Based on material in chapters 7 and 8 ¡  When metals bond with nonmetals, electrons are transferred from the metal to the nonmetal ¡  The metal becomes a cation and the nonmetal becomes an anion. ¡  The attraction between the cation and the anion results in an ionic compound. ¡  In Lewis Theory, we can show this by moving electron dots from the metal to the nonmetal ¡
Lewis Structures: ¡
When Potassium and Chlorine bond: §  Potassium transfers its valence electron to chlorine §  This transfer gives chlorine an octet and leaves potassium with an octet in the previous principal shell (now its valence shell) §  The Lewis dot structure of the anion is usually written with brackets with the charge in the upper right corner (outside the brackets) §  The positive and negative charges attract one another, resulting in the compound KCl ¡  Lewis Structures: ¡  Magnesium loses its 2 valence electrons, forming a 2+ charge. Oxygen gains these two electrons forming a 2-­‐ charge and acquiring an octet ¡  Lewis theory can help predict the correct chemical formulas for ionic compounds. ¡  Ex. Lewis theory predicts 1 K atom for every Cl atom, forming KCl. ¡  Lewis Structures: ¡  Sodium loses its 1 valence electron to get an octet in the previous principal shell ¡  Sulfur must gain 2 electrons to get an octet ¡  The compound that forms between sodium and sulfur requires 2 sodium atoms for every one sulfur atom ¡  Correct Lewis Structure: ¡  Formula: Na2S ¡
Crystalline SOLIDS at room temperature §  Ions are arranged in a repeating 3-­‐D pattern §  Minimizes like-­‐charge repulsions §  Results in a very stable compound §  MASSIVE repeating structure! Makes it hard to melt, and relatively harder to dissolve (requires agitation, etc.) ¡
High melting point §  Strong attractive forces due to the ions §  The large structure makes it hard to melt, and can make it relatively harder to dissolve (requires agitation, etc.) ¡  Conduct electric current when melted, or dissolved in water §  This destroys the stable crystalline structure §  Ions are free to move to the oppositely charged electrode. §  See next slide for image ¡  Metallic bonding: metals bound with other metals ¡  Metals are made of closely packed cations rather than atoms. ¡  Valence electrons are modeled as a “sea of electrons” that move freely among the metal ¡  This makes metals have these properties: §  High ductility (can be pulled into wires) §  High malleability (can be pounded into flat sheets) ¡  Mixtures composed of two or more elements ¡  Often stronger than their individual components §  Steel (iron, carbon, etc) §  Bronze (copper and tin) ¡  When nonmetals bond with other nonmetals, a molecular compound is formed. ¡  Molecular compounds contain covalent bonds, in which electrons are shared between atoms rather than transferred ¡  In Lewis theory, we represent covalent bonding by allowing atoms to share some of the valence electrons in order to attain octets (or duets for hydrogen) ¡
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We know that Noble Gases like Neon and Argon are unreactive. This was explained by Gilbert Lewis who came up with the Octet Rule in 1916 as part of the Lewis Theory of Bonding Octet Rule – Atoms tend to achieve the electron configuration of a noble gas when forming compounds Octet – set of 8 electrons in the highest energy orbital §  Atoms with an octet are stable ¡
Hydrogen, lithium, Beryllium and Helium are exceptions – their most stable electron configuration is known as a duet (2 electrons in their highest energy level) ●He● ¡  Lewis Structures: ¡  In water, hydrogen and oxygen share their electrons so that each hydrogen atoms gets a duet and each oxygen atom gets an octet ¡  The shared electrons (those appearing in the spaces between the two atoms) count toward the octets (or duets) of both atoms ¡  Electrons that are shared between two atoms are called bonding pair electrons ¡  Electrons that are not shared but belong to only one atom are called lone pair electrons or unshared pairs. ¡  In water, there are 2 bonding pairs of electrons (one between each Hydrogen and the oxygen atom) and 2 lone pair electrons (both on the oxygen atom) ¡  Bonding pair electrons are often represented by dashes to emphasize that they form a chemical bond. ¡  Single covalent bond – a bond in which two atoms share a pair of electrons ¡  When we represent compounds this way, it is known as a structural formula. ¡  Structural formula for water: ¡
Ammonia (NH3 ) §  # of Bonding Pairs: ▪  3 §  # of Lone Pairs: ▪  1 ¡
Methane (CH4 ) §  # of Bonding Pairs ▪  4 §  # of Lone Pairs ▪  0 ¡
