Ch. 6 - Chemical Bonds Chemical reactivity depends on electron configuration. Remember the Stable Octet rule: when the highest energy level occupied is filled with electrons (8 electrons for most atoms), then the atom is stable and not likely to react. Electron dot diagrams are models of an atom with each dot representing a valance electron. Always place the first two electrons on the same side and then rotate counterclockwise around the atom. Do electron dot diagrams for the following: Ca O Cl Construct your own electron dot diagram Choose one element & drag the correct number of VALENCE Br electrons around it. ELECTRONS Ca Li S O H N C Cl P Na Mg 1
Elements without complete sets of valence electrons tend to react (by gaining or losing electrons). Ionic bonds: the transfer of electrons from one atom to another, resulting in a more stable electron configuration and the formation of oppositely charged ions which are then attracted to each other (between metals and nonmetals!) Anion = ion formed when atom gains electrons; to name anions use part of the element name plus the suffix -ide. (A Neg ION) Cation = ion formed when atom loses electrons; to name these, just name the element! 2
When anions and cations are close together, a chemical bond will form between them. Chemical bond = force that holds atoms or ions together as a unit Ionic bond = force that holds cations and anions together As soon as the electrons are transferred from one atom to another, ionic bonds will form because of the oppositely charged ions being so close together. Ionization energy = amount of energy used to remove a valence electron from its ground state energy level The lower the ionization energy, the easier it is to remove an electron from an atom. The ionization energy tends to increase as you move across a period and decrease as you move down a group Electronegativity patterns are the same as that of ionization energy. The greater the electronegativity of an atom, the easier it is for the atom to gain an electron. As you move across a period, electronegativity increases. As you move down a group, electronegativity decreases. This is why the most reactive non-metals are found at the top of each group! 3
Ionic Compounds Ionic compounds = represented by chemical formulas which show the elements in the compound and the ratio of the atoms or ions in this compound; to determine the correct ratio, you can look at the electron dot formulas and see how many of each will be needed to complete the required electron transfers. Na + Cl Na +1 + Cl -1 NaCl Mg + Cl Mg +2 + Cl -1 + Cl -1 MgCl 2 The 2 is called a subscript and represents the number of chloride ions present in the ionic compound. Ionic compounds are often arranged in a lattice structure called a crystal. The arrangement of the ions depends on the ratio of ions and their relative sizes. Properties of ionic compounds (these are a result of the strong attractions among ions within a crystal lattice) include: 1. solid at room temperature 2. relatively high melting point 3. crystalline 4. brittle (hammering pushes similarly charged ions close together, they repel each other and the crystal breaks easily) Good conductors of electricity must have charged particles that can move freely from one place to another. When salt is a solid, the electrons are held in fairly fixed positions so it is a poor conductor of electricity. However, when you heat it, the ions gain enough kinetic energy to break the lattice apart and it becomes an excellent conductor of electricity. 4
Writing Formulas for Binary Ionic Compounds 1. Write the symbol for the metal (cation) first, nonmetal (anion) second Ex: Na Cl 2. Use the periodic table and write the charge as a superscript after the symbol Ex: Na +1 Cl -1 3. Criss-cross the number, not the sign, so they become subscripts after the other element Ex: Na +1 Cl -1 Na 1Cl 1 drop the sign from the formula do not use a subscript of 1 (the symbol itself represents 1) Na 1Cl 1 NaCl 4. Reduce subscripts if possible Draw the electron dot diagrams for each and figure out how they will combine to form an ionic compound: Mg Cl Pull 5
Al S Ca O Writing Formulas for Ionic Compounds with Polyatomic Ions 1. Write the symbol for the polyatomic, metal or nonmetal (cation first, anion second) Ex: Mg PO 4 2. Use the periodic table and polyatomic ions list (see list in packet) and write the charges as superscripts Ex: Mg +2 (PO 4) -3 3. Criss-cross the number (not the sign) and write as subscripts Ex: Mg 3(PO 4) 2 if the polyatomic ion has a subscript >1 put the polyatomic ion in parentheses 4. Reduce subscripts if possible Ex: NH 4 + Cl (NH 4) +1 Cl -1 NH 4Cl Mg + PO 4 Mg +2 (PO 4) -3 Mg 3(PO 4) 2 Ca + O Ca +2 O -2 Ca 2O 2 CaO Writing Formulas for Compounds Containing Transitional Metals If the ion is a transitional metal, you will have been given the charge via the Roman numeral in the name. EX: Copper (II) oxide = Cu +2 + O -2 Cu 2 O 2 CuO Iron (III) oxide = Fe +3 + O -2 Fe 2 O 3 Chromium (III) permanganate = Cr +3 + (MnO 4 ) -1 Cr(MnO 4 ) 3 6
Naming Compounds Ionic compounds always include a cation (written first) and an anion (written second) 1) Binary ionic compounds = refer to chart of anions & cations NaCl Sodium chloride MgBr 2 Magnesium bromide CaI 2 Calcium iodide Fr 2 O Francium oxide Li 3 P Lithium phosphide (note that the subscripts represent the number of anions or cations in the compound) 2) Compounds of metals with multiple ions (transition metals): You must include a Roman Numeral to indicate the ion being used. CuO copper (II) oxide (it takes only 1 copper to balance the -2 charge on the oxygen so it must be a +2 copper ion) Cu 2 O copper (I) oxide (it takes 2 copper ions to balance the -2 charge on the oxygen so each of them must be a +1 ion) FeN Iron (III) nitride Pb 2 S 3 Lead (III) sulfide (since you know sulfur has a -2 charge then the total charge for sulfur is -6 and since you know you only have 2 leads, each must have a -3 charge!) 3) Ionic compounds with polyatomic ions Polyatomic ion = atoms joined by covalent bonds to form a charged unit acting as a single ion Most polyatomic ions are anions. To name them, simply use the name of the cation along with the polyatomic ion: NaOH Sodium hydroxide Mg(OH) 2 Magnesium hydroxide (note that parentheses are used if there is more than 1 polyatomic ion) Fe 2(SO 4) 3 Iron (III) sulfate 7
