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Chapter 4. Chemists have found it convenient to represent elements, especially when discussing chemical bonding, using a system devised by G.N. Lewis,

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Presentation on theme: "Chapter 4. Chemists have found it convenient to represent elements, especially when discussing chemical bonding, using a system devised by G.N. Lewis,"— Presentation transcript:

1 Chapter 4

2 Chemists have found it convenient to represent elements, especially when discussing chemical bonding, using a system devised by G.N. Lewis, called Lewis Dot Symbols

3 Chemical Bonds – Force that holds atoms together in a molecule. Types of Chemical Bonds: 1.Ionic Bonds – Assume that every atom wants to have a filled valence level (2 electrons for the first and 8 for each of the others). Then each atom will try to add or lose electrons to achieve this. This is called the Octet Rule.

4 Loss or gain depends on which is easier. (The number of valence electrons in an atom is the # on top of the column (before the A or B). Except for H and He, if the # of valence electrons is less than 4, it will lose that number of electrons. If it is greater than 4, it will gain enough to make it eight. H either gains or loses 1 electron, depending on the other atom. More about this later. He does not do anything. It is chemically inert (it does not bond to any other atom)

5 A) Ion – Any charged atom (i.e., one that has gained or lost electrons B) Cation – A + charged ion C) Anion – A - charged ion

6 An atom cannot gain or lose electrons without other atoms near to accept or give these electrons. Hence when a cation is formed, an anion is also formed (or perhaps more than one of each type). These then attract each other (opposite charges attract) and form the force we call an Ionic Bond. Ionic bonds between atoms can only occur between a metal and a non-metal atom.

7 Covalent Bonds – Chemical Bonds in which 2 atoms share exactly 2 electrons. An example is F 2. Sometimes 2 atoms will share 4 electrons. This is 2 covalent bonds between the same atoms, always referred to as a double bond (the word covalent is not necessary, because there is no such thing as a double ionic bond). Sometimes 6 electrons are shared, which forms a triple bond. 6 is the maximum # of electrons that can be shared by 2 atoms. Covalent bonds occur between 2 non-metal atoms or sometimes between a metal and a non-metal. Two metal atoms never bond together. Why?

8 G.N. Lewis also developed symbols to illustrate chemical bonds. A dash is used to represent a single bond (i.e. 2 electrons) and dots to represent the remaining valence electrons not involved in bonding). ═ is used to represent a double bond and ≡ is used to represent a triple bond.

9 Two Types of Covalent Bonds: 1. Non-polar – When 2 identical atoms bond together, each has an equal attraction for the shared electrons, therefore there is no net gain or loss of electrical charge by either atom. One end of the bond is identical to the other.

10 2. Polar – Whenever 2 different atoms covalently bond together, one atom will have a stronger attraction for the shared electrons than the other, thus one side appears to gain some negative charge and the other side seems to lose negative charge (become positive). The 2 ends of the bond are different in charge and, similar to a magnet, we say that each end is a pole (in this case a positive and negative pole (like a battery), hence the name polar covalent bond.

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12 Sometimes molecules that have polar bonds are, as a whole non-polar compounds, because of symmetry. We won’t discuss how to identify these except in a couple of very important examples. Any compound in which all bonds are non-polar, will be non-polar as a molecule.

13 There is a special class of compounds that we will discuss in more detail later in the semester, called hydrocarbons. These are compounds that contain C and H and nothing else. All bonds are either non-polar C  C bonds or slightly polar C  H bonds. Nevertheless, all these compounds (and there are thousands of them) are non- polar, because of symmetry.

14 Compounds composed of only covalent bonds are considered to be molecular compounds and compounds composed of ionic bonds are considered to be ionic compounds. Ionic compounds can either be simple binary compounds composed of ionic bonds between atoms or more complicated molecules, where one or both ions is actually composed of 2 or more atoms covalently bonded but forming an ion. These are called polyatomic ions. Some of the more common polyatomic ions are listed in table 4.4 on page 107: 2 elements

15 04_T04.JPG

16 We will learn a few of the most common ( hydronium, ammonium, hydroxide, nitrate, sulfate, cyanide, carbonate, bicarbonate and phosphate). Note that all, except two, are anions.

17 A compound between 2 elements is called a binary compound. For compounds between a metal and non-metal, we can predict the formula of the compound that will form. We simply assume each will form the most likely ion (remember the Octet Rule). Then the value of the charge on the cation (without the sign) becomes the subscript for the anion in the compound formula, while the value of the charge on the anion becomes the subscript for the cation. We simply cross-exchange. If both subscripts can be divided evenly by the same number, we do that. Anytime the subscript is one, we don’t write it.

