Atoms Combine for One Reason: To Fill Their Outermost Shell
Learn covalent bonding as shared electron pairs, ionic bonding as transferred electrons, and how to write the formula of any compound by criss-crossing valencies, including polyatomic ions.
Why do atoms combine at all?
To reach a full outermost shell, which is the arrangement they are most stable in.
The noble gases — helium with electrons, neon with and argon with — already have full outer shells. They need nothing, so they form almost no compounds. Every other atom has an incomplete outer shell, and combining is how it completes one.
There are exactly two ways to do it. An atom can share electrons with another atom so that both count the shared pair, which is a covalent bond. Or it can transfer electrons outright, so that one atom loses and the other gains, which is an ionic bond.
That is the mechanism Dalton's theory could not supply, as the previous part of this chapter noted. It states that atoms combine; this page explains how, and it covers the second part of the CBSE Class 9 Science chapter on the atomic foundations of matter.
The noble gases — helium with electrons, neon with and argon with — already have full outer shells. They need nothing, so they form almost no compounds. Every other atom has an incomplete outer shell, and combining is how it completes one.
There are exactly two ways to do it. An atom can share electrons with another atom so that both count the shared pair, which is a covalent bond. Or it can transfer electrons outright, so that one atom loses and the other gains, which is an ionic bond.
That is the mechanism Dalton's theory could not supply, as the previous part of this chapter noted. It states that atoms combine; this page explains how, and it covers the second part of the CBSE Class 9 Science chapter on the atomic foundations of matter.
How does a covalent bond work, and how do you draw its dot structure?
Two atoms share a pair of electrons, and both count that pair towards their own outer shell.
One shared pair is one single bond. Two shared pairs make a double bond and three make a triple bond. Covalent bonds form typically between two non-metals.
In an electron dot structure, the valence electrons of each atom are drawn as dots around its symbol. A shared pair sits between the two symbols; the pairs not involved in bonding are lone pairs.
How to check a structure in one step. Add up the valence electrons of all the atoms, divide by two to get the number of pairs, and confirm that your drawing has exactly that many pairs — bonding and lone together.
**Worked example 1 — hydrogen, .** Each hydrogen has valence electron, so the total is , giving one pair. The two atoms share it, and each now counts electrons — a complete first shell, called a duplet. One single bond, no lone pairs.
**Worked example 2 — chlorine, .** Each chlorine has valence electrons, so the total is , giving seven pairs. One pair is shared and each atom keeps three lone pairs. Each chlorine now counts electrons. One single bond.
**Worked example 3 — water, .** Valence electrons: , giving four pairs. Oxygen shares one pair with each hydrogen — two bonding pairs — and keeps two lone pairs. Oxygen counts and each hydrogen counts .
**Worked example 4 — ammonia, .** Valence electrons: , giving four pairs. Nitrogen forms three bonds and keeps one lone pair.
**Worked example 5 — methane, .** Valence electrons: , giving four pairs, all of them bonding. Carbon forms four single bonds and has no lone pairs.
**Worked example 6 — oxygen, .** Each oxygen has and needs , so they share two pairs — a double bond — and each keeps two lone pairs.
**Worked example 7 — nitrogen, .** Each nitrogen has and needs , so they share three pairs — a triple bond — and each keeps one lone pair.
**Worked example 8 — carbon dioxide, .** Valence electrons: , giving eight pairs. Carbon forms a double bond to each oxygen, using four bonding pairs, and each oxygen keeps two lone pairs — four lone pairs in all. Eight pairs accounted for.
A double bond is two shared pairs, not one bond drawn twice. That is the distinction the counting check enforces: has one double bond and therefore four electrons between the atoms, while has one single bond and two.
Hydrogen is the exception to the octet. Its first shell holds only , so hydrogen completes a duplet and never reaches eight. A dot structure showing hydrogen with eight electrons is wrong on its face.
