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Chemistry 20 – Unit A – Diversity of Matter and Chemical Bonding Section 3.1 structural formula valence electron quantum mechanics orbital valence orbital bonding electron lone pair octet rule Lewis symbol electronegativity covalent bond ionic bond metallic bond Section 3.2 bonding capacity coordinate covalent bond empirical formula molecular formula Lewis formula structural formula stereochemical formula polar covalent bond bond dipole Section 3.3 stereochemistry VSEPR theory polar molecule non-polar molecule polar covalent bond bond dipole Section 3.4 intermolecular bond intramolecular bond van der Waals force dipole-dipole force London force momentary dipole isoelectronic molecule hydrogen bond Section 3.5 crystal lattice covalent network 1 Chapter 3–Understanding Chemical Compounds- Chemical Bonding: page 76 Bonding Theory: A Little History Bonding theory begins prior to Mendeleyev’s periodic table Edward Frankland (1852) -stated each atom had a certain bond capacity Friedrich Kekule (1858) – first to use a dash between atom to illustrate a bond (structural formulas) Jacobus van’t Hoff and Joseph Le Bel (1874) – extended structural formulas to 3 dimensions. Richard Abegg (1904) – theorized that bonding capacity was related to the atoms electronic structure. Also suggested that bonding occurred such that an atom could obtain a “noble gas like” structure. He was the first to bring up the idea of an electron transfer due a chemical reaction to achieve this stability. Atoms would then be held together by having opposite charges. (Ionic bond) Gilbert Lewis (1916)- proposed that atoms could achieve a “noble gas like” stable structure by sharing electrons as well as by transferring them. The electron involved were called valence electrons. (Covalent bond) Linus Pauling (1939)explained why certain electron arrangement were more stable covering both ionic and covalent bonding. Section 3.1 Bonding Theory and Lewis Formulas: energy level: a specific energy an electron can have in an atom or ion orbital: region of space around an atom’s nucleus where an electron may exist. An orbital is not small particles in orbit but may be thought as a 3-D space where the electron might be. valence orbital: volume of space that can only be occupied by electrons in an atom’s highest energy level valence electron: electrons that occupies the outermost energy level of an atom or ion. The number of valence electrons can be determined by referring to the group number (Roman Numeral ‘A’) on the periodic table. 1. Determine the number of valence electrons for the following atoms or ions: a. chlorine f. helium b. strontium g. carbon c. aluminum h. sulfur d. lithium i. krypton e. nitrogen j. chloride 2 Lewis Symbols: The valence electrons are shown as dots around the central symbol, with the entire diagram showing (for an atom) a net charge of zero. Drawing Lewis symbols for main group atoms: Write the element symbol to represent the nucleus and filled energy levels of the atom Add a dot to represent each valence electron Start by placing valence electrons singly into each of four valence orbitals (represented by the four sides of the element symbol). Element can only have 4 orbitals. The only exception is H and He which have only one orbital. Each orbital must be occupied before electrons can be paired up. (Each orbital can only hold a max of 2 e-. More than 2 result in too high of a repulsion force If additional locations are required for electrons once each orbital is half-filled; start filling each of the four orbitals with a second electron until up to eight valence electrons have been represented by dots. (known as the octet rule) Period 2 elements of the periodic table: Group Number Atom Number of valence electrons Li Be B C N O F Ne Lewis Symbol An atom with a valence orbital occupied by a single electron can theoretically share that electron with another atom. Such an electron is, therefore, called a bonding electron. ( ) A full valence orbital, occupied by two electrons, has a repelling effect on electrons in any nearby orbitals. Two electrons occupying the same orbital are called a lone pair. ( ) Number of bonding electrons Number of lone pairs 3 2. Draw a Lewis Symbol for each of the following: a. iodine (I) f. argon (Ar) b. barium (Ba) g. tin (Sn) c. boron (B) h. selenium (Se) d. bismuth (Bi) i. chloride (Cl1-) e. helium (He) j. sulfide (S2-) 3. Complete the table for the following atoms: Atom Valence electrons Lewis Symbol Bonding electrons Lone Pairs K Ca B Si N O F Ar Rb As P C 4 4. Draw the Lewis Symbol of the following atoms and indicate how many bonding electrons are present: a bromine e silicon b lithium f oxygen g polonium h germanium c magnesium d bismuth Electronegativity: page 81 atoms exhibit (have) different abilities to attract negative valence electrons the farther away from the nucleus that negative electrons are, the weaker their attraction to the nucleus (which is positive due to the presence of protons) inner electrons (those closer to the nucleus) shield the valence electrons from the attraction of the positive nucleus the greater the number of positive protons in the nucleus, the greater the attraction for negative electrons combining these three points, it was possible to assign a value to any atom the term electronegativity is used to describe the relative ability of an atom to attract a pair of bonding electrons in its valence level (energy level) electronegativity is usually assigned on a developed scale – according to this scale, fluorine has been given the highest assigned electronegativity of 4.0 – cesium has the lowest assigned electronegativity of 0.8 metals tend to have low electronegativities non-metals tend to have high electronegativities 5. Using your periodic table, determine the electronegativity of the element a-h of table for question 4. a. ____ b. ____ c. ____ d. ____ e. ____ f. ____ g. ____ h. ____ 5 Bonding: page 82 A chemical bond is: Covalent Bonding: page 82 According to Science 10: bond formed when non-metallic atoms share electrons; atoms in a molecule are bound