Acids, bases and salts
- Neutralisation as an ionic equationUse for every acid plus alkali reaction. The hydrogen ion from the acid and the hydroxide ion from the alkali combine to form water, and the spectator ions are left over as the dissolved salt.
Key concepts: **Acids, bases and alkalis defined by their ions**: An *acid* produces hydrogen ions, $\text{H}^+$, when dissolved in water, and this single ion causes every acidic property. A *base* is a metal oxide or metal hydroxide that neutralises an acid. An *alkali* is a base that is soluble in water, releasing hydroxide ions, $\text{OH}^-$. Every alkali is a base, but an insoluble base such as copper(II) oxide is not an alkali., **Basic and acidic oxides**: A *basic oxide* is a metal oxide that reacts with acids to give a salt and water, such as $\text{CaO}$ and $\text{CuO}$. An *acidic oxide* is a non-metal oxide that reacts with bases to give a salt and water, such as $\text{CO}_2$ and $\text{SO}_2$. As a rule, metal oxides are basic and non-metal oxides are acidic., **Solubility rules for common salts**: All sodium, potassium and ammonium salts are soluble, and all nitrates are soluble. Most chlorides are soluble except silver chloride and lead(II) chloride. Most sulfates are soluble except barium sulfate, calcium sulfate and lead(II) sulfate. Most carbonates and most hydroxides are insoluble., **The pH scale and indicator colours**: The pH scale runs from below 0 to 14: below 7 is acidic, 7 is neutral, above 7 is alkaline, and pH falls as the hydrogen-ion concentration rises. Litmus is red in acid and blue in alkali; thymolphthalein is colourless in acid and blue in alkali; methyl orange is red in acid and yellow in alkali., **The three reactions of dilute acids**: A dilute acid reacts in three set patterns. Acid plus a reactive *metal* gives salt plus hydrogen. Acid plus a *base* (metal oxide or hydroxide) gives salt plus water. Acid plus a *carbonate* gives salt plus water plus carbon dioxide. Learn the patterns and the products follow automatically.
Exam tips
- Decide first whether the salt is soluble or insoluble. A *soluble* salt is made by reacting an acid with an excess insoluble base, carbonate or metal and then crystallising, or by titration when the base is a soluble alkali. An *insoluble* salt is made by precipitation, mixing two soluble solutions and filtering off the solid.
- The acid decides the salt's family and the other reactant supplies the metal. Hydrochloric acid gives a *chloride*, sulfuric acid gives a *sulfate* and nitric acid gives a *nitrate*. Name the salt by combining the metal with the acid's ending.
Atoms, elements and compounds
- Number of neutrons in an atomUse to find the neutron count from the nucleon number $A$ and the proton number $Z$. Protons and neutrons both sit in the nucleus, so subtracting the protons from the nucleon total leaves the neutrons.
- Relative atomic mass from isotopic abundancesUse when an element exists as two or more isotopes with known percentage abundances. The result is a *weighted* mean, so it always lies closer to the more abundant isotope and need not be a whole number.
Key concepts: **Covalent bonding**: A *covalent bond* is a shared pair of electrons between two atoms. It forms between non-metal atoms, where each shared pair lets both atoms count those electrons towards a full outer shell. A *simple molecular* substance is made of small separate molecules held to each other by weak intermolecular forces., **Element, compound and mixture**: An *element* is a substance made of only one type of atom and cannot be broken down into anything simpler by chemical means. A *compound* is two or more elements chemically combined in a fixed ratio, with new properties of its own. A *mixture* is two or more substances that are not chemically joined, present in any ratio, each keeping its own properties., **Giant covalent structures**: A *giant covalent* (macromolecular) structure is a continuous network of atoms joined throughout by strong covalent bonds, with no small separate molecules. Diamond, graphite and silicon(IV) oxide are the required examples. Breaking the lattice means breaking many strong covalent bonds, so melting points are very high., **Ions and ionic bonding**: An *ion* is a charged particle formed when an atom loses or gains electrons. *Ionic bonding* is the strong electrostatic attraction between oppositely charged ions. It forms between a metal, which loses electrons to become a positive ion, and a non-metal, which gains them to become a negative ion, building a giant ionic lattice., **Metallic bonding**: *Metallic bonding* is the strong electrostatic attraction between a lattice of positive metal ions and a sea of delocalised electrons that move freely throughout the structure. Those delocalised electrons explain why metals conduct electricity even when solid and why their melting points are high., **The three subatomic particles**: A proton has relative mass $1$ and charge $+1$, a neutron relative mass $1$ and no charge, and an electron negligible mass (about $\frac{1}{1840}$) and charge $-1$. Protons and neutrons form the central nucleus while electrons occupy shells around it, so a neutral atom contains equal numbers of protons and electrons., **What isotopes are**: *Isotopes* are atoms of the same element with the same number of protons but different numbers of neutrons. They have identical electronic configurations, so their chemical properties are the same; they differ only in physical properties that depend on mass, such as density.