One shared pair is one single bond. Two shared pairs make a double bond and three make a triple bond. Covalent bonds form typically between two non-metals.
In an electron dot structure, the valence electrons of each atom are drawn as dots around its symbol. A shared pair sits between the two symbols; the pairs not involved in bonding are lone pairs.
How to check a structure in one step. Add up the valence electrons of all the atoms, divide by two to get the number of pairs, and confirm that your drawing has exactly that many pairs — bonding and lone together.
**Worked example 1 — hydrogen, .** Each hydrogen has valence electron, so the total is , giving one pair. The two atoms share it, and each now counts electrons — a complete first shell, called a duplet. One single bond, no lone pairs.
**Worked example 2 — chlorine, .** Each chlorine has valence electrons, so the total is , giving seven pairs. One pair is shared and each atom keeps three lone pairs. Each chlorine now counts electrons. One single bond.
**Worked example 3 — water, .** Valence electrons: , giving four pairs. Oxygen shares one pair with each hydrogen — two bonding pairs — and keeps two lone pairs. Oxygen counts and each hydrogen counts .
**Worked example 4 — ammonia, .** Valence electrons: , giving four pairs. Nitrogen forms three bonds and keeps one lone pair.
**Worked example 5 — methane, .** Valence electrons: , giving four pairs, all of them bonding. Carbon forms four single bonds and has no lone pairs.
**Worked example 6 — oxygen, .** Each oxygen has and needs , so they share two pairs — a double bond — and each keeps two lone pairs.
**Worked example 7 — nitrogen, .** Each nitrogen has and needs , so they share three pairs — a triple bond — and each keeps one lone pair.
**Worked example 8 — carbon dioxide, .** Valence electrons: , giving eight pairs. Carbon forms a double bond to each oxygen, using four bonding pairs, and each oxygen keeps two lone pairs — four lone pairs in all. Eight pairs accounted for.
A double bond is two shared pairs, not one bond drawn twice. That is the distinction the counting check enforces: has one double bond and therefore four electrons between the atoms, while has one single bond and two.
Hydrogen is the exception to the octet. Its first shell holds only , so hydrogen completes a duplet and never reaches eight. A dot structure showing hydrogen with eight electrons is wrong on its face.
How does an ionic bond form, and which ions does each atom make?
One atom transfers electrons to another, and the oppositely charged ions then attract each other.
An ionic bond is the electrostatic attraction between a positive and a negative ion. It forms typically between a metal and a non-metal, because metals have few outer electrons to lose and non-metals need only a few to gain.
- Metals lose electrons and become cations, which are positive
- Non-metals gain electrons and become anions, which are negative
The ions formed by common elements, with the configuration each ends up with:
- Sodium loses to give with
- Magnesium loses to give with
- Aluminium loses to give with
- Chlorine gains to give with
- Oxygen gains to give with
- Nitrogen gains to give with
Every one of those ions has the configuration of a noble gas. , , and all end with , matching neon, while ends with , matching argon. That is the whole point of the transfer.
Worked example 1 — sodium chloride. Sodium has one outer electron it can give away; chlorine needs exactly one. Sodium hands its electron over, becoming , and chlorine becomes . The two ions attract, and one of each pairs up:
Worked example 2 — magnesium oxide. Magnesium gives two electrons and oxygen needs exactly two, so again one of each:
Worked example 3 — magnesium chloride. Magnesium gives two electrons, but each chlorine can take only one. So two chlorine atoms are needed for each magnesium:
The electrons transferred must balance exactly. No electrons go missing and none appear from nowhere, so the total lost equals the total gained — which is why the numbers of ions come out as they do. Checking that the total positive charge equals the total negative charge is the same check in a different form.
An ionic compound has no molecules. The transfer produces ions, and every positive ion attracts every negative ion around it, so the solid is a giant lattice of alternating ions rather than a collection of separate particles. That is why its properties differ so sharply from a covalent compound's — the subject of the next part of this chapter.
