IB CHEMISTRY Topic 4 Bonding Higher level.

IB CHEMISTRY Topic 4 Bonding Higher level

4.1 Ionic bonding and structure
OBJECTIVES • Positive ions (cations) form by metals losing valence electrons. • Negative ions (anions) form by non-metals gaining electrons. • The number of electrons lost or gained is determined by the electron configuration of the atom. • The ionic bond is due to electrostatic attraction between oppositely charged ions. • Under normal conditions, ionic compounds are usually solids with lattice structures. • Deduction of the formula and name of an ionic compound from its component ions, including polyatomic ions. • Explanation of the physical properties of ionic compounds (volatility, electrical conductivity and solubility) in terms of their structure.

Ionic bonding Ionic Bond – electrostatic attraction between oppositely charged ions Ions with a negative charge are called anions. Ions with a positive charge are called cations.

Ionic bonding – other points
Non directional bond - strength of bond equal in all directions Conducts electricity when molten or in solution (aq) High melting point and boiling point hard solids Low volatility (tendency of a substance to vaporize) Brittle

Formation of ions - cations
Elements with 1, 2, or 3 valence electrons (metals or ‘losers’) react chemically by losing their valence electrons and become positive ions (cations) Element takes on the noble gas electron configuration/structure. Example: Na has one valence electron but once removed it’s electron structure becomes that of Ne. 2.8.1  or s22s22p63s1  1s22s22p6

Lewis structure - formation of sodium ion
Sodium has 1 electron in its outer shell If it loses this it will have no partially filled shells. Na Na Loses 1 electron Sodium atom (2.8.1) Sodium 1+ ion (2.8.0)

Formation of ions - cations
Group 1 Group 2 Group 3 +1 +2 +3 Li+ Na+ K+ Mg2+ Ca2+ Al3+

Formation of ions - anions
Elements with 5, 6, or 7 valence electrons (non-metals) react chemically by gaining enough electrons to become negative ions (anions) and take on the next noble gas electron structure. Result: Element takes on the next noble gas electron configuration/ structure. Example: Cl has seven valence electrons but once one is gained it’s electron structure becomes that of Ar. 2.8.7  or 1s22s22p63s23p5  1s22s22p63s23p6

Lewis structure - formation of chloride ion
Chlorine has 7 electrons in its outer shell. If it gains 1 electron it can achieve a full outer electron shell. It is, therefore, going to be able to accept the electron that the sodium wants to lose. Cl Cl Gains 1 electron (from sodium) Chlorine atom (2.8.7) Chlorine 1- ion (2.8.8)

Formation of ions - anions
Group 5 Group 6 Group 7 -3 -2 -1 N3- P3- O2- S2- F- Cl- Br-

Formation of ions – transition metals
Transition elements can form more than one ion. Example: Fe can form Fe2+ and Fe3+ Ground state: 1s22s22p63s23p64s23d6 Loss of 2 e-s: 1s22s22p63s23p63d6 Loss of 3 e-s: 1s22s22p63s23p63d5

Nonpolar covalent bond
Ionic bond formation Use the periodic table (metal + non-metal) or from the electronegativity values in your data booklet. The degree of polarity can be determined by looking at the difference in electronegativity. Roughly it breaks down like this: 0.0 – 0.4 Nonpolar covalent bond 0.4 – 1.8 Polar covalent bond >1.8 Ionic bond

Electronegativity and ionic bonding
The ability of atoms in a molecule to attract electrons to itself. On the periodic chart, electronegativity increases as you go… …from left to right across a row. …from the bottom to the top of a column.

Some polyatomic ions nitrate NO3- phosphate PO43- hydroxide OH-
Common polyatomic ions formed by non-metals in period 2 and 3 that you must know! Chemical name Chemical formula nitrate NO3- phosphate PO43- hydroxide OH- ammonium ion NH4+ sulfate SO42- hydrogen carbonate HCO3- carbonate CO32-

Describing the lattice structure
Example: Small sodium chloride ionic lattice

Describing the lattice structure
Endlessly repeating lattice of ions…depends on the size of the crystal lattice...although arrangement of ions stays the same The coordination number is (6,6) meaning each Na+ is surround by 6 Cl-, each Cl- is surrounded by 6 Na+.

4.2. Covalent bonding OBJECTIVES • A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei. • Single, double and triple covalent bonds involve one, two and three shared pairs of electrons respectively. • Bond length decreases and bond strength increases as the number of shared electrons increases. • Bond polarity results from the difference in electronegativities of the bonded atoms. • Deduction of the polar nature of a covalent bond from electronegativity values.

Covalent Bonding A covalent bond is the electrostatic attraction between a shared pair of electrons and positively charged nuclei These electrons are called bonding electron pairs. Electrons not involved in bonding are called non-bonding electron pairs.