Chemical energetics
- Enthalpy change from bond energiesUse to calculate an enthalpy change from a table of bond energies. The reactant bonds broken take energy in and are added; the product bonds formed give energy out and are subtracted. The subtraction supplies the correct sign of $\Delta H$ automatically.
- Sign convention for the enthalpy changeUse to read the type of reaction directly from the sign of $\Delta H$. A negative value means the chemicals lose stored energy and release it (exothermic); a positive value means the chemicals gain stored energy from the surroundings (endothermic).
Key concepts: **Bond breaking and bond making**: *Bond breaking* requires energy to pull bonded atoms apart, so it is always endothermic. *Bond making* releases energy as atoms join, so it is always exothermic. Whether the whole reaction is exothermic or endothermic depends on which of these two energy transfers is larger., **Bond energy**: A *bond energy* is the energy needed to break one mole of a particular bond, which is exactly equal to the energy released when one mole of that same bond forms. One value therefore describes both breaking and making the bond, which is why a single table of bond energies can predict an enthalpy change., **Common exothermic and endothermic changes**: Common *exothermic* changes are the combustion of fuels, the neutralisation of an acid by an alkali, and the reaction of a reactive metal with an acid. Common *endothermic* changes are the thermal decomposition of a carbonate and the dissolving of ammonium nitrate or ammonium chloride in water., **Enthalpy change**: The *enthalpy change* $\Delta H$ is the energy transferred to or from the surroundings per mole of reaction, measured in kJ/mol. The symbol $\Delta$ means *change in* and $H$ is the enthalpy, the energy stored in the chemicals. An exothermic reaction releases stored energy and a thermal decomposition is a common endothermic process., **Exothermic and endothermic reactions**: An *exothermic* reaction transfers thermal energy to the surroundings, so the temperature of the surroundings rises. An *endothermic* reaction takes thermal energy in from the surroundings, so the temperature of the surroundings falls. The energy change is always judged from the surroundings, most usefully the water or solution holding the thermometer.
Exam tips
- Energy cannot be seen, so classify a reaction by the surroundings, not the chemicals. If the surroundings get hotter the reaction is exothermic; if they get colder it is endothermic. A mixture that goes cold has not released *cold*; it has absorbed energy from its surroundings.
Chemical reactions
- Average rate of reactionUse to find the mean rate over a measured interval, for example in $\text{cm}^3$ of gas per second or grams lost per second. Read the quantity from a gas syringe, a balance or a graph, then divide by the time elapsed.
- Oxidation number sum ruleUse to find an unknown oxidation number. The oxidation numbers of all atoms add up to zero in a neutral compound, or to the overall charge in an ion. Treat oxygen as $-2$ and hydrogen as $+1$ in their usual compounds, and an uncombined element as $0$.
Key concepts: **Collision theory**: *Collision theory* states that particles can only react when they collide, and the collision must have energy at least equal to the *activation energy* and the correct orientation. Most collisions are too gentle or badly aligned, so only a small fraction are successful; the rate of reaction is set by the number of successful collisions per second., **Conservation of mass**: In any chemical reaction atoms are only rearranged, never created or destroyed, so the total mass of the products equals the total mass of the reactants. An apparent gain in mass means a gas has been taken in from the surroundings, and an apparent loss means a gas has escaped into them., **Dynamic equilibrium**: A *reversible reaction*, shown by the sign $\rightleftharpoons$, reaches *dynamic equilibrium* in a closed system when the forward and reverse reactions occur at equal rates. At equilibrium the concentrations of reactants and products stay constant, though usually not equal, and both reactions are still taking place., **Oxidation and reduction (OIL RIG)**: A *redox reaction* is one in which oxidation and reduction happen together. *Oxidation* is the gain of oxygen, the loss of electrons, or an increase in oxidation number; *reduction* is the loss of oxygen, the gain of electrons, or a decrease in oxidation number. The mnemonic *OIL RIG* records the electron view: Oxidation Is Loss, Reduction Is Gain., **Physical and chemical changes**: A *physical change* rearranges the same particles without making any new substance, for example melting, boiling or dissolving, and it is usually easily reversed with no change in mass. A *chemical change* breaks and remakes bonds to form one or more *new substances* with different properties, signalled by a colour change, a gas, a precipitate or a permanent energy change, and it cannot be reversed by simple cooling or evaporation.
