The breaking of bonds and the forming of bonds occur during chemical reactions.
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1 Chemical Bonding
2 The breaking of bonds and the forming of bonds occur during chemical reactions.
3 Aspirin The formula for a molecule of aspirin is C 9 H 8 O 4 Is it an ionic or covalent (molecular) compound? covalent compound (made of all nonmetals - no ions) What do we call the things that hold a molecule of aspirin together? Bonds
4 Aspirin C 9 H 8 O 4 covalent compound (made of all nonmetals - no ions)
5 The Attachment Between Atoms atoms combine to form ionic bonds covalent bonds (M + NM) (NM + NM) chemical bond a mutual electrical attraction between the nuclei and valence electrons of two atoms that binds the atoms together
6 Ionic Bonding ionic bond electrical attraction between cations and anions; when electrons are taken by one atom from another atom metal and a nonmetal NaCl cation and anion (The charges are hidden to make a neutral compound.)
7 Ionic Bonding: taking of electrons Na 11e - F 9e - Na + 10e - STABLE!!! F - 10e - STABLE!!! 3s 3s 2p 2p 2s 2s 1s 1s
8 I m Positive! A metal ion atom A nonmetal atom ion I m Negative
9 I m Positive! A metal ion A nonmetal ion I m Negative When a metal and a nonmetal atom are around each other there is the opportunity for. the transfer of electrons producing ions that would like to cling to each other. Ionic bonding!!!
10 The simplest ratio of the packed ions is called: cubic shape The Formula Unit Ex: NaCl
11 Ions Metals form cations. (metals lose e - ) Nonmetals form anions. (nonmetals gain e - )
12 Ions cations (+) anions (-) monatomic ions ions formed from one atom Na + or O -2 Examples: polyatomic ions - ions formed from two or more atoms bonded together Examples: NH 4 + or SO 4-2
13 Naming Ions monatomic ions cations named like the atom, only add ion to it» Example: Na + is the sodium ion anions remove the ending to the atom name and add ide and ion to it» Example: Cl - polyatomic ions is the chlorine +ide ion or the chloride ion You do not determine their names, you memorize them
14 Ionic Compounds solid at room temperature (forming crystals) high melting points (thus are usually solid at RT) formula unit represents the lowest ratio of ions that combine to form a neutral compound when dissolved in water, the ionic compounds will break up into ions (dissociate) the solutions of ionic compounds will conduct electricity (electrolytes)
15 Dissociation NaCl(s) HO(l) Na (aq) + Cl (aq) solid placed in water hydrated ions (surrounded by water) dissociation when an ionic compound dissolves to break apart into hydrated ions
16 Dissociation
17 Electrolytes and Nonelectrolytes When an ionic compound dissolves to produce ions, it is called an electrolyte because it conducts electricity in water. When an compound does not dissolve to produce ions, it is called a nonelectrolyte because it does not conduct electricity in water.
18 Electrolytes or salt? Both - Same thing in this example
19 Check for Understanding 1. What kinds of atoms form ionic bonds? 2. What is a polyatomic ion? 3. Name 5 things you learned about ionic compounds.
20 You Try It. Do the Dissociation Equations worksheet.
21 Covalent Bonding covalent bond when electrons are shared between two atoms the electronegativity difference between the two atoms is less than 1.9 usually two nonmetals NO ions formed! (no electrons are taken just shared)
22 When a nonmetal and another nonmetal atom are around each other there is the opportunity for. the sharing of electrons producing molecules in which the atoms like to cling to each other. Covalent bonding!!!
23 The formation of a bond between two nonmetal atoms. Atoms sufficiently far apart to have no interaction
24 Covalent Compounds Also called molecular compounds solid, liquid, or gas at room temperature low melting points molecular formula represents the actual ratio of atoms that combine to form a neutral compound when dissolved in water, the molecular compounds DO NOT break up into ions (NO dissociation) the solutions of molecular compounds DO NOT conduct electricity (nonelectrolytes)
25 Pure Covalent The two fundamental types of bonds. Ionic
26 There is Pure Nonpolar Covalent another type of bond, not purely Polar Covalent covalent and not purely ionic. Ionic
27 Sharing of Electrons How would you know if an electron is going to be taken by one atom from another? Is there ever a time in which the electron is not taken but shared? Is the electron always shared equally?
28 Electronegativity electronegativity a measure of the ability of an atom in a chemical compound to attract electrons from another atom in the compound The difference in electronegativity values for two atoms will indicate whether the two atoms form an ionic bond (e - taken) or a polar or nonpolar covalent bond (e - shared). How is this different from electron affinity?
29 Electronegativity Differences 0.0 to 0.4 nonpolar covalent 0.5 to 1.9 polar covalent 2.0 and up ionic These ranges are flexible, although the general rule is a metal and nonmetal will form an ionic bond and two nonmetals will form a covalent bond. (Learn these values!)
