Chemical Bonding Parts A and B PDF

Explaining the Diversity of Matter: Chemical Bonding  Periodic Properties of Elements Dmitri Mendeleev published a periodic table of elements. He arranged the elements in increasing atomic numbers (number of protons). Elements arranged further by putting them in groups (columns...

Explaining the Diversity of Matter: Chemical Bonding  Periodic Properties of Elements Dmitri Mendeleev published a periodic table of elements. He arranged the elements in increasing atomic numbers (number of protons). Elements arranged further by putting them in groups (columns) with similar properties: Alkali metals Alkaline-earth metals Halogens Noble gases Transition elements Actinides Lanthanides (rare earth) Descriptions of these groups (columns) can be found in your Nelson textbook on page 14.  There are certain trends in reactivity within the periodic table. The metals in the bottom left corner are the most reactive, while the nonmetals in the upper right corner are the most reactive. The noble gases (group 18) seldom react and are considered non-reactive. They are also known as the inert gases. Electronegativity Chemists believe that atoms have different abilities to attract electrons. The term electronegativity is used to describe the relative ability of an atom to attract a pair of bonding electrons in its valence level. Fluorine has the highest electronegativity (4.0), while francium has the lowest electronegativity (0.7). Please note that these are the most reactive metal and nonmetal because of their bonding capabilities.  The electronegativities of atoms are recorded on the periodic table. Metals tend to have low electronegativities, while nonmetals tend to have high electronegativities. Let’s Try Some! Exercise 1: Look on your periodic table and list the electronegativities for the following: Na Cs I F Answers!!! Na = 0.9 Cs = 0.8 I = 2.7 F = 4.0 When two atoms are in close proximity, the electronegativities of the two atoms determine the nature of the bond. Hydrogen (white) and Flourine (green) Valence Electrons, Valence Energy Levels & Valence Orbitals Periodicity of elements is related to the number of valence electrons an element has. Valance electrons are those electrons occupying the highest energy level of an atom. The valence energy level is the outermost energy level of an atom. Ex. The magnesium atom has 12 protons and 12 electrons. The maximum number of electrons in the first energy level is 2, the second energy level has 8 electrons, which leaves 2 electrons to occupy the valence energy level. Exercise 2: How many electrons are in the valence energy level of each of the following? (Please draw out an energy level diagram) Oxygen Argon Chlorine Exercise 2: How many electrons are in the valence energy level of each of the following? (Please draw out an energy level diagram) Oxygen Argon Chlorine 6 e¯ 8 e¯ 7 e¯ 2 e¯ 8e¯ 8e¯ 8 p+ 2 e¯ 2 e¯ O 18 p+ 17 p+ Oxygen atom Argon atom Chlorine atom 6 valence 8 valence 7 valence electrons electrons electrons  Orbital: a region in space around an atom’s nucleus in which an electron may exist. Valence Orbitals: volumes of space that can be occupied by electrons in an atom’s highest energy level. These are the only orbitals we are concerned with when studying bonding. The first energy level has room for only one orbital with a maximum of 2 electrons. Hydrogen has one electron and helium has the maximum number of two electrons. Energy levels above the first have room enough for 4 orbitals. Helium Neon  A valence orbital may contain 0, 1 or 2 electrons. In other words, there may not be more than two electrons in a valence orbital. Electrons will occupy all valence orbitals before forming electron pairs. A maximum of 8 electrons may occupy a valence energy level. This is known as the octet rule.  An atom with a valence orbital that is occupied by one single electron can share that electron with another atom. Therefore, it is called a bonding electron. A full valence orbital, occupied by 2 electrons, repels electrons in nearby orbitals. Two electrons occupying the same orbital are called lone pairs. For example, chlorine has 7 valence electrons. It has 3 lone pairs and one bonding electron. : Cl..  