Bonding spectrum 100% covalent 100% ionic A B A B A+ B- Increasing DEN
Increasing polarity Transfer

Problem 1: Which is more ionic, NaCl, or LiCl?
Δχ(NaCl) =│χa - χb│ = 0.9 – 3.2 = 2.3 Ʃχ(NaCl) = (χa + χb)/2 = 4.1/2 = 2.05 Δχ(LiCl) =│χa - χb│ = 1.0 – 3.2 = 2.2 Ʃχ(LiCl) = (χa + χb)/2 = 4.2/2 = 2.1 Therefore NaCl is more ionic.

polarity

Polarity The partial positive charge is written as δ+ and the partial negative charge is written as δ-. The overall charge is called a dipole moment and is represented by a vector as shown opposite.

Polarity Just because a molecule possesses polar bonds it does not mean that the molecule as a whole will be polar.

Polarity The molecule is polar if the electron densities are not symmetrical such that there is a net dipole moment.

Polarity

Bonds strengths

< < Bonds strengths Each bond contains 2 electrons.
Triple bonds are shorter and stronger than double bonds. Double bonds are shorter and stronger than single bonds. < <

Bonds strengths A strong attraction between two nuclei for the electrons creates a short bond length. Therefore short bond lengths have high bond dissociation enthalpy. Type of bonding Ionic Covalent Metallic Nature of bonding Electrostatic attraction between positive and negative ions. Electrostatic attraction of a shared pair of electrons between atoms. Electrostatic attraction between lattice of positive metal ions and delocalised outer shell electrons. Strength of bonds The smaller the ions and the greater the charge on the ions, the stronger the attraction between the ions (usually). This is due to a greater charge density within the structure. The shorter the bond, the stronger the bond (usually). Double bonds are stronger than single bonds, while triple bonds are stronger than double bonds. The smaller the metal ions, the greater the charge on the ions, and the more delocalised outer shell electrons there are, the stronger the attraction between the ions and electrons (usually). This is due to a greater charge density within the structure.

4.3 Covalent structures OBJECTIVES • Lewis (electron dot) structures show all the valence electrons in a covalently bonded species. • The “octet rule” refers to the tendency of atoms to gain a valence shell with a total of 8 electrons. • Some atoms, like Be and B, might form stable compounds with incomplete octets of electrons. • Resonance structures occur when there is more than one possible position for a double bond in a molecule. • Shapes of species are determined by the repulsion of electron pairs according to VSEPR theory. • Carbon and silicon form giant covalent/network covalent structures. • Deduction of Lewis (electron dot) structure of molecules and ions showing all valence electrons for up to four electron pairs on each atom. • The use of VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three and four electron domains. • Prediction of bond angles from molecular geometry and presence of non-bonding pairs of electrons. • Prediction of molecular polarity from bond polarity and molecular geometry. • Deduction of resonance structures, examples include but are not limited to C6H6, CO32- and O3 . • Explanation of the properties of giant covalent

Lewis structures

Covalent Bonding – Lewis structures
Lewis structures are representations of molecules showing all electrons, bonding and nonbonding. Each bond contains 2 electrons. Elements pair up according to the octet rule – elements gain or lose electrons in order to have the electron configuration of a noble gas. Cl2

How to draw Lewis structures deductively
1. Draw out the atoms with the least occurring atom as the central atom (if applicable). 2. Draw in the valence electrons. 3. Connect up one bond from each of the outside atoms to the central atom.

4. Count up the electrons to see if all the atoms follow the octet rule (here oxygens only have 7).
5. If not, then add extra bonds until they do.

6. If this cannot be done, a coordinate covalent bond – where one atom donates both electrons to the bonding pair, may need to be added. Eg. CO (Free radicals are the other option) 7. Clean up the structure making it easy to read (electrons in pairs, o and x). Structural formula: O=O Structural formula:

Exceptions to the Octet: Boron
Boron forms stable compounds with just 3 valence electrons These compounds are highly reactive:

Exceptions to the Octet: Beryllium
Beryllium forms stable compounds with just 2 valence electrons Compounds that break the octet rule are often toxic and dangerous, ready to react so the octet rule is obeyed!

O N2 CO HCN O C O

C2H4 (ethene) C2H2 (ethyne)
CO NH4+ H3O+ Draw this one now or Draw this one now

VSEPR theory

Shapes of molecules Determined by number of valence electrons of the central atom 3-D shape a result of bonded pairs and lone pairs of electrons VSEPR theory (valence-shell-electron-pair repulsion) states the best arrangement of a given number of electron domains (negative charge centres) is the one that minimizes the repulsions among them.

Shapes of molecules Simply put, electron pairs, whether they be bonding or nonbonding, repel each other. By assuming the electron pairs are placed as far as possible from each other, we can predict the shape of the molecule.