Exam tips
- When explaining any rate change, first decide which lever has moved: the *frequency* of collisions (how often particles meet) or the *proportion* of collisions that are successful (how many reach the activation energy). Concentration, pressure and surface area change only the frequency; temperature changes both; a catalyst changes only the proportion.
Chemistry of the environment
- Complete combustion of carbonUse as the model for burning a fuel in plenty of oxygen, where the only carbon product is carbon dioxide. A sulfur impurity burns in the same way, $\text{S} + \text{O}_2 \rightarrow \text{SO}_2$, releasing the sulfur dioxide that goes on to cause acid rain.
- Percentage by mass of an elementUse to compare fertilisers by the share of a nutrient they carry, most often the percentage of nitrogen. Count *every* atom of the element in the formula, so $\text{NH}_4\text{NO}_3$ contributes two nitrogen atoms, a mass of 28, not 14.
Key concepts: **Chemical tests for the presence of water**: Two colour tests show that a liquid contains water. Anhydrous copper(II) sulfate turns from white to blue, and anhydrous cobalt(II) chloride turns from blue to pink. Both confirm only that water is *present*, not that it is pure, because any aqueous solution gives the same change., **Composition of clean, dry air**: Clean, dry air is about 78% nitrogen, 21% oxygen and 1% argon by volume, with about 0.04% carbon dioxide and traces of other gases making up the rest. Nitrogen and oxygen together account for roughly 99%, so the remaining gases, though tiny in amount, include the carbon dioxide that drives climate change., **NPK fertilisers**: Fertilisers replace the elements that crops take from the soil. An *NPK fertiliser* supplies nitrogen for leafy growth, phosphorus for roots, and potassium for flowers and fruit. A single salt rarely holds all three, so an NPK product is usually a mixture, for example ammonium phosphate combined with potassium chloride., **The greenhouse effect and global warming**: The Earth's surface radiates thermal energy, *greenhouse gases* such as carbon dioxide and methane absorb it and re-emit some back towards the surface, and this keeps the lower atmosphere warmer than it would otherwise be. Rising concentrations of these gases trap more of this energy, raising the average global temperature, which is *global warming*., **Treating water to make it safe to drink**: The domestic supply is treated in three stages in a fixed order: *sedimentation* lets large insoluble particles settle out, *filtration* removes the finer suspended solids, and *chlorination* adds a small, controlled amount of chlorine to kill harmful microbes. The result is safe to drink but still holds dissolved salts, so it is not pure.
Exam tips
- Nitrogen makes up 78% of air yet is *not* a greenhouse gas, while carbon dioxide is only about 0.04% yet is one of the main causes of global warming. A gas warms the climate by absorbing the thermal energy the Earth radiates, so judge it by that property, never by how common it is.
Electrochemistry
- Anode half-equation for a halide ionUse at the anode, where a halide ion loses an electron and is oxidised to a halogen *molecule*. Two ions are always needed because the product is diatomic, for example $\text{Br}_2$ or $\text{Cl}_2$.
- Cathode half-equation for a metal ionUse at the cathode, where a metal cation gains electrons and is reduced to the neutral metal. The number of electrons equals the charge on the ion, so an $\text{Al}^{3+}$ ion needs three electrons while a $\text{Pb}^{2+}$ ion needs two.
- Overall fuel-cell reactionUse to summarise the overall change in a hydrogen-oxygen fuel cell, in which hydrogen and oxygen react to form water. It is exactly the reverse of the electrolysis of water, $2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2$.
