Class 11 Chemistry (India)
- Ionic bonds
- Ionic bonds: Reaction of metals & Non-metals
- Covalent bonds
- Single and multiple covalent bonds
- Metallic bonds
- Drawing Lewis diagrams
- Predicting bond type (metals vs. nonmetals)
- Worked example: Lewis diagram of formaldehyde (CH₂O)
- Worked example: Lewis diagram of the cyanide ion (CN⁻)
- Exceptions to the octet rule
- Worked example: Lewis diagram of xenon difluoride (XeF₂)
- Lewis diagrams
Single and multiple covalent bonds
Octet rule - Matter always wants to be in the most stable form. For any atom, stability is achieved by following the octet rule, which is to say all atoms (with a few exceptions) want 8 electrons in their outermost electron shell (just like noble gases). The electrons present in the outermost shell of an atom are called valence electrons.
Exceptions to the octet rule include hydrogen (H) and helium (He) that follow the duet rule instead. They are the first two elements of the periodic table and have a single electron shell which accommodates only 2 electrons. Other exceptions include some group 3 elements like boron (B) that contain three valence electrons. Theoretically, boron can accommodate five more electrons according to the octet rule, but boron is a very small atom and five non-metal atoms (like hydrogen) cannot pack around the boron nucleus. Thus, boron commonly forms three bonds, BH, with a total of six electrons in the outermost shell. This also results in some anomalous properties for boron compounds because they are kind of “short of electrons”. It should be thus noted that covalent bonding between non-metals can occur to form compounds with less than an octet on each atom.
In general, achieving the octet configuration (i.e. 8 electrons in the outermost shell) is the driving force for chemical bonding between atoms. Take a look at the outer shell configuration (i.e. number of valence electrons) of three atoms – sodium (Na), chlorine (Cl) and neon (Ne):
Outer shell configuration diagrams of sodium (Na), chlorine (Cl) and neon (Ne)
Ionic and covalent bonds
Let’s look at the following two scenarios A and B. There are two kids, Emily and Sarah. They both are very good friends.
Scenario A relationship diagram
Scenario B relationship diagram
Now let’s apply the above analogy to chemical bonding. Assume that Emily and Sarah represent two atoms, and the blanket symbolizes their valence electrons. In scenario A, atom Emily is willing to donate her electrons (blanket) to atom Sarah because by doing so both achieve an octet configuration of 8 electrons in their respective outer shells, making them both happy and stable. This donation of electrons is called ionic bonding.
Example of an ionic bond
Example of an ionic bond
In scenario B, both the atoms Emily and Sarah are equally electronegative. So, neither Emily nor Sarah is ready to part with her electrons (blanket), and they instead share their valence electrons with each other. This is called a covalent bond. Electronegativity is a measure of how strongly an atom attracts electrons from another atom in a chemical bond and this value is governed by where the particular atom is located in the periodic table (francium is the least electronegative element while fluorine is the most electronegative).
Example of a covalent bond
Example of a covalent bond
Polar and Non-polar covalent bond
Let’s go back to Emily and Sarah:
Scenario C relationship diagram
Scenario D relationship diagram
Let’s apply the above analogy to a covalent bond formation. In scenario C, both Emily and Sarah are equally cold (in our analogy this translates to them having the same electronegativity). Because they have the same electronegativity, they will share their valence electrons equally with each other. This type of a covalent bond where electrons are shared equally between two atoms is called a non-polar covalent bond.
Example of a non-polar covalent bond
Example of a Non-polar covalent bond
In scenario D, Emily is cold but Sarah is much colder (no doubt mild hypothermia from playing outside in the rain too long)! Together they share the blanket, but Sarah has a tendency to keep pulling the blanket from Emily in order to warm up more. In the atomic world, one atom (Sarah) is more electronegative than another atom (Emily), and naturally pulls the shared electrons towards itself. This pulling of electrons creates slight polarity in the bond. Covalent bonds where electrons are not shared equally between two atoms are called polar covalent bond.
Example of a polar covalent bond
Example of a polar covalent bond
As shown above, the electrons in a covalent bond between two different atoms (H and Cl in this case) are not equally shared by the atoms. This is due to the electronegativity difference between the two atoms. The more electronegative atom (Cl) has greater share of the electrons than the less electronegative atom (H). Consequently, the atom that has the greater share of the bonding electrons bears a partial negative charge (δ-) and the other atom automatically bears a partial positive charge (δ+) of equal magnitude.
Properties of non-polar covalent bonds:
- often occurs between atoms that are the same
- electronegativity difference between bonded atoms is small (<0.5 Pauling units)
- electrons are shared equally between atoms
Properties of polar covalent bond:
- always occurs between different atoms
- electronegativity difference between bonded atoms is moderate (0.5 and 1.9 Pauling units)
- electrons are not shared equally between atoms
Methane (CH) is an example of a compound where non-polar covalent bonds are formed between two different atoms. One carbon atom forms four covalent bonds with four hydrogen atoms by sharing a pair of electrons between itself and each hydrogen (H) atom. The electronegativity value for carbon (C) and hydrogen (H) is 2.55 and 2.1 respectively, so the difference in their electronegativity values is only 0.45 (<0.5 criteria); the electrons are thus equally shared between carbon and hydrogen. So we can conveniently say that a molecule of methane has a total of four non-polar covalent bonds.
