Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 1

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1 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 1 Bonding Unit 3: Chemistry I In this unit all students must be able to Understand that the structure of molecules is the result of nonmetals sharing electrons in order to form stable outer-energy-level configuration. This sharing of electrons creates what is called a covalent bond. One way to represent the sharing of electrons is by using the orbital notation system that you learned and used in the last unit. You can write the orbital notation for different elements and use those illustrations to show how the valence level orbitals overlap in covalent bonding. Examples: Hydrogen has 1 electron, with an orbital notation of: H: Oxygen has 8 electrons, with an orbital notation of: 2p Before oxygen can bond covalently, it must hybridize. In hybridization, the 2 kinds of valence level orbitals change into one kind of valence level orbital. The valence level is the highest primary electron energy level. This case, the highest primary energy level is 2. The hybridized electron configuration of oxygen looks like this: Unhybridized orbital notation for oxygen Unhybridized orbital notation for oxygen O: 2s O: O: 2s 2p 2 These unhybridized valence orbitals combine and form these 4 equal hybridized orbitals If 2 hydrogen atoms are combined with one oxygen atom, then each of the halffilled (single arrow) hydrogen electron orbitals are combined with each of the half-filled (single arrow) oxygen electron orbitals, creating a set of filled valence level orbitals: o Draw Lewis dot structures for simple molecules Lewis dot structures or Lewis dot formulas are drawings that show the arrangement of electrons in 2 dimensions. Molecules are formed when atoms bond together covalently, which means they bond in a way in which they share This illustration shows modified orbital notations for hydrogen and oxygen before they share electrons (arrows show where in the orbitals the electrons will be shared and will overlap). H: H: 2 O: This illustration shows a combine modified orbital notation for hydrogen and oxygen after they share electrons. Notice how the orbitals overlap. H 2 O: O: 2 Figure 3.1. The illustration above shows an orbital notation illustration of how sharing of electrons occur to form covalent bonds. :H :H

2 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 2 electrons. The sharing of electrons allows atoms to have enough electrons to fill their valence level. The valence level is the highest primary electron energy level of an atom. You can also think of the valence level of an atom as the outer-most primary energy level of an atoms electron cloud. Lewis dot structures, Lewis dot diagrams, Lewis dot formulas Lewis dot structures* show the shared pairs of electrons and also show the lone pairs (sometimes called unshared pairs) of electrons. A Lewis structure can be drawn for any covalently-bonded molecule, any polyatomic ion, and any coordination compound. We will only be learning about covalently-bonded molecules in this unit. Later, we will learn how Lewis dot structures can be also be used to illustrate ionic bonding. First, however, I will give you a method for creating these Lewis dot structures. *Note: These structures are also called formulas, diagrams, or drawings. Most of the time, atoms share electrons in a way that gives each atom 8 electrons in the valance level. Hydrogen (H) is an exception; it only wants to get 2 electrons (2 electrons will give hydrogen the noble gas electron configuration of helium). Remember that valence electrons are those in the highest primary level. You can figure out how many electrons there are in the highest primary level by drawing the orbital notation or writing the electron configurations. H: 1 H 1 or H or H or H:. Orbital notation for hydrogen Electron configuration for hydrogen Lewis dot formula for hydrogen (Note: It doesn t matter on which side you place the single electron) Hydrogen (symbol H) has only has 1 valence electron. This electron is represented by the one arrow in the orbital notation. This electron is also represented by the superscript 1 in the orbital notation. It is also represented by the single dot around the hydrogen symbol in the Lewis dot formula. O: 2s 2p Orbital notation for oxygen O: 2 2s 2 2p 4 Electron configuration for oxygen Definitions: A shared pair of electrons is a) any circled any pairs of electrons between atoms in a sharing arrangement, b) and electrons in between atoms in a Lewis dot formula, or c) any lines representing pairs of electrons in a Lewis structure or formula. A lone pair of electrons is any pair of electrons in the valence level that is NOT involved in bonding. In a Lewis dot formula, or Lewis structure or formula lone pairs are pairs of dots that are placed around an atomic symbol but are NOT between 2 atomic symbols. O H Lewis dot formula for oxygen Note: It doesn t matter on which side the lone electrons or the electron pairs in the Lewis dot formula are placed. So, the Lewis dot formula above is just of 4 possible arrangements.