together by covalent bonds A covalent bond is: The key ideas of the Lewis theory of bonding are: Atoms and ions are stable if they have a noble gas-like electron structure; i.e. a stable octet of electrons. Electrons are most stable when they are paired. Atoms form chemical bonds to achieve a stable octet of electrons. A stable octet may be achieved by an exchange of electrons between metallic and nonmetallic atoms. A stable octet of electrons may be achieved by the sharing of electrons between nonmetallic atoms. The sharing of single valence electrons results in a covalent bond. 6 Ionic Bonds: page 83 Table salt is made up of two elements that can be very dangerous, namely, sodium and chlorine. Combined they form a safe compound. Sodium and chlorine combine during a reaction that will transfer one valence electron from sodium to chlorine. After the transfer sodium and chlorine are no longer neutral. The sodium and chlorine have becomes ions (atoms with unequal number of protons and electrons). The chlorine now has a net charge of -1 and the sodium ion has a charge of +1. Since objects of opposite charge attract each other, the two ions will be held tightly together. This type of attraction is called an ionic bond. In a crystal of sodium chloride, each ion will have six neighboring ions of opposite charge. This makes ionic bonding a very strong type of interaction between atoms. These strong bonded ionic compounds have a high melting point and boiling point. Ionic bonding type of bond formed when electrons transfer between metals and nonmetals electrostatic attraction between oppositely charged ions (positive and negative ions) in the crystal lattice of a salt metals ---------- lose electrons ---------- change into a positive ion nonmetals ---------- gain electrons ---------- change into a negative ion due to unlike charges being present : positive ion ------ is attracted by the ----- negative ion Bonding of NaCl: Lattice Structure and Empirical Formulas 7 Metallic Bonding: page 83 Summary: Bonding Theory (page 84) the formation of a chemical bond involves competition for bonding electrons occupying valence orbitals. if the competing atoms have equal electronegativities, the electrons are shared equally. electrons sharing between atoms with high electronegativities results in covalent bonding. electrons sharing between atoms with low electronegativites, often results in metallic bonding. if the competing atoms have unequal electronegativities, the electrons are unequal covalent bonding; if unequal enough that electrons transfer, the result will be ionic bonding. bonding theory was created by chemists to describe, explain, and predict natural events and observed properties. 8 Section 3.2 Explaining Molecular Formulas: page 85 Molecular Elements: page 85 hydrogen H2 bromine Br2 nitrogen N2 iodine I2 oxygen O2 astatine At2 fluorine F2 phosphorus P4 chlorine Cl2 sulfur S8 Molecular Compounds: page 86 Table from page 87: Bonding Capacities of Some Common Atoms Atom Number of valence electrons Lewis Symbol Number of bonding electrons Bonding capacity Electronegativity carbon nitrogen oxygen fluorine hydrogen 9 Covalent bonds occur when there is a sharing of single valence (bonding) electrons between two nonmetallic atoms in order to form an octet (8) of valence (bonding) electrons (exception to this rule is hydrogen, which will have two valence (bonding) electrons). single covalent bond – is formed when 1 pair of single valence (bonding) electrons is shared double covalent bond - is formed when 2 pairs of single valence (bonding) electrons are shared triple covalent bond - is formed when 3 pairs of single valence (bonding) electrons are shared coordinate covalent bond – a covalent bond in which one of the atoms donates both electrons . hydrogen gas, H2(g) oxygen gas, O2(g) nitrogen gas, N2(g) Types of Formulas: page 87 Ionic Compound Formulas: Lewis formula for ionic compounds: sodium chloride potassium fluoride calcium chloride barium oxide 10 Molecular Compound Formulas: Table 2 Names of Types of Formulas for Molecular Compounds page 88 Example: acetic acid (CH3COOH) Empirical Formula An empirical formula shows the simplest whole-number ratio of atoms in the compound. Empirical formulas are rarely useful for molecular compounds. Molecular Formula A molecular formula shows the actual number of atoms that are covalently bonded to make up each molecule. A molecular formula often has the atom symbols written in a sequence that helps you determine which atoms are bonded to which. Lewis Formula * A Lewis formula is also commonly called a Lewis diagram, or an electron dot diagram. It uses Lewis symbols to show electron sharing in covalent bonds, electron transfer in ionic bonds, and the formation of stable valence octets of electrons in molecules. Structural Formula * A structural formula is also commonly called a structural diagram. As well as showing which atoms are bonded, the type of covalent bond is represented by the number of lines drawn between atomic symbols. Stereochemical A stereochemical formula is a structural formula drawn to try to represent the Formula three-dimensional molecular shape. For much larger molecules, this style of representation often becomes too complex to be practical. * It is strongly recommended that you memorize this terminology information. 11 Determining Lewis Formulas: page 89 Lewis formula predictions are limited to entities with only one central atom. The central atom is the atom to which all the other atoms – peripheral atoms – are bonded to. Simple Lewis formulas can be predicted with a series of five steps. Determine the Lewis Formula and the structural formula for _____________________, ________________________ and _________________________. 1. Count the total valence electrons in the entity by adding the valence electrons of each atom. If the entity is a polyatomic ion, add (usually) or subtract valence electrons to account for the net charge, one for each unit of charge. 2. Arrange the peripheral atom symbols around the central atom symbol, and place one pair of valence electrons ( ) between each peripheral atom and the central atom (bond pairs). 