Key concepts: **Electrolysis and the electrolytic cell**: Electrolysis is the breakdown of an ionic compound, when molten or in aqueous solution, by the passage of an electric current. The liquid decomposed is the *electrolyte*, and the current enters and leaves it through two *electrodes*. The *cathode* is joined to the negative terminal and the *anode* to the positive terminal. An *inert* electrode, such as carbon (graphite) or platinum, carries the current without itself reacting., **Oxidation at the anode, reduction at the cathode**: At the cathode, cations gain electrons, so reduction always occurs there; at the anode, anions lose electrons, so oxidation always occurs there. A useful memory aid is *red cat* (reduction at the cathode) and *an ox* (oxidation at the anode). Charge travels through the metal wires as moving electrons and through the electrolyte as moving ions., **Products from a molten binary compound**: A solid ionic compound does not conduct, because its ions are fixed in the lattice. Once molten the ions are free to move, so electrolysis can occur: the metal is discharged at the cathode and the non-metal at the anode. For molten lead(II) bromide this gives lead at the cathode and bromine at the anode., **The hydrogen-oxygen fuel cell**: A hydrogen-oxygen fuel cell combines hydrogen and oxygen to produce electrical energy directly, with water as the only chemical product. Hydrogen is the *fuel* that is oxidised and oxygen is the *oxidising agent*. Unlike an ordinary battery, the reactant gases are supplied continuously from outside, so the cell runs as long as they are fed in.
Exam tips
- A correct half-equation balances both the atoms and the total charge. Add electrons to whichever side makes the charges equal: electrons appear on the right for an oxidation (loss) and on the left for a reduction (gain). Then check that the number of electrons matches the charge that has changed.
Experimental techniques and chemical analysis
- Moles from concentration and volumeUse to find the amount in moles $n$ from a concentration $c$ in $\text{mol/dm}^3$ and a volume $V$ that has first been converted to $\text{dm}^3$. This is the opening step of every titration calculation.
- Retardation factorUse to identify a component on a chromatogram. Measure the substance distance from the pencil baseline to the centre of its spot, and the solvent distance to the solvent front. The value has no units and always lies between $0$ and $1$.
Key concepts: **Apparatus for measuring volume**: A *burette* delivers any chosen volume of liquid accurately, read to $0.1\ \text{cm}^3$, and provides the variable volume in a titration. A *volumetric pipette* delivers one fixed accurate volume such as $25.0\ \text{cm}^3$. A *measuring cylinder* is far less precise and is used only for approximate volumes. A balance measures mass, not volume, so it can never answer a volume question., **Choosing a separation technique**: *Filtration* separates an insoluble solid from a liquid, leaving the solid as the residue and the liquid as the filtrate. *Crystallisation* recovers a dissolved solid by evaporating to a saturated solution then cooling slowly. *Simple distillation* recovers a pure solvent from a solution, and *fractional distillation* separates two miscible liquids with different boiling points., **Flame tests for metal cations**: A clean wire dipped in the sample and held in a hot flame gives a characteristic colour: lithium red, sodium yellow, potassium lilac, calcium orange-red and copper(II) blue-green. Sodium yellow and calcium orange-red are deliberately set as look-alikes, as are potassium lilac and lithium red, so the five colours must be learned exactly., **Tests for common gases**: Hydrogen gives a squeaky pop with a lighted splint; oxygen relights a glowing splint; carbon dioxide turns limewater milky; ammonia turns damp red litmus blue; chlorine bleaches damp litmus paper; and sulfur dioxide turns acidified potassium manganate(VII) from purple to colourless., **What a titration measures**: A *titration* finds the volume of one solution that exactly reacts with a fixed volume of another, so an unknown concentration can be calculated. The *end-point* is where the indicator changes colour, showing the acid and alkali have exactly reacted. Only *concordant* titres, agreeing to within about $0.10\ \text{cm}^3$, are averaged; the rough first run is discarded.
Exam tips
- Concentration is measured in $\text{mol/dm}^3$, so each volume must be changed from $\text{cm}^3$ to $\text{dm}^3$ by dividing by $1000$ before it is used in $n = c \times V$. Substituting a volume in $\text{cm}^3$ gives an answer one thousand times too large.
