Chem 2202 Unit 2 Notes 2018 PDF

1 Unit 2: CHEMICAL BONDING IN MATTER Chemical Bonds: are attractive electrostatic forces that hold atoms or ions together in a substance. Molecular compounds have covalent bonds Ionic compounds have ionic bonds Metallic compounds have metallic bonds We will be looking at each type of chemical bo...

1 Unit 2: CHEMICAL BONDING IN MATTER Chemical Bonds: are attractive electrostatic forces that hold atoms or ions together in a substance. Molecular compounds have covalent bonds Ionic compounds have ionic bonds Metallic compounds have metallic bonds We will be looking at each type of chemical bond in detail in this unit. Bohr's Model of the Atom - protons and neutrons are located in the nucleus - electrons are located in orbits or specific energy levels outside the nucleus. Energy Level Diagrams - # of e- in an atom is equal to its atomic number - # of energy levels is equal to the period number - Maximum number of electrons in the first three energy levels is 2, 8, 8 - # of electrons in the outermost energy level is equal to the group number of the element. Examples: Draw Energy Level diagrams and energy level diagrams for the following atoms: sodium and chlorine Valence Energy Level: the outermost energy level of an atom that contains electrons. Valence Electrons: electrons that occupy the valence energy level. Note: The number of valence electrons of an element in Group A is equal to the group number in which it is found: Group # IA IIA IIIA IVA VA VIA VIIA VIIIA # of Valence 1 2 3 4 5 6 7 8 Electrons  2 Covalent Bonding: Covalent Bonds are the electrostatic force of attraction that occurs between 2 non-metallic elemental nuclei and a shared pair of electrons. Molecular compounds form from covalent bonds. In order to understand how bonding works in molecular compounds we use Electron Dot Diagrams (also known as Lewis structures) When two atoms approach each other their outermost energy levels (valence shells) come in contact first. Therefore, to show formation of any chemical bonds we need only show valence electrons. This eliminates the need to draw all electrons as in an Energy Level Diagram. This was proposed by Gilbert N. Lewis. To draw the Electron Dot Diagram for an element, (1) Write the symbol of the element. This represents the nucleus and non-valence electrons. (2) Use the periodic table to determine the number of valence electrons for the atom. (3) Place dots (representing electrons) in four areas (orbitals) around the atom (one on each side). DO NOT form pairs, until each of the four areas has at least one dot. Example: Draw an electron dot diagram for Sulfur: Bonding Electron (or 'unpaired electron') is a single electron in a valence orbital of an atom. Bonding electrons are always involved in bonding. Lone Pairs (non-bonding electrons) are electrons found as pairs in filled valence orbitals of an atom. They are rarely involved in bonding. Bonding Capacity is the maximum number of bonds an atom can form. It equals the number of bonding electrons in the atoms. Examples: Complete the following table: Symbol Dot Diagram # lone pairs # bonding bonding electrons capacity Li B C N F Ne  3 Lewis Dot Diagrams for Molecular Compounds Central Atoms are elements that are covalently bonded to more than one other element. They must have a bonding capacity more than 1, but that does not guarantee that the element will be a central atom! When you pair up 2 electrons in a molecule you have created a covalent bond! Example: HCl A covalent bond can be a single bond where one electron pair are shared, or a multiple bond - double or triple bonds. A double bond is when TWO electron pairs are shared between two atoms while a triple bond is when THREE electron pairs are shared between two atoms. Example: O2 and N2 To draw a dot diagram for any molecule: 1) Draw a dot diagram for each element. 2) Elements are bonded together by pairing up bonding electrons 3) Bond together all elements that have a bonding capacity higher than 1 first. These will be the possible central atoms for your molecule. 4) Count the number of unbonded electrons around your central atom. Subtract the number of unbonded elements you have left (ones with bonding capacity =1). a) If your answer is 3 you need a triple bond. Create this now. b) If your answer is 2 you need a double bond. Create this now. c) If your answer is 0 you do not need a multiple bond. 5) Bond the remaining elements to the central atoms by pairing electrons. You should have no unbonded electrons left. If there are any, you have made a mistake. 