HL IB Chemistry - The Covalent Model PDF

Head to www.savemyexams.com for more awesome resources HL IB Chemistry Your notes The Covalent Model Contents Covalent Bonds Lewis Formulas Multiple Bonds Coordinate Bonds Shapes of Molecules Bond P...

Head to www.savemyexams.com for more awesome resources HL IB Chemistry Your notes The Covalent Model Contents Covalent Bonds Lewis Formulas Multiple Bonds Coordinate Bonds Shapes of Molecules Bond Polarity Molecular Polarity Giant Covalent Structures Intermolecular Forces Physical Properties of Covalent Substances Chromatography Resonance Structures (HL) Benzene (HL) Expansion of the Octet (HL) Formal Charge (HL) Sigma & Pi Bonds (HL) Hybridisation (HL) Page 1 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Covalent Bonds Your notes Covalent Bonds What are covalent bonds? Covalent bonding occurs between two non-metals A covalent bond involves the electrostatic attraction between nuclei of two atoms and the electrons of their outer shells No electrons are transferred but only shared in this type of bonding When a covalent bond is formed, two atomic orbitals overlap and a molecular orbital is formed Covalent bonding happens because the electrons are more stable when attracted to two nuclei than when attracted to only one Diagram to show the formation of a covalent bond in a hydrogen molecule The positive nucleus of each atom has an attraction for the bonding electrons shared in the covalent bond In a normal covalent bond, each atom provide one of the electrons in the bond. A covalent bond is represented by a short straight line between the two atoms, H-H Covalent bonds should not be regarded as shared electron pairs in a fixed position; the electrons are in a state of constant motion and are best regarded as charge clouds Hydrogen Molecular Orbital Diagram Page 2 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes A representation of electron charge clouds. The electrons can be found anywhere in the charge clouds Non-metals are able to share pairs of electrons to form different types of covalent bonds Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas This makes each atom more stable The octet rule refers to the tendency of atoms to gain a valence shell with a total of 8 electrons In some instances, the central atom of a covalently bonded molecule can accommodate more or less than 8 electrons in its outer shell Being able to accommodate more than 8 electrons in the outer shell is known as ‘expanding the octet rule’ Accommodating less than 8 electrons in the outer shell means than the central atom is ‘electron deficient’ Some examples of this can be found in the section on Lewis structures Examiner Tip Covalent bonding takes place between two nonmetal atoms. Remember to use the periodic table to decide how many electrons are in the outer shell of a nonmetal atom. Page 3 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Lewis Formulas Your notes Lewis Formulas Lewis formulas are simplified electron shell diagrams and show pairs of electrons around atoms. A pair of electrons can be represented by dots, crosses, a combination of dots and crosses or by a line. For example, chlorine can be shown as: Different Lewis Formulas for chlorine molecules Note: Cl–Cl is not a Lewis formula, since it does not show all the electron pairs. The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons Steps for drawing Lewis Formulas 1. Count the total number of valence electrons 2. Draw the skeletal structure to show how many atoms are linked to each other. 3. Use a pair of crosses or dot/cross to put an electron pair in each bond between the atoms. 4. Add more electron pairs to complete the octets around the atoms ( except H which has 2 electrons) 5. If there are not enough electrons to complete the octets, form double/triple bonds. 6. Check the total number of electrons in the finished structure is equal to the total number of valence electrons Page 4 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Worked example Your notes Draw a Lewis formula for CCl4. Answer: Steps in drawing the Lewis formula for CCl4 Further examples of Lewis formulas Follow the steps for drawing Lewis structures for these common molecules Total number of Molecule Lewis formula valence electrons C + 4H CH4 4 + (4 x 1) = 8 Page 5 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources N + 3H NH3 Your notes 5 + (3 x 1)=8 2H + O H 2O (2x 1) + 6 =8 C + 2O CO2 4 + (2 x 6) = 16 H+C+N HCN 1+ 4 + 5 = 10 Incomplete Octets For elements below atomic number 20 the octet rule states that the atoms try to achieve 8 electrons in their valence shells, so they have the same electron configuration as a noble gas However, there are some elements that are exceptions to the octet rule, such a H, Li, Be, B and Al H can achieve a stable arrangement by gaining an electron to become 1s2, the same structure as the noble gas helium Li does the same, but losing an electron and going from 1s22s1 to 1s2 to become a Li+ ion Be from group 2, has two valence electrons and forms stable compounds with just four electrons in the valence shell B and Al in group 13 have 3 valence electrons and can form stable compounds with only 6 valence electrons There are two examples of Lewis structures with incomplete octets you should know, BeCl2 and BF3: Molecule Total number of valence electrons Lewis formula Be + 2Cl = BeCl2 2 + ( 2 x 7) = 16 B + 3F= BF3 3 + (3 x 7) = 24 Test your understanding of Lewis diagrams in the following example: Page 6 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Worked example Your notes How many electrons are in the 2-aminoethanoic acid molecule? A. 18 B. 20 C. 28 D. 30 Answer: The correct option is D because: You must count the lone pairs on N and O as well as the bonding pairs. There are 5 ‘hidden’ pairs of bonding electrons in the OH, CH2 and NH2 groups Hydrogen does not follow the octet rule Examiner Tip Lewis formulas are also known as electron dot or Lewis structures. Page 7 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Multiple Bonds Your notes Multiple Bonds Non-metals are able to share more than one pair of electrons to form different types of covalent bonds Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas This makes each atom more stable It is not possible to form a quadruple bond as the repulsion from having 8 electrons in the same region between the two nuclei is too great Covalent Bonds & Shared Electrons Table Number of Type of covalent bond electrons shared Single (C − C) 2 Double (C = C) 4 Triple (C ≡ C) 6 Bond energy The bond energy is the energy required to break one mole of a particular covalent bond in the gaseous states Bond energy has units of kJ mol-1 The larger the bond energy, the stronger the covalent bond is Bond length The bond length is internuclear distance of two covalently bonded atoms It is the distance from the nucleus of one atom to another atom which forms the covalent bond The greater the forces of attraction between electrons and nuclei, the more the atoms are pulled closer to each other This decreases the bond length of a molecule and increases the strength of the covalent bond Triple bonds are the shortest and strongest covalent bonds due to the large electron density between the nuclei of the two atoms This increase the forces of attraction between the electrons and nuclei of the atoms As a result of this, the atoms are pulled closer together causing a shorter bond length Page 8 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources The increased forces of attraction also means that the covalent bond is stronger Diagram to show bond lengths for carbon Your notes Triple bonds are the shortest covalent bonds and therefore the strongest ones Examiner Tip Remember: Single covalent bonds are the longest and weakest Triple covalent bonds are the shortest and strongest Page 9 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Coordinate Bonds Your notes Coordinate Bonds What are coordinate bonds? In simple covalent bonds the two atoms involved share electrons Some molecules have a lone pair of electrons that can be donated to form a bond with an electron- deficient atom An electron-deficient atom