China 高一 Chemistry

Chapters: 8

1. Ch.1 Substances and their changes

Classification and transformation; colloids · Ionic reactions; electrolytes · Redox reactions

Colloids and Dispersed Systems – A dispersed system is one substance spread as small bits (the dispersed phase) through another (the dispersion medium). By particle size it is a true solution (below about 1 nm), a colloid (about 1–100 nm) or a suspension (above about 100 nm). Colloid particles are too small to see or to settle, but big enough to scatter light, so a beam passing through shows its path: the Tyndall effect. Colloids are grouped by the states of the two phases (fog, smoke, milk, foam, gel). Colloid particles often carry the same charge and repel each other; adding an electrolyte cancels the charge, so they clump and settle: coagulation.
Ionic Equations and Ion-Exchange Reactions – Salts, strong acids and alkalis split into ions in water: they are electrolytes. When two such solutions mix, a reaction happens only if some ions are removed from the solution as a precipitate, a gas or water. A full ionic equation shows every ion; crossing out the ions that do not change (spectator ions) gives the net ionic equation, for example Ag⁺ + Cl⁻ → AgCl↓. Charges and atoms must balance on both sides.
Redox Reactions: From Oxygen to Electron Transfer – Oxidation first meant adding oxygen or removing hydrogen. Reduction meant the opposite. Today we use a bigger idea: oxidation is losing electrons and reduction is gaining electrons. Both always happen together, so we call them redox reactions. A more active metal gives electrons to the ion of a less active metal.

2. Ch.2 Sodium and chlorine

Sodium and compounds · Chlorine and compounds · Amount of substance; molar concentration · Lab 1: prepare a solution of given molarity

Alkali Metals and Sodium Compounds – Alkali metals (lithium, sodium, potassium, rubidium, caesium) are group 1 of the periodic table. Each atom has one outer electron that it loses easily, forming M⁺ ions. They are soft, light and very reactive: they react with water to make hydrogen and an alkali, and they burn in air. Reactivity rises down the group. Sodium gives Na₂O, Na₂O₂, Na₂CO₃ (washing soda) and NaHCO₃ (baking soda), and each metal colours a flame in its own way.
Halogens: The Group 17 Elements – The halogens are fluorine, chlorine, bromine, iodine and astatine, in group 17 of the periodic table. Each atom has 7 outer electrons, so it gains 1 electron to form a halide ion (X⁻). Down the group, atoms get bigger, colours get darker, melting and boiling points rise, and reactivity falls. A more reactive halogen displaces a less reactive one from its salt.
Mole Concept, Molar Mass and Chemical Formulas – Atoms are far too small to count one by one, so chemists count them by weighing. Atomic masses are given in u, where 1 u is one-twelfth the mass of a carbon-12 atom. One mole is 6.022 × 10²³ particles, and its mass in grams equals the formula mass in u. With moles we can find the percentage of each element in a compound, and work back from percentages to the empirical and molecular formulas.
Lab: Prepare a Solution of Given Molarity – To prepare a solution of a known molarity (mol/L) you first calculate the mass of solute from mass = molarity × volume in litres × molar mass. You weigh it, dissolve it in a little water in a beaker, pour it through a funnel into a volumetric flask, rinse the beaker three times into the flask, and then add water until the bottom of the meniscus touches the ring mark, with your eye level with the mark. A stopper and 10 inversions mix the solution. Errors come from reading the mark wrongly, overshooting the mark, losing solute in transfer, using a hot solution or using a wet or dirty flask.

3. Ch.3 Iron; metal materials

Iron and compounds · Metal materials · Lab 2: properties of iron compounds

Iron and Its Compounds – Iron (Fe, atomic number 26) is a grey, magnetic transition metal. Its atom can lose 2 or 3 electrons, so it makes two families of compounds: iron(II) with Fe²⁺ and iron(III) with Fe³⁺. Iron reacts with dilute acids to give Fe²⁺ and hydrogen, with chlorine to give FeCl₃, with steam to give Fe₃O₄, and with damp air to give rust. The three oxides are FeO (black, Fe²⁺), Fe₂O₃ (red-brown, Fe³⁺) and Fe₃O₄ (black, magnetic, both ions). The two hydroxides, Fe(OH)₂ (white-green) and Fe(OH)₃ (red-brown), do not dissolve in water; Fe(OH)₂ turns into Fe(OH)₃ in air. An oxidising agent changes Fe²⁺ into Fe³⁺ and a reducing agent changes Fe³⁺ back into Fe²⁺.
Corrosion, Its Prevention and Alloys – Corrosion is the slow eating away of a metal by air, moisture or chemicals around it. Iron rusts only when both air (oxygen) and water are present, forming hydrated iron oxide. It can be prevented by painting, oiling, galvanising, chrome plating, anodising or making alloys. An alloy is a uniform mixture of a metal with other metals or a non-metal; alloys are usually harder, stronger and more corrosion-resistant than the pure metal.
Lab: Properties of Iron Compounds (Testing Fe²⁺ and Fe³⁺) – In this practical you identify Fe²⁺ and Fe³⁺ in solution. Sodium hydroxide gives a white-green solid with Fe²⁺ (turning brown in air) and a red-brown solid with Fe³⁺. Potassium thiocyanate (KSCN) turns Fe³⁺ blood red but leaves Fe²⁺ unchanged. Hydrogen peroxide oxidises Fe²⁺ to Fe³⁺ (then KSCN turns red), and iron powder reduces Fe³⁺ to Fe²⁺ (then KSCN gives no red). The lab shows that iron(II) is reducing, iron(III) is oxidising, and each can change into the other.