Shapes – electron domains
An electron domain is the number of bonds/electron pairs there are on the central atom Eg. SO2 has 3 electron domains CH4, NH3, NH4+ all have 4 electron domains

Using the VSEPR Model 1. Draw the electron-dot structure
2. Identify the central atom 3. Count the total number of electron pairs around central atom 4. Predict the electron shape 5. Predict the shape of the molecule using the bonding atoms

Rule of thumb: Let each non-bonding electron pair bend by a further 2-3⁰

Shapes (if no non-bonding electron pairs)
Molecules, or ions, possessing ONLY BOND PAIRS of electrons fit into a set of standard shapes. All the bond pair-bond pair repulsions are equal. All you need to do is to count up the number of bond pairs and chose one of the following examples... C A covalent bond will repel another covalent bond BOND BOND PAIRS SHAPE ANGLE(S) EXAMPLE 2 LINEAR º BeCl2 3 TRIGONAL PLANAR 120º AlCl3 4 TETRAHEDRAL º CH4 5 TRIGONAL BIPYRAMIDAL 90º & 120º PCl5 6 OCTAHEDRAL 90º SF6

Shapes (with non-bonding electron pairs)
If a molecule, or ion, has lone pairs on the central atom, the shapes are slightly distorted away from the regular shapes. This is because of the extra repulsion caused by the lone pairs. BOND PAIR - BOND PAIR < LONE PAIR - BOND PAIR < LONE PAIR - LONE PAIR O O O As a result of the extra repulsion, bond angles tend to be slightly less as the bonds are squeezed together.

Methane CH4 109.5° TETRAHEDRAL H C H C BOND PAIRS 4 LONE PAIRS 0 C H
Carbon - has four electrons to pair up Hydrogen - 1 electron to complete shell Four covalent bonds are formed C and H now have complete shells BOND PAIRS 4 LONE PAIRS 0 109.5° H C BOND ANGLE... SHAPE... 109.5° TETRAHEDRAL

Ammonia NH3 N H N H Nitrogen has five electrons in its outer shell
BOND PAIRS 3 LONE PAIRS 1 TOTAL PAIRS 4 H Nitrogen has five electrons in its outer shell It cannot pair up all five - it is restricted to eight electrons in its outer shell It pairs up only three of its five electrons 3 covalent bonds are formed and a pair of non-bonded electrons is left As the total number of electron pairs is 4, the shape is BASED on a tetrahedron

Ammonia NH3 N H N BOND PAIRS 3 LONE PAIRS 1 TOTAL PAIRS 4 H The shape is based on a tetrahedron but not all the repulsions are the same LP-BP REPULSIONS > BP-BP REPULSIONS The N-H bonds are pushed closer together Lone pairs are not included in the shape 107° H N H N N H ANGLE ° SHAPE... PYRAMIDAL H H

Ammonia NH3 N H N BOND PAIRS 3 LONE PAIRS 1 TOTAL PAIRS 4 H

Water H2O O H O H Oxygen has six electrons in its outer shell
BOND PAIRS 2 LONE PAIRS 2 TOTAL PAIRS 4 H Oxygen has six electrons in its outer shell It cannot pair up all six - it is restricted to eight electrons in its outer shell It pairs up only two of its six electrons 2 covalent bonds are formed and 2 pairs of non-bonded electrons are left As the total number of electron pairs is 4, the shape is BASED on a tetrahedron

Water H2O O H O BOND PAIRS 2 LONE PAIRS 2 TOTAL PAIRS 4 H The shape is based on a tetrahedron but not all the repulsions are the same LP-LP REPULSIONS > LP-BP REPULSIONS > BP-BP REPULSIONS The O-H bonds are pushed even closer together Lone pairs are not included in the shape O 104.5° H O O H H ANGLE ° SHAPE... ANGULAR H H H

Ammonium NH4+ NH4+ N+ BOND PAIRS 4 LONE PAIRS 0 TETRADHEDRAL

Hydronium ion H3O+

Bromine triflouride BrF3
BOND PAIRS 3 LONE PAIRS ’T’ SHAPED ANGLE <90° F Br

:O:: S:O: S Sulphur dioxide SO2
S has 2 bonded atoms , 1 lone pair (electron cloud), 120°, bent, V-shaped :O:: S:O: S .. O O

Ethane C2H3, Ethene C2H4, Ethyne C2H2

Carbon dioxide CO2 O C O C O O C
The shape of a compound with a double bond is calculated in the same way. A double bond repels other bonds as if it was single e.g. carbon dioxide O C O C O Carbon - needs four electrons to complete its shell Oxygen - needs two electron to complete its shell The atoms share two electrons each to form two double bonds DOUBLE BOND PAIRS 2 LONE PAIRS 0 O C 180° Double bonds behave exactly as single bonds for repulsion purposes so the shape will be the same as a molecule with two single bonds and no lone pairs. BOND ANGLE... SHAPE... 180° LINEAR

By Deduction Resonance Structures

Resonance When a Lewis structure allows for the same arrangement of atoms but a different but equally valid arrangement of electrons, resonance occurs eg. Ozone, O3 The resonance hybrid is actually the dotted line below. The electrons are said to be delocalized, and are spread over two bonding orbitals. These electrons are located in the p orbitals. The overlapping p orbitals create a new shape called a pi bond () in a normal bonding situation. In resonance these  bonds then all overlap as well.