3 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 3 Oxygen (symbol O) has only has 6 valence electrons. These electrons are represented by the 2 arrows above 2s and the 4 arrows above 2p in the orbital notation. These electrons are also represented by the superscript 2 in 2s 2 and the superscript 4 in 2p 4 in the orbital notation. It is also represented by the 6 dots around the hydrogen symbol in the Lewis dot formula. Recall that you put one dot on each of the 4 sides of the imaginary square around the oxygen symbol and then pair up dots if necessary until you reach the required number of dots. Now then, let s put these 2 Lewis dot formulas next to each other and see how we can share electrons to give hydrogen the 2 electrons (or dots) that it needs to have a full valence level and give oxygen the 8 electrons (or dots) that it needs to have a full valence level. You should write this down with me as I draw it. H Let s be clear about a couple of terms. A lone pair is a pair of electrons drawn on the same side of the atomic symbol. Hydrogen does not have a lone pair. Oxygen has 2 lone pairs. Oxygen also has 2 lone or single electrons. Now, let s draw a circle around the lone electron on hydrogen and one of the lone electrons around oxygen. This is how we predict the covalent bonding of hydrogen and oxygen. Hydrogen now has 2 electrons and that s all that it wants. Oxygen on the other hand now has 7 electrons. The problem is that oxygen wants 8 electrons. To get another electron, we need another hydrogen atom, like this: See how the 2 nd hydrogen has shared its electron? Now, oxygen has 8 electrons. This drawing is now called the Sharing arrangement for this molecule, but it s not yet the Lewis dot structure for the molecule. Now sharing arrangement is not some special scientific term. It s just my description used so that you and I both know of what is being spoken. We re ready to write the Lewis dot structure. See figure 3.1. All we do is to squeeze the dots we drew circles around between the atomic symbols. O H H O O Figure 3.1. This is an illustration of the Lewis dot structure for water. The dots H O represent electrons. Between each of the 2 hydrogen atoms and the oxygen there are 2 shared electrons or dots (point to these H dots that are between the H and the O). In addition, there are dots that represent electrons that are not shared. These unshared pairs of electrons are also called lone pairs (point to these pairs of dots on the top and on the right side of O). Notice that this arrangement gives oxygen 8 electrons and the 2 hydrogen atom s 2 electrons. This arrangement stabilizes all of these atoms. OK, let s try another one. This time let s write the Lewis dot formulas for an atom of hydrogen and a Lewis dot formula for an atom on chlorine. First, here is the orbital notation, the electron configuration, and the electron dot notation for chlorine. H

4 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 4 Cl: 2s 2p 3s 3p Cl: 2 2s 2 2p 6 3s 2 3p 5 Cl Here is an illustration of how carbon and chlorine combine to make a combination of carbon and chlorine to make carbon tetrachloride (CCl 4 ): Orbital notation for chlorine Electron configuration for chlorine Lewis dot formula for chlorine So, the correct electron dot notation for hydrogen and chlorine is below. H Cl Now I will write the correct sharing arrangement and you can check your work. H Cl Now, let s write a Lewis dot formula for the covalent bonding of these 2 electron dot notations for these 2 atoms. Here is the correct Lewis dot formula for the covalent bonding of these 2 elements and you can check your work. H Cl Now, we complete this whole process for a combination of H N H nitrogen and oxygen. The H N H correct sharing arrangement and Lewis dot formula for the covalent bonding of these 2 H elements is seen below. H Cl Cl C Cl Cl Cl Cl C Cl Cl

5 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 5 Lewis structures, Lewis diagrams, Lewis formulas In Lewis structures the valence electrons that are lone pairs are represented as dots and lines are used to represent shared pairs in a single, double, or triple bond. Lewis structures show each atom and its position in the structure of the molecule using its chemical symbol. Lines are drawn between atoms that are bonded to one another. One line represents 2 electrons that are shared in a bond. Excess electrons that form lone pairs are represented as pairs of dots, and are placed next to the atoms. Students should be able to determine the number of valence electrons. In the previous unit, students also learned how to determine the number of valance electrons from the orbital configuration and the electron configuration of the element. Students should also know that for the representative elements ONLY the number of valence electrons can be determined from the Roman numeral at the top of each column in which the element appears in the periodic table. Single bonds: H O Figure 3.2. An illustration of the Lewis structure or Lewis formula for water (compare to figure 3.2). By definition a single bond is a pair of electrons shared between the nuclei of two atoms bonded together. In a Lewis dot formula, a single bond is represented by a pair of electron dots between the atomic symbols for the bonded atoms. In a Lewis formula, a single bond is represented by a single line between the atomic symbols for the bonded atoms. Double bonds: By definition a double bond is two pairs of electrons shared between the nuclei of two atoms bonded together. In a Lewis dot formula, a double bond is represented by two pairs of electron dots between the atomic symbols for the bonded atoms. In a Lewis formula, a double bond is represented by a two parallel lines between the atomic symbols for the bonded atoms. Triple bonds: By definition a double bond is three pairs of electrons shared between the nuclei of two atoms bonded together. In a Lewis dot formula, a triple bond is represented by three pairs of electron dots between the atomic symbols for the bonded atoms. In a Lewis formula, a triple bond is represented by a three parallel lines between the atomic symbols for the bonded atoms. H Lines represent 2 electrons that are shared between 2 atoms (in this case hydrogen and oxygen) and dots represent in shared electrons (in this case the unshared electrons are in pairs so that are also called lone pairs ).