is an atom that has an unfilled outer orbital So both electrons are from the same atom This type of bonding is called dative covalent bond or coordinate bond An example of a dative bond is in an ammonium ion The hydrogen ion, H+ is electron-deficient and has space for two electrons in its shell The nitrogen atom in ammonia has a lone pair of electrons which it can donate to the hydrogen ion to form a coordinate bond Dative covalent bonding ammonium ion Page 10 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Ammonia (NH3) can donate a lone pair to an electron-deficient proton (H+) to form a charged ammonium ion (NH4+) Your notes More examples of coordinate bonding can be found in the section on Lewis Structures Examiner Tip Coordinate bonds are also referred to as coordination bonds or dative covalent bonds. Page 11 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Shapes of Molecules Your notes Shapes of Molecules What is Valence Shell Electron Pair Repulsion Theory? When an atom forms a covalent bond with another atom, the electrons in the different bonds and the non-bonding electrons in the outer shell all behave as negatively charged clouds and repel each other In order to minimise this repulsion, all the outer shell electrons spread out as far apart in space as possible Molecular shapes and the angles between bonds can be predicted by the valence shell electron pair repulsion theory known by the abbreviation VSEPR theory VSEPR theory consists of three basic rules: 1. All electron pairs and all lone pairs arrange themselves as far apart in space as is possible. 2. Lone pairs repel more strongly than bonding pairs. 3. Multiple bonds behave like single bonds These three rules can be used to predict the shape of any covalent molecule or ion, and the angles between the bonds The regions of negative cloud charge are known as domains and can have one, two or three pairs electrons Two electron domains If there are two electron domains on the central atom, the angle between the bonds is 180o Molecules which adopt this shape are said to be LINEAR Examples of linear molecules include BeCl2, CO2, and HC≡CH Diagram to show molecules with two electron domains Beryllium chloride, carbon dioxide and ethyne all have two electron domains Three electron domains If there are three electron domains on the central atom, the angle between the bonds is 120o Molecules which adopt this shape are said to be TRIANGULAR PLANAR or TRIGONAL PLANAR Page 12 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Examples of three electrons domains which are all bonding pairs include BF3 and CH2CH2 and CH2O Diagram to show molecules with three electron domains Your notes Boron trifluoride, ethene and methanal all have three electron domains If one of these electron domains is a lone pair, the bond angle is slightly less than 120o due to the stronger repulsion from lone pairs, forcing the bonding pairs closer together. E.g. SO2 The bond angle is approximately = 118o The shape of sulfur dioxide Sulfur dioxide is an example of a molecule that 'expands the octet' as you will see there are 10 electrons around the sulfur atom which is possible for 3rd period elements and above This shape is no longer called triangular planar as the shape names are only based on the atoms present, this molecule is BENT LINEAR Four electron domains If there are four electron domains on the central atom, the angle between the bonds is approx 109o. E.g. CH4, NH4+ Page 13 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Molecules which adopt this shape are said to be TETRAHEDRAL Diagram to show molecules with four electron domains Your notes Methane and ammonium ions have four electron domains If one of the electron domains is a lone pair, the bond angle is slightly less than 109o, due to the extra lone pair repulsion which pushes the bonds closer together (approx 107o). E.g. NH3, The shape of ammonia Molecules which adopt this shape are said to be TRIANGULAR PYRAMIDAL or TRIGONAL PYRAMIDAL Page 14 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources If two of the electron domains are lone pairs, the bond angle is also slightly less than 109o, due to the extra lone pair repulsion (approx 104o). E.g. H2O Molecules which adopt this shape are said to be BENT or ANGULAR or BENT LINEAR or V-shaped Your notes (when viewed upside down) The shape of water Lone pairs are pulled more closely to the central atoms so they exert a greater repulsive force than bonding pairs Diagram to show repulsion between different electron pairs The order of electron pair repulsion is lone pairs > lone pair: bonding pair > bonding pairs Summary table of electron domains and molecular shapes These are the domains and molecular geometries you need to know for Standard Level: Bonding pairs Lone pairs Total pairs Domain geometry Molecular geometry Bond angle Page 15 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources 2 0 2 linear linear 180° Your notes 3 0 3 trigonal planar trigonal planar 120° 2 1 3 trigonal planar bent linear 118° 4 0 4 tetrahedral tetrahedral 109.5° 3 1 4 tetrahedral trigonal pyramid 107° 2 2 4 tetrahedral bent linear 104.5° Examiner Tip Be careful to distinguish between molecular shape and electron domain shape as it can be easy to confuse the two. Sometimes they are the same as is the case of methane, but other times they can be different like ammonia which has a tetrahedral domain shape, but triangular pyramid molecular shape. Always draw the Lewis structure before you attempt to deduce the shape and bond angle as you could easily miss some lone pairs Page 16 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Bond Polarity Your notes Bond Polarity What is electronegativity? Electronegativity refers to the ability of an atom to draw an electron pair towards itself in a covalent bond Different atoms have different electronegativities, shown by the Pauling scale below The higher the value, the more electronegative the element is The Pauling Scale First three rows of the periodic table showing electronegativity values In diatomic molecules the electron density is shared equally between the two atoms Eg. H2, O2 and Cl2 Both atoms have the electronegativity value and have an equal attraction for the bonding pair of electrons leading to formation of a covalent bond The covalent bond is nonpolar Diagram to show the electron distribution in a chlorine molecule Page 17 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes The two chlorine atoms have identical electronegativities so the bonding electrons are shared equally between the two atoms and the bond is nonpolar What is meant by a polar bond? When two atoms in a covalent bond have different electronegativities the covalent bond is polar and the electrons will be drawn towards the more electronegative atom As a result of this: The negative charge centre and positive charge centre do not coincide with each other This means that the electron distribution is asymmetric The less electronegative atom gets a partial charge of δ+ (delta positive) The more electronegative atom gets a partial charge of δ- (delta negative) The extend of polarity in a covalent bond varies, depending on how big a difference exists in the electronegativity values of the two bonded atoms The bigger the difference in electronegativity, the higher the polarity of the covalent bond Diagram to show the electron distribution in an HCl molecule Cl has a greater electronegativity than H causing the electrons to be more attracted towards the Cl atom which becomes delta negative and the H delta positive What is a dipole? Page 18 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources The dipole moment is a measure of how polar a bond is The direction of the dipole moment is shown by the following sign in which the arrow points to the partially negatively charged end of the dipole: Your notes The sign shows the direction of the dipole moment and the arrow points to the delta negative end of the dipole Worked example The electronegativity values of four elements are given. C = 2.6 N = 3.0 O = 3.4 F = 4.0 What is the order of increasing polarity of the bonds in the following compounds? A. CO < OF2 < NO < CF4 B. NO < OF2 < CO < CF4 C. CF4 < CO < OF2 < NO D. CF4 < NO < OF2 < CO Answer: The correct option is B because: You have to calculate the difference in electronegativity for the bonds and then rank them from smallest to largest: NO (3.4 - 3.0 = 0.4) OF2 (4.0 - 3.4 = 0.6) CO (3.4 - 2.6 = 0.8) CF4 (4.0 - 2.6 = 1.4) Page 19 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Molecular Polarity Your notes Molecular Polarity To determine whether a molecule with more than two atoms is polar, the following things have to be taken into consideration: The polarity of each bond within the molecule How the bonds are arranged in the molecule (i.e the geometry of the molecule) Some molecules have polar bonds but are overall not polar because the polar bonds in the molecule are arranged in such way that the individual dipole moments cancel each other out A polar molecule There are four polar covalent bonds in CH3Cl which do not cancel each other out causing CH3Cl to be a polar molecule; the overall dipole is towards the electronegative chlorine atom A nonpolar molecule Page 20 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes Though CCl4 has four polar covalent bonds, the individual dipole moments cancel each other out causing CCl4 to be a nonpolar molecule Examiner Tip When the pulls of the atoms are not equal and opposite, there is a net pull so the molecule is polar. Page 21 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Giant Covalent Structures Your notes Giant Covalent Structures Covalent lattices Covalent bonds are bonds between nonmetals in which electrons are shared between the atoms In some cases, it is not possible to satisfy the bonding capacity of a substance in the form of a molecule; the bonds between atoms continue indefinitely, and a large lattice is formed. There are no individual molecules and covalent bonding exists between all adjacent atoms Such substances are called giant covalent substances, and the most important examples are C and SiO2 Graphite, diamond, buckminsterfullerene and graphene are allotropes of carbon Diamond Diamond is a giant lattice of carbon atoms Each carbon is covalently bonded to four others in a tetrahedral arrangement with a bond angle of 109.5o The result is a giant lattice with strong bonds in all directions Diamond is the hardest substance known For this reason it is used in drills and glass-cutting tools Diagram to show the tetrahedral structure of diamond Page 22 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources The structure of diamond Graphite Your notes In graphite, each carbon atom is bonded to three others in a layered structure The layers are made of hexagons with a bond angle of 120o The spare electron is delocalised and occupies the space in between the layers All atoms in the same layer are held together by strong covalent bonds, and the different layers are held together by weak intermolecular forces Diagram to show the layered structure of graphite The structure of graphite Buckminsterfullerene Buckminsterfullerene is one type of fullerene, named after Buckminster Fuller, the American architect who designed domes like the Epcot Centre in Florida It contains 60 carbon atoms, each of which is bonded to three others by single covalent bonds The fourth electron is delocalised so the electrons can migrate throughout the structure making the buckyball a semi-conductor It has exactly the same shape as a soccer ball, hence the nickname the football molecule Diagram to show the interlocking hexagons and pentagons that make up the structure of Buckminsterfullerene Page 23 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes The structure of buckminsterfullerene Graphene Some substances contain an infinite lattice of covalently bonded atoms in two dimensions only to form layers. Graphene is an example Graphene is made of a single layer of carbon atoms that are bonded together in a repeating pattern of hexagons Graphene is one million times thinner than paper; so thin that it is actually considered two dimensional Diagram to show the two dimensional structure of graphene The structure of graphene Page 24 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Silicon The silicon atoms in silicon have a tetrahedral arrangement, just like that of the carbon atoms in Your notes diamond Each silicon atom is covalently bonded to four other silicon atoms Silicon has a giant lattice structure Diagram to show the tetrahedral arrangement in silicon The structure of silicon Silicon(IV) oxide Silicon(IV) oxide is also known as silicon dioxide, but you will be more familiar with it as the white stuff on beaches! Silicon(IV) oxide adopts the same structure as diamond - a giant structure made of tetrahedral units all bonded by strong covalent bonds Each silicon is shared by four oxygens and each oxygen is shared by two silicon atoms This gives an empirical formula of SiO2 Diagram to show the tetrahedral units in silicon(IV) oxide Page 25 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes The structure of silicon dioxide Properties of Giant Covalent Structures Different types of structure and bonding have different effects on the physical properties of substances such as their melting and boiling points, electrical conductivity and solubility Giant covalent lattices have very high melting and boiling points These compounds have a large number of covalent bonds linking the whole structure A lot of energy is required to break the lattice The compounds can be hard or soft Graphite is soft as the forces between the carbon layers are weak Diamond and silicon(IV) oxide are hard as it is difficult to break their 3D network of strong covalent bonds Graphene is strong, flexible and transparent which it makes it potentially a very useful material Most compounds are insoluble with water Most compounds do not conduct electricity however some do Graphite has delocalised electrons between the carbon layers which can move along the layers when a voltage is applied Graphene is an excellent conductors of electricity due to the delocalised electrons Buckminsterfullerene is a semi-conductor Diamond and silicon(IV) oxide do not conduct electricity as all four outer electrons on every carbon atom is involved in a covalent bond so there are no free electrons available Page 26 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Characteristics of Giant Covalent Structures Table Buckminster- Silicon Your notes Diamond Graphite Graphene Silicon fullerene dioxide Melting and Very high Very high Very high Low High Very high boiling point Transparent Grey-white Transparent Appearance Grey solid Transparent Black powder crystal solid crystals Electrical Non- Non- Good Very good Poor Poor conductivity conductor conductor Thermal Good Poor Very good Poor Good Good conductivity Piezoelectric Thinnest and —produces Hardest Good Other Soft and strongest Light and electric known natural mechanical properties slippery material to strong charge from substance strength exist mechanical stress Examiner Tip Although buckminsterfullerene is included in this section it is not classified as a giant structure as it has a fixed formula, C60. Page 27 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Intermolecular Forces Your notes Intermolecular Forces There are no covalent bonds between molecules in molecular covalent compounds. There are, however, forces of attraction between these molecules, and it is these which must be overcome when the substance is melted and boiled These forces are known as intermolecular forces There are three main types of intermolecular forces: London(dispersion) forces Dipole-dipole attraction Hydrogen bonding London (dispersion) forces The electrons in atoms are not static; they are in a state of constant motion It is therefore likely that at any given time the distribution of electrons will not be exactly symmetrical - there is likely to be a slight surplus of electrons on one side of the atoms Diagram to show how London (dispersion) forces arise Page 28 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes London (Dispersion) forces This is known as a temporary dipole It lasts for a very short time as