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Unit Four Regents Chemistry Mr Grodski III III. A. Introduction - We have learned in Unit One that a compound is a substance made up of two or more elements bonded together. This chemical bond is the result of the valence electrons of two or more elements interacting together. Also in Unit One we learned that potential energy is stored in chemical bonds. The energy that is released from these bonds is called chemical energy. Remember that a compound is decomposed by chemical changes only. Chemical changes are due to the breaking and the formation of chemical bonds in compounds. Our bodies obtain energy from food by creating and breaking chemical bonds in the food we eat. -Generally two atoms that compose or form a chemical bond combine and..... This is due to the fact that most composing chemical reactions are exothermic in nature. Release Energy -Generally when two elements are being decomposed or the bonds are being broken and the atoms of the compound are being separated, the compound........ This again is due to the fact that most decomposing chemical reactions are endothermic in nature. Absorb Energy In most chemical reactions there is a number of different compounds being decomposed and composed. For instance in the combustion reaction: H2 + O2 = H2O + Heat (Energy) The Hydrogen and Oxygen are decomposed to compose water and the overall net change in energy is exothermic. Stability- Stability is a measure of potential energy and whether a substance is present for a long period or for a short period. If something is stable it does not change easily while something that is not stable is always changing. Sort of like relationships! In chemistry stability means that the substance, whether that be an electron, atom, or compound is at its lowest energy state. A lower energy state is stable because the substance does not have a lot of energy to do something with. Its kind of like a hyperactive kid that just ate a lot of sugar. This kid has a lot of energy and needs to burn it off and we can say this kid is unstable because he can not control his crazy active behavior. He is unstable because we do not know what he is going to do next. On the other hand another kid is sluggish and tired. This kid is stable because he does not have the energy to change or do anything but sleep. Atoms are no different. Atoms that are reactive have higher energy states when they are alone and are unstable but when bonded to another atom of their liking they 1 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III achieve lower energy states and thus become stable. They do not have the energy to change. Remember Mr. Stable and Mr. Unstable?! STABLE = LOWER ENERGY STATES – Atoms bonded in compounds UNSTABLE = HIGHER ENERGY STATES – Free reactive Atoms not bonded Notice that when atoms in a bond are being pulled apart energy must be absorbed. Thus energy must be added to the atoms and this extra energy raises their energy states. This in turn makes the individual atoms unstable because they have energy to change or do something. Well, then the two elements bonded now have energy to change and pull apart and become unstable and have a higher energy state, when they absorb energy. STABLE = BONDED ATOMS UNSTABLE = FREE REACTIVE ATOM The following energy diagram illustrates this principle: The boy and girl are now married and happy giving off love. When Atoms Bond they release their extra energy and become more stable (happy)! The boy and girl have met and are nervous on their first date. Atoms when they become in close proximately repel each other initially thus adding to the overall increase in energy and becomes more unstable.. The boy and girl are single and are alone and are unhappy and have a lot of love to give but no one to share it with.. Awww! Atoms who are alone are “unhappy” (unstable) and have a lot of energy to give but no one to give -Boy and Girl represent two free atoms that are not connected or bonded. Reading from left to right you can see that the two elements are not bonded and have a higher energy state than when bonded (Boy-Girl). Thus, the free atoms are not as stable as the bonded atoms. One can also observe that the most unstable state or the state at which the two elements have the highest energy is when they are coming together but not yet bonded . The reason for this high energy at this point is that the electrons from each atom at first repel each other. This initial repulsion is due in part by the negatively charged electrons in the outermost shells in each atom repelling (negatives repell). With enough activation energy (match striking), the two atoms can get past this initial repulsion and chemically Bond. Energy Added H–H + O–O H H + O O → H2O + Energy → Energy released 2 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III B. Ionic Bonds - This type of bond represents a connection between the valence electrons of one or more elements with another, where electrons are transferred from one atom to another. All Ionic Bonds occur between a Metal and a Nonmetal!!!!!!!! -Since an Ionic bond is a transfer of electrons from one atom to another it makes sense that it requires a nonmetal and metal. Remember that metals have low ionization energies and low electronegativities and nonmetals have high