6 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 6 o Understand that crystalline structure is the result of the ionic bonding of positive ions and negative ions, forming a neutral compound. o Understand that ionic bonding of positive ions and negative ions forms a neutral compound. The sum of the oxidation numbers (or charge numbers) in the formula of any neutral compound is zero Ions are positively or negatively charged, atomic or molecular-sized particles. Ionic crystals Ions or ionic particles arrange themselves in such a way that positive charges are as close to negative charges as possible. The result is a regular, repeating structure called a lattice. When large enough for us to see, these ionic lattices are known as ionic crystals. Even though ionic compounds are made of charged particles, the positive and negative charges equal out so that the overall charge is 0 or neutral. Understand that metallic atoms can form positive monatomic ions by losing electrons in order to achieve a stable outer energy level electron structure. Metal ions typically have a small number of electrons in their outer or valence energy level. Representative element metals most often lose all their valence electrons, leaving the energy level that was below the valence level as the new outer level. This level has full s and p sublevels. The resulting ion has an electron configuration that is more stable than the atom that previously existed. Creating charged particles, however, creates a new kind of instability due to the charges that are formed. A charged particle is unstable because that charge seeks to establish equilibrium through neutralization. This charge instability problem can be solved by combining opposite charges in a ratio that neutralizes the charges. Note: Creating a charged particle, one that has an unequal number of protons and electrons, is called ionization Figure 3.3. Illustration of a small section of a crystal lattice. Note that most crystal lattices are more complex than this. Figure 3.4. This illustration of a periodic table outline shows the location of metals, non-metals and metalloids. Metalloids are elements that sometime behave like metals and sometimes behave like non-metals. Recall that electrons have a negative charge. Each time that an electron is removed from an atom or molecular structure the atom or structure is left with a + charge. This is because an atom has a neutral charge from having an equal number of positively charged protons and negatively charged electrons. Removing 1 electron creates a + charge, removing 2 electrons creates 2+ charge, removing 3 creates a 3+ charge, and so on. Taking electrons from an atom creates a

7 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 7 monoatomic ion and taking electrons from a molecule creates a polyatomic ion. Positive ions are called cations. Likewise, each time that an electron is added to an atom or molecular structure the electron brings with it a charge to that atom or molecular structure. Adding electrons to an atom creates a monoatomic ion and adding more electrons than protons to what would otherwise be a molecule creates a polyatomic ion. Negative ions are called anions. In other words, a monoatomic ion is a single atom with too many or too few electrons to have a neutral charge. Examples include: Na +, Mg 2+, Al 3+, P 3. O 2, and Cl all of which can be found on your periodic table of oxidation numbers in your test references. A polyatomic ion is a covalently bonded group of atoms with too many or too few electrons to have a neutral charge. Examples include: IO 2, IO 3, NO 3, NO 2, BrO 4, ClO 4, IO 4, BrO 3, BrO 2, ClO 3, CO 2 3, PO 3 4, and PO 3 3 all of which are found in the Common ions chart in your test references. Positively charged ions are called cations. Examples include: Na +, Mg 2+, Al 3+, and NH + 4. Negatively charged ions are called anions. Examples include: IO 2, IO 3, NO 3, NO 2, BrO 4, ClO 4, IO 4, BrO 3, BrO 2, ClO 3, CO 2 3, PO 3 4, and PO 3 3 aio 2, IO 3, NO 3, NO 2, BrO 4, ClO 4, IO 4, BrO 3, BrO 2, ClO 3, CO 2 3, PO 3 4, PO 3 3, P 3, O 2, and Cl. Sometimes a metal will become more stable by simply loosing all of its valence p-sublevel electrons, leaving a full and more stable s sublevel in the valence level. These electrons don t just disappear. They go somewhere. When thinking about ionic bonding between metal and non-metal atoms, it s helpful to think about the electrons lost by the metal being gained by the non-metal. See the next item. o Understand that nonmetal atoms can form negative monatomic ions by gaining electrons in order to achieve a stable outer energy level electron structure. Non-metal ions typically have a large number of electrons in their outer or valence energy level. Non-metals often gain enough extra electrons to fill their valence level (a full valence level for most atoms has 8 electrons). This valence level, with the extra electrons needed to reach 8 in the valence level, will also have full s and p sublevels and full orbitals. All these create additional stability. The resulting ion, therefore, is more stable than the atom that previously existed. When you write the electron configuration of calcium after it has lost 2 electrons, you MUST include the 2+ charge as a superscript on the calcium symbol (Ca 2+ ). Likewise, when you write the electron configuration of sulfur after it has gained 2 electrons, you MUST include the 2 charge as a superscript on the sulfur symbol (S 2 ).