6) Count to ensure each atom has an octet - 8 electrons around it - the exceptions are hydrogen (2) & boron (6) Examples: Draw Electron Dot Diagrams Substance workings Electron Dot structural diagrams Diagrams H2O CH4  4 Substance Workings Electron Dot structural diagrams Diagrams CH2O C3H8 C2H2 CO2 CH3OH CH3NH2 Structural Diagrams: Drawing all these dots is convenient for determining what types of bonds are in your molecular compounds. However, these get very messy in larger molecules and do not allow you to truly see the overall bonding picture. To do this we draw structural diagrams. To draw a structural diagram from a dot diagram: a) Replace bonding pairs of electrons with a stick (-). Double bonds get 2 sticks while triple bonds get 3 sticks. b) All lone pair dots are just dropped. Draw structural diagrams for the dot diagrams in above table.  5 Molecular Shapes All molecules have a definite 3-D shape. Stereochemistry is a study of the shape of chemicals. VSEPR (Valence Shell Electron Pair Repulsion) is used to predict the shape of molecules. In order to use VSEPR theory we will use Electron dot diagrams to show how many and what types of covalent bonds the molecules have and use this information to draw a 3D diagram. 3D diagrams use different bond drawings to illustrate the three dimensions the molecule has. These new bonds are draw as follows: atom is in same plane as central atom atom goes behind the plane (away from you) atom goes in front of plane (towards you) One important thing to remember is that only central atoms can have a shape - if an element is only bonded to one other element it does not have a shape. Each central atom in a molecule has a shape, but they do not always have to be the same. You are also required to know the approximate angles bonds make in each shape. There are many different molecular shapes, but we only deal with 5 in this course: 1. Tetrahedral Molecule Lewis Dot Bp and LP around central Shape & atom Bond Angles CH4 2. Pyramidal NBr3  6 3. V-shaped or Bent Molecule Lewis Dot Bp and LP around central Shape & atom Bond Angles H2O 4. Trigonal Planar BF3 CH2O 5. Linear C2H2 Practice: For each of the following draw the Lewis Dot Diagram and the molecular shape. Indicate the shape around each central atom. Molecule Lewis Dot Bp and LP around central Shape & atom Bond Angles SiH2Cl2  7 Molecule Lewis Dot Bp and LP around central Shape & atom Bond Angles PH2Br SI2 C2H2Cl2 C2HBr CH3OH Electronegativity Draw an electron dot formula for HCl. The formula suggests that the pair of electrons, which constitutes a covalent bond, is shared equally between the hydrogen and chlorine. This is not the case. In hydrogen chloride, the chlorine atom exercises a stronger attractive force on the bonding pair than the hydrogen atom does. We say chlorine is more electronegative than hydrogen.  8 Electronegativity (EN) is the tendency of an atom to attract electrons to it when chemically combined with another atom. Linus Pauling proposed numerical values for the electronegativities of the elements. These values are located below the atomic number on your periodic table. As this value increases the element exerts a greater pull on electrons within a chemical bond. There are two periodic trends in Electronegativity: 1. Increase in EN from left to right within a period. 2. Decrease in EN from top to bottom within a group. *The two trends combine to give fluorine the highest EN at 4.0 and cesium the lowest EN at 0.7. Notes: 1) EN of metals is low. Metals have very little attraction for their own electrons or those of other atoms. They hold their valence electrons very loosely and will lose them easily to other atoms. 2) EN of nonmetals is high. Nonmetals have a strong attraction for their own electrons and those of other atoms. They hold their valence electrons very strongly and will even gain more from other atoms. Molecular Polarity: Bond Polarity is defined as the unequal sharing of electrons between two elements in a covalent bond. A non-polar covalent bond occurs when both atoms involved in a covalent bond have the same EN. The bonding electron pair is shared equally and is uniformly found between the nuclei of two atoms. EX: A polar covalent bond occurs when the two atoms involved in a covalent bond have different EN. The bonding electron pair is unequally shared between the two atoms. This causes an "electron shift" within the bond. EX: Notes: 1) The more electronegative atom will become partially negative (δ-). Since the shared electrons are pulled closer, it appears to have more electrons surrounding it than it has protons in its nucleus. 