Electron Dot Diagrams Electron dot diagrams can represent atoms. These diagrams ONLY show the atom’s valence electrons! These are the only electrons involved in a chemical reaction. Follow these steps: 1) Write the atomic symbol for the atom. This symbol represents the nucleus and the inner energy levels that do not participate in the chemical bonding. 2) Dots () represent the electrons in the valence energy level of the atom. Arrange these dots around the atomic symbol. a) One dot must be placed in each of the four orbitals before any electron pairing occurs. b) Begin with the fifth electron to make lone pairs. c) There are four orbitals within the valence energy level, which allows a maximum of eight electrons in the valence energy level. Let’s try some… Exercise 3: Calcium Oxygen Bromine Carbon Answers!!! Calcium Ca Oxygen :O Bromine :Br Carbon C  Metalloids The staircase line on the periodic table gives a general separation between metals and nonmetals. Elements near this line have some properties of both metals and nonmetals - known as metalloids. Metalloids are unique because they have many bonding electrons compared to both metals and nonmetals. They are generally non-conductive, hard, and have high melting and boiling points. The metalloids are Boron, Silicon, Germanium, Arsenic, Antimony, Tellurium, Polonium and Astatine. Ionic and Covalent Bonding An ionic bond is the attraction that results from a positive ion (metal) and a negative ion (nonmetal). A positive ion (cation) is formed when an atom loses electrons to become similar to the nearest noble gas. The negative ion (anion) is formed when an atom gains electrons to become similar to the nearest noble gas. The two ions are then attracted to each other. A transfer of electrons occurs in an ionic bond. Na2O  A covalent bond is the bond that results from the electrostatic attraction between the electrons of one atom to the nucleus of another and vice versa. Two atoms approach each other close enough to allow the nucleus of each atom to attract the electrons of the other. This results in a mutual sharing of electrons between the two nuclei. Covalent bonds are classified as single, double, or triple bonds depending on the number of electrons shared between the two nuclei.  In a single bond, one pair of electrons is shared between the 2 atoms. In a double bond, two pairs of electrons are shared between the 2 atoms. In a triple bond, three pairs of electrons are shared between the 2 atoms. This helps to explain the diatomic elements. Recall that the diatomic elements are N2, O2, H2, and the halogens. “The gens” (group 17). Name That Bond!!!! Name That Bond!!!! This is a covalent bond as This is an ionic bond as the electrons are shared the electron (s) transfer between the 2 atoms. from one atom to the other. Bond Polarity Linus Pauling He explained that the polarity of a covalent bond is the difference in electronegativities of the bonded atoms. If the bonded atoms have the same electronegativities, they will share electrons equally and form a nonpolar covalent bond. If the atoms have different electronegativities, they will share the electrons unequally and form a polar covalent bond.  Electronegativities can be used to determine the type of bond formed between atoms. The difference between the electronegativities between two atoms can be used to describe the nature of the bond. Electronegativity Difference Type of Bond Between the Descriptions of Electrons in the Between Two Atoms. Atoms Bond Transfer of electrons between  1.7 Ionic metal and nonmetal Electrons shared unequally < 1.7 Polar covalent between unlike atoms Electrons shared equally 0 Nonpolar covalent between identical nonmetals  Examples: 1. Na-Cl Na 0.9 difference 2.3 ionic bond Cl 3.2 -electron transfers 2. N-O N 3.0 difference 0.4 polar covalent O 3.4 - electrons shared unequally 3. C-C C 2.6 difference 0 nonpolar covalent C 2.6 - electrons shared equally Ionic Bonding An ionic bond is the attraction that results from a positive ion (metal) and a negative ion (nonmetal). A positive ion (cation) is formed when an atom loses electrons to become similar to the nearest noble gas. The negative ion (anion) is formed when an atom gains electrons to become similar to the nearest noble gas. The two ions are then attracted to each other. A transfer of electrons occurs in an ionic bond. Na2O Electron Dot Diagrams for Ionic Compounds Unpaired electrons called bonding electrons are available for bonding. Electrons are transferred from the metal to the nonmetal - results in a net negative charge for the nonmetal and a net positive charge for the metal. NaCl(aq) MgCl2(aq) CaO(s) Electron Dot Diagrams for Covalent Compounds Here are some steps to follow… The Lewis-Dot diagram for Water (H2O) 1. Find the number of atoms in one molecule. Two hydrogen and one oxygen 2. If necessary, draw the electron-dot symbols for each of the different type of atoms in the molecule. 