Dipole moment of a water molecule : Water has a very large dipole moment which results from the two polar H—O bonds oriented at an angle of The bond dipoles add up to create a molecular dipole indicated by the green arrow. Privacy Policy. Skip to main content. Basic Concepts of Chemical Bonding. Search for:.
Learning Objectives Apply the rules for assigning oxidation numbers to atoms in compounds. The atom with higher electronegativity, typically a nonmetallic element, is assigned a negative oxidation number, while metallic elements are typically assigned positive oxidation numbers. Key Terms electronegativity : A chemical property that describes the tendency of an atom to attract electrons or electron density toward itself. It indicates of the degree of oxidation of an atom in a chemical compound.
Learning Objectives Apply knowledge of bond polarity and molecular geometry to identify the dipole moment of molecules. Key Takeaways Key Points When non-identical atoms are covalently bonded, the electron pair will be attracted more strongly to the atom that has the higher electronegativity. This results in a polar covalent bond. Polarity refers to a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment.
A polar molecule acts as an electric dipole that can interact with electric fields that are created artificially, or that arise from nearby ions or polar molecules. The ability of an atom to attract electrons in a chemical bond is called the electronegativity of the atom. The electronegativity of an atom is related to its electron affinity and ionization energy.
Electron affinity is the energy liberated by a gaseous atom when an electron is added to it. Ionization energy is the minimum amount of energy necessary to remove the most weakly bound electron from a gaseous atom. Electronegativity level is normally measured on a scale that was created by Linus Pauling.
On this scale, the more electronegative elements are the halogens, oxygen, nitrogen, and sulfur. Fluorine, a halogen, is the most electronegative with a value of 4. The less electronegative elements are the alkali and alkaline earth metals. Of these, cesium and francium are the least electronegative at values of 0. Elements with great differences in electronegativity tend to form ionic bonds. Atoms of elements with similar electronegativity tend to form covalent bonds. Pure covalent bonds result when two atoms of the same electronegativity bond.
Intermediate differences in electronegativity between covalently bonded atoms lead to polarity in the bond. The positively charged protons in the nucleus attract the negatively charged electrons.
As the number of protons in the nucleus increases, the electronegativity or attraction will increase. Therefore electronegativity increases from left to right in a row in the periodic table. This effect only holds true for a row in the periodic table because the attraction between charges falls off rapidly with distance.
The chart shows electronegativities from sodium to chlorine ignoring argon since it does not does not form bonds. As you go down a group, electronegativity decreases.
If it increases up to fluorine, it must decrease as you go down. The chart shows the patterns of electronegativity in Groups 1 and 7. Consider sodium at the beginning of period 3 and chlorine at the end ignoring the noble gas, argon. Think of sodium chloride as if it were covalently bonded. Both sodium and chlorine have their bonding electrons in the 3-level.
The electron pair is screened from both nuclei by the 1s, 2s and 2p electrons, but the chlorine nucleus has 6 more protons in it. It is no wonder the electron pair gets dragged so far towards the chlorine that ions are formed. Electronegativity increases across a period because the number of charges on the nucleus increases. That attracts the bonding pair of electrons more strongly. As you go down a group, electronegativity decreases because the bonding pair of electrons is increasingly distant from the attraction of the nucleus.
Consider the hydrogen fluoride and hydrogen chloride molecules:. The bonding pair is shielded from the fluorine's nucleus only by the 1s 2 electrons.
In the chlorine case it is shielded by all the 1s 2 2s 2 2p 6 electrons. But fluorine has the bonding pair in the 2-level rather than the 3-level as it is in chlorine. If it is closer to the nucleus, the attraction is greater.
At the beginning of periods 2 and 3 of the Periodic Table, there are several cases where an element at the top of one group has some similarities with an element in the next group. Three examples are shown in the diagram below. Notice that the similarities occur in elements which are diagonal to each other - not side-by-side. For example, boron is a non-metal with some properties rather like silicon. Unlike the rest of Group 2, beryllium has some properties resembling aluminum.
And lithium has some properties which differ from the other elements in Group 1, and in some ways resembles magnesium. There is said to be a diagonal relationship between these elements. There are several reasons for this, but each depends on the way atomic properties like electronegativity vary around the Periodic Table.
So we will have a quick look at this with regard to electronegativity - which is probably the simplest to explain. Electronegativity increases across the Periodic Table. So, for example, the electronegativities of beryllium and boron are:. Electronegativity falls as you go down the Periodic Table. So, for example, the electronegativities of boron and aluminum are:. So, comparing Be and Al, you find the values are by chance exactly the same.
The increase from Group 2 to Group 3 is offset by the fall as you go down Group 3 from boron to aluminum. Something similar happens from lithium 1. In these cases, the electronegativities are not exactly the same, but are very close.
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