracking is useful, and why alkanes are unreactive

Cracking turns heavy fractions of crude oil 原油 into more useful, lower-$M_r$ alkanes and alkenes. (A fraction 馏分 is a group of molecules with a similar boiling-point range.)

Alkanes are generally unreactive, especially towards polar reagents. This is because the C–H and C–C bonds are strong and have little polarity 极性, so there is no charge to attract an attacking species.

Environmental effects

Burning alkanes in an internal combustion engine gives off carbon monoxide, oxides of nitrogen and unburnt hydrocarbons. A catalytic converter removes these by turning them into harmless gases.

Cracking heavy fractions at a refinery produces alkenes for plastics and fuels.

Alkenes have a C=C double bond (general formula $\text{C}_n\text{H}_{2n}$). The double bond is the functional group, so alkenes are reactive.

Making alkenes

• elimination 消去 of $\text{HX}$ from a halogenoalkane 卤代烷, using $\text{NaOH}$ dissolved in ethanol, with heat.
• dehydration 脱水 of an alcohol 醇, using a hot $\text{Al}_2\text{O}_3$ catalyst or concentrated sulfuric acid.
• cracking of a longer-chain alkane.

Reactions of alkenes

Most reactions are electrophilic addition 亲电加成 across the double bond:

There are also two oxidation reactions with acidified $\text{KMnO}_4$:

• cold, dilute $\text{KMnO}_4$ adds two $\text{–OH}$ groups to give a diol 二醇.
• hot, concentrated $\text{KMnO}_4$ breaks the C=C bond apart. The products (such as carboxylic acids or carbon dioxide) tell you where the double bond was.

Test for a C=C bond

Shake the compound with orange bromine water 溴水. An alkene decolourises it (turns it colourless) by electrophilic addition. An alkane does not.

The bromine-water test: an alkene decolourises the orange bromine water, while an alkane leaves it orange

Addition polymerisation

In addition polymerisation 加成聚合, many alkene molecules join into one long chain, with no other product. Ethene gives poly(ethene). The long-chain product is a polymer 聚合物.

Addition polymerisation: many ethene molecules open their double bonds and join into the long chain of poly(ethene)

The mechanism and Markovnikov's rule

In electrophilic addition (for example bromine with ethene), the electron-rich C=C attracts the electrophile. This forms a positive intermediate called a carbocation 碳正离子, which the negative part then attacks.

Electrophilic addition of bromine to ethene: the C=C attacks Br$^{\delta+}$, a carbocation forms, then Br$^-$ attacks it

Alkyl groups push electrons towards the positive carbon — this is the inductive effect 诱导效应. So a carbocation with more alkyl groups is more stable: tertiary is more stable than secondary, which is more stable than primary. When $\text{HBr}$ adds to propene, the more stable carbocation forms, so hydrogen adds to the carbon that already has more hydrogens. This pattern is Markovnikov's rule 马氏规则.

Carbocation stability rises from primary to tertiary as more alkyl groups push electrons in — the basis of Markovnikov's rule

A halogenoalkane 卤代烷 is an alkane with one or more halogen atoms in place of hydrogen (a C–X bond, where X is a halogen).

Volatile halogenoalkanes were once widely used as refrigerants and aerosol propellants

Making halogenoalkanes

• free-radical substitution 自由基取代 of an alkane with $\text{Cl}_2$ or $\text{Br}_2$ in ultraviolet light.
• electrophilic addition 亲电加成 of an alkene with a halogen $\text{X}_2$ or a hydrogen halide $\text{HX}$.
• substitution of an alcohol 醇, for example by $\text{HX}$, by $\text{PCl}_5$, by $\text{PCl}_3$ with heat, or by $\text{SOCl}_2$.

Three classes

A halogenoalkane is primary 伯, secondary 仲 or tertiary 叔, depending on how many carbon atoms are joined to the carbon that holds the halogen (one, two or three).

