0620 · 2.5Simple molecules and covalent bonds

Covalent bonding and electron diagrams

0:00Water can turn into vapour without becoming hydrogen and oxygen. Its molecules move apart, but the atoms within each molecule stay joined. That distinction explains much of the behaviour of simple molecular substances. In this video, we will connect electron sharing, molecular diagrams, melting, boiling and electrical conductivity. We need one starting idea about atoms. Electrons occupy shells around the nucleus, and the outer shell is especially important in bonding.

0:27The first shell is full with two electrons. For the atoms we will use, a full outer shell beyond the first contains eight electrons. Noble gases have full outer electron shells, which are particularly stable arrangements. Hydrogen needs one more electron to achieve the electron arrangement of helium. Carbon, nitrogen, oxygen and chlorine can also achieve noble gas electronic configurations through bonding. In these molecules, they do this by sharing electrons rather than transferring them.

0:56A covalent bond is formed when a pair of electrons is shared between two atoms, leading to noble gas electronic configurations. Each hydrogen atom starts with one electron. When two hydrogen atoms share their electrons, both have access to the same pair. Each therefore has two electrons in its first shell. Sharing does not create extra electrons or permanently give the pair to just one atom. The shared electrons count towards the outer shell of both bonded atoms.

1:22They are attracted to both positively charged nuclei, helping hold the atoms together. One shared pair forms one single covalent bond. A molecule is a group of atoms held together by covalent bonds. A hydrogen molecule contains two hydrogen atoms joined by one bond. It is still an element because it contains only one type of atom. A water molecule contains different elements joined together, so water is a compound.

1:47Dot-and-cross diagrams show the electrons involved in bonding. Dots represent electrons from one atom, and crosses represent electrons from another. They are not different kinds of electron. For our diagrams, we show outer shells only, with shared pairs between atoms and unshared electrons around their original atoms. Chlorine has seven outer-shell electrons, so each atom needs access to one more. Two chlorine atoms contribute one electron each to a shared pair.

2:15Each chlorine now counts six unshared and two shared electrons, making eight. The molecule therefore contains one single covalent bond. A pair of outer-shell electrons that is not shared in a bond is called a lone pair. Each chlorine atom has three lone pairs. These electrons must appear in a complete outer-shell dot-and-cross diagram. Showing only the bonding pair would leave out most of the electrons we are supposed to represent.

Single bonds and lone pairs

2:40Hydrogen chloride follows the same sharing principle, even though its atoms are different. Hydrogen and chlorine each give one electron to their bond. Hydrogen counts two electrons, while chlorine counts eight. A full shell does not mean eight for every atom, because hydrogen only needs two. Oxygen has six outer-shell electrons, so it needs access to two more. In water, it shares one electron with each of two hydrogen atoms.

3:06This produces two single covalent bonds. Oxygen also keeps four unshared electrons, arranged as two lone pairs, giving it eight outer-shell electrons altogether. Check each hydrogen separately rather than counting every electron in the molecule around every atom. Each hydrogen shares only the pair between itself and oxygen. That gives each hydrogen two electrons. Oxygen counts both bonding pairs and both lone pairs because all four pairs belong to its outer shell.

3:35Carbon has four outer-shell electrons and needs access to four more. In methane, it forms a single bond with each of four hydrogen atoms. With these four shared pairs, the carbon atom now has eight outer-shell electrons. Each hydrogen has two, and there are no lone pairs on carbon in this molecule. Nitrogen begins with five outer-shell electrons, so it needs access to three more. In ammonia, it forms three single bonds with three hydrogen atoms.

4:02One lone pair remains on nitrogen. Its three bonding pairs and one lone pair give eight electrons, while each hydrogen again counts two. These examples reveal a useful pattern for the molecules in this topic. Hydrogen forms one bond, chlorine forms one, oxygen forms two, nitrogen forms three, and carbon forms four. These numbers follow from how many electrons each atom needs to complete its outer shell. They are a checking tool, not a substitute for showing the electrons.

Double and triple bonds

4:30This next part is for the Extended paper. Atoms can share more than one pair with the same neighbour. Two shared pairs form a double covalent bond, and three shared pairs form a triple covalent bond. Each shared pair still contains two electrons and counts towards both bonded atoms. In oxygen, each atom begins with six outer-shell electrons. Sharing just one pair would give each oxygen access to only seven electrons.

4:56Instead, each contributes two electrons, producing two shared pairs. The double bond and two lone pairs give each oxygen a complete outer shell of eight electrons. In nitrogen, each atom needs access to three additional electrons. Each contributes three electrons to the shared region, creating a triple bond. One lone pair remains on each nitrogen atom. Six shared and two unshared electrons give each atom eight, and no electrons are created or lost.

5:24Carbon dioxide contains one carbon atom and two oxygen atoms. Carbon forms a double bond with each oxygen, giving four shared pairs around carbon. Each oxygen has two bonding pairs and two lone pairs. All three atoms therefore achieve eight outer-shell electrons, and carbon has no lone pairs. Two single bonds would not give the correct diagram for carbon dioxide. They would leave carbon and both oxygen atoms short of a complete outer shell.

5:51Always check the electron count around every atom, not just whether the atoms are connected. A plausible-looking arrangement is not enough if the electrons do not add up. Ethene contains two carbon atoms and four hydrogen atoms. Each carbon forms single bonds with two hydrogen atoms, leaving it needing two more shared pairs. The carbon atoms therefore form a double bond with each other. Each carbon has four bonding pairs altogether, and every hydrogen has one.

