CHEM 121 Introduction to Fundamental Chemistry. Summer Quarter 2008 SCCC. Lecture 5.

Transcription

1 CHEM 121 Introduction to Fundamental Chemistry Summer Quarter 2008 SCCC Lecture 5

2 Forces Between Particles Noble Gas Configurations Ionic Bonding Ionic Compounds Naming Binary Ionic Compounds The Smallest Unit of Ionic Compounds Covalent Bonding Polyatomic Ions

3 It has been recognized for a long time that the noble gases have great chemical stability. With few exceptions they are unreactive or inert. The noble gases have 8 valence electrons with the exception of He which has 2. He 1s 2 Ne 1s 2 2s 2 2p 6 Ar 1s 2 2s 2 2p 6 3s 2 3p 6 Kr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 Xe 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6

4 The electronic configuration of the noble gases is described as being energetically stable. We can draw a Lewis diagram to illustrate the number of valence electrons an atom has. In a Lewis diagram valence electrons are represented by dots placed above, below and to the left and right of the atoms symbol. e.g. element with 4 valence electrons E

5 There are two simple rules to keep in mind when drawing Lewis diagrams: Place one dot in each of the four locations before doubling up. There can be only a maximum of 2 dots in any one location. E E E E

6 What is the Lewis diagram for H? 1. First write the electron configuration: 1s 1 2. Identify the number of valence electrons. 1 valence electron. H For a representative element it is easy to identify the number of valence electrons as this is equal to the group number.

7 What is the Lewis diagram for S? 1. First write the electron configuration: [Ne]3s 2 3p 4 2. Identify the number of valence electrons. 6 valence electrons S Alternatively you can recognize that S is in group VIA so has six valence electrons

8 The octet rule states that: Atoms interact in order to obtain a stable octet of eight valence electrons The octet rule works extremely well at describing the interactions of the representative elements.

9 One way in which atoms can interact to satisfy the octet rule is by transferring electrons between each other. Transferring of electrons results in the atoms acquiring net positive and negative charges. When an atom loses or gains electrons a simple ion is formed.

10 Consider a Na atom what happens if it loses one electron? Na Na + + 1e - [Ne]3s 1 Consider a Cl atom would you expect it to lose or gain electrons? Cl [Ne]3s 2 3p 5 + 1e- Cl -

11 Metals tend to lose electrons forming positively charged ions called cations. A representative metal will lose its group number of electrons to obtain a stable octet. Na Na + + 1e - Mg Mg e - What would the charge be of the ion formed by a Li atom? +1 The ion formed would be Li +

12 Non-metals tend to gain electrons forming negatively charged ions called anions. A representative non-metal will gain (8 - group number) electrons to obtain a stable octet. O + 2e - O 2- S + 2e - S 2- What would the charge be of the ion formed by a I atom? -1 The ion formed would be I -

13 Ions do not form in isolation. They form when electrons are transferred from a metal to a non-metal, that is they are always formed in pairs. M + nx M y+ + nx (y/n)- The ratio of the number of ions formed is always such that the sum of the charges adds up to zero. e.g. (y+) + (n(y/n)-) = (y+) + (y-) = 0

14 The anions and cations formed when a metal and non-metal interact will be attracted to each other. M y+ + nx (y/n)- MX n This attraction is called an ionic bond. A compound containing one type of cation and one type of anion is called a binary ionic compound. In the formula of ionic compounds we normally write the metal atom first.

15 We write a simple ion by first writing the element symbol followed by the charge of the ion as a superscript. e.g. Li +, Mg 2+, Al 3+, Fe 2+, Cu 2+, Cl -, S 2-, O 2-,etc Cations are named by giving the name of the parent metal atom and adding the word ion. e.g. Li + lithium ion, Mg 2+ magnesium ion, Al 3+ aluminum ion, etc

16 Some metal atoms from the d block and groups IIIA-VA can form more than one ion. e.g. Element Iron Chromium Cobalt Copper Common Ionic Forms Fe 2+, Fe 3+ Cr 2+, Cr 3+, Cr 6+ Co 2+, Co 3+ Cu +, Cu 2+ When we write the names of these ions we need to specify the charge by putting it in Roman numerals after the element name. Cr 2+ Chromium(II) ion, Co 3+ Cobalt(III) ion, etc

17 Anions are named by adding the suffix ide to the stem name of the parent non-metal atom and adding the word ion. The stem name is typically the first syllable of the atom name (see table 4.2). e.g. Br - bromide ion, Cl - chloride ion, O 2- oxide ion, S 2- sulfide ion and P 3- phosphide ion.