3. Place more pairs of valence electrons (lone pairs) on all the peripheral atoms, to complete their octets. Recall that a hydrogen atom’s energy level is completed with only two valence electrons. 4. Place any remaining valence electrons on the central atom as lone pairs. 5. If the central atom’s octet is not complete, move a lone pair from a peripheral atom to a new position between that peripheral atom and the central atom. Repeat until the central atom has a complete octet. For a molecule, this completes the Lewis formula. If the entity is a polyatomic ion, place square brackets around the entire Lewis formula, and then write the net charge outside the bracket on the upper right. To show the structural formula, omit all lone pairs and replace every bond pair with a line. Example 1 Example 2 Example 3 E.g. Draw the Lewis diagram for the carbonate ion 12 Draw Lewis diagram for CO321. total e- = C( 4 ) + 3O ( 3 x 6 ) + 2- ( 2 ) = 24 e- O 2. C O 3. O 4. C O O O O 6. O C O O C O 5. O 2- C O O O Resonance structure : When 2 or more Lewis diagram can describe a formula, these diagrams are called resonance structures of that formula. The three resonance structure of CO32- is given below O O 2- O C C O 2- O O 2- C O O O Draw Lewis Formulas and Structural Formulas for the following 13 fluorine gas, F2(g) hydrogen fluoride oxygen difluoride water ammonia methane sulfur hexafluoride phosphine, PH3(g) NH3BH3 5. Draw a Lewis formula of the following molecules: Lewis Symbol Lewis Symbol a HBr(g) b carbon disulfide c HOCl d sulfur dichloride 14 e H2Se(g) f NH2Br g CH3F h carbon dioxide i Silicon tetrafluoride j HI(g) k PBrIF l diiodine sulfide m CH3SH o nitrogen tribromide 15 Draw a Lewis formula to represent the following ions: Write the formula of the ion before drawing the Lewis formula. a. hydrogen ion b. nitrate ion c. hydroxide ion d. carbonate ion e. hydrogen sulfide ion f. hydrogen carbonate ion* * Hint: A hydrogen ion (proton) bonds to any lone pairs from answer (d). 16 Rules for Drawing Lewis Symbols For Molecules That Have More Than One Central Atom. Step 1 Determine the atom with the largest bonding capacity. Use The group number “Roman Numeral” to determine bonding capacity. This is the central atom. Step 2 Draw the central atom. Draw the valence electrons around the central atom. Unpaired (bonding) electrons are available for bonding Step 3 All other atoms are drawn around the central atom pairing up bonding electrons such that each atom is surrounded by an octet of electrons (except hydrogen) Step 4 Work though the chemical formula from left to right ensuring that all atom groups remain intact in the Lewis model. Step 5 If any electrons are left unpaired after all atoms have been drawn, determine if a double bond (sharing of two pairs of electron) or a triple bond (three pairs of electrons) is required to complete the diagram Example a. CH3Cl b. N2Cl2Br2 c. C2H5OH d. CH3OCH3 e. CH2O CH3COCH2OH Draw the Lewis Formula and Structural Formula of the following: a. Lewis Symbol / Structural Formula hydrogen peroxide b. Lewis Symbol / Structural Formula O2F2 17 c. ethane d. C2H4 e. ethanol f. methanethiol CH3SH(g) g. methanol h. C2H2 i. N2H3OF j. HCHO k. CH3NHCH3 l. CH3CHOHCH3 n. N2H4 m. CH3OCH3 18 o. CH2FCH2I p. propane q arsenic trihydride r PHO Structural Formulas: Color Atom represented hydrogen oxygen nitrogen carbon halogens sulfur white ball red ball orange ball black ball green ball yellow ball Bonding capacity forms 1 bond forms 2 bonds forms 3 bonds forms 4 bonds forms 1 bond forms 2 bonds Using a molecular model kit – construct models of each of the following molecules and draw their structural formula. O2 Cl2 HCl H2O NH3 CH4 H2O2 N2H4 19 C2H6 CH3OH C2H5OH CH3OCH3 C2H5OCH3 C2H4 CO2 H2CO Bonding Capacity 1 Example Chlorine 2 Oxygen 3 Nitrogen 4 Carbon Possible combinations Draw the structural formula of the following without the aid of the molecular model kit. 1 CH3NHCH3 2 CH3CH2OCOCHO 20 3 C2H5CCNH2 4 C2H2 5 CH3OH 6 CH3CHOHCONH2 7 CH3COOC2H5 8 CH3CCCH3 9 HOF 10 C2H5COOH 11 C2H4 12 C4H10 13 HCHO 14 CS2 15 C2H5CONH2 16 C2H5COCH3 21 17 HNO 18 CH3CHOHCONH2 19 C3H4 20` PCl3 21 CH3CH2COONH2 22 BrCP 23 CHONHOH 24 CO2 25 C3H7OCH3 26 HCC2CH 27 CO(NH2)2 28 C2H2Cl2 29 CH2CHCH2OH 30 CH2COHOCH3 22 Predicting molecular formulas of nonmetallic compounds Write a partial chemical equation show only the reactants.(ie. reactant + reactant ) Predict the formulas of the following products using the fewest number of atoms and the least number of double or triple bonds. Show the Lewis formula of each type of atom involved and then build Lewis formula of the simplest molecule. From the Lewis formula draw the structural formula. e.g. #1. oxygen gas + fluorine gas e.g. #2. phosphorus + hydrogen gas 1. oxygen gas + hydrogen gas 2. iodine + bromine 3. phosphorus + chlorine gas 4. oxygen gas + chlorine gas 5. carbon + sulfur 6. sulfur + oxygen gas 23 Section 3.3 Molecular Shapes and Dipoles page 91 Valence shell electron pair repulsion theory ( VSEPR theory ) The shapes of molecules are very important because many of their physical and chemical properties depend upon the three dimensional arrangements of their atoms. e.g , The functioning of enzymes, which are substances that control how fast biochemical reactions occur, requires that there be a very precise fit between one molecule and another. Enzymes are large molecule of protein, because of their shape will react only with a specific molecule much like a key will fit into a lock. Even slight alterations in molecular geometry can destroy this fit and deactivate the enzyme, which in turn prevents the biochemical reaction involved from occurring. Very difficult to predict shapes of large molecules 1957 a simpler theory was developed to help explain and predict the stereochemistry of certain chemical elements and compounds. Called VSEPR theory. Valence shell electron pair repulsion theory ( VSEPR ) is the best theoretical explanations of molecular shapes. The general theory is based on the idea that valence shell electron pairs (bonding electrons and nonbonding lone