3. Find and calculate the total number of valence electrons in all the atoms combined. O: 1 x 6e- = 6e- + H: 2 x 1e- = 2e- total 8e- 4. Arrange the atoms so that it forms a model of the structure. Each element should have eight electrons surrounding it except for the hydrogen atom. To determine which atom is located in the center of the structure, usually, it the atom that can form the most bonds, in this case, it is the oxygen. 5. Count the number of electrons in the Lewis dot diagram. The sum of the dots should equal the amount of valence electrons counted in the third step. Let’s try some… NH3 SBr2 CH3OH Don’t Panic!!!! Lewis dot diagrams can take some time to figure out. If you cannot immediately determine the right configuration, take a deep breath and try again. Think of these as puzzles, keep on working with the pieces until they fit together Explaining the Diversity of Matter: Chemical Bonding –Part B Structural Formulas For every pair of bonding electrons in a covalent bond, a line can be drawn representing a bond. Examples: can be represented by Cl-Cl to show a bond. single bo can be represented by H-Br to show a single bond. can be represented with Find the Lewis and Structural Diagrams Draw Lewis formulas for the following: CS2 C2H4  More Lewis Formulas  Draw Lewis formulas for the following: CS2 C 2H 4.... H H :S::C::S: ∙∙ ∙∙ C :: C ∙∙ ∙∙ H H  Bonding Capacities of some Common Atoms Atom # of # of Bonding Valence Bonding Capacity Electrons Electrons Carbon 4 4 4 Nitrogen 5 3 3 Oxygen 6 2 2 Halogens 7 1 1 Hydrogen 1 1 1  You Want What? What is the difference between a Lewis dot diagram, a chemical formula, and a molecular formula? CH4 ____________ ____________ _____________  VSEPR Theory Valance Shell Electron Pair Repulsion Theory Stereochemistry – the study of the 3-D spatial configuration (shape) of molecules and how this effects their reactions. VSEPR Theory is very effective at predicting the shape of molecules. The theory is based on the repulsion of bonded and unbonded electron pairs in a molecule. The pairs of electrons in the valence shell of an atom stay as far apart as possible because of the repulsion of their negative charges. The following table shows the various molecular shapes. See pages 92-95. Note how to draw the stereochemical formulas on page 95.  VSEPR Class Shape Model Example AX Linear HCl AX2 Linear CO2 AX3 Trigonal BH3 Planar AX4 Tetrahedral CH4 : AX3 Pyramidal NH3.. V-shaped H2O :AX2 Multiple Bonds VSEPR Theory is also able to describe the shape of molecules containing double and triple bonds. Treat the multiple bond as one bond, just as you would a single bond. For example, C2H4(g) There are 3 bonds around each H H carbon atom (2 single and 1 \ / and 1 double) and no lone pairs. C = C This is an AX3 arrangement / \ which means a trigonal planar H H shape around each C atom. Polarity of Molecules The molecular shape and bond polarity are used to predict the overall polarity of a molecule. Complete the following steps: 1. Draw the Lewis formula for the molecule. 2. Use the number of electron pairs and VSEPR rules to determine the shape around the central atom. 3. Use electronegativities to determine the polarity of each bond. 4. Add the bond dipole vectors to determine whether the final result is zero (nonpolar molecule) or nonzero (polar molecule). A bond dipole is the charge separation that occurs when the electronegativity difference of two bonded atoms shifts the shared electrons, making one end of the bond partially positive (+) and the other partially negative (-). The arrow representing the bond dipole points from the lower to higher electronegativity. Example #1: O=N=O nitrogen dioxide.... : O :: N::O: Linear shape - ← + → - Resulting bond dipole vector O=N=O arrows cancel each other to 3.4 3.0 3.4 produce a zero total. Therefore, the resulting molecule is nonpolar.  Example #2: Water molecule HOH(l) or H2O(l)  Example #2: Water molecule HOH(l) or H2O(l).... H : O :H angular shape - 3.4 - Here the bond dipoles do not O cancel each other. There is + / \ + an overall polarity in the H H molecule. It is polar. 