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The Big Picture

I have been helping my daughter revise chemistry for her Gateway Science exam this morning and was appalled at the incoherence of the subject. In the end I had to write out the “big picture” for her:

• Chemistry is all about electrons and energy.
• Negatively charged electrons occupy 'shells' around the nucleus.
• The first shell has 2 electrons.
• The next two shells have 8 electrons.
• The nucleus consists of positively charged protons and neutral neutrons.
• The 'atomic number' of an element is the number of protons in the nucleus.
• The 'atomic weight' is the atomic number plus the number of neutrons.
• There are four main types of elements at GCSE:
  1. hydrogen
  2. noble gases
  3. metallic
  4. non-metallic
• Non-metallic elements don't include metallic elements and noble gases are ignored.
• Non-metallic atoms can share electrons between their outer shells in covalent bonds.
• There are four type of materials at GCSE:
  1. molecular
  2. salts
  3. metals
  4. noble gases
• Covalent-bonded atoms are called either 'molecular ions' or 'molecules', depending on their electric charge.
• In molecules, the total numbers of protons and electrons are the same and so they are neutral.
• In molecular ions, the total numbers of protons and electrons are different and so they carry a charge.
• Molecular ions generally have more electrons than protons and so are negatively charged.
• Metal is a substance consisting of metallic ions and loosely associated electrons.
• Salt is a substance consisting of metallic ions and negatively charged ions.
• Each type of salt crystal or metallic crystal has a repeated arrangement of ions.
• Different types of crystal can have different arrangements.
• When atoms or ions are 'bonded', their relative positions are fixed.
• There are 3 main types of bond studied in GCSE chemistry:
  1. Covalent bonding within molecules and molecular ions.
  2. Ionic bonding in solid salts.
  3. Metallic bonding in solid metals.

I have several general rules that bring together the different topics in GCSE chemistry. These contradict much of the muddled teaching.

• Electron shells must be full.
• Only noble gas atoms (e.g. helium) have the same number of protons and electrons.
• Metallic atoms have more protons in their nucleus than electrons in their shells and are positively charged 'metallic ions'.
• Metallic ions are always positively charged.
• Metallic ions don't gain, lose or share electrons.
• When a metallic element is shown as a non-charged substance in an equation, it means the metal substance and not neutral atoms.
• Non-metallic atoms have more electrons than protons.
• A single, non-metallic atom is a negatively charged atomic ion.
• Only covalent bonds are 'chemical' bonds within molecules and molecular ions.
• Ionic bonds and metallic bonds are 'physical' bonds between ions in crystals.
• Crystals are solid.
• Only covalent reactions (changing covalent bonds) enable the gain/loss of electrons to/from their shells.
• Ionic compounds (salts) are not molecules even though chemical equations suggest they are.
• Electrolysis does not convert metallic ions to atoms - it supplies loosely associated electrons to form a metal substance.

Exceptions to these general rules can be ignored at GCSE.

Of the four material types, metals and noble gases can consist of single elements. However, salts and molecules are compounds of more than one element. Together, salts and molecules form practically everything around us in nature.

Salts make up most of the earth's crust and many of the hard, brittle things around us:

• rock
• glass
• china

Molecules make up the atmosphere and the water that we need for life. Most of our bodies are made of molecules as is most of what we eat and drink.

Covalent Bonding

• Covalent bonds bind atoms together in molecules or molecular ions.
• Non-metallic atoms can share electrons between their outer shells.
• This is a very strong 'chemical' bond, binding the atoms together to form either molecular ions or molecules.
• Making or breaking covalent bonds requires a 'chemical' reaction.
• (van der Waals) attraction between molecules is much weaker than the bonding within the molecules.
• Molecular substances generally have low melting points.
• Molecular ions take part in ionic bonding.
• Small molecules often exist as gases. Larger molecules tend to have higher melting and boiling points (except water – see Hydrogen Bonding).
• Covalent crystals are rare but are very hard due to the strong covalent bonds forming the crystal lattice (e.g. diamonds).