30 Coordinate Covalent Bonds A coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons. The shared electron pair comes from one of the bonding atoms. Once formed, a coordinate covalent bond is like any other covalent bond. In a structural formula, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them. For example, Carbon monoxide (CO) is an example of a type of covalent bonding different from that seen in water, ammonia, methane, and carbon dioxide. A carbon atom needs to gain four electrons to attain the electron configuration of neon. An oxygen atom needs two electrons. Yet it is possible for both atoms to achieve noble-gas electron configurations by a type of bonding called coordinate covalent bonding. To see how, begin by looking at the double covalent bond between carbon and oxygen. With the double bond in place, the oxygen has a stable configuration but the carbon does not. As shown below, the dilemma is solved if the oxygen also donates one of its unshared pairs of electrons for bonding.
31 The structural formula of carbon monoxide, with two covalent bonds and one coordinate covalent bond is shown below.
32 Ionic, Polar Covalent, or Nonpolar Covalent? What kind of bond would each pair form? 1. N and S 2. S and C 3. Mg and Cl 4. C and F 5. Ba and O Which one of these bonds has the least ionic character?
33 Valence Electrons valence electrons the electrons in the highest energy level Na: 1s 2 2s 2 2p 6 3s 1-1 valence e - O:? Ne:? Al:? He:?
34 Octet Rule octet rule most atoms will gain or lose electrons to have 8 valence electrons (e - in the highest energy level) Exceptions: H, He, Li, Be, B, and some atoms P and higher on the periodic table Why is an atom like Ca more stable once it becomes an ion? How many valence electrons would calcium have to lose to have 8?
35 Double and Triple Bonds Sometimes atoms bond by sharing more than one pair of electrons. Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. A bond that involves two shared pairs of electrons is a double covalent bond. A bond formed by sharing three pairs of electrons is a triple covalent bond as shown below.
36 Resonance A resonance structure is a structure that occurs when it is possible to draw two or more valid electron dot structures that have the same number of electron pairs for a molecule or ion. Resonance structures are simply a way to envision the bonding in certain molecules. For example, the ozone molecule has two possible electron dot structures. Notice that the structure on the left can be converted to the one on the right by shifting electron pairs without changing the positions of the oxygen atoms. As drawn, these electron dot structures suggest that the bonding in ozone consists of one single coordinate covalent bond and one double covalent bond. Because earlier chemists imagined that the electron pairs rapidly flip back and forth, or resonate, between the different electron dot structures, they used double-headed arrows to indicate that two or more structures are in resonance.
37 Metallic Bonds Metals are made up of closely packed cations rather than neutral atoms. The valence electrons of metal atoms can be modeled as a sea of electrons. That is, the valence electrons are mobile and can drift freely from one part of the metal to another. Metallic bonds consist of the attraction of the free-floating valence electrons for the positively charged metal ions. These bonds are the forces of attraction that hold metals together. The sea-of-electrons model explains many physical properties of metals. For example, metals are good conductors of electrical current because electrons can flow freely in them. As electrons enter one end of a bar of metal, an equal number leave the other end. Metals are ductile meaning they can be drawn into wires. Metals are also malleable, which means that they can be hammered or forced into shapes.
38 Both the ductility and malleability of metals can be explained in terms of the mobility of valence electrons. A sea of drifting valence electrons insulates the metal cations from one another. When a metal is subjected to pressure, the metal cations easily slide past one another like ball bearings immersed in oil. In contrast, if an ionic crystal is struck with a hammer, the blow tends to push the positive ions close together. They repel, and the crystal shatters as shown in Figure 7.12.
39 Alloys Although every day you use metallic items, such as spoons, very few of these objects are pure metals. Instead, most of the metals you encounter are alloys. Alloys are mixtures composed of two or more elements, at least one of which is a metal. Brass, for example, is an alloy of copper and zinc. Alloys are important because their properties are often superior to those of their component elements. Sterling silver (92.5% silver and 7.5% copper) is harder and more durable than pure silver but still soft enough to be made into jewelry and tableware. Bronze is an alloy generally containing seven parts of copper to one part of tin. Bronze is harder than copper and more easily cast. Nonferrous (non-iron) alloys, such as bronze, copper-nickel, and aluminum alloys, are commonly used to make coins. The most important alloys today are steels. The principal elements in most steel, in addition to iron and carbon, are boron, chromium, manganese, molybdenum, nickel, tungsten, and vanadium. Steels have a wide range of useful properties, such as corrosion resistance, ductility, hardness, and toughness. Alloys can form from their component atoms in different ways. If the atoms of the components in an alloy are about the same size, they can replace each other in the crystal. If the atomic sizes are quite different, the smaller atoms can fit into the interstices (spaces) between the larger atoms. In the various types of steel, for example, carbon atoms occupy the spaces between the iron atoms.