An ionic bond is the electrostatic attraction between a positive and a negative ion. It forms typically between a metal and a non-metal, because metals have few outer electrons to lose and non-metals need only a few to gain.
- Metals lose electrons and become cations, which are positive
- Non-metals gain electrons and become anions, which are negative
The ions formed by common elements, with the configuration each ends up with:
- Sodium loses to give with
- Magnesium loses to give with
- Aluminium loses to give with
- Chlorine gains to give with
- Oxygen gains to give with
- Nitrogen gains to give with
Every one of those ions has the configuration of a noble gas. , , and all end with , matching neon, while ends with , matching argon. That is the whole point of the transfer.
Worked example 1 — sodium chloride. Sodium has one outer electron it can give away; chlorine needs exactly one. Sodium hands its electron over, becoming , and chlorine becomes . The two ions attract, and one of each pairs up:
Worked example 2 — magnesium oxide. Magnesium gives two electrons and oxygen needs exactly two, so again one of each:
Worked example 3 — magnesium chloride. Magnesium gives two electrons, but each chlorine can take only one. So two chlorine atoms are needed for each magnesium:
The electrons transferred must balance exactly. No electrons go missing and none appear from nowhere, so the total lost equals the total gained — which is why the numbers of ions come out as they do. Checking that the total positive charge equals the total negative charge is the same check in a different form.
An ionic compound has no molecules. The transfer produces ions, and every positive ion attracts every negative ion around it, so the solid is a giant lattice of alternating ions rather than a collection of separate particles. That is why its properties differ so sharply from a covalent compound's — the subject of the next part of this chapter.
How do you write the formula of a covalent compound from valencies?
Criss-cross the valencies into subscripts, then simplify.
The method, in three steps:
- Write the two symbols side by side, with the less electronegative element first — the metal, or the element further left in the periodic table
- Write each element's valency above its symbol
- Cross them over so that each valency becomes the other element's subscript, then simplify any common factor
A subscript of is never written.
Worked example 1 — water. Hydrogen has valency and oxygen . Crossing over puts under hydrogen and under oxygen:
Worked example 2 — methane. Carbon and hydrogen give , which is written
Worked example 3 — ammonia. Nitrogen and hydrogen give .
Worked example 4 — where simplification is needed. Carbon and oxygen cross over to give . Both subscripts share a factor of :
Worked example 5 — the same situation with silicon. Silicon and oxygen give , which simplifies to .
Worked example 6 — more practice. Hydrogen with sulphur gives . Phosphorus with chlorine gives . Carbon with chlorine gives .
Always look for a common factor before writing the answer down. is not a compound anyone writes, and leaving a formula unsimplified is the commonest mark lost here.
Check the formula by counting valencies. In , carbon uses its valency of against four hydrogens of valency each, so . In , . The total valency on each side must balance, and that is a second verification costing one line.
A few compounds are genuinely not the simplified form. Hydrogen peroxide is and not , because its molecule really does contain two of each atom. So the criss-cross method gives the simplest ratio, and for the compounds in this syllabus that ratio is the formula — but it is worth knowing the method has a limit.
The method, in three steps:
- Write the two symbols side by side, with the less electronegative element first — the metal, or the element further left in the periodic table
- Write each element's valency above its symbol
- Cross them over so that each valency becomes the other element's subscript, then simplify any common factor
A subscript of is never written.
Worked example 1 — water. Hydrogen has valency and oxygen . Crossing over puts under hydrogen and under oxygen:
Worked example 2 — methane. Carbon and hydrogen give , which is written
Worked example 3 — ammonia. Nitrogen and hydrogen give .
Worked example 4 — where simplification is needed. Carbon and oxygen cross over to give . Both subscripts share a factor of :
Worked example 5 — the same situation with silicon. Silicon and oxygen give , which simplifies to .