pairs), being negatively charged, stay as far apart from each other as possible so that the repulsion between them are at a minimum. Stereochemistry: is the study of the 3-D spatial configuration of molecules and how this affects reactions. According to VSEPR theory: only the valence electrons of the central atom(s) are important for molecular shape valence electrons are paired in a molecule or polyatomic ion bonded pairs of electrons and lone pairs of electrons are treated approximately equally Valence electron pairs repel each other electrostatically the molecular shape is determined by the positions of the electron pairs when they are a maximum distance apart 24 Using VSEPR to predict the shapes of of the hydrogen compounds of period 2. BeH2(s), BH3(g), CH4(g), NH3(g), H2O(l), HF(g) Table 1: Geometry of beryllium dihydride (pages 92 – 93) Lewis formula Bond pairs Lone pairs Total pairs General formula Explanation: Consider BeH2 molecule. The Lewis structure is H Be Electron pair arrangement Stereochemical formula H How are these atoms arranged ? Are they in a straight line or at some angle ? The answer to this question depends on the arrangement of electron pairs around the central atom. The central atom is Berylium and there are 2 shared pairs of electrons around Beryllium. The minimum repulsion will occur, when the electron pairs are on opposite sides of the nucleus. We can represent this as: H– Be – H the molecule should be linear Be Table 2: Geometry of boron trihydride Lewis formula Bond pairs Lone pairs Total pairs General formula Electron pair arrangement Stereochemical formula Explanation: H •• The molecule BH3 . The Lewis structure is H B H The central atom B has 3 bonding pairs, the arrangement which will lead to minimum repulsion is a triangle with boron in the center. B The shape of the molecule is trigonal planar 25 Table 3: Geometry of methane Lewis formula Bond pairs Lone pairs Total pairs General formula Electron pair arrangement Stereochemical formula Total pairs General formula Electron pair arrangement Stereochemical formula Total pairs General formula Electron pair arrangement Stereochemical formula Explanation: Table 4: Geometry of ammonia Lewis formula Bond pairs Lone pairs Explanation: Table 5: Geometry of water (page 94) Lewis formula Bond pairs Lone pairs Explanation: 26 Table 6: Geometry of hydrogen fluoride Lewis formula Bond pairs Lone pairs Total pairs General formula Electron pair arrangement Stereochemical formula Explanation: Shapes of Molecules VSEPR theory describes, explains, and predicts the geometry of molecules by counting pairs of electrons that repel each other to minimize repulsion. The process for predicting the shape of a molecule is summarized below: Step 1: Draw the Lewis formula for the molecule, including the electron pairs around the central atom. Step 2: Count the total number of bonding pairs (bonded atoms) and lone pairs of electrons around the central atom. Step 3: Double bonds and triple bonds can be treated as though they are single bonds between neighbouring atoms. Step 4: When applying VSEPR, we general consider the shape around the central atom. If there are more than one central atom, there are more than one shape. Step 5: Refer to Table 7, and use the number of pairs of electrons to predict the shape of the molecule. Eg. Construct the Lewis structure and predict the shape of each of the following. a. SiF4 b. SO42c. PH3 d. ClO4- 27 Table 7: Using VSEPR Theory to Predict Molecular Shape (page 95) General formula* AX2 Bond pairs 2 Lone pairs 0 Total pairs 2 Geometry** Stereochemical formula Examples linear (linear) AX3 3 0 3 trigonal planar (trigonal planar) AX4 4 0 4 tetrahedral (tetrahedral) AX3E 3 1 4 trigonal pyramidal (tetrahedral) AX2E2 2 2 4 angular (tetrahedral) AXE3 1 3 4 linear (tetrahedral) * A is the central atom; X is another atom; and E is a lone pair of electrons ** The electron pair arrangement is in parenthesis The Multiple Bond in VSEPR Models: page 96 28 Draw structural formulas for the following and determined the shape around their central atom(s) E.g. a. C2H4 (ethane) b. NO2- (Nitrite ion) c. Cl2CO d. C2H2 e. CH3OH f. CH3COCH3 Summary Number of electron pairs 2 pairs Molecular shape BeCl2 X 3 pairs Example Cl – Be – Cl linear planar triangular (trigonal planar) 120 BF3 F o B F F 4 pairs tetrahedral CH4 , NH4+ H 109.5o C H H H 29 Bonding pair 4 Lone pair 0 Name & example Tetrahedral CH4 Shape H 109.5o C H H H 3 1 Pyramidal NH3 N H H H 2 2 107.3o Bent H2O O H 104.5o Molecular Polarity: Dipole Theory H page 98 Electronegativity and Bond Polarity Recall that electronegativity of elements increases from left to right within a period and bottom to top within a group or family on the periodic table Linus Pauling explained that the difference in electronegativity between atoms was the reason for the polarity of a covalent bond. Covalent bonds in which the bonding electrons are unequally shared and unsymmetrical distributed between the nuclei of two bonded atoms are called covalent bonds. Such bonds occur between atoms of different electronegativity. The greater the electronegativity difference, the more polar the bond will be. 30 Pauling reasoned that when two atoms of different electronegativity share electrons, the attraction for these electron are unequal, therefore the electrons will spend more of their time closer to the nucleus of higher electronegativity. In this situation, the bond formed is called a polar covalent bond. The presence or absence of polar bonds within molecule plays a very important part in determining the chemical and physical properties of molecules. Some of these properties are: melting points, boiling points, viscosity and solubilities in solvents. eg. Cl2 and HCl ∙∙ ∙∙ ׃Cl ׃Cl׃ ¨ ¨ H ׃Cl׃ 3.2 2.2 3.2 ∙∙ ¨ 3.2 in the middle nonpolar covalent bond closer to the Cl polar covalent bond Both molecules are formed by sharing one pair of electrons. In the chlorine molecule, the electrons are shared and distributed equally by both chlorine atoms since both atoms have the same attraction for the shared pair of electrons(ie. same electronegativity). However