electronegativities and high ionization energies. Thus metals have weak hold of electrons and nonmetals have a strong attraction for electrons. This very important because in order to have an ionic bond one element must be able to pluck an electron(s) from another. Well it should seem obvious that the atom that does the plucking is the element that has the strongest attractions for electrons, which is the nonmetal! And the atom that loses its electron or gets plucked is the metal because it does not have a strong hold of its electrons. For Example: Sodium Chloride, NaCl, table salt, is made of Cl and Na held together by an ionic bond. Na is the metal, Cl is the nonmetal. If we look at the electronegativity of Cl and Na from Table S it easy to see why an electron in transferred and an ionic bond is established. Cl = 3.2 Na = 0.9 *2.3* Its easy to see that the Cl is more than 3 times the electronegativity of Na. Because of this large difference in electronegativity its easy see that Cl will pluck an electron from Na. Also one can see that the difference between the two atoms electronegativity is 2.3 . *This is important to note because if the difference between the two atoms elements is 1.7 or greater then the bond is Ionic! -We already know Na is a metal that likes to lose one electron because of the oxidation state in the upper right corner of the periodic table (+1) and because it has one valence electron to lose. Thus it likes to be Na+1. Notice if Na loses its valence electron it has the same STABLE electron configuration of Ne (Noble Gas). -Cl is nonmetal that likes to gain an electron, one of its oxidation states is -1. Thus it likes to be Cl-1 . Notice when Cl gains an electron it has the same STABLE electron configuration of Ar (Noble Gas) The driving force behind chemical bonds is stable lower energy electron configurations that are attained by differences of atomic radii. 3 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III 1. Lewis Diagrams with Ionic compounds.- A Lewis structure of any compound is simply the electron dot diagrams of the atoms in a compound where all valence electrons are accounted for. Remember that only the valence electrons are available for bonding. The lewis diagram For NaCl is : The two are bonded because of the opposite charges of the two ions that attract for one another. Notice that a total number of eight valence electrons are used . This is called the octet rule! The octet rule means that each atom participating in a ionic bond must achieve eight electrons in its outermost energy level, thus eight dots around each element. It’s easy to see that Cl has attained eight and has achieved a charge of -1 because it plucked an electron from Na. Na attained eight by losing an electron. Remember when Na loses an electron it has the same electron configuration as Ne (a noble gas !) A noble gas is stable because it has eight electrons in its outermost energy shell ( filled energy level ). When an atom attains a stable octet of electrons it is at a lower energy state. When two atoms bond they do so at lower energy states. What is the lewis structure of K and I bonded: Chemical Formula: What is the lewis structure of Ca and Cl bonded: Chemical Formula: 2. Writing Chemical Formulas of Ionic Compounds H20 type of atom subscript = number of Hydrogen atoms Oxygen does not have a subscript following it which implies one. A.-Balance oxidation states- when writing chemical formulas for ionic substances its important to balance the oxidation states to zero. For instance in the case of Calcium Chloride.... Ca has a oxidation state of +2 (from upper right corner of the Periodic Table) and Cl has an oxidation state of -1. Thus in order for the compound to have a zero or neutral charge two Cl are needed for every Ca. 2 Cl = -2 1 Ca = +2 0 Thus: CaCl2 When writing chemical formulas for ionic substances the metal always goes first! 4 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III B.-Kriss Cross Method - this method simply takes the oxidation states of each atom and kriss crosses them as subscripts. For example: Sc+3 O-2 Sc2O3 Notice if you did the math that this compound has a zero or neutral charge. The Lewis Structure for the above ionic compound is : 3. Characteristics of Ionic Compounds - Compounds that have ionic bonds are very strong because of the electrostatic attraction between the metal positive ion and the nonmetal negative ion. Strong Attractions always lead to higher melting points or boiling points. (Remember in vapor pressures the liquid with the highest intermolecular forces of attraction had the highest boiling point!) Thus ionic solids have high melting points! -Also ionic compounds in the solids state are in the fixed geometric patterns or crystal lattice. In the geometric structure of the solid ionic crystal, ions form the crystal lattice and are held in relatively fixed positions by electrostatic attraction. They are fixed because the ions are in the most stable or low energy state! Because of the fixed or stable state of the ions in a crystal lattice: Ionic solids do not conduct electricity or heat! -When these crystal lattices are heated (absorb heat) or dissolved in water (aq) the ions of the metal and the nonmetal are non longer fixed (or stable) and the crystal lattice becomes destroyed. When this happens the ions are free to move independently and are no longer fixed in one position. This free movement of ions in the liquid state (l) or a aqueous (aq) state allows for electrical conductivity! Lightning in a pool! Hair dryer in tub of water! Thus ionic compounds in the liquid or aqueous state will conduct electricity! Try these: 1. 