Nitrate NO3- When writing Lewis structures for species like the nitrate ion, we draw resonance structures to more accurately reflect the structure of the molecule or ion.

In reality, each of the four atoms in the nitrate ion has a p orbital.
The double bond on N=O is normally a  bond. The p orbitals on all three oxygens overlap with the p orbital on the central nitrogen.

This means the  electrons are not localized between the nitrogen and one of the oxygens, but rather are delocalized throughout the ion.

Benzene C6H6

The organic molecule benzene has a p orbital on each carbon atom.

In reality the  electrons in benzene are not localized, but delocalized.
The even distribution of the  electrons in benzene makes the molecule unusually stable.

Carbonate CO32-

Carbon Structures

Allotropes Allotropes are compounds of the same element that differ in structure Carbon can bond with itself in at least three major different inorganic ways giving us 3 very different materials 1. Diamond 2. Graphite (graphene and pyrolytic carbon) 3. Fullerenes - Buckyballs and nanotubes

Diamond Each carbon bonded to 4 other carbons
HL: Carbons are bonded via sp3 hybridization to 4 other carbon atoms forming a giant network covalent compound.

Properties of Diamond High melting point due to strong directional covalent bonds (3550 C) Extremely hard because it is difficult to break atoms apart or move them in relation to one another No electrical conductivity because electrons are localized in specific bonds Insoluble in polar and non-polar solvents because molecular bonds are stronger than any intermolecular forces

Graphite Each carbon bonded to 3 other carbons forming sheets of graphene. These graphene sheets slide over each other. HL: Carbon atoms are bonded via sp2 hybridization. Carbon atoms form sheets of six sided rings with p-orbitals perpendicular from plane of ring

Graphite Structure Carbon has 4 valence e- to bond with. 3 are used for closest atoms in rings. 1 is delocalized in p-orbitals The presence of p-orbitals allows for stronger London forces that hold the sheets together

Properties of Graphite
Different from Diamond Conducts electricity because of delocalized electrons, use as a conductor, lack of vibration between layers however make it an insulator not a conductor of heat Slippery can be used as lubricant, sheets can easily slip past each other (think of a deck of cards), use as pencil ‘leads’ Same as Diamond High melting point (higher actually because of delocalized e-, 3653C), use as carbon fibre Insoluble (same reason)

Fullerenes Buckyballs: spherical Nanotubes: tube shaped
Both have very interesting properties Super strong Conduct electricity and heat with low resistance Free radical scavenger

Buckyballs or Buckminsterfullerene
Carbon atoms bond in units of 60 atoms (C-60) forming a structure similar to a soccerball with interlocking six sided and five sided rings. HL: sp2 hybridization Extra p-orbitals form pi bonds resulting in... Properties: Electrical conductivity Stronger covalent bonds, therefore stronger materials

QUART Silicon Structures

Silicon dioxide (SiO2) SiO2 repeating unit as every oxygen links in with a second Si atom Structurally think tetrahedral, SiO4

Difference between Si and C (diamond)
Si is larger than C so the Si-O bond length is greater The greater the bond length, the lower the bond enthalpy (energy) is This means it is easier to break Therefore Si is more reactive than C (diamond)

4.4 Intermolecular forces
OBJECTIVES • Intermolecular forces include London (dispersion) forces, dipole-dipole forces and hydrogen bonding. • The relative strengths of these interactions are London (dispersion) forces < dipole-dipole forces < hydrogen bonds. • Deduction of the types of intermolecular force present in substances, based on their structure and chemical formula. • Explanation of the physical properties of covalent compounds (volatility, electrical conductivity and solubility) in terms of their structure and intermolecular forces.

Intermolecular forces (IMF)
Intermolecular forces – attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding (Van der Waal’s forces include only 1 & 2 interactions) 1. London (dispersion) forces - the interactions between non-polar molecules from temporary dipoles (induced dipoles). 2. Dipole-dipole forces – from permanent dipoles of polar molecules. 3. Hydrogen bonding – from bonds with H and either N, O or F only.

London Forces Weakest of the three
also called the Dispersion forces (see animation) Caused by the motion of electrons around a nucleus Electrons sometimes are asymmetrical about an atom, leading to temporary dipoles also called instantaneous induced dipoles The formation of a dipole in one molecule can cause an opposite dipole to form in a nearby molecule Dispersion forces increases with atomic radii

Increasing London forces:
Large radius contains more electrons Size of electron cloud ie. with longer molecules Less branching to allow for more surface area interaction.

