2.2.1 Metals, non-metals and the Periodic Table

Metals and non-metals have contrasting physical properties

1. Metals are usually shiny, strong and dense, with high melting points.
2. Metals are good conductors of electricity and heat.
   1. Copper is used for electrical wiring because it conducts well and is ductile.
3. Aluminium can be rolled into thin foil because it is malleable.
4. Non-metals are usually dull, brittle when solid and poor conductors, with low melting and boiling points.
   1. Graphite is an exception, because it is a non-metal that conducts electricity.

Metals lose electrons while non-metals gain or share them

1. When they react, metal atoms lose electrons to form positive ions, such as ${Na}^{+}$ and ${Mg}^{2+}$.
2. Non-metal atoms gain electrons to form negative ions, such as ${Cl}^{-}$ and $O^{2-}$, when they react with metals.
3. Non-metal atoms can also share electrons with other non-metals to form molecules, such as $H_{2}O$.
4. Metal oxides are basic, and those that dissolve in water form alkaline solutions.
5. Non-metal oxides are often acidic, so sulfur dioxide and carbon dioxide form acidic solutions in water.

Electron arrangement gives an element's group and period

1. Elements are placed in order of atomic number, so each element has one more proton and one more electron than the one before.
2. Electrons fill the shells in order, with up to $2$ in the first shell and up to $8$ in each of the next two.
3. For Groups 1 to 7, the group number equals the number of electrons in the outer shell.
4. The period number equals the number of occupied shells.
   1. Sodium, atomic number $11$, has the arrangement $\mathrm{2.8.1}$, so it is in Group 1 and Period 3.

Few outer electrons place metals left and centre

1. Atoms of metals usually have 1, 2 or 3 outer electrons, which they lose relatively easily.
   1. This places most metals on the left-hand side and in the centre of the table.
   2. Aluminium, with the arrangement $\mathrm{2.8.3}$, is a metal in Group 3.
2. Atoms of non-metals usually have 4 to 7 outer electrons, or a full outer shell in Group 0.
   1. This places the non-metals on the right-hand side of the table.
3. A staircase line running from boron down to astatine roughly divides the metals from the non-metals.

2.2.2 Types of chemical bonding and dot and cross diagrams

Chemical bonds form by transferring or sharing electrons

1. Atoms bond in ways that give them a full outer shell, like the outer shell of a noble gas.
2. A metal bonding with a non-metal transfers electrons, forming ionic bonds.
3. Non-metals bonding with each other share electrons, forming covalent bonds.
4. Metal atoms bonding together release their outer electrons into a shared pool, forming metallic bonds.
5. In every case the bond is an electrostatic attraction between opposite charges.

Ionic bonds hold oppositely charged ions together

1. When sodium reacts with chlorine, each sodium atom transfers its one outer electron to a chlorine atom.
   1. Sodium becomes ${Na}^{+}$, with the arrangement $[2.8{]}^{+}$, and chlorine becomes ${Cl}^{-}$, with the arrangement $[\mathrm{2.8.8}{]}^{-}$.
2. The ions attract each other in all directions, so they build a giant structure.
3. A grain of salt contains millions of ions, so the formula NaCl gives the ratio of ions rather than a molecule.
4. In magnesium oxide, magnesium transfers two electrons to oxygen, forming ${Mg}^{2+}$ and $O^{2-}$ ions.

Covalent bonds form molecules, giant structures and polymers

1. A covalent bond forms when two non-metal atoms each put one electron into a shared pair.
   1. The shared pair is attracted to both nuclei, which holds the two atoms together.
2. In simple molecules such as $H_2$, ${Cl}_2$, $H_{2}O$ and ${CH}_4$, a small, fixed number of atoms are joined by covalent bonds.
   1. Separate molecules are held to one another only by weak intermolecular forces, which are much weaker than the covalent bonds inside them.
3. In giant covalent structures such as diamond and silicon dioxide, huge numbers of atoms are joined by covalent bonds in a continuous network.
4. Polymers such as poly(ethene) are very long molecules in which thousands of atoms are joined in chains by covalent bonds.
   1. Neighbouring polymer chains are held together by intermolecular forces, which are stronger than between small molecules because the chains are so long.

Metallic bonding involves a sea of delocalised electrons

1. In a metal, the atoms are packed closely in a regular arrangement.
2. Each atom gives up its outer electrons, leaving positive metal ions.
3. The released electrons spread through the whole metal as a sea of electrons.
4. This strong electrostatic attraction between positive ions and delocalised electrons is called metallic bonding.
5. Ionic, giant covalent and metallic substances are giant structures, while simple molecules and polymers are made of separate molecules.