2.2 + 2.2  Polarity A polar molecule is a molecule that has an overall charge separation - one end of the molecule is positive and the other end of the molecule is negative. If a molecule is polar it is said to have a molecular dipole. A nonpolar molecule has no net charge separation. A molecular dipole is a result of unequal sharing of electrons in a molecule – from a difference in electronegativities Intermolecular Forces Intermolecular forces: The weak forces or bonds betweem molecules. Intramolecular bonds: attractions within a molecule (covalent bonds that hold the molecule together. Intermolecular bonds involve the electrostatic attractive forces between molecules. Ionic substances do not form molecules. Therefore, intermolecular bonding only occurs in substances that form covalent bonds.  Types of Intermolecular Bonding These forces that exist between molecules are referred to as Van der Waals forces. Van der Waals forces can be divided into three different types, namely, dipole-dipole forces, London dispersion forces, and hydrogen bonding. 1) Dipole-Dipole Forces Polar molecules have dipoles (oppositely charged sides). Attraction between dipoles is called the dipole-dipole force. The positive end of one molecule will be attracted to the negative end of a neighboring molecule. This will extend in all directions.  If a molecule is nonpolar, no permanent dipoles exist  The intermolecular force that holds these molecules together in the liquid and solid states is called the London dispersion force. 2) London Dispersion Forces result from the movement of the electrons in the molecule which generates temporary positive and negative regions in the molecule The weak attractive forces that result when the electrons of one molecule are attracted to the positive nuclei of a nearby molecule.  London Dispersion Forces are due to the electrostatic attraction between temporary dipoles in neighboring molecules. Let us consider the Chlorine molecule, Cl2(g). This molecule is nonpolar since there is no difference in electronegativity between the atoms. At a particular instant, we may find that the two electrons that form the bond may be closer to one nucleus than the other. Results in a temporary dipole with one end more negative than the other. This temporary dipole may be attracted to another temporary dipole in a neighboring molecule. In a London dispersion force, there is no net shift of electrons. A temporary dipole-dipole interaction results from the temporary shift of electrons. General Information: London forces are present between all molecules, whether or not any other types of attractions are present. The more polar the molecule, the stronger the dipole – dipole forces. The more electrons that are present in the molecule, the stronger the London forces. forces The greater the distance apart from each other the molecules are, the weaker the London forces. The stronger the London forces and dipole – dipole forces are, the higher the boiling point. point More energy is required to separate the molecules. 3) Hydrogen Bonding Occurs when hydrogen is bonded to a highly electronegative element (fluorine, oxygen and nitrogen) – chemistry is FON!!! In these cases, the bond is strongly polar. The highly electronegative atom pulls hydrogen’s electron away from its nucleus. Another molecule’s lone pair of electrons can now approach the nucleus closely on the side away from its covalent bond. The hydrogen end of the bond takes on a strong positive charge because of the exposed positive nucleus, while the other element takes on a strong negative charge. This positive hydrogen will be attracted to nearby negative atoms. It appears as though the hydrogen atom bonds to different molecules. Hydrogen bonding is the strongest of the intermolecular bonds!!!! Three structures show hydrogen bonding: 1. HF(g) 2. Any molecule with an –OH group in any part of its structure. 3. Any molecule with an –NH group in any part of its structure. - Hydrogen bonding causes very high boiling points as much energy must be added in order to break these bonds holding the molecules together. - Hydrogen bonds explain why ice is less dense than water. The hydrogen bonds hold the water molecules farther apart in the solid state. Ice is less dense and floats on water. - Hydrogen bonds hold the DNA double helix together. If covalent bonds held it together, they would be too strong to allow the helix to open up during DNA replication. - Water has unusually high melting and boiling points because the hydrogen bonds are so strong.  