Covalent Bonding Bonding between non-metals

Ionic Bonding So far, we have studied compounds where one element (a metal) donates one or more electrons to another element (a non-metal). 2Na + Cl 2  2NaCl These ionic compounds are crystalline solids with high melting and boiling points.

What about combinations of elements that do not tend to completely transfer electrons? often composed of one or more non metals the compounds have low melting and boiling points often liquids or gases at room temperature are molecules rather than formula units

How do non-metals complete their outermost energy levels? These atoms share electrons – electrostatic attraction does not explain their bonding. Electron sharing always occurs in pairs – two atoms can share up to 3 pairs of electrons in covalent bonding. There are far more covalent compounds than ionic compounds

Metals and Non-Metals Non-metals are few in number, but many are very abundant compared to metals. There are many more metals than non-metals, but quite a few are relatively rare. Non-metals make many more compounds with each other through sharing electrons than metals do. Often there are multiple compounds of two non-metals, like hydrocarbons.

Using electron dot diagrams, a diatomic molecule of hydrogen can be represented as The two dots between the hydrogen represents a covalent bond – a pair of shared electrons Both atoms attain the He noble gas configuration by sharing the electrons

Similarly, for some other simple covalent compounds: Shared pairs can also be represented as dashes: H-HF-FH-Cl F2F2 HCl

Sharing Electrons In each case, sharing of electrons occurs if the atoms attain the electron configuration of a noble gas. This is the octet rule for covalent bonding. Note that for F and Cl, not all of the pairs of valence electrons are shared – each has 3 unshared pairs of electrons. These are sometimes called lone pairs or non-bonding pairs of electrons.

Multiple Covalent Bonding Some atoms are capable of sharing more than one pair of electron at a time. The air we breathe is composed of two such atoms, O 2 and N 2. In the case of O 2, two pairs of electrons should be shared or O=O However, experimental evidence shows that O 2 only forms a single bond plus 2 unpaired electrons – an exception to the octet rule.

Nitrogen In the case of N 2, three pairs of electrons are shared::N:::N: Nitrogen does follow the octet rule. Each atom has 5 valence electrons and shares 3 electrons with the other to attain the Ne electron configuration.

Electronegativity This term, coined by Linus Pauling in 1930, refers to the ability of atoms to attract electrons in chemical bonds. Electronegativity is not a formal measurement. It is a convenient rating scale that is used to indicate the degree to which electrons are shared by atoms.

The electronegativity scale The scale goes from 0.7 to 4.0 Metals have low electronegativity. The most active metals have electronegativities less than 1.0 Non-metals have high electronegativity, and have values between 2 and 4. Electronegativity increases across periods, decreases down groups (like ionization energy)

Electronegativity and the PT

The Degree of Sharing in Covalent Bonds Diatomic elements (H-H, N=N, O=O, F-F, etc.) share electrons equally. A nonpolar covalent bond – occurs when two atoms of the same element share their bonding electrons.

What about when the bonding atoms are not the same? Different elements have different electronegativities (the ability to attract bonding electrons) A polar covalent bond – occurs when two different atoms of different elements and electronegativities unequally share their bonding electrons.

Unequal Sharing of Electrons The atom with the stronger electronegativity will attract the electrons more stronger than the weaker atom. The more highly electronegative atom will have slight negative charge The less electronegative atom will have a slight positive charge

Showing Bond Polarity Note that the Greek letter delta is used to denote partial positive and negative charges, far less than +1 or –1. Polarity can also be represented by an arrow pointed towards the more electronegative atom.,

Differences in the electronegativities of the bonding atoms can be used to classify the type of bond involved.

When the electronegativity differences are greater than 2.0, the bond is considered to be predominantly ionic. Less than 2.0 is considered primarily covalent. A bond with a small electronegativity difference (less than 0.5) is considered non- polar covalent.

Using Table S of your reference tables, find the electronegativity difference between elements in the following covalent bonds. BondC-BrC-ClC-FC-I Electronegativity Difference Each element that is bonded to carbon is a member of the halogen family (group 17). What happens to electronegativity within this family?

BondC-HN-HO-HF-H Electronegativity Difference What happens to electronegativity as you go across period 2? What happens the polarity of the bonds as you go across the period?

BondN-OC-OS-OP-O Electronegativity Difference Which of these bonds will be the most polar? List these bonds in order of increasing ionic character (largest difference in electronegativity Mg-Cl Mg-F Ca-Br Ca-I Al-Cl Ca-ClK-ClH-Cl Al-HGe-HSb-HAs-H

Classify the following bonds as ionic, nonpolar covalent, or polar covalent Bond Electronegativity Difference Polar, non- polar, ionic? H and S K and Br N and O Cl and Cl S and O Si and Cl

Polar Molecules When a molecule has a polar bond, one end is somewhat positive and the other somewhat negative. These polar molecules are called dipoles. HCl is an obvious example of a dipole.

Symmetrical Molecules When a molecule is made of more than two atoms, the geometry of the molecule becomes important. When polar bonds are distributed uniformly throughout the molecule, the molecule is nonpolar.

Symmetrical Molecules: Methane Consider methane – what is the electronegativity difference between C and H? The symmetrical distribution of the C-H bonds cancels the effects of the slight polarity of the bonds.

Methane Relatives: CH 3 Cl and CCl 4 A relative of CH 4, chloromethane, is a polar molecule because of the presence of the highly electronegative Cl atom. What about CCl 4 ? Polar or non-polar?

Symmetrical Molecules: CO 2 Carbon dioxide features two highly polar C=O bonds. Yet the molecule is a gas at standard temperature and pressure (1 atm, 273K) Bond polarity cancels because the distribution of electrons is in opposite directions (symmetrical charge distribution)

Polar Molecule: Water The electronegative oxygen pulls the bonding electrons away from the hydrogen, gaining a slightly negative charge. The hydrogens become slightly positive. This molecule is highly asymmetrical

Properties of Polar Molecules The marvelous properties of water are due to its polarity. Water has the ability to dissolve many ionic compounds. In general, the presence of positive and negative ends in polar molecules produces forces of attraction between similar nonpolar molecules. This can result in higher melting and boiling points for polar molecules than non-polar ones.

Hydrogen “Bonds” In compounds like water, ammonia and hydrogen fluoride, the hydrogen atoms are bonded to small atoms of high electronegativity. These hydrogen atoms form polar covalent bonds which result in a limited share of the electron pair that forms the bond.

Hydrogen Bonds These hydrogen atoms are still hungry for electrons, and behave much like an exposed proton. They can be attracted to and form another weaker bond with a neighboring atom of high electronegativity. These kind of bonds, about 5% of the strength of a true covalent bond, are called hydrogen bonds.

The strength of a hydrogen bond Increases with the electronegativity of the atom bonded to hydrogen Decreases with increasing size of the bonded atom. Negative charge is more concentrated in smaller atoms and thus leads to stronger H bonds. Any molecule where hydrogen is bonded to highly electronegative elements such as F, O, or N will produce hydrogen bonding.

What effect does hydrogen bonding have on the physical properties of compounds?

N, O, F have high electronegativities, and significant hydrogen bonding. Because of this, these substances have low vapor pressures have high heats of vaporization or melting have high melting and boiling points require large inputs of energy to break the many hydrogen bonds between molecules, breaking down large clusters of molecules into separate molecules

This causes the boiling points of NH 3, H 2 O and HF to be higher than their counterparts in periods 3-5, which have lower electronegativity.

Hydrogen Bonding in DNA

Hydrogen Bonding in Ice

Hydrogen Bonding in Polymers