kinds of chemical bonds

7.0 7.0 The Chemical Bond Types of Bonds There are two kinds of chemical bonds, ionic and covalent. In an ionic bond one or more electrons is transferred from one atom to another, typically from an atom on the left of the periodic table to one on the right. For example sodium will lose an electron to chlorine forming Na+, a sodium cation, and Cl-, a chloride anion. Together they form NaCl, a salt. The atoms are held together by electrostatic forces, typically in a solid crystal structure. If electrons are shared between two atoms rather than transferred from one atom to another the bond is covalent. In a covalent bond two electrons, one from each atom, are held about half way between the nuclei. Atoms from the same region of the periodic table tend to form covalent bonds. For example, the hydrogen molecule, H2, is 100% covalent. Some bonds, for example the OH bond in water, are partially ionic and partially covalent. These are polar bonds and will have some ionic character. Water is a good solvent for salts and other polar and ionic materials because of its polar character. On the other hand it is not a good solvent for non polar hydrocarbons such as fat, oil, or grease. 7.1 Ionic bond; Common anions and cations An atom can be negatively or positively charged. The charge (in the form of electrons) gained by one atom is lost by another. If electrons are transferred from one atom to another rather than shared, the bond is ionic. This is the bond formed by a metal on the left of the chart with an element in groups 7 or 6 or with one of the common anions listed below. The elements on the left easily lose electrons and those on the right gain electrons. Typically enough electrons are gained and lost to satisfy the octet rule or to give the inert gas configuration at each ion. The ions are held together by electrostatic forces. The formation of an ionic bond can be represented by the following, .. .. + . + . Cl . Na Na Cl . .. .. Chemistry can be viewed as an extended discussion about the transfer of electrons (charge) between atoms. Chemical names reflect the loss or gain of electrons. 7.2. Nomenclature (Cations, Anions, Oxidation, Reduction) Cations A charged atom is called an ion. There are two kinds of ions, cations and anions. Cations are positively charged. The size and polarity of the charge is written as a superscript to the right of the atomic symbol. A plus 1 charge means that there is one more proton than electron, the atom has lost one electron. Plus 2 means the loss of two electrons, etc. Some cations are . . .. H+ Na+ K+ +1 NH4+ Cu+ Be2+ +2 Sr2+ Fe2+ +3 Al3+ Fe3+ +4 Ti4+ Mg2+ Ba2+ Ca2+ Cu2+ The process whereby an atom (or anything else) losses electrons and becomes a cation is called oxidation. For example, Na Ca Fe2+ Na+ + Ca2+ + Fe3+ + e2ee- are all oxidations. The charge on the cation is called the oxidation state. The oxidation state is frequently indicated by roman numerals. Fe is iron, so Fe+3 is described as iron (III) where the III indicates a +3 oxidation state, or a +3 charge on the iron atom. Elements on the left side of the periodic chart lose electrons more easily than elements on the right, and thus tend to form cations more readily. Elements in group one lose one electron, in group 2 lose two electrons, in group three lose three electrons, and in group 4 lose four electrons. The name of the element is also used to name the cation. The charge is indicated by a roman numeral. Anions Some elements, especially those on the right of the chart, pick up electrons and form negatively charged anions. Common anions include -1 OH- (hydroxide) Cl- (chloride) Br- (bromide) F- (fluoride) I- (iodide) NO3- (nitrate) -2 SO42- (sulfate) CO32-(carbonate) O2- (oxide) S2- (sulfide) -3 PO43- (phosphate) Elements in group seven (halogens) pick up one electron, while elements in group six (O, S, Se, etc.) pick up two electrons. The process of picking up electrons is called reduction. O2 F2 + + 4e2e2O22F- The reduced form is indicated by a change in the name ending to -ide. fluorine (F) becomes fluoride (F-) chlorine (Cl) bromine (Br) iodine (I) oxygen (O) sulfur (S) nitrogen (N) " " " " " ‘ chloride (Cl-) bromide (Br-) iodide (I-) oxide (S2-) sulfide (S2-) nitride (N3-) There are many complex anions, some of which are listed above, and many more in chemistry textbooks . There is one common complex cation, the ammonium ion, NH4+. For example ammonium chloride, NH4Cl, and diammonium sulfate, (NH4)2SO4. 7.3 Oxidation state The oxidation state of an element is the formal charge on it. The oxidation state of L1+ is +1, of O2- is -2. The complexitiy occurs in the case of compounds, for example NaCl or Na2SO4. In these cases there are a set of simple rules for assigning oxidation states. 1. The oxidation state of an element in its elemental state, uncombined with other elements in a compound, is zero. For example the oxidation states of the elements in O2, H2, S, Al, N2, Cl2 are all zero. In compounds, 2. 3. 4. 5. 6. 