40 Summary Review Intramolecular force - is the force that holds together the atoms making up a molecule or compound. There are three main types of intramolecular force, distinguished by the type of bond and the behavior of the electrons. a. Ionic - forms between a metal and nonmetal, such as sodium and chlorine in NaCl. Electrons in an ionic bond tend to be found mostly around one of the two atoms due to the large electronegativity difference between the two atoms; this is often described as one atom giving electrons to the other. In the case of NaCl, sodium would give an electron to chlorine. b. Covalent - generally forms between two nonmetals. Electrons in a covalent bond are shared between the atoms. There are two types of covalent bonds. The polarity of a covalent bond is determined by the electronegativities of each atom. i. nonpolar covalent bonds - the electrons are evenly shared. ii. polar covalent bonds - the electrons are more likely to be found around one of the two atoms. A polar covalent bond usually creates a dipole moment. c. Metallic - generally forms within a pure metal or metal alloy. Metallic electrons are generally delocalized resulting in a large number of free electrons (an electron sea) around the positive nuclei.
41 Intermolecular forces Intermolecular forces - are forces of attraction or repulsion which act between neighboring particles (atoms, molecules or ions). They are weak compared to the intramolecular forces (ionic, covalent, metallic) which keep a molecule together. Dipole moment - the tendency of electrons to move from their free state positions to a localised position around a more electronegative atom. The three major intermolecular forces are: 1. Hydrogen bond - is the attraction between an electronegative atom and a hydrogen atom that is bonded to a highly electronegative atom, such as nitrogen (N), oxygen (O) or fluorine (F). Strongest of the intermolecular forces. 2. Dipole-dipole interactions - are attractive interactions between permanent dipoles in molecules. These interactions tend to align the molecules to increase attraction. Weaker than a hydrogen bond. 3. Dispersion Forces (London dispersion forces) - arise from the movement of electrons which cause an instantaneous dipole in nonpolar molecules that do not have a permanent dipole. These are the weakest intermolecular force. The strength of the force increases as the molar mass increases, corresponding to a higher boiling point as the size of the molecule increases.
42 Hydrogen Bonds Hydrogen bonds are attractive forces in which a hydrogen covalently bonded to a very electronegative atom is also weakly bonded to an unshared electron pair of another electronegative atom. This other atom may be in the same molecule or in a nearby molecule. Hydrogen bonding always involves hydrogen. It is the only chemically reactive element with valence electrons that are not shielded from the nucleus by other electrons. Remember for a hydrogen bond to form, a covalent bond must already exist between a hydrogen atom and a highly electronegative atom, such as oxygen, nitrogen, or fluorine. The combination of this strongly polar bond and the lack of shielding effect in a hydrogen atom is responsible for the relative strength of hydrogen bonds. A hydrogen bond has about 5% of the strength of an average covalent bond. Hydrogen bonds are the strongest of the intermolecular forces. They are extremely important in determining the properties of water and biological molecules such as proteins and enzymes.
43 For example, the dipole interactions in water produce an attraction between water molecules. Each O H bond in the water molecule is highly polar, and the oxygen acquires a slightly negative charge because of its greater electronegativity. The hydrogen atoms in a water molecule acquire a slightly positive charge. The positive region of one water molecule attracts the negative region of another water molecule, as illustrated in the figure below. This attraction between the hydrogen of one water molecule and the oxygen of another water molecule is strong compared to other dipole interactions. This relatively strong attraction, which is also found in hydrogen-containing molecules other than water, is called a hydrogen bond.
44 The strong attractions between water molecules cause the water to bead together into small drops rather than spread over the surface of the flower.
45 Van der Waals Forces The two weakest attractions between molecules are collectively called van der Waals forces, named after the Dutch chemist Johannes van der Waals ( ). Van der Waals forces consist of dipole-dipole interactions and dispersion forces. Dipole-dipole interactions occur when polar molecules are attracted to one another. The electrical attraction involved occurs between the oppositely charged regions of polar molecules, as shown in Figure The slightly negative region of a polar molecule is weakly attracted to the slightly positive region of another polar molecule. Dipole interactions are similar to but much weaker than ionic bonds. For example, a dipole-dipole interaction can be seen in hydrogen chloride (HCl) where the positive end of the molecule will attract the negative end of the other molecule.
46 Dispersion Forces (London dispersion forces) - arise from the movement of electrons which cause an instantaneous dipole in nonpolar molecules that do not have a permanent dipole. These are the weakest intermolecular force. The strength of the force increases as the molar mass increases, corresponding to a higher boiling point as the size of the molecule increases. For example, fluorine and chlorine are gases at room temperature, whereas bromine is a liquid and iodine is a solid at room temperature.
47 Network Solids Network Solid - is a substance made up of an array of repeating covalently bonded atoms. In a network solid there are no individual molecules and the entire crystal may be considered a macromolecule. For example, diamonds are network solids made of carbon atoms. Silicon dioxide or quartz has a continuous three dimensional network of SiO 2 units resulting in a network solid. Properties Electrical conductivity - Poor, as there are no delocalized electrons, unlike ionic compounds and metallic bonds. Melting point - High, due to the large amount of energy required to rearrange the covalent bonds. Hardness - Hard, due to the strong covalent bond throughout the lattice Solubility - Generally insoluble in any solvent due to the difficulty of solvating a very large molecule.
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