the electrons are constantly moving Temporary dipoles are constantly appearing and disappearing Page 29 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Consider now an adjacent atom. The electrons on this atom are repelled by the negative part of the dipole and attracted to the positive part and move accordingly This is a temporary induced dipole Your notes There is a resulting attraction between the two atoms, and this known as London (dispersion) forces, after the German chemist, Fritz London London (dispersion) forces are present between all atoms and molecules, although they can be very weak They are the reason all compounds can be liquefied and solidified London (dispersion) forces tend to have strengths between 1 kJmol-1 and 50 kJmol-1. The strength of the London( dispersion) forces in between molecules depends on two factors: the number of electrons in the molecule Surface area of the molecules Number of electrons The greater the number of electrons in a molecule, the greater the likelihood of a distortion and thus the greater the frequency and magnitude of the temporary dipoles The dispersion forces between the molecules are stronger and the melting and boiling points are larger The enthalpies of vaporisation and boiling points of the noble gases illustrate this factor: Graph to show the effect of number of electrons on enthalpy of vaporisation and boiling point Page 30 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources As the number of electrons increases more energy is needed to overcome the forces of attraction between the noble gases atoms Your notes Surface area The larger the surface area of a molecule, the more contact it will have with adjacent molecules The greater its ability to induce a dipole in an adjacent molecule, the greater the London (dispersion) forces and the higher the melting and boiling points This point can be illustrated by comparing different isomers containing the same number of electrons: Diagram to show the effect of surface area on intermolecular forces Boiling points of molecules with the same numbers of electrons but different surface areas Dipole-dipole attractions Temporary dipoles exist in all molecules, but in some molecules there is also a permanent dipole In addition to the London (dispersion) forces caused by temporary dipoles, molecules with permanent dipoles are also attracted to each other by permanent dipole-dipole bonding Diagram to show permanent dipole-dipole interactions Page 31 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes The delta negative end of one polar molecule will be attracted towards the delta positive end of a neighbouring polar molecule This is an attraction between a permanent dipole on one molecule and a permanent dipole on another. Dipole-dipole bonding usually results in the boiling points of the compounds being slightly higher than expected from temporary dipoles alone it slightly increases the strength of the intermolecular attractions The effect of dipole-dipole bonding can be seen by comparing the melting and boiling points of different substances which should have London(dispersion) forces of similar strength Comparing butane and propanone For small molecules with the same number of electrons, dipole-dipole attractions are stronger than dispersion forces Butane and propanone have the same number of electrons Butane is a nonpolar molecule and will have only dispersion forces Propanone is a polar molecule and will have dipole-dipole attractions and dispersion forces Therefore, more energy is required to break the intermolecular forces between propanone molecules than between butane molecules The result is that propanone has a higher boiling point than butane Diagram to show the structures of butane and propanone Page 32 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes Comparing substances with permanent and temporary dipoles in smaller molecules with an equal number of electrons Dipole-induced dipole attraction Some mixtures might contain both polar and nonpolar molecules, for example HCl and Cl2 The permanent dipole of a polar molecule an cause a temporary separation of charge on a non-polar molecule This force is called dipole-induced dipole attraction This force acts in addition to the London dispersion forces that occur between nonpolar molecules and the dipole-dipole forces between polar molecules Diagram to show dipole-induced dipole attraction The polar HCl molecule causes a separation of charge on the nonpolar chlorine molecule Hydrogen bonding Hydrogen bonding is the strongest type of intermolecular force Hydrogen bonding is a special type of permanent dipole – permanent dipole bonding For hydrogen bonding to take place the following is needed: A species which has an O or N or F (very electronegative) atom with an available lone pair of electrons A hydrogen attached to the O, N or F When hydrogen is covalently bonded to an electronegative atom, such as O, N or F, the bond becomes very highly polarised Page 33 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources The H becomes so δ+ charged that it can form a bond with the lone pair of an O, N or F atom in another molecule Your notes Diagram to show polarisation of the H–O/N/F bond The electronegative atoms O or N have a stronger pull on the electrons in the covalent bond with hydrogen, causing the bond to become polarised Hydrogen bonds are represented by dots or dashes between H and the N/O/F element The number of hydrogen bonds depends on: The number of hydrogen atoms attached to O, N or F in the molecule The number of lone pairs on the O, N or F Diagram to show hydrogen bonding in ammonia Ammonia can form a maximum of one hydrogen bond per molecule Diagram to show hydrogen bonding in water Page 34 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes Water can form a maximum of two hydrogen bonds per molecule Van der Waals' forces The term Van der Waal's forces is used to include: London dispersion forces Dipole-induced dipole attractions Dipole-dipole attractions These forces occur between molecules (intermolecularly), as well within a molecule (intramolecularly) Diagram to show the difference between intermolecular and intramolecular forces The polar covalent bonds between O and H atoms are intramolecular forces and the permanent dipole – permanent dipole forces between the molecules are intermolecular forces Page 35 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Examiner Tip Your notes The term “London (dispersion) forces” refers to instantaneous induced dipole induced dipole forces that exist between any atoms or groups of atoms and should be used for non-polar species. Page 36 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Physical Properties of Covalent Substances Your notes Physical Properties of Covalent Substances The physical properties of molecular covalent compounds are largely influenced by their intermolecular forces If you know the type of intermolecular forces present you can predict the physical properties like melting and boiling point, solubility, and conductivity Melting and boiling point When covalent molecular substances change state you are overcoming the intermolecular forces The stronger the forces the more energy need to break the attraction Intermolecular forces are much weaker than covalent bonds, so many covalent substances are liquid or gases at room temperature Substance with a low melting and boiling point are said to be very volatile The strength of the intermolecular forces increases with the size of the molecule the increase in the polarity of the molecule Drawing the structure of the molecule helps identify and rank molecules according to boiling point as the following example shows Page 37 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Worked example Your notes Place these three molecules in the correct order from lowest to highest boiling point and explain your reasoning: CH3CH2CH2OH CH3COCH3 CH3CH2CH2CH3 Answer: Step 1: The first thing to do is find the approximate relative molecular mass: CH3CH2CH2OH = 60 CH3COCH3 = 58 CH3CH2CH2CH3 = 58 This tells you the molecules are approximately the same size so the dispersion forces will be