Dot and cross diagrams show outer electrons

1. Electrons from one atom are drawn as dots and electrons from the other atom as crosses, so it is clear where each came from.
2. For a binary ionic compound, draw the outer shell of each ion inside square brackets, with the charge written outside at the top right.
   1. In $MgO$, magnesium's two outer electrons appear as crosses in the outer shell of oxygen, giving $[Mg{]}^{2+}$ and $[O{]}^{2-}$ with eight outer electrons.
3. For a simple covalent molecule, draw the outer shells overlapping, with each shared pair of one dot and one cross inside the overlap.
   1. In $H_{2}O$, oxygen shares one pair with each hydrogen atom, leaving oxygen with two pairs of outer electrons that are not shared.
4. Finish by counting the outer electrons around each atom to check that every outer shell is full.

2.2.3 Limitations of bonding models

Every representation of a substance is a model

1. Chemists represent substances with dot and cross diagrams, displayed formulae, ball and stick models and space-filling models.
2. Each representation is chosen to show one feature clearly, such as which electrons form a bond or how the atoms are arranged.
3. No single model shows everything, because real particles are not balls, sticks or lines on paper.
4. Knowing what a model leaves out helps you choose the right one and avoid false conclusions.

Dot and cross diagrams show electrons, not shape

1. A dot and cross diagram shows which electrons are transferred or shared and which atom they came from.
2. It is two-dimensional, so it cannot show the three-dimensional shape of a molecule.
   1. Methane is drawn flat, but its four hydrogen atoms really point away from the carbon atom in three dimensions.
3. It does not show the relative sizes of atoms or ions.
4. Electrons from different atoms are drawn differently, although all electrons are identical.
5. For an ionic compound it shows one pair of ions rather than the giant lattice.

Displayed formulae and ball and stick models show bonds

1. A displayed formula makes it easy to see which atoms are bonded to which.
2. It is flat, so the bonds appear at right angles even though real molecules are three-dimensional.
3. It does not show the electrons in each bond or the outer electrons that are not shared.
4. A ball and stick model shows the 3D shape of a molecule much better.
5. Its sticks suggest large gaps between atoms and make bonds look like solid rods, whereas bonded atoms are really very close together.
6. The balls are often the wrong relative sizes, and the electrons are not shown at all.

Space-filling and lattice diagrams hide the bonding

1. A space-filling model shows the atoms touching at roughly their correct relative sizes.
   1. It shows the overall shape of a molecule well, but the bonds and bond angles are hard to see.
2. A lattice diagram of sodium chloride shows how the ions repeat in three directions.
   1. It shows only a tiny part of the lattice, which really continues in every direction for millions of ions.
3. A diagram of a metal shows positive ions in a sea of electrons, but it cannot show the electrons moving.

2.2.4 Reactivity, atomic number and the modern Periodic Table

Same outer electrons make group elements react similarly

1. Chemical reactions involve the outer electrons of atoms, which are lost, gained or shared.
2. Atoms in the same group have the same number of outer electrons, so they react in similar ways.
   1. Lithium, sodium and potassium each have one outer electron, and all react with water to give hydrogen and an alkaline solution.
3. The atomic number fixes the number of electrons, so it also fixes the electron arrangement and the way an element reacts.
4. Group 0 elements have full outer shells, so they have no tendency to lose, gain or share electrons and are very unreactive.

Elements near period ends lose or gain electrons

1. Atoms with one or two outer electrons, in Groups 1 and 2, react by losing them to form positive ions.
2. Atoms with six or seven outer electrons, in Groups 6 and 7, react by gaining electrons to form negative ions.
3. This is why elements from opposite sides of the table, such as sodium and chlorine, react vigorously with each other.
4. Elements with four outer electrons, such as carbon, tend to share electrons instead.

Mendeleev ordered by atomic mass but grouped by properties

1. In 1869, Dmitri Mendeleev arranged the 60 or so elements then known in order of increasing atomic mass.
2. He placed elements with similar properties in the same columns, even when this broke the mass order.
   1. He put iodine after tellurium so that it sat with the other halogens, although iodine has the lower atomic mass.
3. He left gaps for elements that had not yet been discovered.
   1. He predicted the properties of the missing elements, and when gallium and germanium were found their properties closely matched his predictions.

Ordering by atomic number explains Mendeleev's swaps

1. In 1913, Henry Moseley showed that each element has its own atomic number, which was later explained as the number of protons in its nucleus.
2. Ordering by atomic number puts tellurium, $52$, before iodine, $53$, with no swap needed.
   1. Atomic masses can run out of order because of isotopes, since an element with many heavier isotopes has a higher average mass.
3. Atomic number also sets the electron arrangement, which explains why elements in a group behave alike.
4. The noble gases, unknown to Mendeleev in 1869, fitted in as a whole new group, Group 0.