Noble gases do not join other atoms to form compounds. They seem to be most stable just as they are.

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1 UNIT 3: TE NATURE MATTER: MLECULES There are fewer than one hundred naturally occurring elements on the earth, but there are billions of compounds made of those elements. In this unit, we will examine how atoms join together to form molecules. Ionization energy and bonding A neutral atom has the same number of protons and electrons, but an atom may gain or lose one or more electrons to form an ion. The energy required to remove one electron from a neutral atom is called the first ionization energy (IE 1 ). Metal atoms have low first ionization energies and tend to lose electrons easily; nonmetal atoms have higher first ionization energies and do not lose electrons easily (nonmetals tend to gain electrons). This is a fundamental difference between metals and nonmetals, and explains much about their bonding behavior. In general, IE 1 tends to increase as you go across a period (less tendency to lost electrons) and decrease as you go down a group or family (greater tendency to lose electrons). Valence electrons and bonding ow do atoms join together to make molecules? ne clue to how atoms bond is their ionization energies. The first ionization energy (IE 1 ) is the energy to remove one electron from a neutral atom; by adding more energy, we can remove more electrons. The energy needed to do this is shown at right for Al. As you might guess, removing IE kj more electrons takes more energy, but look at the dramatic increase in energy to remove IE the fourth electron! A neutral aluminum atom has 13 electrons. After removing one kj IE electron (IE 1 ), it has 12 left; removing another (IE 2 ) leaves 11, and removing a third kj IE (IE 3 ) leaves 10 electrons just like neon, a noble gas. nce the aluminum atom has 4 11,600 kj achieved the same number of electrons as a noble gas, it becomes much more difficult to remove any more electrons. We see the same behavior for all main group elements: once they achieve the same number of electrons as a noble gas, the energy to remove any more electrons increases dramatically. Another clue: elements in a family share similar properties, including the kinds and formulas of the compounds they form. The diagram below illustrates three chemical families arranged a little differently than you are used to seeing them in the periodic table. Noble gases do not join other atoms to form compounds. They seem to be most stable just as they are. halogens noble gases alkali metals 2 e 3 Li 9 10 Ne 11 Na Ar 19 K 35 Br 36 Kr 37 Rb 53 I 54 Xe 55 Cs The halogens (,, Br, and I) each have one less electron than a noble gas element. These elements are extremely reactive, tending to acquire 1 electron to become 1 ions. When it gains one electron, a halogen atom achieves the same number of electrons as a noble gas. The elements in the alkali metal family (Li, Na, K, Rb, and Cs) each have one more electron that a noble gas element. They, too, are an extremely reactive family of elements, tending to lose 1 electron to become +1 ions. When it loses one electron, an alkali metal atom achieves the same number of electrons as a noble gas.

2 It seems that when an atom achieves the number of electrons identical to a noble gas element, it is particularly stable. That number of electrons is the key to understanding bonding in main-group elements. The two electrons of e form a stable, unreactive core called the helium core. The next element is Li. We can think of Li as having a e core plus one electron, for a total of three electrons. We continue across period 2 of the periodic table, adding one more electron on top of the e core for each element, until we get to the next noble gas, Ne, with 10 electrons. Ten electrons form the neon core. Element 11, Na, can be thought of as having a Ne core plus one electron, and as before we can continue adding electrons on top of the neon core across period 3. The electrons outside the noble gas core are called valence electrons, and they are the electrons that form bonds. or main-group elements, the number of valence electrons is given by the group number. In 1916 an American chemist, Gilbert Lewis, devised a simple way to represent the valence electrons. e used the symbol for the element to represent its nucleus plus its core electrons, then arranged dots around the symbol to represent the valence electrons. The Lewis dot symbol for lithium is Li with one dot. The dot may be placed anywhere around the Li. Li Nitrogen has 5 valence electrons. The Lewis dot symbol for nitrogen is an N with 5 dots. The Lewis dot symbol that most accurately represents bonding behavior is to put one dot on each N side around the symbol (top, right, bottom, left), then if there are more valence electrons continue around, pairing them, until they are all shown. or N, that ends up with two dots paired and three dots not paired. ctet rule The stability of the noble gases and the behavior of halogens and alkali metals suggest that elements form chemical bonds by losing, gaining, or sharing valence electrons in order to achieve the same number of electrons as the nearest noble gas in other words, to acquire 8 outer electrons (except for elements near e, which has just 2 electrons). This is called the octet rule. An atom of has 7 valence electrons and needs to gain just one more to achieve an octet of 8 outer electrons, like the noble gas Ne. An atom of with 6 valence electrons needs to gain 2 more electrons, and an atom of N needs to gain 3 more electrons. Nonmetals will either gain or share valence electrons to achieve an octet. A metal like Na has one valence electron and would need to gain 7 more electrons to achieve an octet like Ar. Gaining 7 electrons would be very difficult, so a better strategy for Na is to lose its valence electron, leaving only the Ne core with its complete octet. Metals lose valence electrons to achieve an octet. Thus, we predict that metals lose valence electrons when forming chemical bonds, and nonmetals gain or share electrons when forming chemical bonds. This is one of the fundamental differences between metals and nonmetals. Transition metals do not bond by the octet rule, so we will not study their bonding in this course. 2