similar Step 2: Draw the structures of the molecules and identify the intermolecular forces present So, the order of boiling from lowest to highest is: CH3CH2CH2CH3 ˂ CH3COCH3 ˂ CH3CH2CH2OH Solubility The general principle is that 'like dissolves like' so non-polar substances mostly dissolve in non-polar solvents, like hydrocarbons and they form dispersion forces between the solvent and the solute Polar covalent substances generally dissolve in polar solvents as a result of dipole-dipole interactions or the formation of hydrogen bonds between the solute and the solvent A good example of this is seen in organic molecules such as alcohols and water: Diagram to show the hydrogen bonding between ethanol and water Page 38 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes A hydrogen bond forms between oxygen atom on the ethanol and the hydrogen atom of the water As covalent molecules become larger their solubility can decrease as the polar part of the molecule is only a smaller part of the overall structure This effect is seen in alcohols for example where ethanol, C2H5OH, is readily soluble but hexanol, C6H13OH, is not Polar covalent substances are unable to dissolve well in non-polar solvents as their dipole-dipole attractions are unable to interact well with the solvent Giant covalent substances generally don't dissolve in any solvents as the energy needed to overcome the strong covalent bonds in the lattice structures is too great Conductivity As covalent substances do not contain any freely moving charged particles they are unable to conduct electricity in either the solid or liquid state However, under certain conditions some polar covalent molecules can ionise and will conduct electricity Some giant covalent structures are capable of conducting electricity due to delocalised electrons but they are exceptions to the general rule Comparing the Properties of Covalent Compounds Table Non—polar covalent Polar covalent Giant covalent Ionic substances substances substances substances Melting and boiling Low Low Very high Very high point Volatility Highest High Low Low Some solubility Solubility in polar Insoluble depending on Insoluble Soluble solvents molecular size Page 39 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Some solubility Solubility in non— Soluble depending on None Insoluble polar solvents Your notes molecular size None — except Electrical Only when molten or None None graphite, conductivity aqueous graphene Worked example Compound X has the following properties: Melting point Electrical conductivity solid molten 1450oC poor poor What is the most probable structure of X? A. Network covalent B. Polar covalent molecule C. Ionic lattice D. Metallic lattice Answer: The correct option is A because: A high melting point is characteristic of a giant structure, which could be metallic, ionic or covalent The poor conductivity as a liquid and solid would match a giant covalent or network covalent structure Page 40 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Chromatography Your notes Chromatography What is chromatography? Chromatography is a separation technique that enables the separation of mixtures and includes: paper chromatography thin-layer chromatography (TLC) These chromatography techniques make use of the principle that components in a mixture when dissolved in a fluid (mobile phase), will flow through another material (stationary phase) at varying rates The rate of separation depends upon how the components in the mixture interact with the stationary phase (their retention) and how soluble they are in the mobile phase Therefore the rate of separation depends on the intermolecular forces present For more information on performing chromatography and other separation techniques, see our revision notes on separating mixtures Examiner Tip Column chromatography (CC) and gas chromatography (GC), sometimes called gas-liquid chromatography (GLC), are other chromatographic techniques you may see in other resources They also work on the same principles as paper chromatography and TLC but with different stationary and mobile phases. These are beyond the scope of this specification. What is paper chromatography? In paper chromatography, the mobile phase is a solvent, and the stationary phase is the chromatography paper A pencil line is drawn on chromatography paper, this is the baseline (or origin), and spots of the sample are placed on it Pencil is used for this as ink would run into the chromatogram along with the samples The paper is then lowered into the solvent container, making sure that the pencil line sits above the level of the solvent so the samples don’t wash into the solvent container The solvent travels up the paper by capillary action, taking the sample with it As the solvent moves up the paper, the components in the mixture are dissolved to different extents depending on their solubility, so will travel with the solvent at different rates The extent of solubility depends on the intermolecular forces present The paper contains cellulose fibres which have hydroxyl (OH) groups along their structure Substances in the mixture that can form hydrogen bonds with the OH groups will be more attracted to the stationary phase than those which form weaker intermolecular forces This attraction to the stationary phase also affects the rate of separation Page 41 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Once the solvent front almost reaches the top of the paper, the paper is removed from the solvent and the solvent front is marked on the paper The separated components will appear as distinct spots on the paper Your notes Paper chromatography A dot of the sample is placed on the baseline and allowed to separate as the mobile phase flows through the stationary phase; the reference compound/s will also move with the solvent and are used to identify the components in the mixture. Examiner Tip If the sample does not travel with the solvent, it is because it is insoluble in that solvent An alternative solvent should be used Sometimes a number of solvents need to be trialled in order to find a suitable one in which the components of the sample are separated sufficiently What is thin layer chromatography (TLC)? TLC works in a similar way to paper chromatography but has a different stationary phase The stationary phase is a thin layer of an inert substance (e.g. silica or alumina) supported on a flat, unreactive surface (e.g. glass) The mobile phase, like paper chromatography, is a solvent Page 42 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Silica and alumina contain OH groups so can form hydrogen bonds with components in the sample The components are adsorbed onto the surface of the stationary phase Depending on the strength of interactions with the stationary phase, the separated components will Your notes travel particular distances through the plate What are retardation factors (Rf) values? The extent of separation of the component molecules in the investigated sample depends on their solubility in the mobile phase and the extent of adhesion to the stationary phase The Rf value is used to quantify the distance a particular component travels relative to the solvent front Rf values for compounds are calculated using measurements from the paper chromatogram or TLC plate and can be calculated using the Rf equation: distance travelled by component Rf = distance travelled by solvent These values can be used alongside other analytical data to deduce the composition of mixtures Calculation of Rf values Rf values can be calculated by taking 2 measurements from a chromatogram Page 43 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Examiner Tip Your notes Rf values are quoted as decimals and have no units as they are a ratio of distances When you divide two lengths measured in the same unit, those units cancel out, leaving you with a unitless number. Rf values will always be less than 1 as the component cannot travel further than the solvent front! Page 44 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Resonance Structures (HL) Your notes Resonance Structures The delocalisation of electrons can explain the structures of some