Lewis Structure of Hydrogen: ¡
Each hydrogen has one electron to share with the other hydrogen atom to form a duet. ¡
Bonding Pairs? Lone Pairs? §  1 bonding pair, no lone pairs ¡
Structural formula: H-­‐H ¡
This shows why hydrogen exists in nature as a diatomic molecule instead of as a single atom ¡  Lewis Structure: ¡  If 2 chlorine atoms are paired together, they each get an octet ¡  Bonding Pairs? Lone Pairs? §  1 bonding pair, 6 lone pairs ¡  Structural Formula: ¡  This shows why halogens exist in nature as diatomic molecules ¡
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In Lewis theory, two atoms may share more than one electron pair to get octets. Ex. Oxygen §  We know oxygen exists as a diatomic molecule §  Lewis Structure: §  If we pair two oxygen atoms up, we don’t have enough electrons to give each O atom an octet §  We can take lone pair electrons and turn them into bonding pair electrons §  Each oxygen atom now has an octet because the additional bonding pair counts toward the octet of both oxygen atoms §  Structural formula: ¡  Double Covalent Bond: A bond in which two electron pairs are shared between two atoms ¡  In general, double bonds are shorter and stronger than single-­‐bonds ¡  Ex. The distance between oxygen nuclei in an oxygen-­‐oxygen double-­‐bond (O=O) is 121 pm §  In a single bond it is 148 pm ¡
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Triple Covalent Bond – A bond in which three electron pairs are shared between two atoms Ex. N2 §  Lewis Structure: §  In order to have enough electrons to satisfy the octet rule for both N atoms, we need to convert two additional lone pairs of electrons into bonding pairs. §  Structural Formula: ¡
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Triple bonds are even shorter and stronger than double bonds Because these bonds are so strong, Diatomic nitrogen is a fairly unreactive molecule in nature. Triple bonds have high bond dissociation energies. This is the energy needed to break a covalent bond. Single bonds < double bonds < triple bonds in regards to bond dissociation energies 1. Write the correct skeleton structure for the molecule §
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Atoms should be in the right positions Hydrogen atoms are always on the ends (terminal atoms) -­‐never in the middle of a molecule Molecules tend to be symmetrical if it contains several atoms of the same type 2. Calculate the total number of electrons for the Lewis structure by adding up the valence electrons for each atom in the molecule §
If writing the Lewis structure for a polyatomic ion, the charge of the ion must be considered when calculating the total number of electrons 3. Distribute the electrons among the atoms, giving octets (or duets for hydrogen) to as many atoms as possible §  Start by placing 2 electrons between each pair of atoms §  Then distribute the remaining electrons, first to the terminal atoms, then to the central atom, giving octets to as many atoms as possible 4. If any atoms lack an octet, form double or triple bonds as necessary to give them octets §  Do this by moving lone electron pairs from terminal atoms to bonding regions between atoms. 1. Correct Skeletal Structure O C O (symmetry) 2. Total number of electrons 6 + 4 + 6 = 16 e-­‐ 3. Distribute electrons 4. Form Double/Triple Bonds ¡  Sometimes atoms don’t share evenly (one for one.) ¡  Example: carbon monoxide (CO) ¡  Oxygen shares more electrons than carbon ¡  Once the sharing occurs, this bond is no different than any other covalent bond ¡  In these cases, concentrate on making sure the octet rule is reached, NOT on where the electrons came from. ¡  Follow the same 4 steps, but pay special attention to the charge of the ion when calculating the number of electrons in the Lewis structure §  Add 1 electron for each negative charge §  Subtract 1 electron for each positive charge ¡  We normally show the Lewis structure for a polyatomic ion within brackets and write the charge of the ion in the upper right corner 1. Skeleton Structure CN 2. # of valence electrons 4 + 5 + 1 = 10e-­‐ 3. Distribute Electrons 4. Double/Triple Bonds 5. Enclose in brackets with charge in the upper right corner ¡  When there is more than one correct Lewis structure for a covalent compound ¡  Ex) ozone (O3) ¡  The actual bonding is a hybrid of the different possible resonance structures that you draw. ¡
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Lewis theory is a simple theory, not sophisticated enough to be correct every time Ex. NO – 11 electrons §  Can exist as: (see next slide) ¡