2. 3. 4. Write the lewis diagram for Li & F. Write the chemical formula for Li & F Write the lewis diagram for Al & Cl Write the chemical formula for Al & Cl 5 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III C. Covalent Bonds- This type of bond represents an attraction for valence electrons between nonmetals atoms that results in the sharing of electrons. *All Covalent bonds occur from a nonmetal atoms bonded to nonmetals!* Also table S rule - Difference in electronegativity is < 1.7 is a covalent bond! Remember Ionic Bonds occur from a nonmetal and metal atoms while Covalent Bonds result from Nonmetals and Nonmetals! 1. lewis structures for covalent bonds -Remember that ionic bonds have a transfer of electrons in that each atom attained the octet rule individually. In a covalent bond each element still obtains eight electrons (octet) for each but does so by sharing! For example: In a Ionic bond the Na has attained 8 by losing one electron (has the same electron configuration as Ne) and Cl when it gains one has 8! In an ionic bond each element attains an octet independently! In a covalent bond a lewis structure of HCL or Hydrochloric Acid is Notice that Cl has attained an octet for itself if you count all the electrons that it shares and it has surrounding it. 2 electrons for H 8 electrons for C l ( H has only 2 electrons in its only principle energy level thus its octet # is 2. Its the only exception that you will deal with!) -Notice that the Hydrogen and Chlorine share a pair of electrons. This pair of electrons represents a covalent chemical bond! This compound is contains only one pair of electrons thus represents a single bond. This compound can be written into a structural formula of -The single dash between the two atoms represents a single bond or a pair of electrons. In a structural formula the extra electrons surrounding Cl are discarded because only the electrons involved in the bond are illustrated as a dash in a structural formula. These extra electrons are very important in determining shape of the structure, which we will see later! Write the lewis structure of NH3 6 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III Write the lewis structure of CBr4 Write the lewis structure of CO2 Notice CO2 has double bonds due to two pairs of electrons being shared! When can use lewis structures to determine the number of bonds present! For example determine the number of bonds in a molecule of H2 Determine the number of bonds in a molecule of Cl2 Determine the number of bonds in a molecule of O2 Determine the number of bonds in a molecule of N2 2. Nonpolar covalent bonds- Covalent bonds implies that electrons are shared between two atoms in a bond. Ionic bonds are those bonds where there is an transfer of electron(s) between two atoms. The reason for a transfer of electrons in an ionic bond is that there is a large difference in electronegativity between the metal atom and the nonmetal atom. The nonmetal atom because of its larger electronegativity plucks the electron(s) from the lower electronegative metal atom and thus a transfer. -But in a covalent (sharing) bond the atoms involved are all nonmetals, which have similar electronegativities. When two nonmetal atoms that are bonded have the same electronegativities then the electrons are “equally shared” and are NONPOLAR! Think of it as a tug of war between the two atoms. If each atom in the bond has the same electronegativity, which is the same attraction for electrons, then the each atom will pull on the electrons equally and no one wins and no poles are produced. -All nonmetal atoms that are bonded to themselves are Nonpolar Covalent Bonds. These molecules represent all the bonds in the HOFBrINCl ‘s. ( H2 , O2 , F2 , Br2 , I2 , N2 , Cl2 ) For instance I2 , one of the HOFBrINCl’s equally shares electrons, its lewis structure: -Notice each Iodine atom has 8 electrons and shares a single bond between them. 7 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III 3. Polar covalent bond- In this covalent bond the electrons that the two nonmetals atoms are sharing unequally. The reason that they are sharing electrons unequally is that the two atoms have different electronegativities. -A covalent bond is a bond between two nonmetals. If those two nonmetals are the same element as in the case of HOFBrINCl’s then its a nonpolar covalent bond. If the two nonmetals are different elements than they must have different electronegativities and thus its a polar covalent bond! -In keeping with the tug of war example, with a polar covalent bond the nonmetal with the strongest attraction for electrons or the highest electronegativity will pull the shared electrons closer to its atom and win! If the shared electrons are closer to one atom than another than one side of the bond will have a partial negative charge and the other side a partial positive charge, thus two poles..and polar! -for example has a polar covalent bond. It is a covalent bond because H and Cl are both