Chemistry lesson 3. Study guide - Questions

Transcription

1 Chemistry lesson 3 Study guide - Questions This lesson builds on the understanding of the atom and electrons into the concepts of the octet rule, Lewis dot structures and chemical bonding. This is where the chemical rubber meets the road. Your weekly program. Develop your own method by all means but this will get you going. Read the notes attached here. All questions and the final assessment are based on what is in the notes. Read the chapter sections in the reading section ( the text has much greater detail than what you are expected to know) especially look at diagrams and figures At this stage not much will make sense but that is OK With your text as reference and this study sheet go through the Power point presentation. Make notes where needed. MANDATORY watch all of the embedded video links (you will need internet access to do this) Optional - Listen to the audio file of a live lesson; be aware that there will be long pauses with not much going on at certain points. Refine your notes / mind maps on the key concepts outlined in these weekly study sheets. Check your understanding Now go through your homework questions and answer those Your study resources Textbook Audio files of the actual lessons Weekly study sheet Power Point slides YouTube links and YouTube as a general resource College forum site Molecule kit 1

2 Textbook reference Bettleheim Edition 9 Please read these sections and note take Chapter 3 sections 3.1 to 3.8. Read 3.10 and Key Concepts to understand Octet rule Bond types Ionic & Covalent Ions Anions and Cations Lewis dot structures for compounds and drawing them Electronegativity Polyatomic ions Ionic and covalent compounds Polar and non Polar covalent bonds Valence shell electron pair repulsion model (VSEPR) Use the links in the power point presentation. Good sources for chemistry videos are Bozeman science Crash course chemistry YouTube Socratic.org - 2

3 Study Notes When chemicals bond they are exchanging / sharing electrons. Specifically this involves valence shell electrons. All of us are familiar with shopping or making a deal. It turns out that when atoms form chemical bonds to each other, they are doing exactly that, making a deal with each other to benefit themselves. Octet Rule Ask yourself, what makes atoms happy? It's really simple - atoms are happiest when they have 8 electrons in their valence shell. This rule is known as the octet rule. If the atom's valence shell happens to be the first shell, as in hydrogen or helium, 2 electrons in that shell would make them happy. As we saw earlier the first shell only holds 2 electrons. The octet rule mainly holds true for the first four periods of element in the periodic table. Covalent Bond A covalent bond is a type of chemical bond in which two atoms share a pair of electrons. Organic chemistry which will be studied in modules 2 & 3 is all based on covalent bonding. Organic chemistry is the chemistry of carbon molecules. Analogy Picture this scenario: Two fluorine atoms are at a party. They each have 7 electrons in their valence shell (we know this because they are in group 17) They're both unhappy atoms because they don't have the 8 valence electrons. But when they run into each other, they decide to make a deal that will benefit both of them. 3

4 One fluorine atom says to the other, "I'll share one of my valence electrons with you and you share one of your valence electrons with me. That way, we both have 8 valence electrons in our valence shells". For the rest of the night at the party the two fluorine atoms stay together, bonded side-byside, so they can share their electrons. The bond they form with each other while they share their electrons is known as a covalent bond. A covalent bond is one of two main types of chemical bonds (the other type is called an ionic bond). When two or more atoms bond together, they form a molecule. When the two fluorine atoms bond together they form a molecule called difluorine. 4

5 Another Covalent Bond Example: Two hydrogen atoms and an oxygen atom are sitting on a couch. The oxygen atom has 6 valence electrons, and the hydrogen atoms each have 1 valence electron (we know this from their periodic group numbers). They are all unhappy because none of them satisfy the octet rule (or in the case of the hydrogen atoms, they don't have the 2 electrons in the valence shell they so much desire). The oxygen atom then comes up with a brilliant idea. He proposes that each of the hydrogen atoms share an electron with him. When the atoms follow through with the deal, the oxygen atom has two more electrons to his name, one from each hydrogen atom. The hydrogen atoms each get the extra electron they desperately need to satisfy having two electrons in their outer shell. What you end up with is a molecule with 2 hydrogens, 1 oxygen and 2 covalent bonds holding them together. The molecule they form, which has the chemical formula H 2 O, is known as water. Bonding and Non-Bonding Electrons Bonding electrons are electrons that are involved in bonds between atoms, they are the electrons that are shared. Non-bonding electrons are electrons that are not involved in bonding. Pairs of non-bonding electrons are often referred to as lone pairs. 5