2) The least electronegative atom will become partially positive (δ+). Since the shared electrons are pulled slightly away, it appears to have less electrons surrounding it than it has protons in its nucleus. 3) These electrons are stilled shared; they are NOT exchanged - only shifted. 4) Polar Covalent Bonds have bond dipoles (a δ+ end and a δ+ end) but are still electrically neutral. 5) Bond dipoles are represented by an arrow with the arrowhead pointing towards the partially negative site.  9 6) The greater the difference in electronegativity between two nuclei in a covalent bond the greater the polarity in the bond. Ex: Molecular Polarity is when the bond dipoles of polar bonds within the molecule interact to create an apparent positive side and a negative side to the molecule in 3 dimensions. A molecule is non-polar if all bond dipoles cancel geometrically. This is similar to vector addition in physics - if the bond dipole arrows "cancel out" there will be no positive and negative sides to the molecule. A molecule is polar if all bond dipoles do not cancel. This creates the positive and negative sides to your molecule you need for molecular polarity. Examples: Draw electron dot and 3d structures for the following and use them to determine if the molecules below are polar or non-polar: Substance Lewis dot Molecular Shape Notes BP& LP Draw dipoles CH4 CH3Br NF3 NCl3 H2O SI2  10 BF3 C2H2F2 C2H2 HCN It is very time consuming to draw diagrams to decide if a molecule is polar or non-polar. However, only tetrahedral, trigonal planar and linear molecules have the geometry available to cancel out dipoles. Therefore, it is possible to follow 2 simple rules to determine molecular polarity: (1) If a molecule contains C or Si - the molecule is polar if it contains more than 2 elements. (2) For all other molecules - the molecule is polar unless all electronegativities are the same Molecule Shape & Polarity Molecule Shape & Polarity CH4 SiH3Cl NCl3 PBr3 OF2 SI2 PH2Cl C2H4  11 Si2Cl2 HCN Solubility of Molecular Compounds In unit 1 we looked at solubility of ionic compounds (solubility table on back of periodic table). In Science 1206, you learned that molecular substances may or may not dissolve in water. For molecular compounds solubility in any solvent depends on polarity. General Rule: like dissolves like: 1) polar substances dissolve in polar solvents 2) non-polar substances dissolve in non-polar solvents 3) polar and non-polar substances do not dissolve in each other. In the table above, which substances can dissolve in water? Write yes or no by the compounds Why do water and oil (a hydrocarbon) not mix? Water is polar and oil is non-polar, therefore difference polarities will not dissolve in each other. How does soap work? Soap breaks up the oil into smaller drops, which can mix with the water. It works because soap is made up of molecules with two very different ends. One end of soap molecules love water - they are hydrophilic. The other end of soap molecules hate water - they are hydrophobic. Hydrophobic ends of soap molecule all attach to the oil. Hydrophilic ends stick out into the water. This causes a drop of oil to form: These drops of oil are suspended in the water. This is how soap cleans your hands - it causes drops of grease and dirt to be pulled off your hands and suspended in water. These drops are washed away when you rinse your hands. Melting/Boiling Points When a substance is heated, the particles that make up the substance move faster and faster. If the particles move fast enough they start to slide around each other (liquid) or fly apart (gas). At what temperature this occurs depends on how tightly the particles are held together. For substances that are solids at room temperature, the particles are held together tightly. For gasses the particles are not held together very strongly. The melting/boiling points of chemicals will depend on the type of compound it is. In molecules, the atoms are held together tightly by covalent bonds. However, each molecule is not directly connected to its neighbouring molecule. These neighbouring molecules are instead held together by intermolecular forces  12 There are two types of Intermolecular Forces: 1. van der Waals Forces - involve dipole interactions a) London Dispersion Forces (LDF) b) dipole-dipole (D-D) 2. Hydrogen Bonding (HB) London Dispersion Forces - Electrons are always moving around the atoms involved in a molecule. This creates fluctuating dipoles in this molecule. These