Worked example 6 — more practice. Hydrogen with sulphur gives . Phosphorus with chlorine gives . Carbon with chlorine gives .
Always look for a common factor before writing the answer down. is not a compound anyone writes, and leaving a formula unsimplified is the commonest mark lost here.
Check the formula by counting valencies. In , carbon uses its valency of against four hydrogens of valency each, so . In , . The total valency on each side must balance, and that is a second verification costing one line.
A few compounds are genuinely not the simplified form. Hydrogen peroxide is and not , because its molecule really does contain two of each atom. So the criss-cross method gives the simplest ratio, and for the compounds in this syllabus that ratio is the formula — but it is worth knowing the method has a limit.
How do you write formulae for ionic compounds with polyatomic ions?
Criss-cross the charges exactly as before, and put brackets round a polyatomic ion whenever more than one is needed.
A polyatomic ion is a group of atoms carrying a single overall charge, and it travels through reactions as one unit.
The polyatomic ions to learn, with their valencies:
- Ammonium — valency
- Hydroxide — valency
- Nitrate — valency
- Hydrogen carbonate — valency
- Carbonate — valency
- Sulphate — valency
- Sulphite — valency
- Phosphate — valency
Worked example 1 — simple cases. with gives . with gives . with gives .
Worked example 2 — sodium carbonate. has valency and has valency . Crossing over needs two sodiums and one carbonate:
No brackets, because only one carbonate group is present.
Worked example 3 — calcium hydroxide. needs two hydroxide ions:
The brackets are essential here. means two oxygens and two hydrogens. Writing would mean one oxygen and two hydrogens — a different substance entirely, and the single most penalised slip in the topic.
Worked example 4 — magnesium nitrate. with two gives .
Worked example 5 — where simplification is needed. with crosses over to , and both subscripts share a factor of :
Worked example 6 — aluminium sulphate. with gives
Check the charges balance: and . They cancel, so the formula is right.
Worked example 7 — ammonium sulphate. Two with one :
Here the positive ion needs the brackets, because two ammonium groups are required.
Worked example 8 — calcium phosphate. with gives . Check: and . Correct.
Verify every ionic formula by balancing the charges. Total positive must equal total negative, and that check catches a mis-crossed valency instantly. It is quicker than re-deriving the formula and it is the habit worth forming, because the same check is used throughout Class 11 and Class 12 chemistry.
A polyatomic ion is a group of atoms carrying a single overall charge, and it travels through reactions as one unit.
The polyatomic ions to learn, with their valencies:
- Ammonium — valency
- Hydroxide — valency
- Nitrate — valency
- Hydrogen carbonate — valency
- Carbonate — valency
- Sulphate — valency
- Sulphite — valency
- Phosphate — valency
Worked example 1 — simple cases. with gives . with gives . with gives .
Worked example 2 — sodium carbonate. has valency and has valency . Crossing over needs two sodiums and one carbonate:
No brackets, because only one carbonate group is present.
Worked example 3 — calcium hydroxide. needs two hydroxide ions:
The brackets are essential here. means two oxygens and two hydrogens. Writing would mean one oxygen and two hydrogens — a different substance entirely, and the single most penalised slip in the topic.
Worked example 4 — magnesium nitrate. with two gives .
Worked example 5 — where simplification is needed. with crosses over to , and both subscripts share a factor of :
Worked example 6 — aluminium sulphate. with gives
Check the charges balance: and . They cancel, so the formula is right.
Worked example 7 — ammonium sulphate. Two with one :
Here the positive ion needs the brackets, because two ammonium groups are required.
Worked example 8 — calcium phosphate. with gives . Check: and . Correct.
Verify every ionic formula by balancing the charges. Total positive must equal total negative, and that check catches a mis-crossed valency instantly. It is quicker than re-deriving the formula and it is the habit worth forming, because the same check is used throughout Class 11 and Class 12 chemistry.