Metals
- Displacement of a less reactive metalUse when a more reactive metal is placed in a solution of a less reactive metal's salt. The reactive metal forms ions and the less reactive metal is deposited. The anion, here $\text{SO}_4^{2-}$, is a spectator and is left out of the ionic equation.
- Reaction of a metal with dilute acidUse for any metal above hydrogen in the reactivity series, for example $\text{Zn} + \text{H}_2\text{SO}_4 \rightarrow \text{ZnSO}_4 + \text{H}_2$. The metal atom loses electrons to form a positive ion. Metals below hydrogen (copper, silver, gold) give no reaction.
Key concepts: **Conditions and product of rusting**: *Rusting* is the corrosion of iron and steel. It needs *both water and oxygen* present; salt and acidity speed it up but are not essential. The product is hydrated iron(III) oxide, $\text{Fe}_2\text{O}_3 \cdot x\text{H}_2\text{O}$, the orange-brown flaky solid. Remove either water or oxygen and rusting stops., **Extraction method from position relative to carbon**: A metal's extraction method follows its position relative to *carbon*. Metals below carbon (zinc, iron, copper) are extracted by *reduction with carbon*, the cheaper route. Metals above carbon (potassium to aluminium) hold their oxygen too strongly and must be extracted by *electrolysis* of the molten compound. Very unreactive metals such as gold are found native., **Metallic bonding and the properties it explains**: A metal is a giant lattice of positive ions in a *sea of delocalised electrons* that move freely through the whole structure. The mobile electrons make a metal a good conductor of electricity and heat. The layers of ions slide without shattering, so a metal is *malleable* and *ductile*, and strong metallic bonding gives high melting and boiling points., **Order of the reactivity series**: Metals are listed in order of how readily their atoms lose electrons to form positive ions: *potassium, sodium, calcium, magnesium, aluminium*, [carbon], *zinc, iron*, [hydrogen], *copper, silver, gold*. Carbon and hydrogen are non-metal reference points. A higher metal reacts more vigorously and is harder to extract., **Uses of aluminium and copper**: *Aluminium* is chosen for aircraft bodies, overhead power cables and food cans because it has a low density and resists corrosion through a self-sealing oxide layer. *Copper* is chosen for electrical wiring and saucepan bases because it is an excellent electrical and thermal conductor, is ductile, and does not react with water., **What an alloy is and why it is harder**: An *alloy* is a mixture of a metal with one or more other elements, usually other metals. It is harder and stronger than the pure metal because the added atoms are a *different size*, so they disrupt the regular layers of ions and stop the layers sliding over one another. The metallic bonding remains, so an alloy still conducts.
Organic chemistry
- General formula of the alcoholsUse to write the formula of any alcohol from its carbon number $n$; for example ethanol has $n = 2$, giving $\text{C}_2\text{H}_5\text{OH}$. The *hydroxyl* group $-\text{OH}$ is the functional group, so every member shows the characteristic alcohol reactions.
- General formula of the alkanesUse to write the molecular formula of any alkane from its carbon number $n$; for example pentane has $n = 5$, so its formula is $\text{C}_5\text{H}_{12}$. Alkanes are *saturated*, containing only single carbon-carbon bonds, which is why they hold the maximum possible number of hydrogen atoms.
- General formula of the alkenesUse to write the molecular formula of any alkene from its carbon number $n$; for example propene has $n = 3$, so its formula is $\text{C}_3\text{H}_6$. Alkenes are *unsaturated*, carrying one carbon-carbon double bond $\text{C}=\text{C}$ and therefore two fewer hydrogens than the matching alkane.
- General formula of the carboxylic acidsUse to write the formula of any carboxylic acid; for example ethanoic acid takes $n = 1$ to give $\text{CH}_3\text{COOH}$. The *carboxyl* group $-\text{COOH}$ is the functional group, and it makes these compounds *weak acids* that still show every typical acid reaction.