18 Na and Cl - Na has 1 valence electron which it loses and becomes Na +1 while Cl has 7 valence electrons, so it gains 1 and becomes Cl -1. The 1 on Na becomes the subscript for Cl while the 1 on Cl becomes the subscript for Na. The formula for the compound is NaCl (1 is never written). Let’s do some examples:

19 Ca and Cl – Ca has 2 valence electron which it loses and becomes Ca +2 while Cl has 7 valence electrons, so it gains 1 and becomes Cl -1 (gains 1 electron to get 8). The 2 on Ca becomes the subscript for Cl and the 1 on Cl becomes the subscript for Ca. Thus the formula for the compound is CaCl 2 1 2

20 Mg and SO 4 -2 – Mg has 2 valence electrons which it loses and becomes Mg +2. the sulfate already has a charge of –2. The 2 from Mg becomes the subscript for SO 4 -2 and the 2 from the sulfate ion becomes the subscript for Mg. Thus the formula for the compound becomes Mg 2 (SO 4 ) 2. Note that when there are more than one of a particular polyatomic ion, its formula is placed inside parentheses. But we are not through yet. The subscript on both the sulfate and the Mg can be evenly divided by 2, so we do so. In each case the result is 1. Thus the final formula is MgSO 4. NOTE: No parentheses.

21 Al and O Let’s try one more: 2 3

22 Naming Binary Compounds or ionic compounds with Polyatomic ions. Binary Compounds – Name the metal first followed by the non-metal name (dropping endings such as, “ygen”, ogen, “ur”, “ine”, “ic” or “orous” and replacing with “ide”. Examples – Sodium chloride, beryllium oxide, aluminum nitride.

23 Ionic compounds with polyatomic ions. Simply name the metal followed by the anion polyatomic name unchanged. If the compound contains the ammonium ion with a non-metal element, follow the non-metal rule above. Examples – sodium sulfate, magnesium carbonate, ammonium phosphate. NOTE: It makes no difference how many of each atom or ion is present in the compound, for these cases. For now, we will not worry about binary compounds involving transition metals.

24 When molecules are mixed together, they tend to attract each other, sometimes strongly and sometimes weakly. With polar and ionic compounds it is easy to understand. The positive end of one molecule attracts the negative end of its neighbor. These attractions are what cause solids and liquids. With ionic compounds this is very strong. Ionic compounds tend to be solid at room temperature and at high temperatures. The following is discussed on pages 156 and 159 in Chapter 6 of your book.

25 In gases, the attractions are extremely weak. The stronger the attractions, the more likely it is that the substance will be a solid at room temperature and the higher the melting and boiling points will be. With polar molecular substances, these attractions are much weaker (called Dipole Forces), than in ionic compounds, hence lower melting and boiling points, with one exception. We will return to this in a minute.

26 With non-polar compounds these attractions are extremely weak, (called Dispersion Forces or London Forces) hence many of these are gases or liquids at room temperature. Ionic Compounds  Ionic bonds & dispersion forces between molecules Polar Covalent Compounds  Dipole forces (attractions) & dispersion forces Non-polar Covalent Compounds  Dispersion forces only

27 The exception to the polar molecular situation arises whenever H is directly bonded covalently to N, O or F in the compound. Then the polarity is so strong that the intermolecular attraction is much stronger than would be expected. This is called Hydrogen Bonding.

28 © 2010 Pearson Prentice Hall, Inc. 6/28 Intermolecular Forces and the States of Matter Hydrogen bonds: When a hydrogen atom is covalently bonded to a highly electronegative atom like nitrogen, oxygen, or fluorine (N,O,F), it can exhibit an additional polar attraction. This attraction is called a hydrogen bond.

29 Much higher MP and BP than expected Solid is less dense than liquid (ice floats) DNA The best example is H 2 O. Very strong attractions:

30 Without these special properties, life on earth as we know it, could not exist. We will discuss this more at the end of the semester.

31 We have mentioned that substances can change state from solid to liquid (melting) or vice versa (freezing) or from liquid to gas (vaporization) or vice versa (condensation). These are phase changes. They are physical changes. One other physical change is worthy of mention. Sublimation occurs when a solid vaporizes directly without ever becoming a liquid. Best example is dry ice. The following can be found on page 154 in your text.


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