Exam tip
Exam tip: count the pairs, and bracket the polyatomic ion
For a dot structure, add the valence electrons and divide by two to get the number of pairs, then check your drawing has exactly that many — bonding plus lone. has electrons, so eight pairs.
Hydrogen completes a duplet, not an octet. Never draw it with eight electrons.
A double bond is two shared pairs and a triple bond three — four and six electrons between the atoms.
Ions take a noble gas configuration: , , and all end ; ends .
Metals lose to form cations; non-metals gain to form anions, and the electrons lost must equal the electrons gained.
For a formula, criss-cross the valencies and then simplify: becomes , and becomes .
Bracket a polyatomic ion whenever more than one is needed: , , . is wrong.
Check the charges balance: gives against .
Learn the polyatomic ions with their valencies — they cannot be worked out, and formulas throughout the rest of the course depend on them.
And never write a subscript of .
Hydrogen completes a duplet, not an octet. Never draw it with eight electrons.
A double bond is two shared pairs and a triple bond three — four and six electrons between the atoms.
Ions take a noble gas configuration: , , and all end ; ends .
Metals lose to form cations; non-metals gain to form anions, and the electrons lost must equal the electrons gained.
For a formula, criss-cross the valencies and then simplify: becomes , and becomes .
Bracket a polyatomic ion whenever more than one is needed: , , . is wrong.
Check the charges balance: gives against .
Learn the polyatomic ions with their valencies — they cannot be worked out, and formulas throughout the rest of the course depend on them.
And never write a subscript of .
Did you know
Why table salt has no molecules
Ask what a molecule of common salt looks like and the honest answer is that there is no such thing.
When sodium transfers its electron to chlorine, the result is a ion and a ion. Neither is attached to one particular partner. A positive ion attracts every negative ion near it, and a negative ion attracts every positive one, so the ions pack themselves into a regular three-dimensional lattice in which each sodium ion is surrounded by chloride ions and each chloride ion by sodium ions.
A grain of salt is therefore one enormous structure held together throughout, not a heap of separate units. The formula states the ratio — one sodium ion to every chloride ion — and nothing more.
That is why chemists call it a formula unit rather than a molecule, and why the next part of this chapter calculates a formula unit mass for ionic compounds and a molecular mass for covalent ones. The two words describe genuinely different things.
It also explains the properties. Melting salt means overcoming the attraction between ions right through the lattice, which takes a great deal of energy — hence the high melting point. Melting a covalent substance such as ice means only separating whole molecules from one another, which takes far less, because the strong bonds inside each molecule are not broken at all.
And it explains the cubic shape of a salt crystal. The regular arrangement of ions at the atomic scale shows up as flat faces and right-angled corners at a scale you can see with a hand lens — the lattice, made visible.
When sodium transfers its electron to chlorine, the result is a ion and a ion. Neither is attached to one particular partner. A positive ion attracts every negative ion near it, and a negative ion attracts every positive one, so the ions pack themselves into a regular three-dimensional lattice in which each sodium ion is surrounded by chloride ions and each chloride ion by sodium ions.
A grain of salt is therefore one enormous structure held together throughout, not a heap of separate units. The formula states the ratio — one sodium ion to every chloride ion — and nothing more.
That is why chemists call it a formula unit rather than a molecule, and why the next part of this chapter calculates a formula unit mass for ionic compounds and a molecular mass for covalent ones. The two words describe genuinely different things.
It also explains the properties. Melting salt means overcoming the attraction between ions right through the lattice, which takes a great deal of energy — hence the high melting point. Melting a covalent substance such as ice means only separating whole molecules from one another, which takes far less, because the strong bonds inside each molecule are not broken at all.
And it explains the cubic shape of a salt crystal. The regular arrangement of ions at the atomic scale shows up as flat faces and right-angled corners at a scale you can see with a hand lens — the lattice, made visible.
Exam relevance
How does chemical bonding in Class 9 feed into JEE and NEET?