Key concepts: **Addition and condensation polymerisation**: A *polymer* is a very large molecule built when many small *monomers* join together. In *addition polymerisation* many alkene monomers join by opening their $\text{C}=\text{C}$ double bonds, forming no other product. In *condensation polymerisation* two monomers, each with two reactive groups, join with the loss of a small molecule (usually water) at every linkage, forming a polyester or a polyamide., **Homologous series and functional groups**: A *functional group* is the atom or group of atoms that gives a molecule its characteristic reactions, such as the $-\text{OH}$ group of the alcohols. A *homologous series* is a family of compounds that share the same functional group and the same general formula, whose consecutive members differ by one $\text{CH}_2$ unit. Within a series the functional group fixes the chemical reactions while the chain length sets the physical properties, so boiling point rises steadily down the series., **How an organic name is built**: An organic name is built from a *stem* that counts the carbon atoms in the longest chain (meth, eth, prop, but) and a *suffix* that names the functional group (*-ane*, *-ene*, *-ol*, *-oic acid*). A *locant* number is inserted when the group could sit on more than one carbon, as in but-2-ene, and is omitted when only one position is possible, as in ethene., **Petroleum and fractional distillation**: *Petroleum* (crude oil) is a mixture of hydrocarbons, mostly alkanes of many different chain lengths, and is almost useless until separated. *Fractional distillation* sorts it into *fractions* of similar boiling point in a column that is hot at the bottom and cool at the top. Down the column the chains are shorter: viscosity falls, the colour lightens and flammability rises., **Testing for unsaturation with bromine water**: An *unsaturated* compound contains at least one $\text{C}=\text{C}$ double bond, while a *saturated* compound has only single carbon-carbon bonds. The test is to shake the compound with *bromine water*: an alkene adds bromine across its double bond and *decolourises* the orange solution, whereas an alkane has no double bond and leaves the colour unchanged.
States of matter
- Relative molecular massUse to find the relative molecular mass of a gas by adding the relative atomic mass of every atom shown in its formula. In this chapter $M_r$ is the quantity that fixes how fast a gas diffuses, so compute it before comparing two gases.
Key concepts: **Changes of state**: A *change of state* is a reversible physical change in which no new substance forms. Melting (solid to liquid) and boiling (liquid to gas) take in energy; freezing (liquid to solid) and condensation (gas to liquid) give out energy. *Sublimation* is the direct change of a solid to a gas, shown by solid carbon dioxide and iodine., **Diffusion and the random motion of particles**: *Diffusion* is the net movement of particles from a region of higher concentration to a region of lower concentration, caused by their random motion. No stirring or draught is needed. Diffusion is fastest in gases, slower in liquids and negligible in solids, because the particles are progressively closer together and slower moving., **Properties of the three states**: Every bulk property follows from the particle picture. A solid keeps a fixed shape because its particles are locked in place; liquids and gases flow because their particles can move. Solids and liquids keep a fixed volume and resist compression because their particles already touch, whereas a gas fills its container, has a low density and is easily compressed because of the large spaces between its particles., **The three states in the kinetic particle model**: The *kinetic particle model* states that all matter is made of tiny particles in constant motion. In a *solid* the particles touch in a regular arrangement and only vibrate about fixed positions. In a *liquid* they touch in a random arrangement and slide past one another. In a *gas* they are far apart in a random arrangement and move quickly in all directions.
Exam tips
- At the same temperature all gas molecules have the same average kinetic energy, so a lighter molecule must move faster. The smaller the relative molecular mass, the faster the gas diffuses. Always *calculate* each $M_r$ rather than judging from the size of the formula, because carbon monoxide ($M_r = 28$) is heavier than methane ($M_r = 16$) despite looking small.
- Place the given temperature on a number line marked with the melting point and the boiling point. Below the melting point the substance is solid, between the two points it is liquid, and above the boiling point it is gas. Watch negative values: a boiling point of $-183$ °C lies below room temperature, so the substance is already a gas at 25 °C.
Stoichiometry
- Amount of substance from massUse to convert between a mass $m$ in grams and an amount $n$ in moles, where $M$ is the molar mass in g/mol. This is the first move of almost every reacting-mass calculation.
- Number of particles from amountUse to find how many atoms, molecules or ions are present, where $N_A = 6.02 \times 10^{23}$ per mole. Read the substance carefully, because one mole of $\text{O}_2$ holds twice as many atoms as molecules.
- Relative atomic mass from isotopic abundancesUse when an element exists as two or more isotopes and you are given the mass and percentage abundance of each. The result is a *weighted* mean, so it always lies closer to the more abundant isotope.