Single and Multiple Covalent Bonds
The number of pairs of electrons shared between two atoms determines the type of the covalent bond formed between them.
|Number of electron pairs shared||Type of covalent bond formed|
Now let’s move on to a couple of examples and try to determine the type of covalent bonds formed
Diagram of single covalent bond being formed
Nitrogen atom can attain an octet configuration by sharing three electrons with another nitrogen atom, forming a triple bond (three pairs of electrons shared)
Diagram of nitrogen bonding into octet configuration
Consider the molecule carbon dioxide (CO). Let’s determine the type of covalent bonds it forms.
Diagram of two double covalent bond being formed
Want to join the conversation?
- How do you know the number of valent electrons an element has(7 votes)
- By counting the columns on the periodic table. The column with hydrogen would be Group 1, which means every element within that column only has ONE valence electron to give away. I hope I helped. Sorry if it's still confusing.(39 votes)
- What is the max no of covalent bonds that an atom can form with other atoms? Like in SF6, Sulfur can bond with 6 fluorine atoms, due to additional d orbitals. So, can an element with even more orbitals form even more covalent bonds?(6 votes)
- Sulfur has six valence electrons in the M shell (1s2, 2s2, 2p6, 3s2, 3p4). In order to understand why the six bonds are possible you need to take a look into hybridization. SF6 is so stable that it is energetically favorable for Sulfur to promote two of its electrons to an excited state, which is in the 3d shell, leaving it with a configuration 3s1, 3p3, 3d2. Notice that every orbital has only one unpaired elecron, making they very likely to form a bond with another electron. A sp3d2 hybrid is then formed leaving sulfur with six orbitals without paired electrons. The ability to use the d subshell is what makes it possible for atoms to go beyond the octet, and it's also why atoms up to the second period cannot do that. For that same reason, six or seven bonds are possible, and Xenon can form 8 covalent bonds in the compound XeO4! Meallic elements can definiely have more than eight valence electrons, however they do not tend to form covalent bonds. Therefore the maximum number of covalent bonds should be said to be 7, with the exception of some noble gases since they are very stable by themselves.(4 votes)
- What's the difference between a Polar Covalent Bond and a Covalent Bond?(1 vote)
- Polar covalent bonds do not share electrons equally between two atoms. Non-Polar covalent bonds share electrons equally.(6 votes)
- So, what determines whether a covalent bond will be double, single, or triple? At first I thought electronegativity had something to do with this, but O2 molecules have similar electronegativities, yet they form double covalent bonds.(3 votes)
- The total number of valence electrons a whole compound would have. (Meaning how many more electrons does each atom have than the noble gas before it, then add up that number of electrons for all the atoms to get total valence electrons.) Since every atom needs an octet, with the exception of atoms with a d and f orbital, you can create a lewis structure by placing 8 electrons next to each atom. Now, create bonds to reduce the value by 2 until you have the amount of electrons you intially found were valence in the atom. If you create a single bond, and there are still too many atoms for the number you found, that's how you decide to add more. Conversely, by ensuring each line counts as two electrons, you can determine if you gave a molecule too many bonds, and it needs to have some floating ones on the Lewis Dot. Hope that helps :)(1 vote)
- Even if the electronegativity difference is < 0.5, if the atoms are different and there is some electronegativity difference, wouldn't the electrons be slightly unequally shared between the two atoms?(1 vote)
- I agree, but this is a negligible amount.
Let me use an example of my weight to clarify...
Last week i weighed in at exactly 220.0 lbs, one week later i weighed in at 219.8 lbs.
Sure, my weight has changed, but not by a significant amount. In other words im not going to freak out over 0.2 lbs. Atoms with an electronegativty less than 0.5 in difference will have a slightly unequal distribution of electrons, but not enough to cause a noticeable change.(5 votes)
- Can there be more than three covalent bonds possible between atoms?(3 votes)
- why double bond is more reactive than triple bond ?(3 votes)
- Triple bonds are actually more reactive than double bonds as the sideway overlap of pi bond can be easily broken by addition reactions. Since double bonds have lesser number of pi electrons, so they are relatively more stable than triple bonds.(0 votes)
- How do you know which atom can have how many bonds For example Be (Beryllium) can have only 2 bonds and H(Hydrogen) can only have 1 bond(1 vote)
- The number of bonds formed by an element can only be decided by the number of valence electrons participating in forming bonds. For example, Beryllium electronic configuration is 1s2, 2s2; here valence electrons are 2 therefore only 2 electrons can participate in bond formation. While hydrogen has 1 valence electron therefore it can form only 1 bond. Also the group number tells of the valency of the element. D block elements show variable valencies because these elements have vacnt orbitals where the electrons can jump to and therfore provide more than one way of bonding. i hope this helps.(3 votes)
- Are ionic bonds the strongest all of bonds? Is there any reference page to study coordinate bonds?(3 votes)
- covaelent bonds are stronger than ionic bonds, as shared electrons are harder to seperate then donated electrons(0 votes)
- Why does each single covalent bond count for TWO electrons towards an atom's octet?(1 vote)
- It is mutual sharing and the minimum number of electrons to share is 1. So if the firt element is sharing one electron the second element should also share atleast one electron. Hence single covalent bond is sharing 1 electron from each element perspective.(2 votes)