Effects of Intermolecular Bonding on Physical Properties Melting and boiling points can be a measure of intermolecular bond strength. If the intermolecular bond is strong, the molecule will be very stable - It will take a lot of energy to break apart the molecules. (high boiling and melting points) Generally, as the molecular mass increases, the melting points and boiling points also increase. Water is an exception to this generalization. Water has unusually high melting and boiling points because the hydrogen bonds in water are extremely strong. Another important point to note is that the more complex the molecule is, the stronger the intermolecular bonds will be.  Bond Energy Objects tend towards a lower, more stable energy state. Ex. a ball rolls down a hill to be a lower, more stable position. Similarly, when a chemical bond is formed, the elements have reached a lower, more stable energy level - the energy of a compound is frequently less than the energy of the individual elements that make up the compound. Consider the methane molecule, CH. It is made up of 4(g) carbon and hydrogen. This reaction gives off energy and is considered exothermic – the reactants have more energy than the products in the following equation. C(s) + 2H2(g) CH4(g) + 74.4 kJ The point at which both atoms are most stable is the point at which the chemical bond is formed. Relative Strength of Various Bonds The following bond scale is an approximation only. In order of increasing strength on a scale of 1-to-100: 1 5 10 10 0 Bond Type and Solubility Generally, “like dissolves like”. Polar molecules will generally dissolve in polar liquids. Nonpolar molecules will generally dissolve in nonpolar liquids. For example, acids (which are polar molecules) will dissolve in water which are also polar molecules.  Structures and Physical Properties of Solids A. Ionic Crystals – crystal lattice structure B. Metallic Crystals – metallic bonding C. Molecular Crystals – molecules arranged in a regular lattice D. Network Covalent Crystals – network of covalent bonds  Crystal Lattices Why are ionic compounds solid at room temperature? Why do they have high boiling and melting points? Why are they good conductors of electricity? Ionic compounds are very rigid because the ions arrange themselves such that they are always surrounded by oppositely charged ions - they will never arrange themselves with a positive ion next to another positive ion. (Like charges repel, and unlike charges attract.) The fact that ionic compounds have high melting and boiling points also indicates that they are strongly bonded together - a lot of energy is needed to break intermolecular bonds. Intermolecular bonds are bonds between compounds. The arrangement of ions for an ionic compound is called its crystal lattice. Metallic Crystals Metals are shiny, silvery, flexible solids that conduct heat and electricity well. They have closely packed structures. The valence electrons are not held strongly by their atoms. This is due to the low electronegativity of metal atoms. The electrons can easily become mobile. The electrons act like a negative glue holding the positive metal atoms firmly together. This produces continuous, closely packed structures. Molecular Crystals They have relatively low melting points and a general lack of hardness. This is due to the fact that London, dipole-dipole, and hydrogen bonding forces are not very strong compared to ionic and covalent bonds. The individual entities are neutral molecules so they can’t conduct an electric current. Covalent Network Crystals These are very hard and brittle, with high melting points. They are insoluble and do not conduct electricity. Examples include diamond, quartz, and silicon carbide. They consist of a network of covalent bonds. bonds The atoms are continuously linked throughout the crystal by strong covalent bonds. Most covalent networks involve carbon or silicon atoms which bond strongly and often to themselves and other atoms. The interlocking structure provides the strength, hardness and high melting points. The electrons are not free to move, which explains the nonconduction of electricity.