nonmetals and if you subtract the two electronegativities obtained from table S - 3.2 - 2.2 = 1.0 1.0 is less than 1.7 thus it must be covalent. The electronegativities are different thus it must be a polar covalent bond. -If we were to draw a lewis structure for the HCL showing the pull of electrons due to the differences in electronegativities it would look something like this: -Notice that the Chlorine atom which is the most electronegative of the two atoms pulls the electrons closer to itself and has a partial negative charge. Hydrogen has its electrons pulled farther away and has a partial positive charge. THIS IS NOT LIKE A IONIC BOND BECAUSE THE ELECTRON ALTHOUGH UNEQUALLY SHARED ARE NOT BEING PLUCKED JUST PULLED UNEQUALLY! - The arrow below the lewis structure above demonstrates a Dipole Moment, which is the direction of the charge. Charge always flows from positive to negative. All polar bonds have a dipole moment. 4. Coordinate covalent bonds- This type of covalent bond is no different than any other covalent bond except in this case one of the two atoms involved in the bond donates a complete pair of electrons to be shared and the other atom brings none. This is an example of a Ammonium Ion -For example NH4+ : shows an example of a coordinate covalent bond between the top H and the Nitrogen. represent electrons from the hydrogen. represent electrons from the nitrogen. Notice Nitrogen has five valence electrons and Hydrogen has one valence electron. The top H and the N is a coordinate covalent bond because both electrons (a bond represents a pair) are from the Nitrogen and the hydrogen did not bring any electrons! 8 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III 5. Polar vs. Nonpolar Molecules- To Be Polar or Not to be Polar That is The Question! To Be Polar: 1. Molecule must have polar bonds!(not HOFBrINCl’s) 2. Must be asymmetrical!! (no symmetry) Is CCL4 , H2 , NH3, H20 Polar or Nonpolar molecules? Knowing our two rules: TO BE POLAR (Molecule) 1. Must have Polar Bonds! 2. Must be asymmetrical! If one of these two are false for a molecule then it must be Nonpolar! A nonpolar molecule is a molecule that can have or Nonpolar Bonds Polar Bonds The reason that CH4 is nonpolar is due to the fact that the second rule is not true. It does have symmetry! Symmetry simply implies that the molecule is shaped in a particular way that cancels out the dipole moment! Remember that the dipole moment is the direction of the charge and charge always flows from the positive pole to the negative pole of the Dipole (polar molecule). Remember that a polar molecule has two poles due to differences in Electronegativity. For example: Electrons are being pulled toward Br thus Br becomes - Electrons are being pulled away from I, thus I becomes + Electronegativities from table S: 2.9 2.7 Because of the differences in the electronegativities in these two atoms there is unequal sharing of electrons in the bond. Br has a stronger attraction for electrons than I which is demonstrated by a higher electronegative value of 2.9. If Br has a stronger attraction for electrons than it pulls the electrons closer toward itself and Br gains a partial negative charge. I has its electrons pulled away from it by the more electronegative Br and gains a partial positive charge. If the BrI molecule has two ends with different charges than its polar (or a Dipole). The dipole moment goes from right to left. It is not canceled out here thus BrI is polar. 9 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III This molecule has polar bonds but is symmetrical thus it is nonpolar Again symmetry means that the dipole moments cancell out. In this case the direction of the charge of all the atoms attached is toward the Carbon thus the dipole moments cancells out! If the dipole moments all went away from the carbon it would also cancell out. A molecule with polar bonds and is nonpolar must be a molecule that must have symmetry( - dipole moments cancel out!) No polar bonds no dice! Again symmetry is dependent upon dipole moments canceling out and not the overall shape of molecule. This molecule is also a Tetrahedral but it is a Polar molecule because it does not have symmetry, which means that the dipole moments do not cancel out! Cl is more electronegative than C thus the direction of charge is different! This molecule is a Tetrahedral and is nonpolar because it is has symmetry which means that dipole moments cancell out! Therefore just because a molecule is Linear, Trigonal planar, or Tetrahedral does not mean they are automatically nonpolar because of their symmetrical shapes! All dipole moments must cancel out in each molecule in order for a molecule to be non polar. If they have different atoms attached to them as shown in the above CH3Cl molecule they will not have symmetry because there dipole moments or the direction of each charge will not cancel out!Bent and Pyramidal shapes are always polar because of there asymmetrical shapes regardless if the atoms attached are all the same or if they are different! Shapes of molecules are determined by the number of valence electrons, bonding and nonbonding electrons. Since valence electrons are all the same in a group or family of elements ( column ) then we can predict for shapes of molecules based on this information. 