in the hydrogen chloride, the shared electrons are not equally shared by both atoms. The chlorine with a stronger electronegativity attracts the shared pair of electrons more strongly than does the hydrogen. Since the shared electrons are displaced towards the chlorine and away from the hydrogen, the chlorine becomes partially negative(δ-) and the hydrogen partially positive (δ+). In this situation, a bond dipole occurs between the hydrogen and chlorine and the bond is said to be polar. Bonds or molecules that have opposite poles are said to be polar. Polar covalent bonds can be said to have a charge separation or bond dipole. A bond dipole can be represented by an arrow pointing to the negative side of the bond. eg. δ+ H—Cl δ- δ – Greek lower case delta represents partial charge 2.1 3.0 31 For each of the following: 1. Draw the structural formulas 2. Draw in the bond dipoles eg H2O a) HF b) NH3 c) CH4 d) CO2 e) CH3OH f) OF2 g) C2H6 h) CH3Cl O H H Polar Molecule A molecule is polar if the dipoles do not cancel each other out. If bond dipoles do not cancel each other out, the molecule is said to be polar. Using this information, determine if the above molecules are polar or nonpolar. Predicting Bond Types: Bond type can be described as belonging to one of three classes: 1. nonpolar covalent: no difference in electronegativity 2. polar covalent: a difference in electronegativity exists 3. ionic: attraction between a metal and a nonmetal 32 Summary of Bonding: Intramolecular force ionic bond Bonding model involves an electron transfer between a metallic atom and a nonmetallic atom, resulting in the formation of cations and anions polar covalent bond cations and anions attract each other involves unequal sharing of pairs of bonding electrons by atoms of two different nonmetallic elements bonds can involve 1, 2, or 3 pairs of bonding electrons, i.e., single (weakest), double or triple (strongest) bonds non-polar covalent involves equal sharing of electrons bond bonds can involve 1, 2, or 3 pairs of electrons, i.e., single (weakest), double or triple (strongest) bonds Eletronegativites Revisted Note the electronegativities of elements from left to right across periods and down groups. The trends observed for each of these is? Across a period: the electronegativities generally increase from left to right across a period with the Group VII element having the highest value for the period. Why? Down a group: the electronegativities generally decrease from top to bottom down a group. Francium is the element with the lowest electronegativity. Why? Electronegativity is a numerical value describing the relative tendency of an atom to attract negative electrons to itself when bonded to another atom. As we move along a period of the periodic table the nuclear charge increases and so electrons are attracted more strongly. However, when we start a new period, the electrons in the filled inner shell are able to shield the outer electrons from the nuclear charge. Consequently, electronegativity decreases as we move down a group. The noble gases are an exception to this; they are not given a value. This is because their exceptional stability makes it difficult to attract electrons (i.e. they form NO bonds). 33 Label the electronegativity, and label the charges (if any) on the ends of the bond. Classify the following bonds as either ionic, polar covalent, or nonpolar covalent. a Al - Si h B - Na b Ba - O i Ca - Cl c C-H j F-S d Li - S k Br – Rb e Ca - P l F-F f H – Cl m N–O g I – Br n Mg – S Polarity Indicate the polarity (polar / nonpolar) of the following bonds: a P – Cl f O–H b C=O g S–C c H–F h I–I d P–H i C–C e N – Br j N=N 1. Molecular Polarity Draw the structural formulas of the following and using bond dipoles, determine if it is polar or nonpolar. 1 H2O 2 CO2 3 HI 4 CH3OH 34 5 F2 6 CH3I 7 NI3 8 C3H8 9 C2H6 10 CS2 11 CCl4 12 H2Se Guidelines for Predicting Polar and Non-polar Molecules Type Polar AB Description diatomic compounds CO(g) HBx any molecule with a single H HCl(g) AxOH any molecule with an OH at one end C2H5OH(l) OxAy any molecule with an O at one end H2O(l), OCl2(g) NxAy any molecule with an N at one end NH3(g), NF3(g) all elements Cl2(g), N2(g) most carbon compounds (including CO2(g), CH4(g) Nonpolar Ax CxAy Examples organic solvents, fats, and oils) 35 Polar Substances: page 103 What empirical test can we do to demonstrate the polarity of molecules? (page 98) 36 Section 3.4 Intermolecular Forces page 105 Types of forces that exist: Intramolecular Forces Intermolecular Forces forces that hold atoms together in a molecule forces of attraction and repulsion that exist between molecules forces of attraction within a molecule between two atoms forces of attraction between molecules ionic covalent Van der Waals Forces transfer or exchange of valence electrons sharing of single valence electron single double triple polar London Dispersion Forces Dipole-Dipole Forces Hydrogen Bonding polar / non-polar Intermolecular Forces Intermolecular forces are the forces of attractions that exist between molecules in a compound. These cause the compound to exist in a certain state of matter: solid, liquid, or gas; and affect the melting and boiling points of compounds as well as their solubilities of one substance in another. Solid: A state of matter in which the matter is not compressible nor does it flow. Liquid: A state of matter in which the matter is not compressible but can flow. Gas: A state of matter in which the matter is compressible and can flow. The melting point of a compound is the temperature at which a compound turns from a solid to a liquid or a liquid to a solid. The boiling point of a compound is the temperature at which a compound turns from a liquid to a gas or a gas to a liquid. This temperature is a true measure of the forces of attractions between molecules as molecules separate from one another when they turn from a liquid to a gas. The stronger the attractions between particles (molecules or ions), the more difficult it will be to separate the particles. When substances melt, the particles are still close to one another but the forces