4. Ch.4 Structure and periodic law

Atomic structure and periodic table; nuclides · Periodic law · Chemical bonds · Lab 3: trends in periods and groups

Periodic Classification of Elements – Scientists sorted elements so that similar ones sit together. Döbereiner made triads, Newlands found that every 8th element repeats (octaves), Mendeleev arranged elements by atomic mass and left gaps for unknown ones. Moseley showed that atomic number is the real key. Modern periodic law: the properties of elements are a periodic function of their atomic numbers. The modern table has 7 periods and 18 groups. Elements with Z > 100 get temporary IUPAC names built from digit roots (nil, un, bi, tri, quad, pent, hex, sept, oct, enn) plus -ium.
Periodic Trends in Properties – Two forces decide almost every trend: the pull of the nucleus (effective nuclear charge) and the distance of the outer shell. Across a period the nuclear pull grows while the shell stays the same, so atoms shrink, ionisation enthalpy rises, electron gain enthalpy becomes more negative and electronegativity rises. Down a group a new shell is added, so atoms grow and these values fall. Cations are smaller and anions bigger than their atoms. Valence follows the outer electrons; metallic reactivity is highest at the bottom left and non-metallic reactivity at the top right.
Kossel-Lewis Approach and the Ionic Bond – Atoms join so that each gets a stable outer shell of 8 electrons (an octet), like a noble gas. Kossel said atoms can give or take electrons to make ions (ionic bond). Lewis said atoms can also share pairs of electrons (covalent bond). Lewis structures show these electrons as dots and lines. Formal charge (V − L − B/2) helps pick the best Lewis structure. Ions pack into a crystal, and the energy released is linked to the lattice enthalpy.
Lab: Trends in a Period and a Group of Elements – In a lab we can see periodic trends with our own eyes. Across period 3 from sodium to chlorine, metal activity falls (sodium fizzes in water, magnesium barely, aluminium not at all), non-metal activity rises (chlorine is the strongest), and oxides change from strong base to strong acid. Down group 1 the metals get more active (Li, Na, K). Down group 17 the halogens get less active, so chlorine pushes bromine and iodine out of their salts, and bromine pushes out iodine.

5. Ch.5 Important non-metals

Sulfur and compounds · Nitrogen and compounds; acid rain · Inorganic non-metallic materials · Lab 4–5: removing ions from crude salt; sulfur conversions

Sulfur and Its Compounds – Sulfur (S, atomic number 16, electrons 2, 8, 6) is a yellow non-metal in group 16 with oxygen. Solid sulfur is made of S₈ rings and has rhombic and monoclinic forms. It burns to sulfur dioxide (SO₂), a gas that bleaches, kills microbes and causes acid rain. In the contact process SO₂ is oxidised to SO₃ over a V₂O₅ catalyst and turned into sulfuric acid (H₂SO₄), a strong acid that also dehydrates and oxidises. Hydrogen sulfide (H₂S) smells of bad eggs. Sulfate ions are found with barium chloride: a white precipitate of BaSO₄ that does not dissolve in dilute HCl.
Nitrogen and Its Compounds – Nitrogen gas (N₂) makes up 78% of air. Its triple bond makes it very unreactive, so plants cannot use it directly. Fixation turns N₂ into compounds: lightning and bacteria do it naturally, and the Haber process makes ammonia (NH₃) industrially. Ammonia is a basic gas that forms ammonium salts. At high temperature N₂ and O₂ form NO, which becomes brown NO₂; NO₂ in rain makes nitric acid and acid rain. Nitric acid is a strong acid and oxidising agent used for fertilisers.
Lab: Purify Crude Salt and Change Sulfur Compounds – Crude salt holds Ca²⁺, Mg²⁺ and SO₄²⁻ impurities. We remove them by precipitation: BaCl₂ takes out sulfate as BaSO₄, NaOH takes out magnesium as Mg(OH)₂, and Na₂CO₃ takes out calcium (and the extra barium) as carbonates. Then we filter, neutralise the extra OH⁻ and CO₃²⁻ with dilute HCl, and evaporate to get pure NaCl crystals. In part two we watch sulfur move between oxidation numbers −2 (H₂S), 0 (S), +4 (SO₂) and +6 (H₂SO₄): going up is oxidation, going down is reduction.