Ionic Bonding

• Ionic bonds bind metallic ions and negatively charges ions together in salt crystals.
• Ionic bonding is caused by the electrostatic attraction between metallic ions and negatively charged ions.
• Metallic ions together with negatively charged ions form ionic compounds called salts.
• Ionic bonds are strong 'physical' bonds and don't form molecules.
• Ionic bonds can be broken without chemical reactions (by melting or by dissolving).
• Salts do not exist as molecules even though chemical equations suggest they do.
• Each metallic ion is associated with several adjacent, negatively charged ions.
• Each negatively charged ion is associated with several adjacent, (positive) metallic ions.
• Solid salts form salt crystals.
• The relative positions of the metallic ions and negative ions are fixed (bonded) in a regular 'salt crystal' lattice.
• The arrangement of ions in one part of a crystal is identical to the arrangement in another part.
• Salt crystals do not conduct electricity as the locations of the ions are fixed.
• Salt crystals have high melting points and are hard and brittle.
• When salt crystals melt, the ionic bonds are broken and they form 'electrolytes'.
• Electrolytes have different chemical properties compared to salt crystals.
• Ions don't have fixed locations in electrolytes and can move with electric current.
• Electrolytes conduct electricity, reacting chemically at the anode and cathode (see Electrolysis).
• Many salts can also dissolve in water to form electrolytes.

Metallic Bonding

• Metallic bonds bind metallic ions together in metallic crystals (solid metal).
• Metallic bonding is the electrostatic attraction between the metallic ions and loosely associated electrons.
• Metallic bonds are strong 'physical' bonds.
• Metallic bonds can be broken without chemical reactions (by melting).
• Metals generally have high melting points and their solids are strong and ductile (not brittle like salts are).
• Metal, as a substance, is a combination of metallic ions and loosely associated electrons.
• Each metallic ion is loosely associated with several electrons at any one time.
• Each electron is loosely associated with several metallic ions at any one time.
• The positions of metallic ions are fixed (bonded) in a regular 'metallic crystal' lattice.
• The positions of loosely associated electrons are not fixed.
• When metallic crystals melt, the metallic bonds are broken and the metallic ions can move about.
• As atomic size increases, metallic bonding gets weaker - lowering melting point (compare this with molecular substances, where melting point increases with molecular size).
• Unlike salts, solid and liquid metals have similar chemical properties.
• Metals don't dissolve in water.
• Some metals react chemically with water (alkali metals).
• Other metals don't react with pure water (H₂O) but will react with impurities in water (iron rusts).
• Loosely associated electrons give metals several properties:
  1. They conduct electricity - electrons flow (move easily) between ions.
  2. They conduct heat - high-energy electrons diffuse rapidly through the metal.
  3. They reflect light - metallic lustre.
  4. They are ductile.

A normal piece of metal consists of many crystals. Some metal components are carefully (expensively) grown as a single crystal to make them stronger.

Note

When two electrolytes are mixed, all the ions mix with each other, unless a precipitate is formed.

Example: If a solution of sodium chloride is mixed with a solution of potassium bromide, the resulting solution is the same as if a solution of sodium bromide is mixed with a solution of potassium chloride.

Note

The following common equation for the formation of table salt suggests the formation of 2 salt molecules (which don't exist):

2 Na + Cl₂ => 2 NaCl

However, it might better be written as:

2 (Na⁺ + e^-)_(metal) + Cl_2(molecule) => 2 Na⁺_(ion) + 2 Cl^-_(ion) => 2 (Na⁺ + Cl^-)_(salt)

The metal provides 2 of its loosely associated electrons to the covalent reaction, breaking the covalent bond to convert a chlorine molecule into 2 chloride ions. The chloride ions then combine with sodium ions, also from the metal, to form a salt. At no point does a neutral sodium atom change to a sodium ion by giving up a single electron from its outer shell. This is because neutral sodium atoms don't exist. Also, at no point is a molecule of sodium chloride formed. This is because salt molecules don't exist.

Obviously, the first form of the equation is much easier to write, which is why it is used. Unfortunately, it does not explain to the poor student what is happening. Remember, chemistry developed as a useful science for centuries without fully understanding how the reactions occur; often by trial-and-error. In practice, most industrial chemists only need to know how to make a reaction happen or how to stop unwanted reactions. However, understanding how reactions occur gives modern chemists a framework to build this knowledge upon.

Hydrogen

• Hydrogen is very special and very important in chemistry.
• Hydrogen can be either just a single proton or have 2 electrons.
• A hydrogen ion can carry either a single positive charge (as a proton) or a single negative charge.
• Hydrogen is not an alkali metal because a positively charge hydrogen ion is just a proton.
• Hydrogen is not a halogen, even though a negatively charged hydrogen ion carries a single negative charge.
• Hydrogen can break away from water molecules as protons.
• Hydrogen causes acids and alkalis to be reactive.
• Hydrogen can exist as neutral molecules, H₂.
• H₂ is a flammable gas.
• Hydrogen gas can be released at the cathode in electrolysis.

Hydrogen Bonding and Water

While this might not be part of Gateway Science, it is essential to understand much else!