Primary, secondary and tertiary halogenoalkanes, set by how many carbons are joined to the carbon bearing the halogen

The C–X bond is polar, so the carbon is slightly positive. A nucleophilic substitution 亲核取代 happens when a nucleophile 亲核试剂 (a lone-pair species) attacks that carbon and replaces the halogen.

Many halogenoalkane reactions are carried out by heating the mixture under reflux

To identify the halogen, warm the halogenoalkane with silver nitrate 硝酸银 in ethanol. A silver halide precipitate 沉淀 forms, and its colour shows which halogen is present (white $\text{AgCl}$, cream $\text{AgBr}$, yellow $\text{AgI}$).

The same halogenoalkane can instead undergo elimination 消去 to form an alkene 烯烃. The conditions decide which reaction wins:

• $\text{NaOH}$ in water → nucleophilic substitution → an alcohol.
• $\text{NaOH}$ in ethanol, heated → elimination → an alkene.

The same halogenoalkane: NaOH in water substitutes to an alcohol, while NaOH in ethanol (with heat) eliminates to an alkene

Nucleophilic substitution can follow two routes:

• S$_\text{N}$2: one step. The nucleophile attacks at the same time as the halogen leaves, passing through a crowded transition state 过渡态 where both are half-bonded. The rate depends on both the halogenoalkane and the nucleophile.
• S$_\text{N}$1: two steps. First the C–X bond breaks to give a carbocation 碳正离子; then the nucleophile attacks it. The rate depends only on the halogenoalkane.

S$_\text{N}$2 is a single step through a crowded transition state; S$_\text{N}$1 is two steps via a carbocation

Alkyl groups push electrons towards the positive carbon (the inductive effect 诱导效应), so they stabilise the carbocation. This is why:

• primary halogenoalkanes mostly react by S$_\text{N}$2.
• tertiary halogenoalkanes mostly react by S$_\text{N}$1 (their carbocation is well stabilised).
• secondary halogenoalkanes use a mixture of the two.

How fast a halogenoalkane reacts depends on the strength of the C–X bond, measured by its bond energy 键能. The C–I bond is the weakest, so iodoalkanes react fastest; the C–Cl bond is the strongest of the three, so chloroalkanes react slowest. So when tested with silver nitrate, an iodoalkane gives its precipitate first — its higher reactivity 反应活性 comes from the weaker C–X bond.

The weaker the C–X bond, the faster the halogenoalkane reacts — so iodoalkanes react fastest and chloroalkanes slowest

An alcohol 醇 has the $\text{–OH}$ (hydroxyl) functional group.

Ethanol, the alcohol in hand sanitiser, kills microbes

Making alcohols

Reactions of alcohols

• combustion: alcohols burn in oxygen to give carbon dioxide and water.
• substitution to a halogenoalkane, for example with $\text{HX}$, $\text{PCl}_5$, $\text{PCl}_3$ and heat, or $\text{SOCl}_2$.
• with sodium: alcohols react with sodium metal to give hydrogen and a sodium alkoxide — like water, but more slowly.
• oxidation 氧化 with acidified $\text{K}_2\text{Cr}_2\text{O}_7$ (or $\text{KMnO}_4$). The product depends on the class of alcohol (see below).
• dehydration 脱水 to an alkene, using a hot $\text{Al}_2\text{O}_3$ catalyst or concentrated acid.
• ester formation: an alcohol reacts with a carboxylic acid (with concentrated $\text{H}_2\text{SO}_4$ catalyst) to make an ester.

Three classes and how oxidation tells them apart

An alcohol is primary 伯, secondary 仲 or tertiary 叔, depending on how many carbons are joined to the carbon holding the $\text{–OH}$. Some molecules have more than one $\text{–OH}$ group.

Oxidation by class: a primary alcohol gives an aldehyde then a carboxylic acid, a secondary gives a ketone, a tertiary is not oxidised

In a quick test, acidified $\text{K}_2\text{Cr}_2\text{O}_7$ turns from orange to green with a primary or secondary alcohol, but stays orange with a tertiary alcohol.