8 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 8 Example: Using orbital notation, here is a way to think about it. The electrons that calcium wants to give up go to sulfur like this: Ca: S: S 2 : 2s 2s 2p 2p 3s 3s 3p 3p Using Lewis dot notation, here is still another way to think about it. Once again the electrons that calcium wants to give up go to sulfur like this: And because electrons have a negative charge the result is: Note the brackets around the sulfide ion. Because the 2 charge is outside the brackets, it communicates to you and me that the charge is shared by the entire structure. Sometimes we don t write it that way (we write it S 2 ) and we expect the reader to understand that the charge is shared over the whole structure. In the case of calcium and sulfur, calcium and sulfur had just the right number of electrons for sulfur. In the case of some other sets of atoms (which become ions) you may need 2, 3, 4 or any other number of metal or non-metal ions to get everything worked out to the most stable form. In all the cases above, the resulting formula unit of calcium sulfide is CaS. Notice when we put the 2 ions together there are no charges. That s because the 2+ charge is canceled out by the 2 charge and the resulting calcium sulfide compound is neutral. 4s And because electrons have a negative charge the result is: Ca 2+ : 2s 2p 3s 3p 2s 2p. C. 3s 3p.. S.. :.. Ca 2+ 2 : S.. :

9 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 9 Unhybridzed p x orbital hybridizes into an hybrid orbital P y p z p x Unhybridzed p z orbital hybridizes into an hybrid orbital p x p z p y Unhybridzed s orbital hybridizes into an hybrid orbital Unhybridzed p y orbital hybridizes into an hybrid orbital Figure 3.5. Recall that s and p orbitals that are not hybridized have the arrangement in the illustration above left. In the valence electron energy level, when electrons with this arrangement of orbitals bond, the orbitals hybridize. If the result of bonding creates a total of 4 single bonds plus lone pairs (# of lone electron pairs) + (# of single bonds) = 4 hybridized orbitals then the kind of hybridization that is illustrated on the right occurs. This kind of hybridization is called. It is called because the 4 orbitals stem from one s orbital and three p orbitals. Figure 3.21 and figure 3.22 illustrate the 2 other kinds of hybridization that you need to understand in a high school chemistry class.

10 Unit 3 Ray Tedder s Chemistry I Test Prep Guide page 10 P y p z p x s p x p z p y The unhybridzed 3s orbital below corresponds to the spherical s orbital illustration above The unhybridzed 3p orbital below corresponds to any of the double balloon shaped p orbital illustrations above The hybridzed 3 orbitals below correspond to any of the single balloon shaped orbital illustrations above Cl: 2s 2p 3s 3p Cl: 2s 2p 3 Unhybridized orbital notation for chlorine Unhybridized orbital notation for chlorine Figure 3.6. Recall that orbital notation is a bookkeeping system and does not show the actual shape of orbitals. The illustrations above show the relationships between orbital notations and the actual shape of valence level orbitals. Valence electrons: Valence electrons are those electrons in the highest primary energy level of an atoms electron cloud. In the illustrations above, the valence electrons are in the 3 rd primary level of chlorine. Valence electrons include those in the s and p sublevels in that highest level, so chlorine has 7 valence electrons.