6:18A double bond involves four electrons, not two. In a displayed formula, one line represents one shared pair, so a double bond uses two lines. Dot-and-cross diagrams show those electrons explicitly instead. Drawing an electron as a dot or a cross does not change what the bond is.

Worked example: methanol diagram

6:36Let us work through methanol as a complete diagram-building example. Its molecule contains one carbon atom, one oxygen atom and four hydrogen atoms. Three hydrogens attach to carbon, while the fourth attaches to oxygen. Start with this arrangement rather than placing all four hydrogens around carbon. First, count the available outer-shell electrons. Carbon supplies four, oxygen supplies six, and the four hydrogen atoms supply one each. That makes fourteen electrons altogether.

7:04There are five single bonds in the arrangement, so ten electrons will be used in shared pairs. Place one shared pair in each bond, using one electron from each joined atom. Carbon contributes an electron to each of its four bonds. Oxygen contributes one to the carbon bond and one to the hydrogen bond. Its remaining four electrons form two lone pairs, accounting for all fourteen electrons. Finally, check the outer shell of every atom.

7:30Carbon has four shared pairs, giving eight electrons. Two shared pairs and two lone pairs also make eight around oxygen. Each hydrogen has one shared pair and therefore two electrons, so every atom has the required noble gas electronic configuration. The same checking method helps with an unfamiliar but similar molecule. Establish which atoms are connected, add the shared pairs, and include any lone pairs. Then count around each atom and check the total number of electrons.

7:58Shared electrons count around both neighbours, but only once when counting the whole molecule.

Properties of simple molecules

8:03The Core property statements apply to simple molecular compounds such as water and methane. They generally have low melting points and boiling points, and they are poor electrical conductors. Their structure consists of separate molecules rather than one continuous bonded structure. Their exact values vary, so at room temperature they may be gases, liquids such as water, or solids. This next part is for the Extended paper. The explanation depends on distinguishing bonds within molecules from forces between molecules.

8:32Covalent bonds hold the atoms together inside each molecule. Intermolecular forces act between separate molecules, and these forces are weak compared with the covalent bonds. The word intermolecular means between molecules. In our diagram, solid lines show the covalent bonds within each water molecule. Dashed connectors represent attractions between neighbouring molecules. Keeping these two features separate prevents the most important mistake in explaining melting and boiling. During melting, molecules become able to move past one another.

9:03During boiling, they separate much further as they enter the gas state. These changes overcome intermolecular attractions, rather than breaking the covalent bonds inside the molecules. Relatively little energy is needed because those intermolecular forces are weak. This explains the generally low melting points and boiling points of simple molecular compounds. The molecules themselves remain intact during these changes of state. Water vapour still contains water molecules, not separate hydrogen and oxygen atoms.

9:31Boiling is therefore not evidence that the covalent bonds in water are weak. Electrical conductivity requires charged particles that can move through the substance. Simple molecular substances have no mobile ions or free electrons to carry charge. Their electrons are involved in bonds or remain associated with particular atoms. Moving neutral molecules through a liquid does not provide an electrical current. Having electrons is therefore not the same as having free electrons.

9:57A lone pair is not part of a bond, but its electrons cannot travel through the substance. Simple molecular substances remain poor conductors when melted because melting does not release charge carriers. Solutions that form ions need separate consideration, which belongs to later topics.

Exam questions and recap

10:14Examiners commonly see diagrams with missing lone pairs or incorrect numbers of shared electrons. Avoid these errors by checking both the outer-shell count and the total electron count. Another mistake is adding square brackets and ion charges to these neutral molecules. Electron sharing here does not represent electron transfer or the formation of ions. An Extended question could ask you to draw the dot-and-cross diagram of carbon dioxide and explain its properties.

10:39The command word draw requires the actual electron diagram, not just a description. Explain requires reasons linked to structure and bonding. Start by checking the outer-shell electrons before writing about the properties. Carbon has four outer-shell electrons, and each oxygen has six. Draw two shared pairs between carbon and each oxygen. Add two lone pairs to each oxygen and none to carbon. Each atom then has eight outer-shell electrons, with sixteen electrons shown altogether.

11:08A clear property explanation gives one creditworthy point in each sentence. Carbon dioxide has a simple molecular structure with weak intermolecular forces. Little energy is needed to overcome these forces, so it has a low boiling point. The covalent bonds within its molecules are not broken. It has no mobile ions or free electrons to carry charge, so it does not conduct electricity. For a Core question asking you to describe the properties, give low melting and boiling points and poor electrical conductivity.

11:37For an Extended question asking you to explain them, include the reasons. Avoid describing the covalent bonds as weak. The low boiling point comes from weak attractions between molecules, not weak bonds within them. The central idea is that sharing pairs of electrons gives the atoms in these molecules noble gas electronic configurations. One pair forms a single bond, two form a double bond, and three form a triple bond.

12:00Lone pairs must also appear in complete outer-shell diagrams. Always check hydrogen for two electrons and the other atoms here for eight. Simple molecular properties depend on what happens between molecules and whether charge can move. Weak intermolecular forces explain their generally low melting points and boiling points. The absence of mobile ions or free electrons explains their poor conductivity. Keeping those ideas separate from the strong covalent bonds lets you explain the behaviour rather than memorise it.