Acidified dichromate turns orange to green with a primary or secondary alcohol, but stays orange with a tertiary alcohol

• distillation removes the aldehyde as it forms, before it can be oxidised further.
• reflux keeps boiling the mixture and returning the vapour, so the alcohol is fully oxidised to the carboxylic acid.

Distillation removes the aldehyde as it forms; reflux keeps boiling and returning the vapour, fully oxidising to the acid

A real distillation set-up: the vapour boils off, cools in the condenser and is collected — the same idea separates ethanol from a fermented mixture

The iodoform test

If you warm an alcohol that contains the $\text{CH}_3\text{CH(OH)}-$ group with alkaline aqueous iodine, you get a pale yellow precipitate of tri-iodomethane 三碘甲烷 ($\text{CHI}_3$) and the ion $\text{RCO}_2^-$. This is a useful test for that group.

Acidity of alcohols

The $\text{–OH}$ group makes alcohols very weakly acidic: they can lose the $\text{H}^+$ to form an $\text{RO}^-$ ion. But their acidity 酸性 is lower than that of water. This is because the alkyl group pushes electron density onto the oxygen, which makes the $\text{RO}^-$ ion less stable, so the alcohol holds onto its $\text{H}^+$ more tightly.

Aldehydes and ketones are carbonyl compounds 羰基化合物 — they contain the C=O carbonyl 羰基 group.

• in an aldehyde 醛 the carbonyl carbon is on the end of the chain (it also carries an H), written $\text{–CHO}$.
• in a ketone 酮 the carbonyl carbon is in the middle, between two other carbons.

Propanone (acetone), the simplest ketone, is the solvent in nail-polish remover

Both have the C=O carbonyl group: in an aldehyde it is on the end of the chain (–CHO), in a ketone it is in the middle

Making aldehydes and ketones

Both are made by the oxidation 氧化 of an alcohol 醇 with acidified $\text{K}_2\text{Cr}_2\text{O}_7$ or $\text{KMnO}_4$:

• a primary alcohol, with distillation 蒸馏, gives an aldehyde.
• a secondary alcohol gives a ketone.

Reactions

• reduction 还原 with $\text{NaBH}_4$ or $\text{LiAlH}_4$ turns a carbonyl compound back into an alcohol (an aldehyde gives a primary alcohol; a ketone gives a secondary alcohol).
• reaction with hydrogen cyanide 氰化氢 ($\text{HCN}$), with $\text{KCN}$ as catalyst and heat, adds $\text{H}$ and $\text{CN}$ across the C=O to make a hydroxynitrile 羟基腈. This adds one carbon to the chain.

The mechanism: nucleophilic addition

The reaction with $\text{HCN}$ is a nucleophilic addition 亲核加成. The carbonyl carbon is slightly positive (oxygen pulls the electrons away). So:

1. the $\text{CN}^-$ ion (a nucleophile) attacks the slightly positive carbon.
2. this breaks the C=O double bond, leaving a negative oxygen ($\text{O}^-$).
3. the $\text{O}^-$ takes an $\text{H}^+$ (from $\text{HCN}$) to finish the hydroxynitrile.

Nucleophilic addition of HCN: CN$^-$ attacks the $\delta+$ carbonyl carbon, the C=O breaks to O$^-$, then O$^-$ takes an H$^+$

Detecting any carbonyl

Add 2,4-DNPH reagent (2,4-dinitrophenylhydrazine). An orange precipitate confirms that the compound is an aldehyde or a ketone.