dipoles can then induce opposite dipoles in a neighbouring molecule creating a small force of attraction. This is constantly occurring in all molecular compounds holding the individual particles together. These are also the forces of attraction that hold together the noble gasses (albeit weakly) Diagram: Factors Affecting the Strength of LDF 1) # of electrons in a molecule - the more electrons a substance has, the more easily an uneven distribution of charge can occur resulting in stronger LDF. 2) Shape of the molecule - the more spherical the shape, the less surface area there is giving less opportunity to induce a charge on a nearby molecule. This gives weaker LDF (more complicated molecules have higher LDF) Q. Compare the boiling points of the halogens with the number of electrons each has. Answer. Halogen Boiling Point (̊C) # of Electrons fluorine, F2 -188 chlorine, Cl2 -34.6 bromine, Br2 58.8 iodine, I2 184  13 Q. Butane, C4H10, has a much higher boiling point than Cl2, thus must have stronger intermolecular forces. Give a reason for the difference in boiling points. Note: Molecules that have the same number of electrons are said to be isoelectronic. Dipole-Dipole Forces: is the simultaneous attraction of a molecular dipole by the surrounding molecular dipoles of opposite charge - occurs in polar molecules only Diagram: Question 1: Explain which of the following compounds, F2 or CH3F, has the higher boiling point.  14 Question 2: Which of the following substances, ICl or CCl4, has the higher boiling point? Question 3: Which of the following substances, I2 or IBr, has the higher boiling point? Note: LDF are a stronger force than dipole-dipole. Question 5: Predict the boiling point of HF given the following data. Hydrogen Halide Number of Electrons Boiling Point (C) HF 10 ? HCl 18 -83.7 HBr 36 -67.0 HI 54 -35.4 Answer: Based on the number of electrons, strength of LDF, one would predict that HF has a lower boiling point than -83.7C. It in fact has the highest boiling point of all at 19.4C. Question: WHY NOT WITH HF?  15 This suggests the existence of an additional intermolecular force greater than van der Waals forces. --- Hydrogen Bonding. Question. So, what do HF, NH3 and H2O have in common? Answer: They each contain a highly polar bond O - H, N - H, H - F, with the positive end being hydrogen. Hydrogen Bonding: simultaneous attraction of a proton (very partially positive hydrogen) by the electron pairs of adjacent N, O, or F atoms Diagram: Hydrogen bonding is represented using dotted lines (represents a weaker bond type than a covalent bond). The hydrogen bonding in water when combined with its shape give it some unique properties: Example: Water (ICE) 1) snowflakes are unique with geometric shapes (usually 6-sided) 2) Water has the hydrogen bonding and the water molecule has a v- shaped, when six water molecules combine to give a hexagonal shape a hole is crated inside the molecule causing ice crystals to take up more space (increase volume) and a lower density which causes ice to float!  16 Question. Which of the following compounds exhibit H-bonding: H2O, CH4, CHF3, NF3, CH3OH? Answer. Draw structural diagrams for each. H2O CH4 CHF3 NF3 CH3OH Properties of Hydrogen Bonded Substances Hydrogen bonding affects both physical and chemical properties of substances. 1. increased melting and boiling points Pair Substance Boiling Point # of hydrogen (̊C) electrons bonding? 1 HF 19.4 10 yes HCl -83.7 18 no 2 H2O 100 10 yes H2S -61 18 no 3 NH3 -40 10 yes PH3 -90 18 no 2. increased solubility of substances mutually involving hydrogen bonding. EX: water, H2O; methanol, CH3OH; and ethanol, C2H5OH are soluble in each other in all proportions. They all exhibit H-bonding. Q: Which of these substances CH3F or CH3OH are more soluble in water? 3. shape and stability of certain structures. Water's interesting properties are discussed above, but the behaviour and stability of protein molecules within the body depends on their shape which is affected by the hydrogen bonding present. DNA, protein synthesis and enzyme behavior all depend (and require) hydrogen bonding!  17 Question 1: List the following substances, C3H8, CH3Cl, C2H5OH in order of increasing boiling point? Question 2: Which substance, C4H10 or C2H5OH has the highest boiling? Question 3: Why is it difficult to determine whether C9H20 or CH3OH has the higher boiling point? Aside: One would always predict the H-bonded substance as having the higher boiling point unless there are way, way more electrons (say 50+ more) Bonding of Molecular compounds - In Science 1206 you learned that molecular Compounds are formed when valence electrons are shared between non-metallic elements. Theses sharing of electrons are called covalent bonds (Intramolecular forces). Properties of Molecular Compounds 1. Can be solids, liquids, or gases at room temperature 2. Are not generally hard  18 3. Relatively low melting and boiling points. 