18 Binary ionic compounds are named by giving the name of the cation first followed by the name of the anion and dropping all occurrences of the word ion. e.g. FeO Iron(II) oxide, NaCl Sodium chloride What is the name of the binary ionic compound formed between magnesium and chlorine? What is its formula? Magnesium chloride MgCl 2

19 In our early lectures we defined a molecule as the limit of physical subdivision. Molecules exist as particles containing the number of atoms specified by their formula. e.g. a water molecule is a particle containing 2 hydrogen atoms and one oxygen atom and has the formula H 2 O.

20 The molecular weight (mass of 1 mole) of a water molecule is equal to the sum of the atomic weight of oxygen plus two times the atomic weight of hydrogen. M = x = gmol -1

21 Ionic compounds do not exist as discrete molecules. Instead they exist as crystals where ions of opposite charges occupy positions known as lattice sites. Ions combine in the ratio that results in zero charge to form ionic compounds. NaCl

22 A quantity called the formula weight may be obtained by adding the atomic weights of the ions in the formula of an ionic compound. e.g. The formula weight of MgCl 2 is FW = x35.45 = gmol -1 Formula weights are used in a similar way to molecular weights.

23 Non-metals may also complete their octets by sharing electrons. This may occur between non-metal atoms of the same type: e.g. H 2, O 2, N 2, Cl 2, F 2, I 2, etc Or between different types of non-metal atoms: e.g. CO 2, H 2 O, CH 4, etc

24 Consider two hydrogen atoms separated by a large distance. Each has 1 electron in a 1s atomic orbital Now lets bring the two atoms together so there orbitals overlap.

25 The atomic orbitals overlap to form a new molecular orbital. This is a stable configuration as each H atom can have a full 1s susbshell (like He) where the electrons spend most of their time shared between the atoms. In this arrangement each nucleus feels an inwards attraction to the two electrons. This is called covalent bonding.

26 We can draw Lewis diagrams showing the arrangement of valence electrons in covalent compounds. In these diagrams we represent each pair of electrons between atoms as a line. So for the H 2 molecule discussed previously the Lewis diagram would be: H H All other electrons are represented by dots as described previously.

27 I use a five step method to draw the Lewis diagram for a covalent molecule formed between non-metals. 1. Add up the valence electrons and write this number down. e.g. For CH 4 C has 4 valence electrons Each H has 1 valence electron 8 valence electrons total.

28 2. Write down the central atom this is the first atom in the formula. e.g. for CH 4 C

29 3. Connect each of the outer atoms to the central atom with single lines (each one electron pair). Subtract the number of electrons used from the total number of electrons. e.g. for CH 4 H H C H H 8-(4x2)=0

30 4. If there are any electrons left over add electrons as lone pairs to the outer atoms until their octets are complete. Subtract these electrons from the total. 5. If there are still electrons left over add these as lone pairs to the central atom. If the octet of the central atom is not complete then we need to add electrons from the outer atoms to form double and triple bonds until its octet is complete.

31 What is the Lewis structure of NH 3? N 5e - 3H 3e - H N H 8e - -6e - 2e - -2e - 0e - H

32 What is the Lewis structure of SO 3? N 6e - 3O 18e - O S O 24e - -6e - 18e e - 0e - O

33 There are actually three possible Lewis structures for SO 3. O S O O S O O S O O O O Each of these three structures is equivalent. We say they are in resonance or that they are resonance structures.

34 We have previously looked at simple ions that are formed when an atom loses or gains electrons: e.g. Cl -, I -, Mg 2+, Cu 2+, Na +, O 2-, P 3- etc Two or more atoms covalently bonded to each other and having an overall charge is called a polyatomic ion. e.g. OH -, CN -, NH 4 +

35 We can determine the Lewis structure of a polyatomic ion in a similar way as for a molecule. We follow the same procedure for drawing the Lewis structure of a molecule except when determining the number of valence electrons (step 1). 1. If we have a polyatomic anion we add electrons equal to the charge. 2. If we have a polyatomic cation we subtract electrons equal to the charge.

36 What is the Lewis structure of NH 4+? N 5e - 4H 4e - H e - 8e - H N H -8e - 0e - H

37 What is the Lewis structure of CN -? N 5e - C 4e - - C N 10e - - 2e - - 1e - 8e - -8e - 0e -

38 What is the Lewis structure of CO 3 2-? C 4e - 3O 18e e - O C O 24e - - 6e - O 18e - -18e - 0e -

39 So far we have been considering how electrons are distributed between atoms in molecules and polyatomic ions. An important question is: How can we predict the shape of molecules and polyatomic ions? Can anyone think why we might want to do this?