Telling an aldehyde from a ketone

Aldehydes are easily oxidised to carboxylic acids, but ketones are not. Two tests use this difference:

Telling an aldehyde from a ketone: an aldehyde gives a brick-red precipitate with Fehling's and a silver mirror with Tollens'; a ketone gives no change

A positive Tollens' test: an aldehyde coats the tube with a shiny silver mirror

The iodoform test

If the compound has the $\text{CH}_3\text{CO}-$ group, warming it with alkaline aqueous iodine gives a pale yellow precipitate of tri-iodomethane 三碘甲烷 ($\text{CHI}_3$) and the ion $\text{RCO}_2^-$.

A carboxylic acid 羧酸 has the $\text{–COOH}$ (carboxyl) functional group. It is a weak acid.

Vinegar is a dilute solution of ethanoic acid, a carboxylic acid

Making carboxylic acids

• oxidation 氧化 of a primary alcohol 醇 or an aldehyde 醛 with acidified $\text{K}_2\text{Cr}_2\text{O}_7$ or $\text{KMnO}_4$, with reflux 回流 (so it is fully oxidised).
• hydrolysis 水解 of a nitrile 腈 with dilute acid or alkali, then acidifying.
• hydrolysis of an ester 酯 with dilute acid or alkali and heat, then acidifying.

Reactions of carboxylic acids

These reactions all show that carboxylic acids are acids:

• with a reactive metal (a redox 氧化还原 reaction): gives a salt 盐 and hydrogen.

• with an alkali (a neutralisation 中和): gives a salt and water.

• with a carbonate 碳酸盐: gives a salt, water and carbon dioxide. The fizzing of $\text{CO}_2$ is a test for a carboxylic acid.

Carboxylic acids behave as acids: a salt with a metal (+ H$_2$), with an alkali (+ water), or with a carbonate (+ water + CO$_2$ — the fizz is a test)

Two reactions change the functional group:

• esterification 酯化 with an alcohol (concentrated $\text{H}_2\text{SO}_4$ catalyst) gives an ester.
• reduction 还原 by $\text{LiAlH}_4$ gives a primary alcohol.

An ester has the $\text{–COO–}$ group. It often smells sweet or fruity.

Esters give many fruits and perfumes their sweet smell

Making esters

An ester forms in a condensation 缩合 reaction between an alcohol and a carboxylic acid, with concentrated $\text{H}_2\text{SO}_4$ as catalyst. A water molecule is lost, and the reaction is reversible:

Esterification is a condensation: the –OH from the acid and the –H from the alcohol leave as water, forming the C–O–C ester link

Hydrolysis of esters

Hydrolysis splits the ester back apart. The conditions change the products:

• dilute acid and heat: reversible. Gives back the carboxylic acid and the alcohol.
• dilute alkali and heat: not reversible. Gives the alcohol and the salt of the carboxylic acid (the carboxylate ion).

Hydrolysing an ester: dilute acid (reversible) gives the carboxylic acid and alcohol; dilute alkali (not reversible) gives the carboxylate salt and alcohol

An amine 胺 has an $\text{–NH}_2$ group (the nitrogen replaces a hydrogen of ammonia).

Making a primary amine

Heat a halogenoalkane 卤代烷 with ammonia dissolved in ethanol, under pressure:

This is a nucleophilic substitution 亲核取代: the lone pair on the nitrogen of ammonia attacks the slightly positive carbon and pushes out the halogen. You use an excess of ammonia, or the amine made can react again.

Making a nitrile

Heat a halogenoalkane with potassium cyanide ($\text{KCN}$) in ethanol:

This is also a nucleophilic substitution, with the $\text{CN}^-$ ion as the nucleophile. It is useful because it adds one carbon to the chain. The product is a nitrile 腈.

Making a hydroxynitrile

Add $\text{HCN}$ (with $\text{KCN}$ as catalyst, and heat) to an aldehyde 醛 or ketone 酮. The $\text{H}$ and $\text{CN}$ add across the C=O bond to give a hydroxynitrile 羟基腈. The reagent is hydrogen cyanide 氰化氢, and the mechanism is nucleophilic addition 亲核加成.