species that don’t seem to fit with a Lewis formula Delocalised electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or one covalent bond The Lewis diagram for the nitrate (V) ion gives a molecule with a double and two single bonds There are three possible Lewis formulas These structures are called resonance structures However, studies of the electron density and bond length in the nitrate (V) ion indicate all the bonds are equal in length and the electron density is spread evenly between the three oxygen atoms The bond length is intermediate between a single and a double bond The actual structure is something in between the resonance structures and is known as a resonance hybrid Resonance structures of the nitrate (V) ion To determine the Lewis formula of the nitrate (V) ion first count the number of valence electrons and then add one electron for the negative charge on the ion Number of valence electrons = N + 3O +1 = 5 + ( 3 x 6) + 1 = 24 electrons Three structures are possible, consisting of a double bond and two singles: Resonance structures in the nitrate ion Dotted lines are used to show the position of the delocalised electrons Page 45 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes Resonance hybrid nitrate (V) ion The criteria for forming resonance hybrids structures is that molecules must have a double bond (pi bond) that is capable of migrating from one part of a molecule to another This usually arises when there are adjacent atoms with equal electronegativity and lone pairs of electrons that can re-arrange themselves and allow the double bonds to be in different positions Other examples that you should know about are the carbonate ion, benzene, ozone and the carboxylate anion Resonance Hybrids Table Below are some other resonance structures and hybrids that you should know: Species Lewis Resonance Formulas Resonance Hybrid Carbonate ion, CO32- Benzene, C6H6 Ozone, O3 Page 46 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes Carboxylate ion, RCOO- Page 47 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Benzene (HL) Your notes Benzene Structure of Benzene The structure of benzene was determined many years ago, by a chemist called Kekulé The structure consists of 6 carbon atoms in a hexagonal ring, with alternating single and double carbon-carbon bonds This suggests that benzene should react in the same way that an unsaturated alkene does However, this is not the case The structure of benzene Like other aromatic compounds, benzene has a planar structure due to the sp2 hybridisation of carbon atoms and the conjugated π system in the ring Each carbon atom in the ring forms three σ bonds using the sp2 orbitals The remaining p orbitals overlap laterally with p orbitals of neighbouring carbon atoms to form a π system Page 48 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources This extensive sideways overlap of p orbitals results in the electrons being delocalised and able to freely spread over the entire ring causing a π system The π system is made up of two ring-shaped clouds of electron density - one above the plane and Your notes one below it Benzene and other aromatic compounds are regular and planar compounds with bond angles of 120o The delocalisation of electrons means that all of the carbon-carbon bonds in these compounds are identical and have both single and double bond character The bonds all being the same length is evidence for the delocalised ring structure of benzene Evidence for delocalisation This evidence of the bonding in benzene is provided by data from: Enthalpy changes of hydrogenation Carbon-carbon bond lengths from X-ray diffraction Saturation tests Infrared spectroscopy Enthalpy changes of hydrogenation Hydrogenation of cyclohexene Each molecule has one C=C double bond The enthalpy change for the reaction of cyclohexene is -120 kJ mol-1 C6H10 + H2 → C6H12 ΔHꝋ = -120 kJ mol-1 Hydrogenation of benzene The Kekulé structure of benzene as cyclohexa-1,3,5-triene has three double C=C bonds: Structural formula of benzene The structural formula of benzene shows the alternating single and double bonds It would be expected that the enthalpy change for the hydrogenation of this structure would be three times the enthalpy change for the one C=C bond in cyclohexene C6H6 + 3H2 → C6H12 ΔHꝋ = 3 x -120 kJ mol-1 = -360 kJ mol-1 When benzene is reacted with hydrogen, the enthalpy change obtained is actually far less exothermic, ΔHꝋ = -208 kJ mol-1 This means that 152 kJ mol-1 less energy is produced than expected Therefore, the actual structure of benzene is more stable than the theoretical Kekulé model Page 49 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Carbon-carbon bond lengths Cyclohexene contains two different carbon-carbon bonds Your notes The single carbon-carbon bond (C-C) has a bond length of 154 pm The double carbon-carbon bond (C=C) has a bond length of 134 pm The Kekulé structure of benzene as cyclohexa-1,3,5-triene has three single C-C and three double C=C bonds It would be expected that benzene would have an equal mixture of bonds with lengths of 134pm and 154 pm Comparing carbon-carbon bond lengths The bond lengths observed in benzene are intermediate of single and double carbon-carbon bond lengths All of the carbon-carbon bond lengths are 140 pm suggesting that they are all the same and also intermediate of the single C-C and double C=C bonds Saturation tests Cyclohexene will decolourise bromine water as an electrophilic addition reaction takes place The Kekulé structure of benzene as cyclohexa-1,3,5-triene has three double C=C bonds It would, therefore, be expected that benzene would easily decolourise bromine water Benzene does not decolourise bromine water suggesting that there are no double C=C bonds Infrared spectroscopy Cyclohexene shows a peak in the range of 1620 - 1680 cm-1 for the double C=C bond within its structure The Kekulé structure of benzene as cyclohexa-1,3,5-triene has three double C=C bonds It would, therefore, be expected to also show a peak at 1620 - 1680 cm-1 for the double C=C bonds Page 50 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Benzene does not show a peak in this range for the double C=C bonds, instead, peaks are seen at around 1450, 1500 and 1580 cm-1 which are characteristic of double C=C bonds in arenes Your notes Page 51 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Expansion of the Octet (HL) Your notes Expansion of the Octet Expansion of the octet Elements in period 3 and above have the possibility of having more than eight electrons in their valence shell This is because there is a d-subshell present which can accommodate additional pairs of electrons This is known as the expansion of the octet The concept explains why structures such as PCl5 and SF6 exist, which have 5 and 6 bonding pairs of electrons respectively, around the central atom Five electron pairs Phosphorus pentachloride, PCl5 An example of a molecule with five bonding electron pairs is phosphorus pentachloride, PCl5 The total number of valence electrons is = P + 5Cl = 5 + (5 x 7) = 40 The number of bonding pairs is 5, which accounts for 10 electrons The remaining 30 electrons would be 15 lone pairs, so that each Cl has 3 lone pairs The completed Lewis formula looks like this: Lewis formula of PCl5 The octet of the central phosphorous atom has been expanded to hold 10 electrons Sulfur tetrafluoride, SF4 The total number of valence electrons is = S + 4F = 6 + (4 x 7) = 34 The number of bonding pairs is 4, which accounts for 8 electrons The remaining 26 electrons would be 13 lone pairs Fluorine cannot expand the octet so each fluorine would accommodate 3 lone pairs, accounting for 24 electrons, leaving 1 lone pair on the sulfur (sulfur has expanded the octet) The completed Lewis formula looks like this: Lewis formula of SF4 Page 52 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes The octet of the central sulfur atom has been expanded to hold 10 electrons Chlorine trifluoride, ClF3 The total number of valence electrons is = Cl + 3F = 7 + (3 x 7) = 28 The number of bonding pairs is 3, which accounts for 6 electrons The remaining 22 electrons would be 11 lone pairs Fluorine cannot