10 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III D. Polyatomic IonsA single atom with a charge is a monatomic atom. Example: Na+ . A compound of two or more covalently bonded atoms with a charge is called a polyatomic ion. Polyatomic Ions are very stable and behave like a monoatomic particle and the bonds that hold them are stronger than those of the metal it ionically bonds with. Therefore, during reactions, the polyatomic ion usually remains intact as it passes from the reactants to the products AgNO3 + NaCl AgCl + NaNO3 -Notice the Nitrate (NO3) stays intact through reactions. Notice that the nitrate ion from table E is -1 and Ag has a +1 oxidation state (from upper righthand of Ag in periodic table) and AgNO3 is neutral. Check NaCl you will see that it also will be neutral. Table E lists all of the polyatomic ions that you will be using in this course. Polyatomic ions are most often have negative oxidation states THUS ionically bond with Metals. Although the bonds which keep the atoms in a polyatomic ion are covalent bonds, the polyatomic ions possess a charge to satisfy the octet rule for the covalent bonds in the ion. The above polyatomic ion is has a negative charge of one because it needs one electron from a metal in order to satisfy the octet on the Left oxygen. The triangle represents the electron from the metal. Notice all octets are satisfied. When Na loses an electron it becomes positively charged and ionically bonds with NO3 and the lewis structure looks like this: ** Polyatomic Ions have both Covalent Bonds and Ionic bonds** When naming polyatomic compounds use the Kriss Cross method and makes sure the charges balance out to zero! Remember all polyatomic ions are listed in Table E! 11 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III E. Metallic Bonding - Metallic bonding occurs between two metals. A metal bond consists of an arrangement of positive ions that are located at the crystal lattice sites and are immersed in a “sea” of mobile electrons. These mobile electrons can be considered as belonging to the whole crystal rather than to individual atoms. Remember most metals are transitional elements that have an abundant amount of electrons in their d orbitals that creates this “sea” and many of metals physical and chemical characteristics. F. Characteristics of Different solids 1. Ionic Compounds – (Ionic Compounds- binary & ternary Salts) ∼Metals bonded to nonmetals∼ 1. high melting points 2. low electrical conductivity in solid state 3. high electrical conductivity in liquid or aqueous state 4. mostly soluble in water (dissolves): called electrolyes. 5. All ionic substances are called salts 2. Molecular solids- (covalently bonded- often referred to as Organic because they predominate in living tissue) ∼nonmetals bonded to nonmetals∼ 1. Soft 2. Good electrical insulators (rubber) 3. poor heat conductors 4. low melting points 5. mostly insoluble in water: does not dissolve in water! 3. Metallic solids- (transitional metals bonded to themselves) ∼metals bonded to metals∼ 1. good conductors of electricity and heat 2. great strength 3. malleability and ductility 4. luster & colored solutions when forming complex ions in water 4. Network Solids- ( Covalently bonded crystals) - All of our crystals on earth are ionic crystals that result from Ionic Compounds (ternary and binary salts) except for a few crystals that are covalently bonded. These substances are called Network Solids because they have a network of arrangement much like the crystal lattice of ionic and metallic compounds. There are only three in existence and they include: diamond (C), silicon carbide (SiC), and silicon dioxide (SiO2). Characteristics include: 1. Hard 2. Poor conductors of heat and electricity 3. High melting points 12 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III G. Naming and Writing Formulas of Chemical Compounds 1. Formulas: a. Empirical Formulas - Simplest atomic ratio in which elements combine to form a compound. For example H2O2 ( Hydrogen peroxide) has a empirical formula of HO. One part Hydrogen to one part Oxygen. b. Molecular Formulas - Covalent structures are comprised of molecules. These molecules sometimes consist of atoms combined chemically in multiples of their simplest ratio but many time they include molecules that have multiples of their smallest ratio. For example H2O2 is the molecular formula for Hydrogen peroxide yet its empirical formula is HO. The molecular formula is often referred to as the “true” formula. 2. Naming: a. Binary compounds (two elements!) 