4. Solubility in H2O: varies from high to low 5. Conductivity: poor to none in all states. Explaining properties of molecular compounds Q: Explain in terms of bonding present why molecular compounds are soft and have low melting points. Answer: The individual molecules of a molecular compound are held together by weak intermolecular forces. Large amounts of energy are not needed to break these forces making them soft and causing them to have low melting points. Q: Explain in terms of bonding present why some molecular compounds are soluble in water and others are not. Answer: Solubility in polar water will depend on the polarity of the molecular compound. If the molecule is polar, then the molecular compound will be soluble in water. If the molecule is non-polar, then it will not be soluble in water. Like dissolves like. Q: Explain in terms of bonding present why molecular compounds do not conduct electricity in either state. Answer: Conducting electricity depends upon the movement of electrons. In order for electrons to be able to move through a substance there must be charged particles able to change position. Molecular compounds do not conduct in either state because there are no charged particles available to move about. There are no ions present, and all electrons are held tightly in a fixed position due to covalent bonding which involves sharing electrons for purpose of the stability of the atoms. These shared electrons are held tightly in orbitals and cannot leave that region of space. However, some molecular compounds have the ability to conduct small amounts (ex: graphite), but these are usually compounds of metalloids. Network Covalent Bonding - Recall that ionic compounds form crystals that are a 3-D array of positive and negative ions in a fixed ratio. - So far we have only looked at covalently bonded atoms as they exist in individual molecules. However, it since certain atoms have a bonding capacity over 2 it is possible for these atoms to form covalent bonds in a giant 3D array as well. A Network Solid is a giant rigid structure in which non-metallic atoms are covalently bonded together in a continuous 3-D array. These covalent bonds are very strong and very directional in a 3-D network forming a single very large molecule called a macromolecule.  19 Examples (YOU MUST MEMORIZE THESE 3 CHEMICALS): 1) diamonds (Cn) Silicon carbide(SiC) Silicon dioxide (SiO2) Graphite (C) Network covalent bonding is the simultaneous attraction of every atom to adjacent atoms by covalent bonds, within a 3-D lattice of atoms. Properties of network solids 1) Strong and hard 2) High melting points 3) Not soluble on water 4) Do not conduct electricity 1) Explain in terms of bonding present why network covalent solids are hard and have very high melting points. Answer: Network covalent solids are hard and have high melting points due to each atom, within the network solid, being bound in 3 dimensions to four other atoms, and so on. This gives a tight 3-D network where each atom has a tetrahedral arrangement about it. All atoms are held in rigid positions, thus making it very difficult to break apart. 2) Explain in terms of the bonding present why network covalent solids do not conduct electricity. Network solids cannot conduct electricity in any state because every atom is covalently bonded to every other atom in a tight compacted 3D array which involves sharing electrons for purpose of the stability of the atoms. Those strong 3D covalent bonds do not allow network solids to conduct electricity because the electrons are held tightly in a fixed position and cannot move freely within the network solids. 3) Explain, why diamonds are the hardest substances on Earth. In diamond, the electrons are shared in 3 dimensions with four other carbon atoms to form very strong covalent bonds resulting in an extremely rigid tetrahedral crystal. It is this simple, tightly-bonded 3D arrangement that makes diamond one of the hardest substances on Earth. 