6. Ch.6 Reactions and energy

Heat and electric energy from reactions; cells · Rate and limit of reactions · Lab 6–7: chemical to electrical energy; rate factors

Chemical Energy: Heat, Bond Energy and Cells – A chemical reaction breaks old bonds (energy in) and makes new bonds (energy out). If more energy goes out than in, the reaction is exothermic and heat is released (ΔH negative). If more goes in, it is endothermic (ΔH positive). With bond energies, ΔH = energy to break bonds − energy released making bonds. If the electrons of a redox reaction are made to travel through a wire instead of jumping directly, the energy comes out as electricity: that is a galvanic cell, such as Zn–Cu giving 1.1 V. Joining cells in series gives a battery. Primary cells are used once, secondary cells are rechargeable, and fuel cells are fed with fuel.
Rate of a Chemical Reaction – The rate of a reaction tells how fast a reactant is used up or a product is made, per unit time. Rate = −Δ[R]/Δt = +Δ[P]/Δt (unit mol L⁻¹ s⁻¹). The rate law, rate = k[A]^x[B]^y, is found by experiment; x + y is the order. Molecularity is the number of particles that collide in one elementary step.
Lab: Make a Simple Cell and Test Rate Factors – Part A: put a zinc plate and a copper plate in dilute acid, join them through a voltmeter and read about 1.0 V. Zinc gives electrons (anode), hydrogen forms at copper (cathode). A different metal pair gives a different voltage, and two identical plates give 0 V. Part B: plot gas collected against time. The slope of the line is the rate. A higher concentration, a higher temperature (roughly doubling per 10 °C), smaller pieces and a catalyst all raise the rate, and a fair test changes only one of them at a time. The final amount of product does not change, only how fast it forms.

7. Ch.7 Organic compounds

Introduction to organic compounds · Ethylene and polymers · Ethanol and acetic acid · Basic nutrients · Lab 8–9: ball-and-stick models; ethanol and acetic acid

Carbon and Its Compounds: Bonding, Hydrocarbons and Naming – Carbon has 4 outer electrons, so it shares electrons (covalent bonds) instead of gaining or losing them. Because it bonds to itself (catenation) and always makes 4 bonds (tetravalency), it forms millions of compounds: chains, branches and rings, saturated or unsaturated. Compounds with the same functional group form a homologous series that differs by –CH₂–, and IUPAC names are built from the number of carbons + a suffix or prefix for the functional group.
Polymers – A polymer is a giant molecule made by joining many small molecules (monomers) into a long chain. In addition polymerisation, monomers with a C=C double bond open up and link with nothing lost (ethene → poly(ethene)). In condensation polymerisation, two kinds of monomer with reactive groups at both ends join and give off a small molecule such as water at every link (nylon, polyester). Nature makes polymers too: starch, cellulose, proteins, DNA, rubber. Separate chains give thermoplastics that melt and can be recycled; cross-linked chains give thermosets that never melt. Most plastics do not rot, so we must reduce, reuse and recycle them.
Chemical Properties of Carbon Compounds, Ethanol and Ethanoic Acid – Carbon compounds burn in oxygen to give CO₂, water, heat and light (combustion); strong oxidising agents turn alcohols into acids (oxidation); unsaturated compounds add hydrogen with a nickel catalyst (addition); saturated compounds swap hydrogen for chlorine in sunlight (substitution). Ethanol reacts with sodium to give hydrogen and dehydrates to ethene; ethanoic acid is a weak acid that makes esters with alcohols and CO₂ with carbonates.
Biomolecules – A cell is mostly water, plus four big families of carbon compounds: proteins, carbohydrates, lipids and nucleic acids. Grinding tissue in acid separates small molecules (acid-soluble pool) from big ones – proteins, polysaccharides and nucleic acids (acid-insoluble pool). Proteins are chains of amino acids folded into four levels of structure. Polysaccharides are chains of sugars; lipids are fatty acids on glycerol; nucleic acids are chains of nucleotides. Enzymes are protein catalysts that bind a substrate at the active site, lower the activation energy, and are affected by temperature, pH, substrate level and inhibitors.

8. Ch.8 Chemistry and sustainability

Use of natural resources · Rational use of chemicals · Environmental protection; green chemistry

Chemistry in Everyday Life: Medicines, Food, Cleaning and Soil – Chemistry is at work in your home every day. Antacids are weak bases that calm extra stomach acid. Medicines are grouped by what they do: painkillers, antibiotics, antiseptics, antihistamines, tranquilisers. Food chemicals keep food safe and tasty: preservatives, antioxidants and sweeteners. Soap cleans because one end of its molecule loves water and the other loves oil. Farmers add lime to acid soil and fertilisers to feed crops. Some metal ions are needed by the body, others are toxic, which is one reason we recycle.
Green Chemistry – Green chemistry means designing chemical products and processes that make less waste, use safer substances, save energy and use raw materials that can be renewed. Key tools are high atom economy, catalysts, safer solvents such as water, renewable feedstocks, and products that break down safely. A life cycle assessment checks the impact of a product from raw material to disposal.