! CHEM!161:!Chapter!8!v1114! ! ! ! ! ! ! page!1!of!19!Chapter 8: Chemical Bonds: What Makes a Gas a Greenhouse Gas? Problems: 8.1-8.40, 8.41a, 8.42a, 8.44-8.56, 8.67-8.74, 8.77, 8.81-8.85, 8.87-8.88, 8.91-8.98, 8.99 (c,d only), 8.100 (a,d only), 8.101-8.105, 8.107-8.110, 8.115-8.120, 8.122-8.137, 8.139-8.143, 8.147-8.149, 8.152, 8.155, 8.160, 8.165-8.166, 8.171-8.174, Chapter 10 Problems: 10.1, 10.9-10.14, 10.100 8.1 TYPES OF CHEMICAL BONDS chemical bond: what holds atoms or ions together in a compound The three types of chemical bonds are ionic bonds, metallic bonds, and covalent bonds. – Metallic bonds hold metal atoms together in metals. – Ionic bonds hold ions together in ionic compounds. – Covalent bonds hold atoms together in molecules. METALLIC BONDS Metals exist as nuclei surrounded by a sea of electrons. → The electrons in a metal are shared among all the nuclei, so the electrons are delocalized (i.e., not fixed to a specific atom). → The electrons can shift throughout the entire metal. → Electrons are free to move throughout the solid → metals’ unique properties. – Pots and pans are usually made of metal because metals conduct heat and electricity as electrons flow through the metal. – Metals are malleable and ductile because electrons act as a glue, holding the positively-charged nuclei together → hammered metal and metallic wire.! CHEM!161:!Chapter!8!v1114! ! ! ! ! ! ! page!2!of!19!THE IONIC BOND Ex. 1 Give the Lewis electron-dot formula below for each of the following atoms and ions: sodium sodium ion chlorine chloride ion Metal atoms transfer valence electrons to nonmetal atoms. → positively charged metal cations and negatively charged nonmetal anions – the charged ions are attracted to each other → ionic compound = 3D network of ions → ionic bond: electrostatic attraction between positively charged cation and negatively charged anion Properties of Ionic Substances: – Ionic compounds exist as networks of ions, with cations surrounding anions, and vice versa. → To melt ionic compounds, every bond between each Na+ ion and the Cl− ions surrounding it must be broken, as well as bonds between each Cl− ion and the Na+ ions surrounding it. → A lot of energy is required to break all of these bonds. → Ionic compounds have relatively high melting points—much higher than water’s (0°C). → All ionic compounds are solids at room temperature. – Ionic compounds do not conduct electricity when solid (ions are fixed in place), but they do conduct in the molten (liquid) and aqueous states (ions move around freely). 10.1 INTERACTIONS BETWEEN IONS (Note this is covered in Chapter 10!) The actual melting point of an ionic compound is determined by the strength of the ionic bonds in the compound, which is defined by Coulomb's law. Coulomb's law : The strength of interactions between ions (given by the lattice energy=E) is directly proportional to the product of the ions’ charges (Q1 and Q2) and inversely proportional to the distance between their nuclei (d). Coulomb's law: E ∝ dQ Q21 where Q1 and Q2 are the charges on the ions and d is the distance between the ions’ nuclei! CHEM!161:!Chapter!8!v1114! ! ! ! ! ! ! page!3!of!19!ION-ION INTERACTIONS Coulomb's law, E ∝ dQ Q21 where E=lattice energy=energy released when free gaseous ions combine to form one mole of a solid ionic compound Note: The value of E will be negative because the product of the charges will be negative (due to the negative charge on the anion). – The stronger the ionic bonds that form, → more energy is released when the ionic compound forms, → the more negative the value for E. Thus, the relative strength of the ionic bonds in an ionic compound is given by: 1. Charges of ions: Higher the charge → the stronger the bond – The bond in CaO (Ca+2 and O-2) is stronger than that in NaCl (Na+ and Cl−) → The melting point of CaO (2927°C) is much higher than NaCl's (801°C). 2. Distance between two ions: Shorter distance → stronger the bond – Na+ and Cl− have smaller radii than K+ and Br−. → NaCl’s melting point (801°C) is higher than KBr's (734°C). Note: The strength of the ionic bond is generally determined foremost by the charges, and only if the charges are similar does one compare the distance between nuclei to determine the relative strength of the bond. Ex. 1 Circle the compound in each pair with the stronger ionic bond: a. NaF or MgO c. strontium sulfide or magnesium oxide b. Al2O3 or BaS d. lithium nitride or calcium sulfide Ex. 2: Given that an ionic compound’s melting point depends on the strength of its ionic bonds, rank the following in terms of increasing melting point: lithium fluoride, magnesium oxide, sodium chloride, barium sulfide, and potassium bromide. __________ < __________ < __________ < __________ < __________ lowest m.p. highest m.p. Ex. 3: Rank the following in terms of increasing melting point: magnesium fluoride, mercury, magnesium oxide, calcium sulfide, and sodium sulfide. __________ < __________ < __________ < __________ < __________ lowest m.p. highest m.p.! CHEM!161:!Chapter!8!v1114! ! ! ! ! ! ! page!4!of!19!8.2 LEWIS STRUCTURES COVALENT BOND In 1916, an American chemist, Gilbert N. Lewis, proposed that atoms bond by sharing electrons. → Thus, a bond where atoms share electrons is called a covalent bond. octet rule: Atoms form covalent bonds in such a way that all atoms get eight electrons (an octet) or 4 pairs of electrons, except hydrogen which only needs 2 electrons. – i.e., atoms bond to get the same number of valence electrons as the Noble gas in the same period. covalent bond: sharing of one or more pairs of electrons between 2 nonmetal atoms, so both atoms can achieve a Noble Gas electron configuration. – Covalent bonds hold atoms together in a molecule. H• + •H → H:H Note: When two H atoms bond to form a hydrogen molecule, both atoms get two electrons, just like a stable helium atom. We can also represent the H2 molecule as follows: H H This overlapping region is