Chapter 3.1 Structures and Properties of Substances. Chemical Bonding

Transcription

1 Chapter 3.1 Structures and Properties of Substances Chemical Bonding

2 The orbitals in the Periodic Table The elements of the periodic table can be classified according to the type of orbital that is being filled. Elements that appear in the s block and the p block are called either the main group elements or the representative elements. These elements are representative of a wide range of physical and chemical properties. Among them there are highly reactive, moderately reactive and non reactive elements. While most are solids at room temperature, roughly one quarter of them are gases, one is a liquid. Elements that appear in the d block are called the transition elements. the f block elements are called the inner transition elements 1 18 (IA) (VIIIA) 1 1s 2 (IIA) 13 (IIIA) 14 (IVA) 15 (VA) (VIA) (VIIA) 2 2s 2p 3 3s 3 (IIIB) 4 (IVB) 5 (VB) 6 7 (VIB) (VIIB) (VIIIB) 11 (IB) 12 (IIB) 3p 4 4s 3d 4p 5 5s 3 (IIIB) 4d 5p 6 6s 4f 5d 6p 7 7s s block (main group elements) 5f f block (inner transition elements) 6d d block (transition elements) p block (main elements)

3 Chemical Bonding A water molecule has a bent shape, carbon dioxide is linear. An ammonia molecule looks like a pyramid, and sulfur hexafluoride is shaped like an octahedron. All molecules in nature have a specific shape, which is important to their chemistry. Each nerve cell in the brain communicates with adjacent nerve cells by releasing molecules called neurotransmitters from one cell to the next. Enzymes are assisting in the chemical breakdown of food in our digestive system. The aroma of cologne is the result of odorous molecules migrating to specific sites in our nasal passages. Each of these situations depends on the ability of one molecule with a specific shape to fit into a precise location with a corresponding shape (a receptor). The properties of substances derive from the ways in which particles bond together

4 Chemical Bonding Of the about 120 elements that occur in nature or that have been produced synthetically, only the noble gases exist naturally as single, uncombined atoms. In nature, systems of lower energy tend to be favored over systems of higher energy. In other words, lower-energy systems tend to have greater stability than higher-energy systems. Bonded atoms, therefore, tend to have lower energy than single, uncombined atoms Defintion Chemical bonds are electrostatic forces that hold atoms together in compounds and involve the interaction of valence electrons.

5 Using Lewis Structures to Represent Atoms To draw the Lewis structure of an atom: 1.replace its nucleus and inner electrons with its atomic symbol 2.add dots around the atomic symbol to symbolize the atom s valence electrons (many chemists place the dots starting at the top and continue adding dots clockwise, at the right, then bottom, then left. then begin again at the top) Drawing a Lewis structure for a molecule lets you see exactly how many electrons are involved in each bond, and helps you to keep track of the number of valence electrons Na Mg Al Si P S Cl Ar

6 Using Lewis Structures to Represent Atoms There are two ways to show the bonding pairs of electrons. Use dots only Show the bonding pairs as lines between atoms. In this case dots are reserved for representing a lone pair (a non-bonding pair) of electrons O C O or O C O 4 lone pairs

7 Ionic Bonding Ionic bonding occurs between atoms of elements that have large differences in electronegativity usually a metal (low electronegativity) and a non-metal (high electronegativity). The units of ionic compounds such as sodium chloride and magnesium fluoride cannot be separated easily by direct heating of the crystal salts. The ions that make up the ionic solid are arranged in a specific array of repeating units. In solid sodium chloride, for example, the ions are arranged in a rigid lattice structure. In such systems, the cations and anions are arranged so that the system has the minimum possible energy Lattice structure of sodium chloride

8 Ionic Bonding Because of the large differences in electronegativity, the atoms in an ionic compound usually come from the s block metals and the p block non-metals. Mg F E.g. magnesium in Group 2 and fluorine in Group 17 combine to form the ionic compound magnesium fluoride (MgF2). The figure shows a repeating unit in the crystal model of magnesium fluoride. The process that results in the formation of ions can be illustrated with Lewis structures F F Mg F Mg 2 F

9 Ionic Bonding Because of the large differences in electronegativity, the atoms in an ionic compound usually come from the s block metals and the p block non-metals. Mg F E.g. magnesium in Group 2 and fluorine in Group 17 combine to form the ionic compound magnesium fluoride (MgF2). The figure shows a repeating unit in the crystal model of magnesium fluoride. The process that results in the formation of ions can be illustrated with the box diagram F 1s 2s 2p F 1s 2s 2p Mg 1s 2s 2p 3s Mg 2 1s 2s 2p F 1s 2s 2p F 1s 2s 2p

10 Ionic Bonding Practice problem Write electron configurations for the following elements: a. Li b.ca 2 c.br d.o 2 Draw Lewis structures for these chemical species Draw orbital diagrams (box) and Lewis structures to show how the following pairs of elements can combine. In each case, write the chemical formula for the product. e.li and S f.ca and Cl g.k and Cl h.na and N

11 Properties of Ionic Solids In general, ionic solids have the following properties: crystalline with smooth, shiny surfaces hard but brittle non-conductors of electricity and heat high melting points many ionic solids are also soluble in water (MgF2 is an exception) The amount of energy given off when an ionic crystal forms from the gaseous ions of its elements is called the lattice energy (e.g. The lattice energy of MgF2 is 2957 kj/mol). The same amount of energy must be added to break the ionic crystal back into its gaseous ions.