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In these cases, we write the best Lewis structure we can Boron – tends to form compounds with only 6 electrons around the B instead of 8 §  Ex. BH3 SF6 and PCl5 have more than 8 electrons around the central atom in their Lewis structure §  We call these expanded octets ¡  Which of these molecules have atoms that have exceptions to the octet rule? Write the Lewis structures to help you determine the answer. §  PCl2F3 (hint: spread the halides around the central atom phosphorus—think symmetry!) §  NF3 §  SF4 §  SCl2 ¡  Which of these molecules have atoms that have exceptions to the octet rule? Write the Lewis structures to help you determine the answer. §  PCl2F3 §  NF3 §  SF4 §  SCl2 ¡  Sometimes 2-­‐D models do not show the entire story! ¡  Valence Shell Electron Pair Repulsion Theory §  According to this theory, the repulsion between electron pairs causes molecular shapes to adjust so that there is much distance as possible between valence electron pairs ¡  In this model, bonded atoms as well as lone pairs repel. There are several shapes. We will look at these in a lab activity. ¡  Polarity is partial charge (think back to biology, water is a polar molecule) ¡  Partial positive side, partial negative side ¡  Ionic compounds are not considered polar…
the electronegative force is so strong that there is transferring of electrons rather than sharing. ¡  Sharing that occurs between atoms with different electronegativity results in an unequal sharing of electrons. This is polarity. ¡  Forces BETWEEN molecules (not the ones holding the molecule itself together) ¡  These are weaker than covalent or ionic bonds, but must be discussed! §  Hydrogen bonds §  Van Der Waals forces ▪  Dipole interactions ▪  Dispersion forces ¡  These forces account for differences in molecule and compound properties. ¡
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Strongest of the IM forces Special to HYDROGEN only. Think back to biology—water is a “famously polar” molecule that experiences hydrogen bonding Bonds between H and a very electronegative atom (only N, O, or F) Results in a VERY polar bond (because of the small size of H) The H is then attracted to the unshared electrons of ANOTHER atom §  This could be part of another molecule (water molecules) §  This could be another part of the same molecule if the molecule is large enough (proteins, nucleic acids) ¡  Dipole interactions (aka Dipole-­‐Dipole) §  Occur between polar molecules that attract each other §  Similar to ionic bonding, but WEAKER! §  This sounds like hydrogen bonding! It’s not. This is for elements other than FON. ¡  This force gets weaker as the molecules move away from each other. ¡  Dispersion forces §  The weakest of all the intermolecular forces §  Caused by electron motion, even in non-­‐polar molecules. §  Due to momentary electron location §  Higher in molecules that have a lot of electrons §  F<Cl<Br (can be liquid at RT) < I (is solid at RT!) §  Sort of like dipole interactions, but momentary ¡
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Much less common than the simple molecular compounds This is not intermolecular forces…this is within the compound itself. Solids in which all of the atoms are covalently bonded to each other Diamonds are a network solid form of carbon. Vaporization point (solid to gas) is 3500 C!!! ¡  Check out this vid to review IM forces. ¡  http://youtu.be/oYd9lq3K9Ek ¡
Melting points are lower than ionic compounds §  Many are gas or liquid at room temp §  This is due to weak intermolecular forces ¡
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Generally poor conductors of heat and electricity Brittle or soft However…the more atoms in a molecule, the higher the MP…so really really big covalent molecules actually will have high MPs! These are called network solids ¡  Impurities in compounds cause MPs to decrease and BPs to increase. ¡  Like dissolves like! §  Polar things dissolve polar things §  non-­‐polar things dissolve non-­‐polar things The property chart will give you information about the MP, BP, volatility, solubility and conductivity of the following types of compounds. You have a hard copy and there is one on the website. Giant ionic Metallic Giant covalent Simple molecular Ionic compound vs. covalent compound definitions, properties and electron dot models ¡  Octet and Duet rules ¡  Covalent compounds ¡
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Octet exceptions Single/double/triple and bond dissociation energy Resonance Polarity Intermolecular forces (dipole, dispersion, hydrogen) Network solids VSEPR models