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Introduction to Bonding Most of the matter on Earth is in the form of compounds. Even when a substance exists as a pure substance, it tends to eventually combine with other elements. For example, if you leave an iron nail out in the rain it will combine with the oxygen in the air to form iron oxide (rust). In this investigation, you will build models of atoms and discover one of the fundamental ideas of chemistry: how electrons are involved in the formation of chemical bonds. You have already learned: A neutral atom has the same number of electrons and protons. The electrons occupy energy levels surrounding the nucleus. The electrons occupy the lowest energy level first. Once an energy level is full, electrons will start to fill the next level. Part I: The Valence Electrons Directions: Using the atom building game, build each element in the following table. You do not have to add protons and neutrons to the models, only the electrons, since they are the only subatomic particles involved in bonding. The first one is done for you as an example. Element Atomic Number B He Li F Ne Na Cl Be K N O 5 Electrons in outermost energy level 3 How does it become stable? Ion formed Lose 3 electrons B3+ Use your table to answer the following questions: A. What do lithium, sodium, and potassium have in common? B. What do fluorine and chlorine have in common? C. The electrons in the outermost energy level of an atom are called valence electrons. These are the electrons involved in chemical bonds. How is this number related to the group number of the element? Part II: Modeling a Chemical Bond Atoms that have a complete outermost energy level are stable. We have been using the yellow marbles and either adding more or taking away marbles to show the atom gaining or losing electrons, but what really supplies those electrons? Another atom either gives or takes away electrons to make the atom stable. This is how chemical bonds are formed. Directions: 1. Take out two atom building game boards and the red, blue and yellow marbles. 2. Find sodium on the periodic table. a. How many protons does sodium contain? ____ Add this many red marbles to the first board. b. How many neutrons does sodium contain? ____ Add this many blue marbles to the first board. c. How many electrons does sodium contain? ____ Add this many yellow marbles to the first board. 3. Find chlorine on the periodic table. a. How many protons does chlorine contain? ____ Add this many red marbles to the second board. b. How many neutrons does chlorine contain? ____ Add this many blue marbles to the second board. c. How many electrons does chlorine contain? ____ Add this many yellow marbles to the second board. Questions: A. What will sodium do to become stable? _______________________________ B. What will chlorine do to become stable? ______________________________ C. Why might these two atoms bond together to form a molecule? An element’s oxidation number is the same as the charge an atom has when it ionizes (in other words, when it gains or loses electrons to become stable). Move one electron from the outer ring of sodium to the outer ring of chlorine. D. Are the sodium and chlorine now stable? _________________ E. What is sodium’s oxidation number? (Hint: what is the charge on it?) ______ F. What is chlorine’s oxidation number? ______ G. Suppose the sodium ion and chlorine ion stayed together as a molecule. How many total positive charges does this molecule have? _______ How many total negative charges does this molecule have? _______ What is the molecule’s overall charge? _____ Compounds overall are electrically neutral. To show this, we add together the oxidation numbers: Oxidation number of sodium ________ + Oxidation number of chlorine ________ Equals: ________ In a compound, the sum of the oxidation numbers is always __________. Part III: Oxidation Numbers and the periodic table Remember, the oxidation number is the charge an atom has when it loses or gain electrons to become stable. It indicates how many electrons are lost or gained. If the atom loses one or more electrons, the number is positive. If the atom gains one or more electrons, the number is negative. We can’t build a model every time we want to find an element’s oxidation number. Luckily, the periodic table shows some trends that we can remember to help us make things easier. Look at the following partial periodic table. We are just going to be working with the “A” groups of elements. 1. On the first line, write the correct Roman numeral for that group. 