12 Covalent Bonding Covalent bonding results from the balance between the forces of attraction and repulsion that act between the nuclei and electrons of two or more atoms. Example In H2 molecule there is an optimum separation for two hydrogen atoms at which their nucleus-electron attractions, nucleusnucleus repulsions, and electron-electron repulsions achieve this balance. This optimum separation favors a minimum energy for the system, and constitutes the covalent bond between the two hydrogen atoms. nucleus electron repulsion optimum separation H2 molecule attraction Unlike ionic bonding, covalent bond involves the sharing of pairs of electrons and the formation of a new orbital, caused by the overlapping of atomic orbitals. overlapping region of increased of atomic electron orbitals density 1s 1s

13 Characteristics of Covalent Bonding Generally, electron-sharing enables each atom in a covalent bond to acquire a noble gas configuration. The period 2 non-metals from carbon to fluorine must fill their 2s and 2p orbitals to acquire a noble gas configuration like that of Ne (octet rule). E.g. In the formation of the diatomic fluorine molecule, F2, the bonding (shared) pair of electrons gives each fluorine atom a complete valence level. lone pairs, are not involved in bonding. bonding pair lone pairs F F or F F

14 Characteristics of Covalent Bonding Some molecules are bonded together with two shared pairs of electrons. These are called double bonds. CO2 is an example of a covalent molecule that consists of double bonds O C O or O C O Molecules that are bonded with three shared pairs of electrons have triple bonds. Nitrogen, N2, another diatomic molecule, is a triple-bonded molecule N N or N N

15 Characteristics of Covalent Bonding Bond energy is the energy required to break the force of attraction between two atoms in a bond and to separate them. Thus, bond energy is a measure of the strength of a bond. The bond energy increases if more electrons are shared between two atoms because there is an increase in charge density between the nuclei of the bonded atoms. Between Carbon Atoms and Between Nitrogen Atoms C C C Bond Bond energy (kj/mol) Bond Bond energy (kj/mol) C C C N N N N N N

16 Properties of Ionic Solids In contrast to ionic solids, covalent compounds typically have the following properties: exist as a soft solid, a liquid, or a gas at room temperature have low melting points and boiling points are poor conductors of electricity, even in solution may not be soluble in water Diamond (C) Quartz (SiO2)

17 Predicting Covalent or Ionic Bonding We can use the electronegativity difference between the bonding atoms to predict the type of bond. E.g. Two atoms with identical electronegativities, such as chlorine ( EN= =0) share their electrons equally. They are bonded covalently. In sodium chloride, chlorine (EN=3.16) attracts an electron much more strongly than sodium (EN=0.93). Therefore, sodium s valence electron has a very high probability of being found near chlorine. A high electronegativity difference is characteristic of ionic compounds. mostly ionic ( EN > 1.7) polar covalent ( EN ) EN For atoms that have EN between 0.4 and 1.7, the bond is polar covalent. A polar covalent bond has an unequally shared pair of electrons between two atoms. This unequal sharing results in a bond that has partially positive and partially negative poles. mostly covalent ( EN < 0.4) The relationship between bonding character and electronegativity difference 0

18 Metallic Bonding About two-thirds of all the naturally occurring elements are metals (see lesson 1). Metals conduct electricity and heat in both their solid and liquid states. Most metals are malleable and ductile (can be easily stretched, bent, and deformed without shattering of the whole solid). In general, metals change state at moderate to high temperatures. Most metals have either one or two valence electrons Lithium Sodium Potassium

19 Metallic Bonding Based on electronegativity differences metals do not form ionic bonds with other metals. Similarly, metals do not have a sufficient number of valence electrons to form covalent bonds with one another. Metals do, however, share electrons. Unlike the electron sharing in covalent compounds, however, electron sharing in metals occurs throughout the entire structure of the metal. Metals are composed of a densely packed core of metallic cations, within a delocalized region of shared, mobile valence electrons (free-electron model). The force of attraction between the positively charged cations and the pool of valence electrons that moves among them constitutes a metallic bond. Na millions of atoms e sea

20 Properties of the Metallic Bonding The free-electron model explains many properties of metals: Conductivity: Metals are good conductors of electricity and heat because electrons can move freely throughout the metallic structure. This freedom of movement is not possible in solid ionic compounds, because the valence electrons are held within the individual ionic bonds in the lattice. Melting and Boiling Points: The melting and boiling points of Group 1 metals are generally lower than the melting and boiling points of Group 2 metals. Because the greater number of valence electrons and the larger positive charge of Group 2 atoms result in stronger metallic bonding forces

21 Properties of the Metallic Bonding The free-electron model explains many properties of metals: Malleability and Ductility: The malleability of metals can be explained as metallic bonds are non-directional. The positive external ions are layered as fixed arrays (like soldiers lined up for inspection). When stress is applied to stress applied a metal, one layer of positive ions can slide over another layer. The layers move without breaking the array. external stress applied deformed metal Figure 4.9 Metals are easily deformed because one layer of positive ions can slide over another. At the same time, the free electrons (shown as a yellow cloud) continue