Although water (H₂0) is a molecular substance, it exhibits some weak ionic characteristics; called hydrogen bonding:

• There is a slight positive charge at one end of each molecule and a slight negative charge at the other end. This is caused by a difference in the distribution of the positive protons and the negative electrons of the hydrogen atoms.
• The angle formed by hydrogen-oxygen-hydrogen nuclei varies but averages about 105°. This shape puts the hydrogen nuclei fairly close together.
• Electrons are concentrated mostly around the oxygen nucleus leaving the hydrogen atoms with a weak positive charge (about one third of each proton's positive charge) at one end of the molecule.
• The other end of the molecule, away from the hydrogen nuclei, carries a weak negative charge (about two-thirds of an electron's charge).
• The attraction between a positive charged end of one water molecule to a negatively charged end of another is called ‘hydrogen bonding’.
• Other molecules also exhibit some hydrogen bonding (e.g. alcohols).
• Different liquids with hydrogen bonding will mix (e.g. water and alcohol).
• If one liquid has strong hydrogen bonding and another liquid does not have hydrogen bonding, the liquids will not mix (e.g. water and oil).
• Water molecules will form a crystalline structure below 0° Celsius - ice. This is much hotter than other molecular substances with even higher molecular weights, e.g. CO₂.
• Without hydrogen bonding, water would be a gas at room temperatures and life would not exist.
• Due to the fixed alignment of molecules, the crystalline form of H₂O is less dense than liquid H₂O - so ice floats on water. (Without hydrogen bonding, solids forming on a large surface normally sink - e.g. metal or wax.)
• Water can dissolve many ionic salts to form electrolytes (e.g. table salt, NaCl).
• Water will form crystals with some ionic salts (e.g. blue coloured copper sulphate, which is white without water).
• Water forms a meniscus. Which way the meniscus goes (up or down) depends on the material of the container.
• Some materials are not attracted to water (hydrophobic) e.g. wax - the meniscus goes down.
• Some materials are strongly attracted to water (hydrophilic) e.g. glass - and the meniscus goes up.
• Pure water does not conduct electricity.
• Electrolytes conduct electricity (electrolysis).
• Most water is impure and will conduct some electricity.

Acids and Alkalis

• The bond between hydrogen and oxygen is weak compared to other covalent bonds.
• In water molecules, electrons spend most of their time around the oxygen nucleus.
• Hydrogen nuclei frequently break away from water molecules as single, positively charged protons.
• The concentration of free protons in water is measured in pH.
• The lower the pH, the higher the concentration of free protons!
• In neutral water (pH 7), hydrogen nuclei average about 1ms with each water molecule before breaking away again.
• In acids and alkalis, the hydrogen nuclei break away more frequently.

Alkali Metals

• Alkali metal atoms have one more proton than electrons and carry a single positive charge.
• Alkali metal salts are generally soluble in water.
• Metallic crystals of alkali metals will dissociate water, releasing hydrogen and forming an alkali hydroxide solution.
• As alkali metals increase in atomic size, the attraction to the loosely associated electrons becomes weaker.
• As this attraction gets weaker, so the dissociation of water gets more vigorous.

Alkali Earth Metals

• Alkali earth metal atoms have two more protons than electrons and carry a double positive charge.
• Many alkali earth metal salts are insoluble in pure water (like 'earth' or 'soil' - e.g. calcium carbonate, CaCO₃, in limestone rock).

Halogens

• Halogen atoms have one more electron than protons.
• Halogens are non-metallic and can exist as negatively charged atomic ions.
• Halogens can form covalent bonds in molecules and molecular ions.
• Pure halogen substances exist as molecules (e.g. Cl₂).
• In pure halogens, melting point increases with atomic weight:
  • Fluorine and chlorine are gases at room temperature.
  • Bromine is liquid at room temperature.
  • Iodine and above are solid at room temperature.
• As halogen ions increase in size with atomic number, so their ionic bonds get weaker and they are less reactive.
• Halogens react with water to form acids.
• Fluorine is the smallest halide and its acid is the strongest (very dangerous).

Electrolysis

When an electric current is passed through an electrolyte, reactions occur at the anode and the cathode.