7. Group I m4etals are +1. Group II metals are +2 Group III metals are +3. Oxygen is -2 except in peroxides where it is -1. Hydrogen is +1 except in metal hydrides where it is -1. Florine is -1. Other elements, for example Cl, can be anything and must be calculated from the charge on the other atoms in the compound and the total charge of the compound, usually zero. For example, what is the charge on Cl in NaClO? in KClO2? Salts, acids, bases, hydrates. Cations and anions typically combine in such a way to make a neutral (zero net charge) ionic salt. Many salts are soluble in water, for example NaCl, table salt. Most nitrate, ammonium, and group one salts are soluble in water. 7.4 There are two special cases. When the cation is a proton, H+, an acid is formed, and when the anion is hydroxide, OH-, a base is formed. If we combine the H+ with OH- we get H2O, H+ + OHH2O The complete acid base reaction will involve additional anions and cations, for example HCL acid + NaOH base H2O + NaCl water salt The acid reacts with a base to from water and a salt. This is characteristic of acid base reactions and distinguishes them from other reaction types. Acids and base can be anhydrous or hydrated. Oxides from the left of the chart are basic and when hydrated are writen as hydroxides. CaO + H2O Ca(OH)2 2NaOH 2Al(OH)3 Na2O + H2O Al2O3 + 3H2O Acids can also be hydrated, SO3 + H2O H2SO4. 7.5 Arrows and Equilibrium An arrow going one way indicates that one is talking about just that one way process, for example the process whereby oxygen picks up electrons to form an oxide ion. O2 + 4e2O2- An arrow pointing both ways refers to an equilibrium process, where both the foreward and back reactions are occurring at the same time. It is analogous to patrons leaving and entering a bar. For example, Ca(OH) CaO + H2O indicates that some of the calcium hydroxide present is losing water to form calcium oxide, and some of the oxide is picking up water to form the hydroxide. At any given moment significant amounts of both the oxide and hydroxide are present. An uneven set of arrows, H+ + OH- H2O means that the equilibrium is favored in the direction of the largest arrow, in this case there is much more water present than either H+ or OH-. 7.6 Covalent Bonding The simplest description of the covalent bond is according to the octet or filled shell rule. An atom with an unfilled shell of orbitals can accept enough electrons from other atoms to fill the shell even though they formally belong to the other atom. As the atoms approach each other, orbitals containing one electron each overlap so that one pair of electrons (one electron from each atom) occupies both orbitals. To picture this situation in the form of a simple cartoon, the electronic structure of each reacting atom is written as a Lewis dot structure, Section 1.9. The valence electrons are represented as dots filling the four (one s and three p) valence orbitals. One bond is formed for each unpaired electron. Consequently hydrogen (H) can form one bond, helium (He) none (no unpaired electrons), lithium (Li) one, Beryllium (Be) forms two, boron (B) three, carbon (C) four, nitrogen (N) three with one lone pair, oxygen (O) two, fluorine (F) one, and neon (Ne) none since it has a closed shell without forming bonds. Atom H Li Be B C N O F No. Bonds 1 1 2 3 4 3 2 1 No. Non Bonded Electron Pairs 0 0 0 0 0 1 2 3 The same situation occurs for period 3. The strength of the bond is a result of the build up of electron density between the two positive nuclei where the orbitals overlap. In the following picture two hydrogen atoms share their unpaired electrons to form a two electron covalent bond. Each hydrogen has two electrons occupying the region of space designated as the 1s orbital. The two electrons are in a stable closed shell in the case of both He and H2. . H. + H. H He 7.7 Covalent Sigma Bonds and the Shape of Molecules The shape of the molecule is determined largely by electron repulsion. Electrons will stay as far from each other as possible. Two electrons will be on opposite sides from each other, three at opposite corners of a triangle, and four at opposite corners of a tetrahedron. Since bonds are electrons, molecular shapes follow the same rules, where the bonding atom is placed where the electrons are. The rules are as follows. H .. . Electrons in two orbitals Electrons in three orbitals Electrons in four orbitals sp sp2 sp3 linear triangular tetrahedral For example six electrons will form a tetrahedron around oxygen in water (write the Lewis dot structure). Two unpaired electrons in two of the orbitals will combine with the unpaired single electron from each of two hydrogens to form water, and the shape will be bent. A lone electron pair is a filled orbital which occupies a hybridized orbital but is not shared. There are two lone electron pairs in water, below. . O. .. + .H H .. O. .. How would one form OH- + H+ from the above water molecule? Double and triple bonds, Sigma (σ) and p Bonds σ For any atom (such as nitrogen, and especially carbon or silicon) or ion (for example -) with four or more electrons in the four s and p orbitals, the s and the three p orbitals Al hybridize into four equivalent sp3 orbitals directed towards the corners of a tetrahedron. This occurs provided no double bond character occurs. A double bond may occur between two adjacent p orbitals directed parallel to each other, Figure 1. In this case the overlap is less and the second p bond is somewhat weaker than the sigma bond. Although the double bond is stronger than one single bond, it is not twice as strong. In the case of a single double bond the remaining s and two p orbitals again form equivalent sp2 orbitals directed toward the vertices of an equilateral triangle. It is possible to form two π bonds, a triple bond, in which case the remaining sp orbitals form two linear sigma bonds. 