Dipole-dipole forces Occurs between polar molecules, also called permanent dipole The partially charged ends of the molecules attract and repel each other Don’t forget London forces are always acting as well

Dipole-dipole forces Polar molecules (polar covalent) have slightly charged ends Opposites attract. Large electronegative difference = stronger attraction. London Forces should be explained as a temporary polarization of moving electron clouds. The causes depolarization in neighbouring atom electron clouds VIA induction. Atoms with more electrons will have more or stronger London Forces due to increased polarization strength.

Hydrogen Bonds Hydrogen bonds are a special form of dipole attraction
A hydrogen bond is the attractive force that forms between an unshared electron pair and a hydrogen atom covalently bonded with a strongly electronegative element (NOF). Hydrogen is the only reactive element without an underlying layer of electrons – this makes a hydrogen bond about 5% as strong as an average covalent bond

Hydrogen Bonds Covalent bond Hydrogen bond Hydrogen Bonding (F, O or N bonded to H) with a free lone electron pair Due to small size and high electronegativity of non metals Creates a large charge difference Basically a super strong dipole-dipole bond London Forces should be explained as a temporary polarization of moving electron clouds. The causes depolarization in neighbouring atom electron clouds VIA induction. Atoms with more electrons will have more or stronger London Forces due to increased polarization strength.

Boiling point trends Phase change when IMF are overcome
Be sure to explain using the words IMF and how they affect the bonds BETWEEN particles. London Forces are ALWAYS present!!! BP tends to increase with increased Molecular Mass Exception: Hydrogen bonding is very strong (O, F, N) which explains the abnormally high BP Strength of H-Bonding due to Large electronegativity Small size allows for close approach of other dipoles CH4 has a very low BP because it lacks any polarity, the only bonds that must be overcome are vaLondon Forces

General physical property trends
London forces: Lowest MP, Non polar Butane (C4H10) Dipole-dipole: Slightly miscible Propanone C3H6O Hydrogen Bonding: Miscible with polar substances H2O Ionic Bonding: Only conducts electricity when liquid or aqueous. (Decomposition when it does) NaCl Metallic Bonding: Conducts electricity, not water soluble, MP regulated by, valance, size and packing. Fe Giant Covalent: Highest MP, Insoluble in both non-polar and polar solvents. Does not conduct electricity except for graphite. Diamond and Graphite (Allotropes) Increasing Melting Point

1. 2. 3. Dipole – Induced dipole Van der Waal’s Forces
(Induced dipole – Induced dipole)

How IMF affect the boiling points of substances
When a liquid turns into a gas the attractive forces between the particles are completely broken so boiling point is a good indication of the strength of intermolecular forces. Covalent macromolecular structures, such as diamond, have extremely high melting and boiling points. Metals and ionic compounds also tend to have relatively high boiling points due to ionic attractions. Hydrogen bonds are in the order of 1/10th the strength of a covalent bond whereas London forces are in the order of less than 1/100th of a covalent bond. The weaker the attractive forces the more volatile the substance.

IMF – comparing compounds
The presence of hydrogen bonding can be illustrated by comparing the boiling points of: HF and HCl H2O and H2S NH3 and PH3 CH3OCH3 and CH3CH2OH CH3CH2CH3, CH3CHO, and CH3CH2OH

4.5 Metallic bonding OBJECTIVES • A metallic bond is the electrostatic attraction between a lattice of positive ions and delocalized electrons. • The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion. • Alloys usually contain more than one metal and have enhanced properties. • Explanation of electrical conductivity and malleability in metals. • Explanation of trends in melting points of metals. • Explanation of the properties of alloys in terms of non-directional bonding.

Metallic bonding Metallic bonds are the electrostatic attraction between a lattice of positive ions and delocalized electrons.

Metallic bonding Strength of metallic bond increases with charge of ion Strength of metallic bond increases with smaller size of the ion

Metallic bonding – electrical conductivity and malleability
For conductivity to occur the substance must possess electrons or ions that are free to move. Metals (and graphite) contain delocalized electrons and are excellent conductors. Metals are malleable which means they can be bent and reshaped under pressure. This due to the close-packed layers of positive ions can slide over each other without breaking more bonds than are made. See difference between covalent brittleness and metals malleability

Alloys Alloys are mixtures of metals and one or more other element (usually also a metal). They produce increased strength, durability, hardness, resistance to corrosion, and magnetic properties. This is due to the different elements being of different sizes preventing the easy sliding between atoms.

Alloy examples Brass is a mixture of copper and zinc. Steel is a mixture of iron and carbon (and others).