expand the octet so each fluorine would accommodate 3 lone pairs, accounting for 18 electrons, leaving 2 lone pairs on the chlorine The completed Lewis formula looks like this: Lewis formula of ClF3 The octet of the central chlorine atom has been expanded to hold 10 electrons Triiodide ion, l3- The total number of valence electrons is = 3I + the negative charge = (3 x 7) + 1 = 22 The number of bonding pairs is 2, which accounts for 4 electrons The remaining 18 electrons would be 9 lone pairs Iodine would accommodate 3 lone pairs, accounting for 12 electrons, leaving 3 lone pairs on the central iodine The completed Lewis formula looks like this: Lewis formula of I3– The octet of the central iodine atom has been expanded to hold 10 electrons Six electron pairs Sulfur hexafluoride, SF6 Page 53 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources An example of a molecule with six bonding electron pairs is sulfur hexafluoride, SF6 The total number of valence electrons is = S + 6F = 6 + (6 x 7) = 48 The number of bonding pairs is 6, which accounts for 12 electrons Your notes The remaining 36 electrons would be 18 lone pairs, so that each F has 3 lone pairs, accounting for all electrons and no lone pairs The completed Lewis formula looks like this: Lewis formula of SF6 The octet of the central sulfur atom has been expanded to hold 12 electrons Bromine pentafluoride, BrF5 The total number of valence electrons is = Br + 5F = 7 + (5 x 7) = 42 The number of bonding pairs is 5, which accounts for 10 electrons The remaining 32 electrons would be 16 lone pairs Fluorine cannot expand the octet so each fluorine would accommodate 3 lone pairs, accounting for 30 electrons, leaving 1 lone pair on the bromine The completed Lewis formula looks like this: Lewis formula of BrF5 The octet of the central bromine atom has been expanded to hold 12 electrons Xenon tetrafluoride, XeF4 The total number of valence electrons is = Xe + 4F = 8 + (4 x 7) = 36 The number of bonding pairs is 4, which accounts for 8 electrons The remaining 28 electrons would be 14 lone pairs Each fluorine would accommodate 3 lone pairs, accounting for 24 electrons, leaving 2 lone pairs on the xenon The completed Lewis formula looks like this: Lewis formula of XeF4 Page 54 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes The octet of the central xenon atom has been expanded to hold 12 electrons Revisiting Valence Shell Electron Pair Repulsion Theory (VSEPR) Molecular shapes and the angles between bonds can be predicted using the three basic rules associated with valence shell electron pair repulsion theory known by the abbreviation VSEPR theory VSEPR theory consists of three basic rules: All electron pairs and all lone pairs arrange themselves as far apart in space as possible. Lone pairs repel more strongly than bonding pairs Multiple bonds behave like single bonds For more information about valence shell electron pair repulsion theory, see our revision note on Shapes of Molecules Molecular geometry versus electron domain geometry It is important to distinguish between molecular geometry and electron domain geometry in exam questions Molecular geometry refers to the shape of the molecules based on the relative orientation of the atoms Electron domain geometry refers to the relative orientation of all the bonding and lone pairs of electrons The Lewis formula for water enables us to see that there are four electron pairs around the oxygen so the electron domain geometry is tetrahedral However, the molecular geometry shows us there are two angled bonds so the shape is bent, angular, bent linear or V-shaped (when viewed upside down) Lewis formula and molecular shape of water Page 55 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources The Lewis formula of water and molecular shape can be used to determine the electron domain and molecular geometries Your notes Five electron domains Table of molecular geometries associated with five electron domains Electron Molecular domain Bonding pairs Lone pairs Shape example geometry geometry trigonal trigonal 5 0 bipyramidal bipyramidal trigonal 4 1 seesaw bipyramidal trigonal 3 2 T-shape bipyramidal Page 56 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes trigonal 2 3 Linear bipyramidal PCl5 is a symmetrical molecule so the electron cloud charge is evenly spread This means that it will be a non-polar molecule as any dipoles from the P–Cl bonds would be cancelled out SF4 and ClF3 are asymmetrical molecules having one or two lone pairs on one side of the central axis making the overall molecule polar Six electron domains Table of molecular geometries associated with six electron domains Electron Molecular domain Bonding pairs Lone pairs Shape example geometry geometry octahedral 6 0 octahedral square based octahedral 5 1 pyramid Page 57 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes octahedral 4 2 square planar SF6 is a symmetrical molecule so the electron cloud charge is evenly spread with 90o between the bonds This means that it will be a non-polar molecule as any dipoles from the S–F bonds would be cancelled out XeF4 is also non-polar despite having two lone pairs The bonding pairs are at 90o to the plane and the lone pairs are at 180o The lone pairs are arranged above and below the square plane resulting in an even distribution of electron cloud charge BrF5 is asymmetrical having a lone pair at the base of the pyramid making the overall molecule polar Worked example What is the electron domain geometry, molecular geometry and F-Xe-F bond angle of xenon difluoride, XeF2? Answer Count the valence electrons = Xe + 2F = 8 + (2 x 7) = 22 There are two bonding pairs, accounting for 4 electrons, so 18 electrons remain Each fluorine should have 3 lone pairs, accounting for 6 pairs or 12 electrons, which leaves 3 lone pairs on the xenon Xenon, therefore, has 2 bonding pairs and 3 lone pairs meaning it has: Electron domain geometry = trigonal bipyramid Molecular geometry = linear The bond angle will be 180o (having the same structure as the triiodide ion) Examiner Tip Lewis structure and Lewis formula may be used interchangeably, but Lewis formula is the preferred term in the specification. Page 58 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Formal Charge (HL) Your notes Formal Charge A limitation of the model of covalent bonding is that when drawing Lewis formulas for molecules, it is sometimes possible to come up with more than one structure while still obeying the octet rule This leads to the problem of deciding which structure is appropriate and is consistent with other information such as spectroscopic data on bond lengths and electron density One approach to determining which is the preferred structure is to determine the formal charge (FC) of all the atoms present in the molecule It is a kind of electronic book keeping involving the bonding, non-bonding and valence electrons Formal charge is described as the charge assigned to an atom in a molecule, assuming that all the electrons in the bonds are shared equally between atoms, regardless of differences in electronegativity The formula for calculating FC is FC= (number of valence electrons) - ½(number of bonding electrons) - (number of non-bonding electrons) or FC= V - ½B - N The Lewis formula which is preferred is the one which: the difference in FC of the atoms is closest to zero has negative charges located on the most electronegative atoms The process of drawing a Lewis formula has been covered previously, but here is a reminder of how to draw the Lewis formula of tetrachloromethane, CCl4, Diagram to show the Lewis formula of carbon tetrachloride Page 59 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes Steps in drawing the Lewis formula for CCl4 To work our the formal charge of the C and Cl atoms in the structure simply apply the FC formula: FC for carbon = (4) - ½(8) - 0 = 0 FC for chlorine = (7) - ½(2) - 6 = 0 Notice that formal charge is calculated for one of each type of atom and does not count the total number of atoms in the molecule Page 60 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Worked example Your notes What is the formal charge on boron in the