Key concepts: **Deducing an ionic formula from charges**: An ionic compound is overall neutral, so its formula is the smallest whole-number ratio of ions whose charges cancel. Combine the ions so the total positive charge equals the total negative charge, then reduce to lowest terms. A polyatomic ion that takes a subscript above one must be bracketed first, as in $\text{Ca(NO}_3)_2$., **Empirical and molecular formula**: The *molecular formula* gives the actual number of atoms of each element in one molecule, for example ethane $\text{C}_2\text{H}_6$. The *empirical formula* gives the simplest whole-number ratio of those atoms, here $\text{CH}_3$. An ionic compound is always written as its empirical formula, such as $\text{NaCl}$, because a giant lattice has no discrete molecules., **Relative atomic, molecular and formula mass**: The *relative atomic mass* $A_r$ is the average mass of an element's atoms on a scale where one carbon-12 atom is exactly 12. The *relative molecular mass* $M_r$ is the sum of the $A_r$ values of every atom in a molecule; for an ionic compound the same sum is called the *relative formula mass*. All three are ratios, so they have no units., **The mole and the Avogadro constant**: The *mole* is the chemist's counting unit: one mole of any substance contains $6.02 \times 10^{23}$ particles, a number called the *Avogadro constant* $N_A$. The mole is the hub of stoichiometry because mass, number of particles, gas volume and concentration are each linked to one another only through the amount in moles.
Exam tips
- To balance a symbol equation, adjust only the *balancing numbers* in front of each formula, never the subscripts inside a formula. The number of atoms of every element must then be equal on both sides. Changing a subscript changes the substance itself and scores no marks.
The Periodic Table
- Group I metal reacting with waterUse for any alkali metal M reacting with water. The products are a soluble metal hydroxide, which makes the solution alkaline and turns universal indicator purple, and hydrogen gas. The gas is hydrogen displaced from the water, never oxygen.
- Halogen displacement reactionUse when a halogen $\text{X}_2$ is added to a solution of a halide $\text{Y}^-$. The reaction proceeds only if X is more reactive than Y, so chlorine displaces bromide and iodide, bromine displaces only iodide, and iodine displaces neither.
- Ion charge predicted from Group numberUse to predict the charge of a simple ion straight from its Group. A metal in Groups I to III loses electrons to give a positive charge equal to the Group number; a non-metal in Groups V to VII gains electrons to give a negative charge equal to eight minus the Group number. Group IV elements rarely form simple ions.
Key concepts: **Group I: the alkali metals**: The Group I elements (lithium, sodium, potassium and below) are the *alkali metals*: soft, low-density, low-melting metals, each with one outer-shell electron. Every Group I atom loses that single electron to form a 1+ ion, so the Group has one fixed valency. They are stored under oil because they react quickly with oxygen and water vapour in air., **Group VII: the halogens**: The Group VII elements (fluorine, chlorine, bromine, iodine) are the *halogens*: reactive non-metals that exist as *diatomic* molecules such as $\text{Cl}_2$. Each atom has seven outer electrons and gains one more to form a 1− *halide* ion such as $\text{Cl}^-$, or shares a pair of electrons to bond covalently., **Metals, non-metals and the dividing line**: Metals occupy the large region on the left of the table and non-metals the smaller region to the upper right, separated by a staircase line. Position predicts the ion sign: a metal has few outer electrons and loses them to form a positive ion, while a non-metal has a nearly full shell and gains electrons to form a negative ion. An element on the line, such as silicon, is a *metalloid* with intermediate properties., **Noble gases: a full outer shell**: The Group VIII (Group 0) elements (helium, neon, argon and below) are the *noble gases*: colourless, *monatomic* gases that are chemically *inert*. Their atoms already have a complete outer shell (eight electrons, or two for helium), so they have no tendency to gain, lose or share electrons., **Reading the Periodic Table: Groups and Periods**: A *Group* is a vertical column; for Groups I to VII every element has the same number of outer-shell electrons, which gives the Group similar chemical properties. A *Period* is a horizontal row, and the Period number equals the number of occupied electron shells. The elements are arranged in order of increasing proton number., **Transition elements: the defining properties**: The *transition elements* form the central block between Group II and Group III. Compared with the Group I metals they are hard, dense and high-melting, and they show three signatures: *variable oxidation states* (for example $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$), *coloured compounds*, and use as *catalysts*, such as iron in the Haber process.