Because formula writing and electron counting are used in every chemistry question for the next three years, and this page is where both are established.
This is the foundation for the Class 11 Chemistry chapter Chemical Bonding and Molecular Structure, examined in both JEE Main and NEET. That chapter keeps the electron dot structure — where it is called a Lewis structure — and adds formal charge, resonance, the octet rule's exceptions, and VSEPR theory for predicting molecular shape. The pair-counting check used above is exactly the first step of drawing a Lewis structure there, and students who never learned to count pairs draw shapes by guesswork.
Lone pairs matter far more later than they appear to here. In VSEPR, the number of lone pairs on the central atom decides the molecule's shape — water is bent rather than linear precisely because oxygen keeps two lone pairs, which is the count made in worked example 3 above. JEE Main shape questions are unanswerable without it.
Ionic bonding is developed with lattice energy and the Born-Haber cycle in Class 11 and Class 12, and the noble gas configuration of each ion becomes the basis of periodic trends in Class 11 Classification of Elements and Periodicity.
Formula writing is used in every equation you balance from here on, and the polyatomic ions listed above reappear in Class 11 Equilibrium, in the whole of inorganic qualitative analysis, and in Class 12 Coordination Compounds.
What the questions look like. For NEET, match-the-column items pairing a molecule with its number of lone pairs or its bond type are common, and assertion-reason questions favour the statement that an ionic compound has no molecules. For JEE Main, the material appears as a step inside a harder question — a shape, a bond angle or a hybridisation that requires the Lewis structure first. Formula-writing errors show up indirectly, by making a stoichiometry answer wrong for a reason that has nothing to do with stoichiometry.
How board and competitive emphasis differ. A board paper asks you to draw the dot structure of three simple molecules and write five formulas by criss-crossing. A competitive paper assumes both and asks what follows from them — the shape, the polarity, the number of bonding pairs. So the speed and accuracy of the dot structure matter more than the ability to describe the method.
The single trap that costs the most marks. Writing for calcium hydroxide. Without brackets the subscript applies only to the hydrogen, giving a formula with one oxygen instead of two — and every mass and stoichiometry calculation built on it is then wrong. Bracket a polyatomic ion the moment you need more than one of it.
A second trap worth naming. Treating a double bond as one shared pair. It is two, so has four electrons between its atoms and has eight in its two double bonds. The pair count catches this: must account for all eight pairs, and a structure with single bonds accounts for only six.
This is the foundation for the Class 11 Chemistry chapter Chemical Bonding and Molecular Structure, examined in both JEE Main and NEET. That chapter keeps the electron dot structure — where it is called a Lewis structure — and adds formal charge, resonance, the octet rule's exceptions, and VSEPR theory for predicting molecular shape. The pair-counting check used above is exactly the first step of drawing a Lewis structure there, and students who never learned to count pairs draw shapes by guesswork.
Lone pairs matter far more later than they appear to here. In VSEPR, the number of lone pairs on the central atom decides the molecule's shape — water is bent rather than linear precisely because oxygen keeps two lone pairs, which is the count made in worked example 3 above. JEE Main shape questions are unanswerable without it.
Ionic bonding is developed with lattice energy and the Born-Haber cycle in Class 11 and Class 12, and the noble gas configuration of each ion becomes the basis of periodic trends in Class 11 Classification of Elements and Periodicity.
Formula writing is used in every equation you balance from here on, and the polyatomic ions listed above reappear in Class 11 Equilibrium, in the whole of inorganic qualitative analysis, and in Class 12 Coordination Compounds.
What the questions look like. For NEET, match-the-column items pairing a molecule with its number of lone pairs or its bond type are common, and assertion-reason questions favour the statement that an ionic compound has no molecules. For JEE Main, the material appears as a step inside a harder question — a shape, a bond angle or a hybridisation that requires the Lewis structure first. Formula-writing errors show up indirectly, by making a stoichiometry answer wrong for a reason that has nothing to do with stoichiometry.