7.8 How does one know weather or not a double bond is present? I most cases one can make a guess from the formula by knowing how many bonds each element must form and the numbers of each. Try the examples below. Hint, assume that H forms one bond and not with itself, oxygen two, not with itself, and go from there. One might follow the following logic; four of more valence electrons implies tetrahedral shape, sp3. One double bond will use one p orbital, changing the symmetry to sp2, planar, and two π bonds or a triple bond gives sp, linear. 7.9 The Representation of Covalent Bonds Covalent bonds are represented by lines. The hydrogen molecule H2 is H - H, water is H-O-H, hydrogen sulfide H2S is H-S-H, carbon dioxide CO2 is O=C=O. Note that oxygen has two bonds and carbon four in accordance with the rules, section 4.1. 7.10 Atomic and Molecular Orbitals H . + H . .. .. Two hydrogen atoms bond together when they approach each other so closely that their 1s atomic orbitals overlap. The two atomic orbitals combine (add) to form a new molecular orbital called a 1 σ (sigma) orbital sometime written as the molecular wave function ψg = φ1 + φ2 where φ1 and φ2 are the 1s atomic orbitals on hydrogen atoms 1 and 2 respectively. ψg is the molecular orbital sometimes written as 1σg. The sigma refers to the symmetry of the orbital in analogy with the 1s hydrogen orbital. There are two electrons in this lowest energy orbital, one from each hydrogen atom. Since the orbitals add mostly between the nuclei in the overlap region, that is where the bonding electrons are found, Figure 7.1. Since there are two hydrogen atoms and two atomic orbitals, there must also be two molecular orbitals. The other molecular orbital is the only other possible linear combination, ψu = φ1 - φ2. Since this orbital is the difference between the two atomic orbitals, it has the shape shown in Figure 7.1. The difference is greatest at the ends of the molecule where the contribution from the furthest atom is zero, and is zero in the middle where the contributions from the two atoms are the same. This is called an anti bonding orbital and is designated as a 1σu*. Electrons in this orbital will be located at the ends of the molecule, away from the bonding electrons, and will have a higher energy. The g used as a subscript is from the German word gerade and, like the use of sigma, refers to the symmetric nature of the bonding (in this case) molecular orbital. The u is from the word ungerade and refers to the anti-symmetric nature of the antibonding orbital. Symmetry plays a major role in physics and in parts of chemistry. The * superscript means that the orbital is anti-bonding. 1σu * 1s 1 σg H-H Energy Figure 7.1 Electrons represented by vertical lines occupy 1s atomic ortitals in seperated hydrogen atoms. As the atoms come closer the atomic orbitals either add to form lower energy bonding orbitals (1σg) where the electorn density builds up between 1s H H the nuclei, or subtract to form higher energy antibonding orbitals (1σu*) where the electon density (if the orbitals were occupied) is at the ends of the molecule. Two electrons will occupy the bonding orbital to form the hydrogen molecule, H-H. Electrons in a bonding orbital stabilize the molecular bond, and electrons in an anti-bonding orbital destabilize the bond. Each electron in a bonding orbital contributes half of a bond, two electrons are one bond, and each electron in an anti bonding orbital subtracts half of a bond. In the case of helium each atom contributes two electrons, so four electrons occupy the two molecular orbitals, two in the 1s bonding orbital and two in the 1s anti bonding orbital. The net bond, sometimes called the bond order, is therefore zero since the two electrons in the bonding orbital add one bond and the two electrons in the anti bonding orbital subtract one bond. Similar considerations can be applied to more complicated molecules involving larger atoms with more orbitals and electrons and more than two atoms. See for example the short book Energy levels in Atoms and Molecules, Richards and Scott, Oxford Chemistry Primers series, Oxford Press, 1994. Problems and Pretest 1. Using these considerations draw the line and dot structures of NH3, water, methane, CO2 , N2H4, N2H2., HCN, CH2NH. 2. Write formulas for aluminum phosphate, aluminum sulfate, iron (II) phosphate, copper (I) sulfide, copper (II) sulfate. What is the oxidation state of P, Al, iron, copper, oxygen and sulfur in each compound? 3. Write the names of the following compounds, Al2O3, FeS, Fe2S3, CaCO3, Na2CO3, KOH, MgSO4, NH4OH and identify the oxidation states. 4. Identify each as an oxidation or reduction. S2 I2 K Mg + + 4e2eK+ + 2S22Ie2e- Mg2+ + V4+ V5+ + e-