2. On the second line, write the number of valence electrons that group has. 3. On the third line, write what the elements must do to become stable. (i.e. gain 2 e- lose 3 e- etc.) 4. On the fourth line, write the charge the elements will have when they become stable. (This is the oxidation number.) Part IV: Naming Ionic Compounds Compounds that are formed from ions are called ionic compounds. 1. Using the periodic table on the previous page, write down the ion formed by each element (symbol and oxidation number). Sodium ______ Sulfur ______ Chlorine ______ Oxygen ______ Magnesium ______ Aluminum ______ Beryllium ______ Nitrogen ______ 2. The cards on the card stock page represent the ions. A peg represents a negative charge (extra electrons that have been gained); a notch represents a positive charge (missing electrons –they have been given away). 4. For each compound below, put the cards together so that each notch is matched by a peg. This will provide equal numbers of positive and negative charges, as in real ionic compounds. 5. Write down the number of each element that you need to create a neutral molecule. Add the oxidation numbers of the elements represented to make sure it is zero. 6. Write the chemical formula in the box provided. The first one has been done for you. Elements Cation Anion sodium and oxygen Na1+ O2sodium and chlorine beryllium and sulfur magnesium and oxygen aluminum and chlorine sodium and nitrogen beryllium and oxygen aluminum and oxygen sodium and sulfur beryllium and nitrogen magnesium and chlorine magnesium and sulfur aluminum and sulfur beryllium and chlorine magnesium and nitrogen aluminum and nitrogen Questions: A. Can you join together sodium and aluminum? Explain. Number of Cation needed 2 Number of Anion needed 1 Chemical Formula Na2O B. Look at your cation, anion, and your chemical formula. Is there a “shortcut” way to get to the formula from the ions? Explain. Writing Formulas for Binary Ionic Compounds Things to remember: Ion – an atom that has gained or lost electrons, so has a positive or negative charge. Cation – an ion with a positive charge; Anion – an ion with a negative charge Oxidation number – the number that indicates how many electrons are lost or gained during bonding. This is the same as the charge on the stable ion of that element. In order to gain stability, the sum of the oxidation numbers for a compound must equal zero. Binary compound – a compound formed between a metal and nonmetal Writing formulas for binary compounds: 1. Write the ion (symbol and oxidation number) of the cation first. 2. Write the ion (symbol and oxidation number) of the anion second. 3. Write subscripts (small numbers near the bottom) is to indicate how many ions of each element is needed to make the sum of the oxidation numbers zero. Do not write the number 1 as a subscript. Examples: Magnesium and Oxygen 1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable. Mg2+ 2. The negative ion will be oxygen (Group VI), since it gains 2 electrons to become stable. Mg2+ O23. The sum of the oxidation numbers is already zero, so we only need one of each ion. Mg1O1 --- we don’t write the ones, so it is MgO Magnesium and Bromine 1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable. Mg2+ 2. The negative ion will be bromine (Group VII), since it gains 1 electron to become stable. Mg2+ Br13. The sum of the oxidation numbers is not zero, so we only need to add more ions until it becomes zero. Mg2+ Br1Br1By adding another bromine ion, we have +2 + -2, which = 0. So, our formula is MgBr2 Practice 1. Sodium and oxygen 2. Calcium and fluorine 3. Iron(II) and sulfur Crisscross Method A shortcut to figure out the number of each ion needed to form a neutral compound is called the crisscross method: 1. Write each ion – cation first, followed by the anion. 2. Cross the number that is a superscript of one to be the subscript of the other. (Just cross the number, not the positive or negative sign. The sign is dropped.) 3. Reduce your subscripts if they are not in the lowest possible ratio. Examples: Magnesium and Bromine 1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable. Mg2+ 2. The negative ion will be bromine (Group VII), since it gains 1 electron to become stable. Mg2+ Br13. Cross the two numbers and drop the positive or negative sign Mg2+ Br1- So, our formula is MgBr2 Magnesium and Oxygen 1. The positive ion will be magnesium (Group II), since it loses 2 electrons to become stable. Mg2+ 2. The negative ion will be oxygen (Group VI), since it gains 2 electrons to become stable. Mg2+ O23. Cross the two numbers and drop the positive or negative sign Mg2+ O2- So, our formula is Mg2O2 But you can reduce 2:2 to 1:1, so the formula becomes MgO. Practice 1. Sodium and sulfur 2. Lithium and iodine 3. Copper(IV) and oxygen Binary Compounds with Polyatomic Ions Polyatomic Ions are covalently-bonded