Type of bonding Covalent Ionic Metallic Hydrogen bonding Dipole-dipole forces London forces Cause Two strong electronegative atoms One strong electronegative and one weak electronegative atom Two weak electronegative atoms An electronegative atom from one molecule and a hydrogen atom from another molecule Two polar molecules Two non-polar molecules DECREASING STRENGTH  Variation of strength within group - The more positive the nucleus the better The more electrons the better The more polar the better More electrons the better Example Diamond NaCl Cu H2O SO2 O2

14.1 Further aspects of covalent bonding and structure (PTO)
Higher level OBJECTIVES • Covalent bonds result from the overlap of atomic orbitals. A sigma bond (σ) is formed by the direct head-on/end-to-end overlap of atomic orbitals, resulting in electron density concentrated between the nuclei of the bonding atoms. A pi bond (π) is formed by the sideways overlap of atomic orbitals, resulting in electron density above and below the plane of the nuclei of the bonding atoms. • Formal charge (FC) can be used to decide which Lewis (electron dot) structure is preferred from several. The FC is the charge an atom would have if all atoms in the molecule had the same electronegativity. FC = (Number of valence electrons)-½(Number of bonding electrons)-(Number of non-bonding electrons). The Lewis (electron dot) structure with the atoms having FC values closest to zero is preferred. • Exceptions to the octet rule include some species having incomplete octets and expanded octets. • Delocalization involves electrons that are shared by/between all atoms in a molecule or ion as opposed to being localized between a pair of atoms.

14.1 Further aspects of covalent bonding and structure (cont...)
Higher level OBJECTIVES • Resonance involves using two or more Lewis (electron dot) structures to represent a particular molecule or ion. A resonance structure is one of two or more alternative Lewis (electron dot) structures for a molecule or ion that cannot be described fully with one Lewis (electron dot) structure alone. • Prediction whether sigma (σ) or pi (π) bonds are formed from the linear combination of atomic orbitals. • Deduction of the Lewis (electron dot) structures of molecules and ions showing all valence electrons for up to six electron pairs on each atom. • Application of FC to ascertain which Lewis (electron dot) structure is preferred from different Lewis (electron dot) structures. • Deduction using VSEPR theory of the electron domain geometry and molecular geometry with five and six electron domains and associated bond angles. • Explanation of the wavelength of light required to dissociate oxygen and ozone. • Description of the mechanism of the catalysis of ozone depletion when catalysed by CFCs and NOx.

FC = valence – ½ bonding – non-bonding
Higher level Formal charge (FC) Formal charge (FC) is: FC = valence – ½ bonding – non-bonding electrons electrons electrons To determine the most favoured Lewis structure we chose: The structure with a FC closest to zero The structure with the negative FC on the most electronegative atom

Higher level Problem 1: Determine the FC on the following compound: FC = valence electrons – ½ bonding electrons - non-bonding electrons FC(N) = 5 - ½ x = 0 FC(C) = 4 - ½ x = 0 FC(S) = 6 - ½ x = -1

Higher level The electronegativity of N is 3.0 and S is 2.5, so the negative FC is best suited to N. Therefore the bottom structure is correct.

Problem 2: Determine which is the preferred structure.
Higher level Problem 2: Determine which is the preferred structure.

Higher level FC(B) = 3 – 1/2x6 – 0 = 0 FC(F) = 7 – ½ x 2 – 6 = 0 FC(B) = 3 – ½ x8 – 0 = -1 FC(F single bond) = 7 – ½ 2 x -6 = 0 FC(F double bond) = 7 – ½ 4 x -4 = +1 The first structure has the most atoms on zero FC, so is the correct structure.

VSEPR with expanded shells
Higher level VSEPR with expanded shells Elements in the third period (Al to Ar) and below can promote an electron to have more than 4 electron domains. This is called an expanded octet. eg. P = [Ne] 3s13p33d1 making PCl5 5 electron domains: non-bonding electrons in the equitorial positions of a trigonal bipyramidal shape 6 electron domains: non-bonding electrons in the axial positions of a octahedron or square bipyramidal.

Trigonal Bipyramidal Electron Domain
Higher level Trigonal Bipyramidal Electron Domain There are two distinct positions in this geometry: Axial Equatorial Non-bonding electron pairs fill the equatorial bonds first because they are only repelled by 2 bonds (axial). Axial bonds are repelled by 3 bonds (equatorial).

Higher level Trigonal Bipyramidal Electron Domain Don’t forget non-bonding electron pairs distort the angles!