of attraction that held the particles rigidly together in the solid state have been sufficiently overcome 37 to allow the particles to move. When substances boil, the particles are completely separated from one another and the attractions between molecules are completely overcome. The energy required to cause substances to melt and to boil, and thus disrupt the forces of attraction, comes from the environment surrounding the material. If you place a piece of ice in your hand, the ice will melt more quickly than if it is placed on a cold counter top. The energy required to melt the ice comes from your hand, your hand gets colder and the ice gets warmer. Intermolecular forces are generally much weaker than covalent bonds Only 41 kJ/mol of energy is required to overcome the intermolecular attraction between H2O molecules in the liquid state (i.e. the energy required to vaporize the sample) Intermolecular forces However, 286 kJ/mol of energy is required to break the covalent bond between the H and O atoms in the H2Ol molecule. Intramolecular Forces (ie covalent bonds) Thus, when a molecular substance changes states, the atoms within the molecule are unchanged The temperature at which a liquid boils reflects the kinetic energy needed to overcome the attractive intermolecular forces (likewise, for the temperature at which a solid melts). Types of attractive forces between neutral molecules Dipole-dipole forces London dispersion forces Hydrogen bonding forces Typically, dipole-dipole and dispersion forces are grouped together and termed van der Waals forces Dipole-Dipole Forces A dipole-dipole force exists between polar molecules Polar molecules attract one another when the partial positive charge on one molecule is near the partial negative charge on the other molecule The polar molecules must be in close proximity for the dipole-dipole forces to be significant In a condensed state such as a liquid, where many molecules are in close proximity, the dipoles find the best compromise between attraction and repulsion. This means the molecules orient themselves to maximize the + & _ interactions and to minimize the + & + and - & - interactions as shown below. 38 London Dispersion Forces Nonpolar molecules would not seem to have any basis for attractive interactions. However, gases of nonpolar molecules can be liquefied indicating that if the kinetic energy is reduced, some type of attractive force can predominate. Fritz London (1930) suggested that the motion of electrons within an atom or non-polar molecule can result in a transient dipole moment A Model To Explain London Dispersion Forces: Helium atoms (2 electrons) Consider the particle nature of electrons The average distribution of electrons around each nucleus is spherically symmetrical The atoms are non-polar and posses no dipole moment The distribution of electrons around an individual atom, at a given instant in time, may not be perfectly symmetrical o Both electrons may be on one side of the nucleus o The atom would have an apparent dipole moment at that instant in time (i.e. a momentary or instantaneous dipole) o A close neighboring atom would be influenced by this apparent dipole - the electrons of the neighboring atom would move away from the negative region of the dipole Due to electron repulsion, a temporary dipole on one atom can induce a similar dipole on a neighboring atom This will cause the neighboring atoms to be attracted to one another This is called the London dispersion force (or just London force) It is significant only when the atoms are close together 39 The ease with which an external electric field can induce a dipole (alter the electron distribution) with a molecule is referred to as the “polarizability” of that molecule The greater the polarizability of a molecule the easier it is to induce a momentary dipole and the stronger the dispersion forces Larger molecules tend to have greater polarizability o Their electrons are further away from the nucleus (any asymmetric distribution produces a larger dipole due to larger charge separation) o The number of electrons is greater (higher probability of asymmetric distribution) In general, we may conclude that the more atoms in a molecule(greater number of electrons), the greater the opportunity to form an instantaneous dipole. The result is to increase the attractive forces and raise the boiling point Dispersion forces are also present between polar/non-polar and polar/polar molecules (i.e. between all molecules) Hydrogen Bonding A hydrogen atom in a polar bond (e.g. H-F, H-O or H-N) can experience an attractive force with a neighboring electronegative molecule or ion which has an unshared pair of electrons (usually an F, O or N atom on another molecule) Hydrogen bonds are considered to be dipole-dipole type interactions A bond between hydrogen and an electronegative atom such as F, O or N is quite polar: The hydrogen atom has no inner core of electrons, so the side of the atom facing away from the bond represents a virtually naked nucleus This positive charge is attracted to the negative charge (lone pairs)of an electronegative atom in a nearby molecule 40 Because the hydrogen atom in a polar bond is electron-deficient on one side (i.e. the side opposite from the covalent polar bond) this side of the hydrogen atom can get quite close to a neighboring electronegative atom (with a partial negative charge) and interact strongly with it (remember, the closer it can get, the stronger the electrostatic attraction) o Hydrogen bonds vary from about 4 kJ/mol to 25 kJ/mol (so they are still weaker than typical covalent bonds. o But they are stronger than dipole-dipole and or dispersion forces. o They are very important in the organization of biological molecules, especially in influencing the structure of proteins Water is unusual in its ability to form an extensive hydrogen bonding network As a liquid the kinetic energy of the molecules prevents an extensive ordered network of hydrogen bonds When cooled to a solid the water molecules organize into an arrangement which maximizes the attractive interactions of the hydrogen bonds o This arrangement of molecules has greater volume (is less dense) than liquid water, thus water expands when frozen o The arrangement has a hexagonal geometry (involving six molecules in a ring structure) which is the structural basis of the six-sidedness seen in snow flakes o Each water molecule can participate in four hydrogen bonds , one with each nonbonding pair of electrons and one with each H atom Diagram 41 Higher IMFs result in a strong attractive force between molecules that requires alot of energy to break resulting in a high boiling point. -A molecule like water possesses all 3 types of IMF has a a high boiling point of 100°C. -A molecule like F2 has only LDF has a much lower boiling point of -188°C. Eg.