1. Binary salts (Ionic compounds) - In a binary compound composed of a metal and a nonmetal, the metallic element is usually named normally and written first. The nonmetal ends in ide For example: NaCl is named sodium chloride CaCl2: LiI : 2. Binary compounds of two nonmetals (covalent) The less electronegative element is usually named and written first normally and the other nonmetal ends in ide . Prefixes and Suffixes are used to indicate the number of atoms of each of the nonmetals. mono- (1), example, CO : carbon monoxide di- (2), example, CO2 : carbon dioxide tri- (3), example, SO3 : sulfer trioxide tetra- (4), example, CCl4 : carbon tetrachloride penta- (5), example, PCl5 : phosphorus pentachloride N2O : dinitrogen monoxide b. Ternary compounds (Salts- Ionic) - contain three different kinds of atoms. Most compounds that contain three or more different kinds of atoms contain polyatomic ions. The names of these compounds with the exception of the acids, which we will learn, later includes just the normal name of the metal and the name of the polyatomic ion. NaNO3 is called sodium nitrate and there is no use of prefixes or suffixes in this case. Mg3(PO4)2 Write the chemical formula for Potassium Chlorite: 13 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III 4. Stock System - This is newer method in naming chemical compounds that utilizes a Roman Numeral to identify a particular oxidation state of an element that has multiple oxidation states. For example Nitrogen has 8 oxidation states. They include -3, 2, -1, +1, +2, +3, +4, and +5. These oxidation states can be found at the upper right hand corner of each element in the Periodic Table. Oxidation states are the ions or charges that a particular element can have (losing or gaining electrons or both). In the case of Nitrogen it can either lose electrons or gain electrons to obtain positive or negative ions. The Stock System eliminates the need to to use prefixes or suffuxes. It does this by using Roman Numerals to identify the Positive Ions of a particular element. For example Nitrogen with its multiple oxidation states will be written as a compound with oxygen under both formats. Formula Name of Compound Stock System Name N2 O NO N2O3 NO2 N2O5 dinitrogen monoxide nitrogen monoxide dinitrogen trioxide nitrogen dioxide dinitrogen pentoxide nitrogen(I) oxide nitrogen(II) oxide nitrogen(III) oxide nitrogen(V) oxide If you check the charges of both elements you would see that they would balance to zero! H. Molecular Attraction - Molecular attractions account for the intermolecular attractions between molecules in liqiuds and gases. These are not chemical bonds but can be very significant in strength. Remember that the strength of the intermolecular attractions affects vapor pressure and boiling points. Strong intermolecular attractions equal lower vapor pressure and higher boiling points! 1. Dipole - Dipole Attractions- This type of molecular attraction utilizes the attraction between two dipoles (two polar molecules). The positive end of one dipole will attract the other end of another dipole. For example in the case of HCl : These attractions can be very strong. The strength of a dipole - dipole interaction is determined by the difference in electronegativities obtained from table K. The larger the difference in electronegativity the larger greater the dipole force. Dipole - Dipole forces of attraction ocurr from two polar molecules!!! 2. Hydrogen Bonding - The strongest type of Molecular Attractions are Hydrogen Bonding or H-Bonding. This type of attraction is the a special type of dipole dipole attraction. Its still requires two polar molecules but one of the molecules must have a hydrogen covalently bonded to a highly electronegative atom. In this way the hydrogen has its pair of sharing electrons pulled away from it to make it very polar. Just like dipole - dipole the strongest H-Bonding occurs when the difference in 14 Grodski Productions Inc. © 1997 Unit Four Regents Chemistry Mr Grodski III electronegativity between the H and the element its bonded to is greatest. Since F is the most electronegative atom than the strongest H - Bonding takes place in molecules in HF. -Water has a relatively high boiling point because of H - Bonding! H -Bond 3. Van der Waals Force (London Forces)- his type of intermolecular attraction occurs between nonpolar molecules. This type of attractions are the weakest because of the absence of electrostatic forces that polar molecules have. Van er Waals forces appear to be due to chance distribution of electrons resulting in momentary dipole attractions. That is even a nonpolar molecule can have have a momentary dipole due to electrons moving to one side of the molecule. If more electrons move to one side of a molecule it can give that side a momentary negative charge which can attract a momentary positive charge from another molecule. The strength of a Van der Waals force is stronger in larger molecules and or heavier, denser atoms because they have more electrons that can produce larger momentary charges, which are stronger. The strength of a van der Waals force also increases as two nonpolar molecules come closer to one another and weaker as they move apart. 4. Molecule Ion Attraction- This type of attraction is one where a molecule like water is attracted to ion. In aqueous solutions, where a chemical is dissolved in water the water molecule surrounds the ion and dissolves it. Dissolving of a particle by water is called hydration. When NaCl(s) is mixed with water, the crystal lattice of the ionic compound is broken down and free ions are released into the solvent (water). Ions have charges and the negative charge from Cl is attracted to the positive H end of the polar water molecule. The positive charge of the Na is attracted to the negative O in water. When something dissolves in water the water molecules surround the ions as shown. if something is not soluble than the solvent (water in this case cannot surround it). Notice that a polar molecule could only dissolve an ionic solid because of its dipoles. 15 Grodski Productions Inc. © 1997