How board and competitive emphasis differ. A board paper asks you to draw the dot structure of three simple molecules and write five formulas by criss-crossing. A competitive paper assumes both and asks what follows from them — the shape, the polarity, the number of bonding pairs. So the speed and accuracy of the dot structure matter more than the ability to describe the method.
The single trap that costs the most marks. Writing for calcium hydroxide. Without brackets the subscript applies only to the hydrogen, giving a formula with one oxygen instead of two — and every mass and stoichiometry calculation built on it is then wrong. Bracket a polyatomic ion the moment you need more than one of it.
A second trap worth naming. Treating a double bond as one shared pair. It is two, so has four electrons between its atoms and has eight in its two double bonds. The pair count catches this: must account for all eight pairs, and a structure with single bonds accounts for only six.
Key takeaways
Covalent bonds, ionic bonds and formulae: quick revision
- Atoms combine to reach a full outermost shell. Noble gases already have one, so they barely react.
- Covalent bonding is sharing electron pairs, typically between two non-metals. One shared pair is a single bond, two a double, three a triple.
- Check a dot structure by adding all valence electrons and dividing by two: has , so eight pairs.
- : one pair, each atom reaching a duplet. : seven pairs, one shared and three lone pairs each.
- : four pairs — two bonding and two lone pairs on oxygen. : four pairs — three bonding and one lone pair. : four pairs, all bonding.
- : a double bond. : a triple bond. : two double bonds plus four lone pairs.
- Hydrogen completes a duplet, never an octet.
- Ionic bonding is electrostatic attraction after transfer, typically between a metal and a non-metal.
- Metals lose to form cations: , , . Non-metals gain to form anions: , , .
- Every ion ends with a noble gas configuration — or .
- pairs one of each; pairs one of each; needs two chlorines because magnesium gives two electrons.
- Electrons lost must equal electrons gained, which is the same as the charges balancing.
- Criss-cross the valencies and simplify: , , , , , . And becomes , becomes .
- Polyatomic ions: , , , with valency ; , , with ; with .
- , , , , , , and simplifying to .
- Bracket a polyatomic ion when more than one is needed — is wrong.
- Balance the charges to verify: gives against .
- An ionic compound has no molecules — it is a giant lattice, and the formula gives only the ratio of ions.
Draw the dot structures of five molecules and check each by the pair count before looking anything up — a structure that accounts for every pair is almost certainly right.
- Covalent bonding is sharing electron pairs, typically between two non-metals. One shared pair is a single bond, two a double, three a triple.
- Check a dot structure by adding all valence electrons and dividing by two: has , so eight pairs.
- : one pair, each atom reaching a duplet. : seven pairs, one shared and three lone pairs each.
- : four pairs — two bonding and two lone pairs on oxygen. : four pairs — three bonding and one lone pair. : four pairs, all bonding.
- : a double bond. : a triple bond. : two double bonds plus four lone pairs.
- Hydrogen completes a duplet, never an octet.
- Ionic bonding is electrostatic attraction after transfer, typically between a metal and a non-metal.
- Metals lose to form cations: , , . Non-metals gain to form anions: , , .
- Every ion ends with a noble gas configuration — or .
- pairs one of each; pairs one of each; needs two chlorines because magnesium gives two electrons.
- Electrons lost must equal electrons gained, which is the same as the charges balancing.
- Criss-cross the valencies and simplify: , , , , , . And becomes , becomes .
- Polyatomic ions: , , , with valency ; , , with ; with .
- , , , , , , and simplifying to .
- Bracket a polyatomic ion when more than one is needed — is wrong.
- Balance the charges to verify: gives against .
- An ionic compound has no molecules — it is a giant lattice, and the formula gives only the ratio of ions.
Draw the dot structures of five molecules and check each by the pair count before looking anything up — a structure that accounts for every pair is almost certainly right.