Covalent bonds = bond formed by atoms sharing a pair of valence electrons; are usually made up of nonmetals bonding with nonmetals because nonmetals tend to have such high ionization energies that the electrons can t escape to form cations. Covalent bonds result in the formation of a molecule. The attractions between the shared electrons and the protons in each nucleus hold the atoms together in a covalent bond. Many nonmetal elements exist as diatomic molecules which are 2 of the same atoms covalently bonded together. Common diatomic molecules are: Bromine (Br 2 ) Oxygen (O 2 ) Hydrogen (H 2 ) (BrOHNClIF) Nitrogen (N 2 ) Chlorine (Cl 2 ) Iodine (I 2 ) Fluorine (F 2 ) Nitrogen (with 5 valence electrons) would need to covalently link up to 3 other nitrogen atoms in order to share a total of 8 electrons. Atoms sharing 2 pairs of electrons = double bond Atoms sharing 3 pairs of electrons = triple bond Atoms sharing electrons don t have equal attractions for the electrons Elements on the right tend to have greater attractions for electrons than elements on the left. Elements at the top of a group have greater attraction for electrons than elements at the bottom of the group. 8
Fluorine is on the far right and also at the top so it has the strongest electron attraction and is the most reactive nonmetal. In diatomic molecules, the electrons are shared equally. In compounds, electrons are often not shared equally. Polar covalent bond = a covalent bond in which electrons are not shared equally; the atom with the greatest electron attraction is denoted with the partial negative charge while the other atom is denoted with a partial positive charge. The above molecule s polarity could be shown by writing: If a molecule has only 2 atoms and they are different, it will be polar. If a molecule has more than two atoms, its polarity depends on the type of atoms and the shape of the molecule. Ex: CO 2 is not polar because the atoms line up horizontally and there are 2 double bonds between each oxygen and the central carbon. O = C = O There is equal pull on the electrons so the molecule is nonpolar. When molecules are near each other, there is an attraction that exists which is strong enough to hold the molecules together in a liquid or a solid. However, attractions between polar molecules are stronger than attractions between nonpolar molecules. Methane (at. mass of 18, b.p. of -161 o C) and water (at. mass of 18, b.p. of 100 o C) have very different boiling points because methane is nonpolar and water is polar. The slightly negative oxygen of one water molecule will be attracted to the slightly positive hydrogen of another water molecule. This attraction will increase the energy required for water molecules to evaporate. Nonpolar molecules have a slight attraction towards each other but it is not near as strong as the attraction between polar molecules. 9
Naming Molecular Compounds 1. Name the first element (usually it is the most metallic of the elements in the compound - furthest to the left on the periodic table or nearest the bottom if both are in the same group) add the correct prefix to indicate the number of atoms of each type that are present (refer to chart) never use the prefix mono- if the element name begins with a vowel drop the final letter in the prefix 2. Name the second element change the ending to -ide add correct prefix to indicate the number of atoms drop the vowel in the prefix if the name begins with a vowel (pentaoxide would be pentoxide!) Number of atoms Prefix 1 Mono 2 di 3 Tri 4 Tetra 5 Penta 6 Hexa 7 Hepta 8 Octa 9 Nona 10 Deca Ex: CO = carbon monoxide P 2 O 5 = diphosphorus pentoxide CCl 4 = carbon tetrachloride Writing Molecular Formulas 1. Write the symbols for the elements in the order they appear in the name EX: carbon dioxide CO 2. The prefix indicates the number of atoms and is written as a subscript (if there is no prefix, there is only 1 atom of that element) EX:carbon dioxide CO 2 dinitrogen tetraoxide N 2 O 4 10
Metallic Bonds If there aren t any nonmetal atoms to accept electrons that metals want to get rid of, there is another way that metals can achieve som sort of electronic stability. Because electrons in metals can move freely among other metal atom (this is why metals are such good conductors!), the metal atom becomes like a cation surrounded by a pool of shared electrons. Metallic bonds are the attraction between a metal cation and the shared electrons that surround it The cations in a metal form a lattice that is held in place by strong metallic bonds between the cations and the surrounding valence electrons. Overall, the metal is neutral. The more valence electrons the stronger the metallic bonds Na has 1 valence electron weak bonds (Na is a soft metal with a low melting point) W has several valence electrons strong bonds (W is a hard metal with a high melting point) Properties of Metals (are a result of the mobility of electrons within the metal lattice): 1. ability to conduct an electric current 2. malleable 3. ductile 4. forms alloys - is a mixture of 2 or more elements with at least one being a metal Bronze copper + tin, harder than brass, weathers slowly Brass copper + zinc, shinier than bronze, weathers rapidly Steel iron + carbon, carbon atoms fill in the gaps between the iron atoms making it stronger Stainless Steel iron + chromium + very little carbon, chromium prevents steel from rusting Aluminum + copper or manganese used for airplanes (aluminum + magnesium is even lighter) 11