4) Why are network covalent solids insoluble in water? Covalent network solids are bonded by strong covalent forces into a three dimensional network where every atom is linked (bonded) to every other atom in a tight compacted 3D array of covalent bonds. There are no possible attractions that could occur between solvent molecules (water) and carbon atoms which could outweigh the attractions between the covalently bonded carbon atoms.  20 Bonding in Ionic Compounds - In Science 1206 you learned that Ionic Compounds are formed when valence electrons are transferred from a metal atom to a nonmetal atom creating ions of opposite charge. These ions are simultaneously attracted to each other through ionic bonds. Cations: - are formed when metals lose one or more valence electrons. Anions: - are formed when nonmetals gain one or more valence electrons. - Ionic bond formation can be shown using the energy level diagrams learned in 1206. The valence electrons are transferred so that the ion has the same electron configuration as the nearest noble gas: Examples: Use Bohr diagrams and energy level diagrams to explain the bonding between sodium and chlorine atoms. - We can also illustrate the ionic bond formation using Lewis dot diagrams, which will be the focus for our course: Ex1: Draw Lewis dot diagrams for the following ionic compounds A) Sodium chloride B) calcium chloride C) magnesium oxide Ex2: Write a Lewis equation for the reaction between each pair of elements. a) Sodium and Oxygen b) Magnesium and nitrogen  21 You could also look at ionic bonds in terms of electronegativity as well - metals have very low electronegativity and therefore do not have a very strong attachment to their valence electrons. Non-metals have high electronegativity and therefore have a strong attachment to valence electrons. Therefore the electron pairs are no longer shared - the non-metals basically "steal" the metal's electrons. Properties of Ionic Compounds: 1. Solids at room temperature. 2. Generally have high hardness and high melting/boiling points. 3. All are soluble to some degree in water. Most have high solubility in H2O, but it does vary. 4. Ionic compounds are brittle when they are solids. 5. Good conductors in liquid state and aqueous solution, but do not conductor in solid state. All of the above properties can be explained using the concept of ionic bonding. 1-2: State, Hardness, Melting Point - To break apart an ionic compound, you have to separate the individual positive and negative ions. - Each ion has several strong ionic bonds to its neighbours - therefore it will require a large amount of energy to break all these bonds making it hard. - Physical energy is not enough to break the bonds which accounts for its hardness and state. - Large amount of heat energy needed to get ions moving fast enough to break the ionic bonds which accounts for the high melting point and boiling point. 3. Why Ionic compounds have are Solubility in Water? Water is a polar molecule meaning it has a partially positive side and a partially negative side. The partially positive (hydrogen) end of the water will be attracted to the negative ions in the ionic crystal while the partially negative (oxygen) end of the water will be attracted to the positive ions in the ionic crystal. These attractions pull the ions away from the crystal. Once in solution, each ion is surrounded by a layer of water molecules. - Ion-Dipole Forces: An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole. Most commonly found in solutions. Especially important for solutions of ionic compounds in polar liquids. A positive ion (cation) attracts the partially negative end of a neutral polar molecule.  22 A negative ion (anion) attracts the partially positive end of a neutral polar molecule. Ion-dipole attractions become stronger as either the charge on the ion increases, or as the magnitude of the dipole of the polar molecule increases. It is important to note that in a homogenous solution the ions and water molecules are evenly dispersed at the particle level. Diagram: 4. Brittle (break easily when bent). - An ionic compound contains ionic bonds which are made up of positive and negative ions. When an ionic crystal is bent or hit, the positive and negative ions in the bond are forced away from ions of opposite charge and closer to ions with the same charge. This shift results in same charges coming in contact with each other creating a repulsive force resulting in the ionic crystal breaks into smaller pieces.  