6 Lewis Dot Structures A Lewis Dot Structure is a simple way of showing how atoms are bound together in molecules. Here are the basic things you need to know about Lewis Dot Structures: 1. Lewis dot structure are mainly drawn for covalent bonded molecules 2. Dots are drawn around an atomic symbol to show valence electrons which are not involved in bonding to other atoms. 3. The dots are drawn in pairs. 4. Lines are drawn between atoms to represent pairs of electrons involved in covalent bonds. Example: Let's look at the Lewis Dot structure of difluorine as an example. As we discussed above, the fluorine atoms each initially have 7 valence electrons. When the two fluorine atoms come together they each share 1 of their electrons to form a covalent bond. This leaves 6 electrons on each fluorine which are non-bonding. Steps to Drawing a Lewis Structure 1. Pick a Central Atom Start your structure by picking a central atom and writing its element symbol. This atom will be the one with the lowest electronegativity. Sometimes it's difficult to know which atom is the least electronegative, but you can use the periodic table trends to help you out. Electronegativity definition - Is a measure of an atoms attraction for the electrons its shares in a bond with another atom. Electronegativity typically increases as you move from left to right across the periodic table and decreases as you move down the table, from top to bottom. You can consult a table of 6

7 electronegativities, but be aware different tables may give you slightly different values depending on how electronegativity is calculated. Once you have selected the central atom, write it down and connect the other atoms to it with a single bond. You may change these bonds to become double or triple bonds as you progress. 2. Count Electrons Lewis electron dot structures show the valence electron for each atom. You don't need to worry about the total number of electrons, only those in the outer shells. The octet rule states that atoms with 8 electrons in their outer shell are stable. This rule applies well up to period 4, when it takes 18 electrons to fill the outer orbitals. 32 electrons are required to fill the outer orbitals of electrons from period 6. However, most of the time you are asked to draw a Lewis structure, you can stick with the octet rule. 3. Place Electrons around Atoms Once you have determined how many electrons to draw around each atom, start placing them on the structure. Start by placing one pair of dots for each pair of valence electrons. Once the lone pairs are placed, you may find some atoms, particularly the central atom, don't have a complete octet of electrons. This indicates there are double or possibly triple bonds. Remember, it takes a pair of electrons to form a bond. Examples 7

8 Nitrogen molecule 8

9 Bond Preferences of Certain Elements It turns out that certain elements have a specific number of bonds they like to make. This is based on the number of valence electrons they have and the number of electrons they need to satisfy the octet rule. Knowing these rules will make drawing Lewis dot structures a lot easier. Atom Number of Covalent Bonds it Makes Hydrogen (H) 1 Carbon (C) 4 Nitrogen (N) 3 Oxygen (O) 2 Elements in group 17 (known as halogens) 1 9

10 Ionic bonds To understand ionic bonding we need to go into some points about electrical charges in atoms. Firstly a quick summary. Ionic bonds form when a metal reacts with a non-metal. Metals form positive ions; non-metals form negative ions. Ionic bonds are the electrostatic forces of attraction between oppositely charged ions. Atomic Charge You learned that protons have positive charges, electrons have negative charges and neutrons have neutral charges. It turns out that an atom can have a charge too. An atom's charge is referred to as atomic charge. Figuring out the charge on an atom is similar to calculating the balance on a bank statement. Atomic charge is worked out by adding all the positive and negative charges together and representing the difference as either + or -. Now, think of a positive charge like a deposit, a negative charge like a withdrawal and the atomic charge as the final balance. Just like you subtract the withdrawals from the deposits to get the final balance, you subtract the negative charges (this is equal to the number of electrons) from the positive charges (this is equal to the number of protons) to get the atomic charge. Example: Let's say an atom has 1 proton and 1 electron. This means that the atom has 1 positive charge and 1 negative charge. 1-1 = 0. Therefore the atomic charge of that atom is 0. Let's say an atom has 6 protons and 5 electrons. That means the atom would have 6 positive charges and 5 negative charges. The atomic charge is equal to 6-5 = +1 (when dealing with charge we always write +1 instead of just 1) Now let's we had an atom with 3 protons and 4 electrons. That means the atom would have 3 positive charges and 4 negative charges and the atomic charge would be 3-4 = -1 It is important to point out that molecules can have a charge too. Neutral When the charge of an atom or molecule is 0, we say that it is neutral. Example: An atom has 1 proton and 1 electron. It's charge is 1-1 = 0. This atom is neutral because it's charge is 0. 10

11 It is important to point out that a neutral atom has the same number of protons as electrons. As you know, the atomic number tells you how many protons an atom has. Therefore, in a neutral atom, the number of electrons it will have will be equal to it's atomic number. Simply put, an atom or molecule that is not neutral is said to be charged. Remember it this way: when we say something is politically charged, for instance like a speech or a commercial, what we really mean is that it is not neutral. When the charge of an atom or molecule is positive, we say that it's positively charged. When the charge of an atom or a molecule is negative, we say that it is negatively charged. Once the electrons have been placed, put brackets around the entire structure. If there is a charge on the molecule, write it as a superscript on the upper right, outside of the bracket. Example: The sulphate ion has the molecular formula of SO 4 and a charge of -2. This can be written as SO 4-2. A molecular ion is referred to as a polyatomic ion. Cations are positively charged and Anions negatively charged. When these ions come together the charges either repel or attract. Where they attract they bond in a manner that the electric charges are balanced. 11

12 In the table below are some common cations and anions. Example: Table salt is composed of a sodium ion 1+ bonded with Chlorine1-. The respective charges attract and they form an ionic bond. 12