3 ctet rule, ions, and ionic chemical bonds When a nonmetal forms a bond with a metal, a simple electron transfer is possible. The metal loses valence electrons to form a positive ion (a cation, pronounced cat-ion), and the nonmetal gains valence electrons to form a negative ion (an anion, pronounced an-ion). Sodium has a neon core plus 1 valence electron. To achieve the same number of electrons as a noble gas, it needs to lose that valence electron. Chlorine has 7 valence electrons. To achieve the same number of electrons as the nearest noble gas, argon, it needs to acquire one more valence electron. Na Na If sodium transfers its single valence electron to chlorine, Na ends up with a Ne core and its complete octet, and ends up with 8 outer electrons (like Ar). Both atoms end up with the same number of electrons as a noble gas, a more stable arrangement. After the electron transfer, however, the sodium and chlorine are no longer neutral atoms. Na has 11 protons but now has only 10 electrons, so it has become a +1 ion. has 17 protons and now has 18 electrons, so it has become a 1 ion. The oppositely charged ions are strongly attracted to each other, and will stay together in what is called an ionic bond. Compounds like Na, made of a metal and a non-metal joined in an ionic bond, are called salts. In a salt, the metal ion does not join exclusively with the nonmetal that received its electrons. Instead, the ions formed by electron transfer collect into a crystal, a regular pattern of positive and negative ions held together by the attraction of the opposite charges. Each ion is equally attracted to all other ions of opposite charge. If a salt melts or dissolves in water, its ions separate in the water and move freely, so the solution can conduct electricity. This diagrams shows the formation of the salt magnesium 2 2+ oxide. ere, two electrons are transferred from Mg to, Mg Mg giving each an octet. Magnesium s octet is its Ne core, left after losing its valence electrons, and oxygen s octet is completed when it accepts the electrons from Mg. Mg2+ and 2 ions form and stick together in an ionic crystal. 1 1 This diagram shows the formation of the salt 2+ Ca Ca calcium chloride. ere, Ca transfers its valence electrons to two chlorine atoms to form one Ca2+ ion and two 1 ions. The formula for the salt is Ca2, indicating a ratio of 1 Ca2+ to 2 1 in the crystal. A cation, having lost its outer electrons, is always smaller than its parent atom, and an anion, having gained one or more electrons, is always smaller than its parent atom. Thus Na1+ is smaller than Na, and 1 is larger than (the numbers in the illustrations are the radii, in picometers). 3