Two ways cyanide adds a carbon: KCN with a halogenoalkane substitutes to a nitrile; HCN with a carbonyl adds to a hydroxynitrile

Hydrolysis of nitriles

Warm a nitrile with dilute acid (or dilute alkali, then acidify). This hydrolysis 水解 turns the $\text{–CN}$ group into a $\text{–COOH}$ group, giving a carboxylic acid 羧酸:

A nitrile can also be reduced by hydrogen and a catalyst to form an amine, which is the reduction 还原 route to a longer-chain amine.

Nitriles are a useful hub: KCN adds a carbon to make the nitrile, which then hydrolyses to a carboxylic acid or reduces to an amine

Reducing a nitrile gives an amine; amines and diacids link up to make polyamides such as nylon — the fibre first made famous in stockings

In addition polymerisation 加成聚合, many small molecules join into one very long chain, with no other product made.

Each small molecule is a monomer 单体. It must be unsaturated — it has a C=C double bond. The double bond opens up so that the monomers can link together. The long chain that forms is the polymer 聚合物.

Poly(ethene), a common addition polymer, is supplied as tiny pellets a few millimetres across; these are later melted and moulded into bottles, bags and other products

Repeat units

The repeat unit 重复单元 is the small part that is copied again and again along the chain. To find it, take the monomer, change the C=C to a single C–C, and draw bonds going out at each end.

• poly(ethene) 聚乙烯 is made from ethene:

• poly(chloroethene) 聚氯乙烯 (PVC) is made from chloroethene:

Finding the repeat unit: change the monomer's C=C to a single bond and draw bonds out at each end; reverse it to find the monomer

Finding the monomer

To go the other way, look at one repeat unit of the polymer, and put the C=C double bond back in. That gives you the monomer.

Poly(alkene)s are very hard to get rid of:

• they are non-biodegradable 不可生物降解 — microbes cannot break them down, so they stay in the ground for a very long time.
• their combustion 燃烧 (burning) can release harmful gases. For example, burning PVC gives off toxic hydrogen chloride.

Most addition polymers are non-biodegradable, so they build up as waste in the environment

Poly(alkene) waste is hard to dispose of: it is non-biodegradable, and burning PVC releases toxic hydrogen chloride

This topic does not add new reactions. Instead it asks you to join up the reactions you already know, so you can build a target molecule in several steps.

Medicines are built up from simple starting materials by multi-step organic synthesis

Identifying functional groups

A molecule may have more than one functional group 官能团. Use the test reactions from the syllabus to identify each one, and then predict how the molecule will behave. For example:

• decolourises bromine water → a C=C double bond (an alkene 烯烃).
• gives a precipitate with silver nitrate → a halogenoalkane 卤代烷.
• orange $\text{K}_2\text{Cr}_2\text{O}_7$ turns green → a primary or secondary alcohol 醇.
• orange precipitate with 2,4-DNPH → an aldehyde 醛 or ketone 酮.
• fizzes with a carbonate → a carboxylic acid 羧酸.

The standard identification tests: each reagent gives a characteristic observation that points to one functional group

A map of the AS reactions

Each row turns one functional group into another. Learn it as a map you can travel around:

A map of the AS reactions: each arrow turns one functional group into another. Work backwards from your target to plan a route

Planning a multi-step route

To devise a synthetic route 合成路线:

1. compare the target with the starting material — what has changed (the functional group, the number of carbons)?
2. work backwards from the target: which single reaction could make it, and from what?
3. repeat until you reach the starting material.
4. write each step with its reagent 试剂 and conditions.

If you need to add a carbon, the $\text{KCN}$ step is the key — it is the only AS reaction that lengthens the chain.

Analysing a route

When you are given a route, for each step state the type of reaction (such as oxidation 氧化, reduction 还原, substitution, addition or elimination) and the reagent used. Also think about possible by-products 副产物 — for example, making an amine from a halogenoalkane also gives a mixture of further-substituted amines, so the yield of the simple amine is low.