#1 The boiling point of argon is –189.4ºC. a. Why is it so low? b. How does this boiling point help prove that London dispersion forces exist? c. The boiling point of xenon is –119.9ºC. Why is it higher than argon? Eg#2 Put the following in order from the lowest to the highest boiling point? Explain your answer. F2, C2H6, NH3, CH3Cl How does structure and polarity affect boiling points? The variation of boiling point is decided by two factors: 1. The molecular size. The larger molecules have a higher boiling point. 2. The intermolecular forces acting upon the molecules. The greater the forces the higher the boiling point. The more attracted the molecules are to one another the less likely the molecules will be able to escape from one another and become vapor molecules. With regard to strength: Hydrogen Bonding > Dipole-dipole Forces > London Dispersion Forces 42 Relation of Boiling Point to the Number of Electrons and to the Type of Intermolecular Forces Complete the following page. The first is given as an example. Note that there are six series or groups of molecules. Molecular Substances with state Polarity of molecule Number of electrons Boiling point (oC) London Dispersion Forces (polar / non-polar) eg F2(g) non-polar 18 -188 1 Cl2(g) -35 2 Br2(l) 59 3 I2(s) 184 4 ClF(g) -101 5 BrF(g) -20 6 BrCI(g) 5 7 ICl(s) 97 8 IBr(s) 116 9 CH4(g) -162 10 C2H6(g) -87 11 C3H8(g) -45 12 C4H10(g) -0.50 13 C5H12(l) 36 14 CF4(g) -129 15 CCl4(l) 77 16 CBr4(s) 189 17 CH3F(g) -78 18 CH3Cl(g) -24 19 CH3Br(g) 3.6 20 CH3I(l) 43 21 CH3OH(l) 65 DipoleDipole Forces Hydrogen Bonding 43 22 C2H5F(g) -38 23 C2H5Cl(g) 13 24 C2H5Br(l) 38 25 C2H5I(l) 72 26 C2H5OH(l) 78 1. Which of the following would you expect to have the highest boiling point BrF(g) or C3H8(g). Account (why) for the difference in boiling points. Check your answer with the table above. 2. The different series of substances give in the pervious table, in general, have increasing boiling points with increasing number of electrons. Explain this trend of number of electrons and strength of intermolecular forces. 3. Methanol, CH3OH(l), and ethanol, C2H5OH(l), each have the least number of electrons but the highest boiling point of their respective series. Account for this. 4. Explain the difference in boiling point between C2H6 and CH3F. 44 PbF2 melts at 855 oC and boils at 1290 oC. PbI2 melts at 402 oC and boils at 954 oC. Why? Which would you expect to have the higher melting and boiling point: Cl2 or O2? ________ Why? Which would have the higher boiling point, HF or HBr? ________ Why? Bonds holding ionic substances (ionic bonds) together are (stronger or weaker) _______________ than bonds holding molecular substances (covalent bonds) together. Physical Properties of Liquids Many physical properties of liquids can be explained by intermolecular forces. This includes surface tension, shape of the meniscus, volatility and capillary action. Surface Tension By definition, the molecules of a liquid exhibit intermolecular attraction for one another. What happens to molecules at the surface in comparison to those in the interior of a liquid? Molecules in the interior experience an attractive force from neighboring molecules which surround on all sides Molecules on the surface have neighboring molecules only on one side (the side facing the interior) and thus experience an attractive force which tends to pull them into the interior 45 The overall result of this asymmetric force on surface molecules is that: The surface of the liquid will rearrange until the least number of molecules are present on the surface The surface molecules will pack somewhat closer together than the rest of the molecules in the liquid o The surface molecules will be somewhat more ordered and resistant to molecular disruptions o Thus, the surface will seem to have a “skin” The “inward” molecular attraction forces, which must be overcome to increase the surface area, are termed the “surface tension” Liquids that have strong IMF also have high surface tensions. Thus, because water has hydrogen bonding, it has considerable more surface tension than other liquids Surface tension determines whether or not droplets of a liquid will bead up, as on a freshly waxed car, or spread out when placed on a flat surface. Cohesive forces bind molecules of the same type together(binding of like molecules) Adhesive forces bind a substance to a surface (binding of unlike molecules) For example, attractive forces (hydrogen bonding) exists between glass materials (silicon dioxide) and water. This is the basis of “capillary” action, where water can move up a thin capillary, against the force of gravity. Surface tension “pulls” neighboring water molecules along. The liquid climbs until the adhesive and cohesive forces are balance by the force of gravity A. When adhesion is greater than cohesion. The liquid (for example water) rises in the capillary tube. B. When cohesion is greater than adhesion, as it is for mercury, a depression of the liquid in the capillary tube results. Note that the meniscus in the tube of water is concave or rounded downward, whereas that in the tube of mercury is convex, or rounded upward. 