Higher level There are four distinct molecular geometries in this domain: Trigonal bipyramidal, Seesaw, T-shaped, Linear

Octahedral Electron Domain
Higher level All positions are equivalent in the octahedral domain. There are three molecular geometries: Octahedral, Square pyramidal, Square planar

Overlapping orbitals (s and p)
Higher level s (sigma) bonds result from the axial overlap of orbitals p (pi) bonds result from the sideways overlap of parallel p orbitals double bonds form by one s (sigma) and one p (pi) bond triple bonds form by one s (sigma) and two p (pi) bonds

σ bonds Sigma bonds (σ) are characterized by Head-to-head overlap.
Higher level Sigma bonds (σ) are characterized by Head-to-head overlap. Cylindrical symmetry of electron density about the internuclear axis.

π bonds Pi bonds (π) are characterized by Side-to-side overlap.
Higher level Pi bonds (π) are characterized by Side-to-side overlap. Electron density above and below the internuclear axis.

Higher level Single bonds are always  bonds, because  overlap is greater, resulting in a stronger bond and more energy lowering. In a multiple bond one of the bonds is a  bond and the rest are  bonds.

Higher level Note the weakness of the extra bond – the electron overlap is weaker, yet it is still an extra bond so the energy required to break them is still a bit stronger. Note the smaller bond lengths have greater bond dissociation enthalpies.

Higher level CASE STUDY CONTINUED Ozone

Resonance and bond order
Higher level Resonance and bond order Bond order is the measurement of the number of electrons involved in bonds between two atoms in a molecule. If the bond order is 0 then there is no bonding. Bond order = 𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑝𝑎𝑖𝑟𝑠 𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛𝑠 Problem 1: Determine the bonding order of ozone. Bond order = 3/2 = 1.5 Bond length is measured by X-ray crystallography

Problem 2: Determine bond order
Higher level Problem 2: Determine bond order Bond order = 𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑏𝑜𝑛𝑑𝑖𝑛𝑔 𝑝𝑎𝑖𝑟𝑠 𝑡𝑜𝑡𝑎𝑙 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛𝑠 = 4 3

Importance of ozone Higher level Ozone at ground level reacts with chemicals to form smog, harms respiratory systems and degrades materials (eg. rubber). Ozone at atmospheric levels is essential for life. It absorbs harmful UV radiation (which causes cancer and inhibition of photosynthesis). Exothermic reactions cause a temperature inversion in the stratosphere – this warm layer prevent convection keeping the layers of Earth’s atmosphere stable.

Ozone formation review
Higher level Review: O2 requires photons with more energy as O3 only has the energy of a 1 ½ bond. O2(g) + UV (λ<242nm)  O•(g) + O•(g) O3(g) + UV (λ<330nm)  O2(g) + O•(g) Bond angle <120⁰ ~ 117⁰

UV and O2/O3 bond calculations
Higher level UV and O2/O3 bond calculations Equations: Ephoton = hf E = energy in one photon (J) h = planks constant (6.63x10-34Js) f = frequency (Hz) v = fλ v = velocity = c = speed of light = (3.00 x 108 m/s) λ = wavelength (nm)

Higher level Problem 1: The bond energy in ozone is 363kJ/mol. Calculated the wavelength of UV radiation needed to break this bond. Ephoton = J/mol / 6.02x1023photons/mol = 6.03x10-19J per photon Ephoton = hf so f = Ephoton / h = 6.03x10-19J / 6.63x10-34Js = 9.09 x 1014s v = fλ so λ = v/f = 3.00 x 108 m/s / 9.09 x 1014s = 3.30 x 10-7m = 330nm

Ozone formation Higher level UV light breaks the bonds in oxygen to create free radicals (atoms/molecules with unpaired valence electrons and so highly reactive): O2(g) + UV (λ<242nm)  O•(g) + O•(g) This then undergoes an exothermic reaction: O•(g) + O2(g)  O3(g) • is the symbol to denote a radical. O• actually has 6 electrons in the valence shell (not one).

Ozone depletion Higher level The reverse also occurs creating an O3 cycle. Depletion is also exothermic. O3(g) + UV (λ<330nm)  O2(g) + O•(g) O3(g) + O•(g)  2O2(g)

Ozone destroying chemicals
Higher level Ozone destroying chemicals CFCs – chloroflourocarbons, catalyze the breakdown of ozone NOx – also catalyze the breakdown of ozone These compounds are free radicals (not complete octets) and hence are highly reactive.