groups of atoms that behave as a unit and carry a charge EX: NO3Nitrate ion C2H3O2acetate ion the charge that is shown is for the whole group of atoms together, not just the last element these ions have very specific names – you will use the following list to find the name and formula for the polyatomic ions you will use in this class. Name Ammonium Acetate Chlorate Hydroxide Nitrate Carbonate Sulfate Sulfite Phosphate Nitrite Chlorite Cyanide Formula NH41+ C2H3O21ClO31OH1NO31CO32SO42SO32PO43NO21ClO21CN1- Use the crisscross method for polyatomic ions, just as you did for regular ions. Never change anything within the polyatomic ion – if you need to add a subscript, put it outside of a parenthesis. Examples: Magnesium Phosphate Mg2+ PO43Mg3(PO4)2 Practice 1. magnesium carbonate 2. calcium hydroxide 3. manganese(IV) sulfite Naming Binary Compounds Ionic compounds have two word names: 1. Write the name of the cation first – this is just the name of the element 2. If the cation is a transition metal, add a Roman numeral in parentheses to indicate the oxidation number of the ion in this compound. a. to find the oxidation number do the crisscross method backwards. b. check the negative ion to see if it has the correct oxidation number. If not, multiply both charges by the appropriate factor until it does. c. remember: the Roman numeral only goes in the name, not into the formula. 3. Write the name of the anion – take the root of the element name and end with –ide. For example: 1. NaF is __________________________________________ 2. MgO is _________________________________________ 3. CuCl2 is _________________________________________ 4. PbO is __________________________________________ Practice 1. Name the following compounds a. ZnS b. K3N c. BaO d. CaO e. AlF3 f. CuI2 Naming Polyatomic Compounds It is important that you recognize the polyatomic. If there are more than 2 elements in the formula, then you know it contains a polyatomic. Note: only one polyatomic is a cation, ammonium (NH4+) All the rest will come after the metal ion. You name polyatomic compounds just like you would other ionic compounds, but do not change the ending of the polyatomic ion. EX: Name ZnSO4 ____________________________ EX: Name NaOH ____________________________ Practice 1. Name these compounds a. CaCO3 b. KClO2 d. Al(OH)3 e. Mg3(PO4)2 Notes: Bonding I. Bonding Remember, a compound consists of 2 or more elements chemically bonded together. Examples: H20, NaCl, Sb3(PO4)5 smallest unit of a compound is called a molecule Only electrons move during bonding. it is the movement of electrons in the outer energy level (valence electrons) that determines how one atom will bond with another. Octet rule – all atoms need 8 valence electrons to be stable (the exception is those atoms that are stable by filling the first energy level, they only need 2) 4 types of chemical bonds: 1. Ionic bonds 2. Covalent bonds 3. Polar covalent bonds 4. Metallic bonds II. Ionic bonds Electrons are transferred from one atom to another. When an atom loses an electron it becomes a positive ion (cation). When an atom gains an electron it becomes a negative ion (anion). Opposites attract, so the cation is attracted to the anion This is the strongest type of bond Ionic bonds form between metals and nonmetals Where are the nonmetals found on the periodic table? ____________ Where are the metals found on the periodic table? _______________ Example: NaCl Lewis dot: Energy levels: 1 III. Covalent Bonds Electrons are shared between atoms, but not transferred The electrons are in the outermost energy level of both atoms at the same time Covalent bonds form between nonmetals Example: NH3 Lewis dot: Energy levels: Ionic vs. Covalent Properties type of elements type of bonding structure melting point hardness conduct electricity? Covalent Bonds Nonmetals only share electrons not rigid – solid, liquid or gas low soft hard no Ionic Bonds Metals and Nonmetals transfer electrons crystalline high very hard yes (when dissolved in water) Practice 1. Determine whether the following are ionic or molecular covalent compounds a. N2O5 b. a liquid at room temperature c. PbNO3 d. a salt that conducts electricity when in solution e. KF f. AgCl g. gasoline h. PCl3 2 IV. Polar Covalent Bonds A special type of covalent bond in which the electrons are shared unequally The electrons spend a little more time with one atom than with the other. This causes the “stronger” atom in the molecule to be slightly negative, while the other atom is slightly positive. Example: H2O Lewis dot: Energy levels: V. Metallic Bonds Bonds between positive ions and free-roaming electrons. These are flexible and are good conductors as solids (electrons are free to move) 3