46 Eg. Which would have a higher surface tension H2O or C6H14? Why? Would the shape of the H2O meniscus in a glass tube, be the same or different than C6H14? Volatility 47 3.5 Structures and Physical Properties of Solids Solids definite shape and volume vary in properties such as hardness, melting point, conductivity, characteristics fracture – the way solids break. May be curved (conchoidal) as in glass of flint or perfectly flat. Are classified according to their chemical properties as ionic, metallic, molecular or covalent network solids. Ionic Solids attraction of a cation to an anion to form the strongest attractive force known in chemistry form very rigid lattices with large lattice energies high melting and boiling points regular crystal of alternating positive and negative charges. Extremely common since they are very stable and unreactive. Most metal ions are found in this form in nature. Formation of Ionic Compounds Lewis symbol equation .....O . Ca 2+ ..O .. 2- .. .. .Ca . Balanced chemical equation 2 Ca(s) + O2(g) 2 CaO(s) A Model for Ionic Compounds The structure of most binary compounds such as NaCl, KCl, and KI can be explained by the close packing of spheres. Generally the larger ion (usually the anion) is in one of the close packing arrangements with the smaller cations fitting within the holes between the closest packed anions. The packing arrangement is done in such a way as to minimize anionanion and cation-cation repulsions. 48 The strong attractions of the positive and negative charges hold the crystal rigidly together. Hitting a metallic crystal with a hammer would give a completely different result compared to hitting an ionic crystal with the same force. In a metal the atoms would shift their positions with no disruption of the metallic bond. However in an ionic crystal, movement of the atoms by a little as one ionic diameter will cause cations previously aligned with anions to be now aligned with other cations and anions to be aligned with other anions. This would cause large repulsion forces, which would shatter the crystal as shown below. Illustration of why crystals shatter. Light circles are cations and dark circles are anions. Metallic Crystals Metals have the physical characteristics of malleability, ductility, conductors of heat and electricity, durability, and high melting points. The simplest model that explains these observations is the electron sea model. For metals the atoms have low electronegativities; therefore the electrons are free to move over (delocalized) over all the atoms. We can think of the structure of a metal as an arrangement of positive atom cores in a sea of electrons To maximize the bonding in a metal it makes sense to pack as many atoms around each other as possible, maximize the number of nearest neighbors The electrons are not bound to any particular atom and are free to move about in the solid. The mobile electrons can conduct heat and electricity and the metal ions can easily be hammered into a sheet or pulled into a wire. 49 Molecular Solids In network solids, the solid can be considered as one giant molecule. In a molecular crystal, the attractive forces between them are van der Waals forces and/or hydrogen bonding. Eg. Sulfur dioxide (SO2), in which the predominant attractive force is a dipole-dipole interaction. Intermolecular hydrogen bonding is mainly responsible for maintaining the three dimensional lattice of ice. Some substances that hold molecular crystals with LDF are Ne, Xe, S8, CH4 etc In general except ice, molecules in a molecular crystal are packed together as closely as their size and shape will allow. Since the IMF holding these crystals together are very weak, relative to covalent and ionic bonds, molecular crystals are more easily broken apart than ionic or covalent crystals. (most melt below 100ºC) Covalent Network Crystals Eg. Diamond and quartz Among the hardest materials on earth Properties include hard, brittle, high melting points (higher than even ionic crystals). Other examples include SiC(s) (silicon carbide) used for sandpaper and WC(s) (tungsten carbide used for tips of saw blades and drill bit Diamond, Analogy to Methane In methane the bonding network of the carbon atoms is terminated by the H atoms. Because the H atoms can only form one covalent bond, which is confined within the molecule, there is no possibility for extended covalent bonding between different molecules. The inter-molecular bonds must therefore rely on weaker dispersion forces for their cohesive forces. In diamond, which can be constructed by replacing the H atoms in methane by other C atoms, the bonding is not terminated and the C atoms are extended throughout the solid in a tetrahedral geometry. These are strong bonds, as a result diamond has a very high melting (>3550°C) and boiling (4827°C) point and is the hardest natural substance known... 50 Methane Crystal Structure of Diamond Silicon dioxide forms a covalent crystal, with each silicon atom forming bonds to four oxygen atoms and each oxygen bonding to two silicon atoms with a tetrahedral geometry. Silicon carbide is another network crystal similar to diamond with alternating tetrahedral silicon and carbon atoms. It is very hard and is sometimes used as an industrial substitute for diamonds. Benzene and Graphite. The similarity between the bonding in methane and diamond allowed us to interpret the insulating properties of diamond; in a similar way it is possible to interpret the electronic properties of graphite. The properties of diamond and graphite are very different, diamond is hard, colorless, and an insulator; graphite is a well known lubricant, black, and a conductor. These differences arise from the different bonding and structure. The structure of graphite is based on layers of connected six-member rings of carbon atoms - the structure of the layers is very similar to that of benzene, see below. While the bonding in the layers is very strong, only weak van der Waals forces bond the layers together (thus the effectiveness of graphite as a lubricant). 51 Benzene Layered structure of Graphite The slipperiness or lubricating qualities of graphite can be explained by noting that graphite has very strong bonding within the layers but very little bonding between the layers giving a structure of flat sheets. This arrangement allows the layer to slide past one another. 52