• Metal can be deposited at the negative cathode (e.g. copper, silver and gold plating).
• If the electrolyte is dissolved in water, the water can dissociate at the cathode releasing hydrogen instead of depositing metal – depending on the position of the metal in the reactivity table.
• The cathode provides the metallic ions with loosely associated electrons and they combine with metallic bonding to form metal. The equation is usually expressed to suggest that metallic ions become metal atoms – this is wrong.
• If the anode is metal, its loosely associated electrons can be removed by electric current and the metallic ions move into the electrolyte – metal is lost from the anode.
• If the anode is carbon (graphite), electrons are removed from negatively charged ions in the electrolyte by covalent reactions. There is more electron sharing by covalent bonding in the resulting molecules. A gas (e.g. oxygen or chlorine but not hydrogen) is often released.
• Ions travel through the electrolyte in a combination of electrostatic force and diffusion.
• The reaction at the cathode is a form of 'reduction' as electrons are supplied.
• The reaction at the anode is a form of 'oxidation' as electrons are removed.

Electrolysis of Sodium Chloride Solution (Sea Water or Brine)

• Carbon electrodes are used at both the anode and the cathode.
• Sodium is above hydrogen in the reactivity table and so water will dissociate at the cathode.
  • 2 H₂O + 2 e^- => 2 OH^- + H₂
• When water dissociates, hydrogen gas is released and hydroxide ions (OH^-) are repelled into the electrolyte.
• Initially, chloride ions (Cl^-) combine to release chlorine gas at the anode.
  • 2 Cl^- => Cl₂ + 2 e^-
• Chloride ions travel through the electrolyte to the anode both by electrostatic attraction and down the concentration gradient.
• The positively charged sodium ions are attracted towards the negatively charged cathode, where they concentrate.
• As the sodium ions concentrate, they also diffuse away from the cathode down the concentration gradient.
• Eventually, the electrostatic attraction towards the cathode is balanced by the concentration gradient away from the cathode.
• Sodium ions continue to move by diffusion through the electrolyte.
• As electrolysis continues, the original neutral sodium chloride solution gradually turns into alkali sodium hydroxide solution.
• When all the chloride ions have been converted to chlorine, a solution of sodium hydroxide remains (see below).

Electrolysis of Sodium Hydroxide Solution

• Oxygen gas is released at the anode.
  • 4 OH^- => 2 H₂O + O₂ + 4 e^-
• Hydrogen gas and hydroxide ions are produced at the cathode (see above).
• Hydroxide ions travel from the cathode, through the electrolyte to the anode.
• Electrolysis of sodium hydroxide solution will dissociate the water solvent, leaving concentrated sodium hydroxide.
• As the concentration increases, sodium hydroxide salt crystals will begin to form.
• Salt crystals do not conduct electricity and will also block the passage of ions, slowing down the electrolysis.
• Eventually, the passage of ions will be blocked completely by the salt crystals (cutting off the electrical current) and electrolysis will stop.

Aluminium Smelting

• Aluminium is above hydrogen in the reactivity table.
• Aluminium would not be deposited if the electrolyte is dissolved in water (rather, water would dissociate).
• Aluminium ore (bauxite) is insoluble in water.
• Bauxite is melted (together with cryolite, which lowers the melting point) to form an electrolyte without water.
• Aluminium ions are supplied with loosely associated electrons from the cathode to form liquid (molten) metal:
  • Al³⁺ + 3 e^- => Al_(metal)
• Aluminium ions are 'reduced' to aluminium metal at the cathode because electrons are supplied.
• Oxide ions combine to form oxygen molecules at the anode:
  • 2 O^2- => O₂ + 4 e^-
• Oxide ions are 'oxidised' to oxygen molecules at the anode because electrons are removed.
• Some of the hot oxygen burns away (oxidises) the carbon anode:
  • C + O₂ => CO₂

Corrosion

• Iron (steel) immersed in sea water will rust.
• Loosely associated electrons are lost to the sea and replaced with oxide ions from the sea.
• The negatively charged oxide ions and the positively charged iron ions form salt crystals of iron oxide (rust).
• Iron oxide crystals are larger than iron crystals and flake off, exposing more iron to the sea to be corroded.
• If not protected, iron objects will corrode completely to rust.
• Iron structures (e.g. ships) can be protected from corrosion by a zinc 'sacrificial anode'.
• Sacrificial anodes dissolve in the electrolyte (sea water) and supply electrons to the iron structure.
• The iron structure becomes a negatively charged cathode.
• Cathodes do not corrode in electrolytes.

Although aluminium reacts more strongly than iron with oxide ions, aluminium oxide does not flake off but forms a protective layer over the aluminium.

If a metal object is immersed in an electrolyte of a metal lower in the reactivity table, metal ions from the object will move into the electrolyte to be replaced by metal ions from the electrolyte - e.g. an iron nail gets coated in copper when immersed in copper sulphate solution.