4 ctet rule and covalent chemical bonds The element chlorine is always found in the form of diatomic 2 molecules. Each chlorine atom has 7 valence electrons, and needs just one more electron to achieve an octet. Because both atoms want to acquire an electron and neither wants to give one up, the solution is to merge their outer valence levels and share a pair of electrons. Look at water, 2. Each hydrogen atom has one valence electron, and could achieve the same number of electrons as e if it acquired one more electron. xygen has 6 valence electrons, and could achieve an octet (like Ne) if it acquired two more electrons. The oxygen atom can share one unpaired electron with one hydrogen atom, and the other unpaired electron with the other hydrogen atom. The water molecule is held together by two single covalent bonds. The shared electrons are bond electrons, and the unshared valence electrons are called lone pairs. Each bond pair can also be shown with a line, so 2 is shown with two single lines for the two covalent bonds, plus two pairs of dots for the lone pairs on oxygen. Sharing of a pair of valence electrons is a covalent bond. When the chlorine atoms share the pair of electrons, each comes closer to the stability associated with the filled outer level of a noble gas. When the chlorine atoms bond together, they form a molecule: the two atoms that are sharing electrons stay together in a single neutral particle. Unlike a salt, a molecular substance is made of neutral molecules, not charged particles, so it does not conduct electricity in the liquid or solid state. Every dot structure you draw for a molecule should conform to the octet rule: the structure you draw must give every atom a filled valence level with eight dots around each atom (including all bond electrons), except which needs just two dots, and no unpaired electrons left over. Now for a more complicated example, 2. As you can see in this Lewis dot structure, the two oxygen atoms cannot achieve octets with a single bond. Each has only 7 dots around it, and still has an unpaired electron. or each in 2 to achieve octet, each oxygen atom must put both unpaired electrons into one bond. They share two pairs of electrons to form a double covalent bond. Each has eight dots around it: 2 lone pairs, and the two shared pairs of the double bond. The double bond is shown with two lines, indicating that the oxygen atoms are sharing two pairs of electrons. Compare this example to the one for water to be sure you understand the difference between a double bond and two single bonds. Can you see how a triple bond could form? 4

5 A strategy for writing covalent Lewis dot structures Pairing up unpaired electrons works for simple molecules, but the octet rule also describes the bonding in more complex molecules. ere is a simple strategy for drawing all covalently bonded dot structures: 1. Add up the total number of valence electrons in the molecule. If the molecule has a negative charge, add that number of electrons to the total. 2. Identify the central atom or atoms, usually the one(s) with the most unpaired electrons in its dot structure. is never a central atom, and C is nearly always the central atom. If two atoms are in the same family, choose the one lower down in the periodic table as the central atom 3. Arrange the other atoms around the central atom(s), as symmetrically as possible. The atoms on the edges of the structure are called terminal atoms. 4. Join the atoms with single bonds. Keep track of how many valence electrons have been used (remember, one single bond represents 2 shared valence electrons). 5. Complete the octets of the terminal atoms by adding lone pairs, then try to complete the octet(s) of the central atom(s). 6. If all the valence electrons are used up before you can complete the octet(s) of the central atom(s), move in one or more lone pairs from a terminal atom to form a double or triple bond. 7. Check your structure: did every atom achieve an octet (remember, that s 8 for every atom except, which needs only 2)? Count electrons again: were all valence electrons used, neither more nor less? Example 1. Draw the Lewis dot structure for P 3 P has 5 valence electrons and each has seven, for a total of 26 dots in the structure. A simple, symmetrical structure puts P at the center, with the s as terminal atoms. P Now we join the atoms with single bonds. Each bond uses 2 of our 26 dots. We have 20 dots left. We add dots to complete the octets on the s. This uses 18 dots; we have 2 dots left. P P P still does not have an octet. We put the last 2 dots on P. All the dots have been used, and every atom has a complete octet. The structure conforms to octet rule. P 5