13 Salt ionic lattice Chemical compounds that consist of cations and anions are called the generic label of salts. The chemical formula is written with the cation first then the anion second. Ionic substances form giant ionic lattices containing oppositely charged ions. Ionic compounds have high melting and boiling points, and conduct electricity when melted or dissolved in water. Here are some examples of polyatomic ions. Become familiar with these but you do not have to remember them. Examples Polyatomic Ion Charge = -1 acetate - C 2 H 3 O 2 - bicarbonate (or hydrogen carbonate) - HCO 3 - bisulfate (or hydrogen sulfate) - HSO 4 - chlorate - ClO 3 - chlorite - ClO 2 - cyanate - OCN - cyanide - CN - dihydrogen phosphate - H 2 PO 4 - hydroxide - OH - nitrate - NO 3-13

14 nitrite - NO 2 - perchlorate - ClO 4 - permanganate - MnO 4 - thiocyanate - SCN - Examples Polyatomic Ion Charge = -2 carbonate - CO 3 2- chromate - CrO 4 2- dichromate - Cr 2 O 7 2- hydrogen phosphate - HPO 4 2- peroxide - O 2 2- sulfate - SO 4 2- sulfite - SO 3 2- thiosulfate - S 2 O 3 2- Examples Polyatomic Ion Charge = -3 borate - BO 3 3- phosphate - PO 4 3- Polar bonds A polar bond is a covalent bond between two atoms where the electrons forming the bond are unequally distributed or shared. This disparity is due to a difference in electronegativity between the atoms. This causes the molecule to have a slight electrical dipole moment where one end is slightly positive and the other is slightly negative. Electronegativity definition The most common definition of electronegativity is that an element's electronegativity is the power of an atom when in a molecule to attract electrons to itself. The first scale of electronegativity was developed by Linus Pauling and his scale ran from about 0.7 (an estimate for francium) to 2.20 (for hydrogen) to 3.98 (fluorine). Bonds on this scale range from covalent (low electronegativity) to ionic (high electronegativity). Electronegativity has no units of measure. Electronegativity in molecules leads to polar bonds Polar bonds are on the dividing line between pure covalent boding and pure ionic bonding. The diagram below outlines an easy way to determine polarity. 14

15 Another factor that determines polarity of a molecule is its shape. A molecule can have polar bonds but due to it shape the polarity will not be expressed in definite areas of positive or negative charge. Carbon dioxide is a good example of this. Water is one of the most important polar molecules for living beings. Its polarity gives water many of it unique properties. We cover this in more detail in lesson 8. Valence Shell Electron Pair Repulsion Theory (VSEPR) Valence shell electron pair repulsion theory, VSEPR, is a technique for predicting the shape or geometry of atomic centres in small molecules and molecular ions. As mentioned above molecule shape influences whether bond polarity is expressed in the molecule as a whole. Molecular shape affects what chemical reactions / interactions can occur. VSEPR theory is based on the idea that the geometry of a molecule or polyatomic ion is determined primarily by repulsion among the pairs of electrons associated with a central atom. Remember same charges repel. The pairs of electrons may be bonding or nonbonding (also called lone pairs). Only valence electrons of the central atom influence the molecular shape in a meaningful way. This theory works for inorganic and organic molecules. 15

16 Basic assumptions of VSEPR 1. Pairs of electrons in the valence shell of a central atom repel each other. 2. These pairs of electrons tend to occupy positions in space that minimize repulsions and maximize the distance of separation between them. 3. The valence shell is taken as a sphere with electron pairs localizing on the spherical surface at maximum distance from one another. 4. A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair. Some typical molecule shapes predicted by this theory. VSEPR geometries Electron pairs Geometry 2 linear 3 trigonal planar 4 tetrahedral 5 trigonal bipyramidal 6 octahedral Two electron pairs or electron clouds Carbon dioxide Linear molecule 16

17 Three electron pairs or electron clouds Sulphur trioxide Trigonal planar Four electron pairs or electron clouds Methane Tetrahedral 17

18 Five electron pairs or electron clouds Phosphorus pentachloride Trigonal bipyramidal 18

19 Six electron pairs or electron clouds Sulphur hexafluoride Octahedral 19

20 Homework 1. Using your model kit make each of the following molecules. Take time to see the molecule shape in all dimensions. Carbon dioxide Oxygen dihydride Water Methane Phosphorus pentachloride Sulphur hexafluoride Note your kit is designed such that the correct molecular shape (VSEPR) forms due to the construction of the bonding points. 2. Define anions and cations plus give three examples of each with their charge 3. Define a polyatomic ion and give three examples 4. What are the two major types of chemical bonds and explain what electrons do in each type 5. Explain the steps in drawing a Lewis structure of a covalent compound using an example (do not copy the text example) 6. Define a polar covalent bond 7. Explain briefly the valence shell electron pair repulsion theory with an example. 8. What are the two factors that determine if a molecule will be polar? 20