23 5. Conductivity - As discussed before, electricity is the flow of charged particles. In order for a substance to conduct, there must be movement of charged particles (electrons, ions) both freely and independently of each other. - In ionic solids, the ions are held in a fixed position within the individual ionic bonds thus the ions are not free to move. Therefore, ions cannot move beyond vibrating in place (cannot change position) - In a liquid state or aqueous solution, the ionic bonds have been broken into its ions. These ions can now move freely and change position carrying charge particles through the solution. Metallic Bonding There are three things true for all metals: 1) They have low electronegativity - they do not hold on to their valence electrons strongly. 2) They have empty valence orbitals - they have completely empty sides in their Lewis Diagrams 3) The loosely held valence electrons are free to move from one valence orbital to the next. The atoms become positive ions because the valence electrons are dispersed or delocalized. The positive ions remain fixed, but the electrons move and are simultaneously attracted to more than one nucleus. A sample of a metal can be thought of as a group of cations in a "sea" of free moving valence electrons. Diagram: Metallic Bond: is the simultaneous attraction of metal cations for the free moving valence electrons.  24 Ex: Draw a metallic bonding for AuCu where you have 70% Au and 30% Cu Properties of Metals: 1) shiny, malleable, ductile 2) conducts in both solid and liquid 3) range of melting points Explaining Properties of Metals 1. Malleable and Ductile Malleable: Able to bend without breaking Ductile: Able to be hammered into thin sheets - When a force is applied to a piece of metal, the positive ions are forced to move. As the metal ions move, the free moving valence electrons move too resulting in no repulsion within the metallic bonds. The cations slide past each other with the 'sea' of electrons acting as a lubricant. Thus, the metallic bonds are not fixed. Bonding in metals is not rigid. As a metal is struck by a hammer, the cations slide through the electron sea to new positions while continuing to maintain their connections to each other. The same ability to reorganize explains why metals can be pulled into long, thin wires. 2. Good Conductors. - Metals are good conductors of electricity because they contain charged particles (valence electrons) that can move freely around in them. In metals the electrons can wander about between the metal ions, these electrons are called free electrons (valence electrons). The more valence electrons there are in a solid the better it will conduct electricity. If we now think of the metal as a wire as shown in the diagram then the tiny electrons show no particular direction of movement  25 But if we now connect a battery to the two ends of the wire the electrons drift down the wire in one direction 3. Metals have a Range of Melting/Boiling Points. - The more valence electrons freely moving about, the larger the number of electrostatic attractions between stationary positive ions and valence electrons, the higher the melting point -Sodium only has one valence electron to contribute to the sea of electrons and melts at 97.8oC. -Magnesium has two valence electrons to contribute to the sea of electrons and melts at 649oC 4. Metals form Alloys. - Alloys are solid solutions of two or more metals. The valence electrons of one metal are simultaneously attracted by another metal. The electrons are free to move from one metal to another regardless of type. The study of metals and how they interact with each other is called Metallurgy. Relative Strength of Chemical Bonds - The different types of chemical bonds we have studied in this unit have different strengths. The stronger the particles are held together, the higher the melting/boiling points. Generally, Network covalent bonds are the strongest, ionic bonds are second with metallic bonds being the weakest. However, you have to remember that covalent bonds are only holding the particles together in Network Covalent solids. In molecular compounds the particles are held together by weak intermolecular forces. This means that when comparing chemicals of different types: Network Solids have the highest boiling points Ionic Compounds Metallic Compounds Molecular Compounds have the lowest boiling points. Remember that when comparing different molecular compounds, we look at the IMF's present. You are not responsible for comparing ionic compounds to other ionic compounds (or metals to metals) Ex 1: List the following in order of decreasing boiling point. a) Na, SiO2 , C3H8 , NaCl , C2H5OH , OF2 25  26 b) Au, MgCl2, H2O, H2S, Cgraphite c) Ca, KF, SiC, Ar 2: Explain which will have the higher boiling point - C2H5OH or NaCl? 26