6 Example 2. Draw the Lewis dot structure for C 2 C has 4 valence electrons, each has one, and the has 6, for a total of 12 dots in the structure. A simple, symmetrical structure puts C at the center, with an on each side and the above. We join them with single bonds, using up 6 dots. We have 6 dots left. Each is satisfied with the two dots in its bond. We add dots to complete the octet on the. This uses 6 dots so we have none left. C still does not have an octet, but we do not have any more dots. We move a lone pair from oxygen into a bond with carbon, making a double bond. Now every atom has a complete octet. The structure conforms to octet rule. C C C Example 3. Draw the Lewis dot structure for CN C has 4 valence electrons, has one, and the N has 5, for a total of 10 dots in C the structure. A simple structure puts C at the center, with the on one side N and the N on the other. We join them with single bonds, using up 4 dots. We have 6 dots left. The is satisfied with the two dots in its bond. We add dots to complete the octet on the N. This uses 6 dots so we have none left. C N C still does not have an octet, but we do not have any more dots. We move one lone pair from nitrogen into the bond with carbon, making a double bond but C still does not have an octet. To give C an octet, we move another lone pair from nitrogen into the bond with carbon, making a triple bond. Now every atom has a complete octet. The structure conforms to octet rule. C N C N Example 4. Draw the Lewis dot structure for 1 has 6 valence electrons, the has one, and we add one more for the 1 charge. That gives a total of 8 dots in the structure. We join the and with a single bond, using up 2 dots. We have 6 dots left. The is satisfied with the two dots in the bond. We add dots to complete the octet on the. This uses 6 dots so we have none left. Both atoms have a complete octet. The structure conforms to octet rule. We bracket the structure and give it a 1 charge overall. 1 6

7 Metallic bonding We have seen that metals bond with nonmetals by electron transfer, to form ions and ionic bonds, and nonmetals bond with other nonmetals by sharing electrons to form covalent bonds. Do metals form bonds with other metals? Metals tend to lose electrons. ne way to imagine how metals bond together to make a solid is to imagine that each metal donates its valence electrons to a common pool or sea of electrons that exists between the atoms. Each metal atom thus becomes a metal cation surrounded by the sea of mobile electrons. Each metal cation is strongly attracted to the negative electron sea, which is why when a metal is struck, the atoms slide past each other but do not separate from the electron sea. The electrons are free to move among the metal cations, which explains why metals conduct electricity even in the solid state. Because each metal cation is attracted to the electron sea, not to the other cations, metals do not form true compounds with other metals. In a true compound, the proportions of elements are constant (for example, one Na1+ and one 1 in Na, or two atoms and one atom in 2). Metals may be mixed in any proportions, and all the metals atoms will donate valence electrons to the electron sea and be held together by it. Metals form mixtures called alloys, instead of true compounds. Three types of chemical bonds You can see now that we classify chemical bonds into three broad categories, based on how the valence electrons are arranged. atoms metal + nonmetal bond ionic character e transferred nonmetal + nonmetal covalent e shared metal + metal metallic e pooled properties forms a salt = crystal of cations & anions brittle solid does not conduct electricity but conducts electricity when melted or dissolved in water composed of neutral molecules brittle does not conduct electricity in any state (solid, liquid, or dissolved in water) lattice of cations embedded in sea of electrons malleable solid and liquid both conduct electricity 7

8 To summarize this brief tour of bonding, 1. Noble gas elements are extraordinarily stable. That number of electrons in an atom forms a stable core. Electrons beyond the noble gas core are called valence electrons. 2. Elements form compounds by either transferring or sharing valence electrons to achieve the same number of electrons as the nearest noble gas (the octet rule). 3. An ionic bond forms when a metal gives up its valence electrons to form a positive ion (a cation), and a nonmetal accepts the electrons to form a negative ion (an anion). The oppositely charged ions are strongly attracted to each other and stay together to form a crystal lattice, called a salt. Because they are made of ions (not molecules), ionic compounds (salts) can conduct electricity when they are melted or dissolved in water, because under those conditions the ions separate and are free to move. 4. A covalent bond forms between two nonmetals, which share valence electrons. Including shared electrons, each element achieves an octet. Compounds in which the elements are joined by covalent bonds are molecular compounds. They are made of molecules, which do not conduct electricity when melted or dissolved in water. 5. A metallic bond forms when metals donate their valence electrons to a common pool (the sea of electrons ) to form a lattice of cations held together by their attraction to the sea of negative electrons. The mobile electrons allow metals to conduct electricity even in the solid state, and allow the cations to slide past each other, making the metal malleable. 6. Even though hydrogen is shown in the same group as the alkali metals, is not a metal and does not make ionic bonds like a metal. ydrogen will share its electron to achieve two outer electrons. 7. The octet rule does not predict all aspects of chemical bonds, so don t be alarmed when you run across strange-looking compounds like S 6, or find you can t predict the charges of the transition metals because they don t conform to octet rule. 8

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