Catalytic ozone destruction
Higher level The balance of the ozone cycle is disrupted by the depletion of ozone from: nitrogen oxides: NO•(g) + O3(g)  NO2•(g) + O2(g) NO2•(g) + O•(g)  NO•(g) + O2(g) Nitrogen oxides slowly diffuse their way from the troposphere. Aircraft fly in the lower stratosphere causing direct injection of NOx. NO can last from 22 to 111 years before breaking down. Nitrogen monoxide free radical (unpaired valence electrons) Nitrogen dioxide free radical (unpaired valence electrons) Nitrogen monoxide is the primary reaction

Catalytic ozone destruction
Higher level 2. chlorofluorocarbons (CFCs): CCl2F2(g)  CClF2•(g) + Cl•(g) by UV radiation Cl•(g) + O3(g)  O2(g) + ClO•(g) ClO•(g) + O•(g)  O2(g) + Cl•(g) ClO•(g) + O3 (g)  2O2(g) + Cl•(g) As with NO, the Cl free radical catalyzes the decomposition reaction thousands of times: Summary: O3(g) + O•(g)  2O2(g)

Higher level 14.2 Hybridization OBJECTIVES • A hybrid orbital results from the mixing of different types of atomic orbitals on the same atom. • Explanation of the formation of sp3, sp2 and sp hybrid orbitals in methane, ethene and ethyne. • Identification and explanation of the relationships between Lewis (electron dot) structures, electron domains, molecular geometries and types of hybridization.

Higher level It’s hard to imagine tetrahedral, trigonal bipyramidal, and other geometries arising from the atomic orbitals we recognize.

Orbital hybridization
Higher level Orbital hybridization Hybridization is the mixing of different types of orbitals to produce new types of orbitals. It is a mathematical procedure. The most common hybrid orbitals are combinations of s and p orbitals that then form sigma bonds. (The number of orbitals mixed equals the number of hybrid orbitals produced. Not all orbitals in a level are hybridized. Count the number of sigma bonds to determine the number of hybridized orbitals.)

Orbital hybridization
Higher level

Two sp orbitals Higher level In triple bonds, as in acetylene, two sp orbitals form a  bond between the carbons, and two pairs of p orbitals overlap to form the two  bonds.

Three sp2 orbitals Higher level In a molecule like formaldehyde (shown at left) an sp2 orbital on carbon overlaps to form a  bond with the corresponding orbital on the oxygen. The unhybridized p orbitals overlap in  fashion.

Four sp3 orbitals Higher level In a molecule like methane (shown at left) the sp3 orbitals on carbon overlap to form a  bond with the corresponding orbitals on hydrogen.

Bonding in CO2 Pi bonds form from overlapping p orbitals
Higher level Pi bonds form from overlapping p orbitals C and O form hybrid orbitals to make the sigma bond

BeF2 Hybridization – orbital diagrams
Higher level BeF2 Hybridization – orbital diagrams Consider beryllium: In its ground electronic state, it would not be able to form bonds because it has no singly-occupied orbitals.

Higher level BeF2 Hybridization – orbital diagrams But if it absorbs the small amount of energy needed to promote an electron from the 2s to the 2p orbital, it can form two bonds. σ bond π bond

Higher level BeF2 Hybridization Mixing the s and p orbitals yields two degenerate orbitals that are hybrids of the two orbitals. These sp hybrid orbitals have two lobes like a p orbital. One of the lobes is larger and more rounded as is the s orbital.

Higher level BeF2 Hybridization These two degenerate orbitals would align themselves 180 from each other. This is consistent with the observed geometry of beryllium compounds: linear.

Higher level BeF2 Hybridization – orbital diagrams σ bond σ bond With hybrid orbitals the orbital diagram for beryllium would look like this. The sp orbitals are higher in energy than the 1s orbital but lower than the 2p.

BF3 Hybridization – orbital diagrams
Higher level BF3 Hybridization – orbital diagrams Using a similar model for boron leads to… σ bond σ bond σ bond

Higher level BF3 Hybridization …three degenerate sp2 orbitals.

CH4 Hybridization – orbital diagrams
Higher level CH4 Hybridization – orbital diagrams With carbon we get… σ bond σ bond σ bond σ bond

CH4 Hybridization Higher level …four degenerate sp3 orbitals.

CHN Hybridization – animation
Higher level CHN Hybridization – animation N C

Higher level C N

Higher level H C N

Hybridization of nitrogen in ammonia
Higher level Hybridization of nitrogen in ammonia Hybrid energy state Lone pair 3 σ bonds Number of electron domains will tell you what hybidization there is. 4 electron domains, therefore 4 sp3 hybyridization

Hybridization of oxygen in CO2
Higher level Hybridization of oxygen in CO2 Number of electron domains will tell you what hybidization there is. Pi bonds form from normal p orbitals π bond Hybrid energy state Lone pair Lone pair σ bond

Problem 1: Analyse the SO42- molecule.
Higher level Problem 1: Analyse the SO42- molecule.

IB CHEMISTRY Topic 4 Bonding Higher level.

Similar presentations

Presentation on theme: "IB CHEMISTRY Topic 4 Bonding Higher level."— Presentation transcript:

Similar presentations

About project

Feedback

Log in

Auth with social network:

IB CHEMISTRY Topic 4 Bonding Higher level.

Similar presentations

Presentation on theme: "IB CHEMISTRY Topic 4 Bonding Higher level."— Presentation transcript:

Similar presentations

About project

Feedback