BH4- ion? Answer Boron is a group 13 element, so has 3 valence electrons. Hydrogen has one valence electron and the charge on the ion is -1, so there are 8 electrons in the diagram. The Lewis formula is therefore: Lewis formula of BH4- The number of bonded electrons is 8 and the number of non-bonded electrons is zero. So the formal charge on B is: FC (B) = (3) - ½(8) - 0 = -1 It is possible to draw three resonance structures for sulfur dioxide, SO2: The three resonance structures of sulfur dioxide The first structure is an illustration of the expansion of the octet as the sulfur has 10 electrons around it Formal charge can be used to decide which of the Lewis formulas is preferred The FC on the first structure is as follows: FC on sulfur = (6) - ½(8) -(2) = 0 FC on oxygen = (6) - ½(4) -(4) = 0 Difference in FC = ΔFC = FCmax- FCmin = 0 The FC on the second (and third) structures is as follows: FC on sulfur = (6) - ½(6) -(2) = +1 Page 61 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources FC on left side oxygen = (6) - ½(2) -(6) = -1 FC on right side oxygen = (6) - ½(4) -(4) = 0 Your notes Difference in FC = ΔFC = FCmax- FCmin = 2 Worked example What is the formal charge on the two resonance structures shown? Resonance structures of carbon dioxide Deduce which is the preferred structure. Answer Structure I FC on carbon = (4) - ½(8) -(0) = 0 FC on oxygen = (6) - ½(4) -(4) = 0 Difference in FC = ΔFC = FCmax- FCmin = 0 Structure II FC on carbon = (4) - ½(8) -(0) = 0 FC on left oxygen = (6) - ½(6) -(2) = +1 FC on right oxygen = (6) - ½(2) -(6) = -1 Difference in FC = ΔFC = FCmax- FCmin = 2 Structure I is the preferred structure as the difference is zero Examiner Tip The term Lewis structure and Lewis formula mean the same thing. Page 62 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Sigma & Pi Bonds (HL) Your notes Sigma & Pi Bonds Bond overlap in covalent bonds A single covalent bond is formed when two non-metals combine Each atom that combines has an atomic orbital containing a single unpaired electron When a covalent bond is formed, the atomic orbitals overlap to form a combined orbital containing two electrons This new orbital is called the molecular orbital The shape of the molecular orbital is dependent on the shape of the atomic orbitals that combined The greater the atomic orbital overlap, the stronger the bond There are two main types of molecular orbital: a sigma (σ) bond and a pi (π) bond What is a sigma bond? Sigma (σ) bonds are formed from the head-on / end-to-end combination or overlap of atomic orbitals The electron density is concentrated along the bond axis (an imaginary line between the two nuclei) s orbitals overlap this way as well as p to p, and s with p orbitals The formation of sigma bonds from s orbitals Sigma orbitals can be formed from the head-on combination of s orbitals The formation of sigma bonds from an s and a p orbital Hydrogen fluoride has sigma bonds between s and p orbitals Page 63 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes The formation of sigma bonds from p orbitals Fluorine has sigma bonds between p orbitals The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond The electrostatic attraction between the electrons and nuclei bonds the atoms to each other A single covalent bond is always a sigma bond What is a pi bond? Pi (π) bonds are formed from the lateral (sideways) combination or overlap of adjacent p orbitals The two lobes that make up the π bond lie above and below the plane of the σ bond This maximises the overlap of the p orbitals A single π bond is drawn as two electron clouds one arising from each lobe of the p orbitals The two clouds of electrons in a π bond represent one bond containing two electrons The electron density is concentrated on opposite sides of the bond axis π bonds are only found within double and triple bonds The formation of a pi bond from p orbitals Page 64 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources π orbitals are formed by the lateral combination of p orbitals Your notes Examples of sigma & pi bonds Hydrogen The hydrogen atom has only one s orbital The s orbitals of the two hydrogen atoms will overlap to form a σ bond The formation of a sigma bond in hydrogen Direct overlap of the 1s orbitals of the hydrogen atoms results in the formation of a σ bond Ethene Each carbon atom uses three of its four electrons to form σ bonds Two σ bonds are formed with the hydrogen atoms One σ bond is formed with the other carbon atom The fourth electron from each carbon atom occupies a p orbital which overlaps sideways with another p orbital on the other carbon atom to form a π bond This means that the C-C is a double bond: one σ and one π bond The formation of a pi bond in ethene Page 65 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Overlap of the p orbitals results in the forming of a π bond in ethene Your notes The formation of sigma bonds and a pi bond in ethene Each carbon atom in ethene forms two sigma (σ) bonds with hydrogen atoms and one sigma (σ) bond with another carbon atom. The fourth electron is used to form a pi (π) bond between the two carbon atoms Ethyne This molecule contains a triple bond formed from two π bonds (at right angles to each other) and one σ bond Each carbon atom uses two of its four electrons to form σ bonds One σ bond is formed with the hydrogen atom One σ bond is formed with the other carbon atom Two electrons are used to form two π bonds with the other carbon atom The formation of sigma bonds and pi bonds in ethyne Page 66 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Your notes Ethyne has a triple bond formed from two π bonds and one σ bond between the two carbon atoms Predicting the Type of Bonds Whether sigma (σ) or pi (π) bonds are formed can be predicted by consideration of the combination of atomic orbitals Page 67 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Worked example Your notes What type of molecular orbitals are found in the following chemicals? 1. Nitrogen, N2 2. Hydrogen cyanide, HCN Answer 1: Nitrogen contains a triple bond and a lone pair on each nitrogen atom Nitrogen atoms have the electronic configuration 1s22s22p3 The triple bond is formed from one σ bond between the two nitrogen atoms and the lateral overlap of two sets of p orbitals on the nitrogen atoms to form two π bonds NOTE: The σ bond between the two nitrogen atoms is formed from the head-on overlap of two sp hybrid orbitals For more information, see our revision notes on Hybridisation The two π bonds are at right angles to each other The triple bond is formed from two π bonds and one σ bond Answer 2: Hydrogen cyanide contains a triple bond One σ bond is formed between the H and C atom A second σ bond is formed between the C and N atom Page 68 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources The remaining two sets of p orbitals of nitrogen and carbon will overlap to form two π bonds at right angles to each other Your notes Hydrogen cyanide has a triple bond formed from a σ bond and the overlap of two sets of p orbitals of nitrogen Page 69 of 73 © 2015-2024 Save My Exams, Ltd. · Revision Notes, Topic Questions, Past Papers  Head to www.savemyexams.com for more awesome resources Hybridisation (HL) Your notes Hybridisation What is hybridisation? The ground state of the electrons in a carbon atom is 1s22s22p2 This can be represented using a spin diagram as shown: Ground state of carbon Orbital spin diagram for carbon in the ground state This electronic structure would imply that carbon forms two covalent bonds using the unpaired 2p electrons Since the 2s electrons are paired there would be no reason for them to be involved in bonding However studies of carbon compounds show that carbon typically forms four covalent bonds that are all equal in energy This puzzle has been explained using the theory of bond hybridisation A half full p-subshell has a slightly lower energy than a partially filled one. The difference in energy between the 2s and 2p subshells is small, so an electron can fairly easily be promoted from the 2s to the 2p giving the new arrangement: Excited state of carbon Orbital spin diagram for carbon in the excited state