S3

3.1: Classification of Elements

(see 3.2 Periodic Trends old syllabus for more detail)

Atomic Radius

• increases down a group and decreases across a period

  • down a group: number of occupied electron levels increase

  • across a period: number of occupied electron levels stay the same but the number of protons increase, increasing the nucleus’ force of attraction to the outer electrons

Ionic Radius

• positive ions are larger than parent atoms due to loss of outer energy level (valence)

• negative ions are smaller than parent atoms due to addition of electrons

  • increased electron repulsion causes electrons to move

• increase in nuclear charge (number of protons) causes ionic radius to decrease

  • increased attraction between outer electrons and nucleus

• ionic radius increases down a group due to increased amount of occupied energy levels

Ionization Energy

• increases across a period and decreases down a group

Electron Affinity

• decreases down a group and increases across a period

Electronegativity

• decreases down a group and increases across a period

  • across → increase in nuclear charge increases attraction between nucleus and bond electrons

  • down → increases distance between nucleus and bond electrons so reduced attraction

Group 1: Alkali metals

• physical properties

  • good conductors of electricity and heat (mobility of outer electrons)

  • low density

  • shiny grey surfaces when freshly cut with a knife

• chemical properties

  • very reactive metals

  • forms ionic compounds with non-metals

  • forms single charged ions (X+)

  • reactivity increases down group (lower IE)

• reaction with water: forms hydrogen and metal hydroxide

  • lithium: floats and reacts slowly (releases hydrogen but keeps shape)

  • sodium: reacts vigorously (heat produced melts the unreacted metal)

  • potassium: reacts more vigorously (heat produced ignites hydrogen)

Group 17: Halogens

• physical properties

  • coloured

  • gradual change from gases (F2, Cl2) to liquid (Br) and solid (I2, At2)

• chemical properties

  • very reactive non-metals

  • reactivity decreases down group (lower attraction)

  • form ionic compounds with metals and covalent compounds with non-metals

• displacement reactions

  • the more reactive halogen displaces the less reactive halogen

• halides

  • halogens produce insoluble salts with silver forming precipitates

Period 3 Oxides

• ionic character of period 3 oxides decrease from left to right

  • electronegativity value approaches oxygen, so the difference is less

• 

• ionic oxides

  • dissolve in water to form alkaline solutions

  • reacts with acid to form a salt and water

• non-metallic oxides

  • reacts with water to form acidic solutions

• amphoteric oxides

  • essentially insoluble (does not affect pH when added to water)

  • shows both basic and acidic behaviour

Acid Rain

• produced by non-metal oxides

• sulfur oxides

  • S(s) + O2(g) → SO2(g) sulfur dioxide

  • H2O(l) + SO2(g) → H2SO3(aq) dissolve in rainwater

  • 2SO2(g) + O2(g) → 2SO3(g) sulfur trioxide

  • H2O(l) + SO3(g) → H2SO4(aq) dissolve in rainwater (acid)

• nitrogen oxides

  • N2(g) + O2(g) → 2NO(g) nitrogen monoxide

  • N2(g) + 2O2(g) → 2NO2(g) and 2NO(g) + O2(g) → 2NO2(g) nitrogen dioxide

  • H2O(l) + 2NO(g) → HNO2(aq) + HNO3(aq) dissolve in rainwater

  • 2H2O(l) + 4NO2(g) + O2(g) → 4HNO3(aq) oxidized

Oxidiation States

• oxidation is:

  • addition of oxygen

  • removal of hydrogen

  • electron loss

  • an increase in oxidation state

• rules to assign oxidation states:

  1. atoms in the free (uncombined) element have an oxidation state of zero

  2. in simple ions, the oxidation state is the same as charge of the ion

  3. oxidation states of all atoms in a neutral compound must add up to zero

  4. oxidation states of all atoms in a polyatomic ion must add up to the charge

  5. usual oxidation state for an element in a compound is the one most commonly found

  6. F (fluorine) has oxidation state of -1 all the time (most electronegative)

  7. O (oxygen) has oxidation state of +2 except in peroxides

  8. Cl (chlorine) has oxidation state of -1 except when bonded to more electronegative ions

  9. H (hydrogen) has oxidation state of +1 except when forming ionic hydrides

  10. oxidation state of a transition metal in a complex ion can be found using the charge on the ligands

Transition Metals

• metals in the d-block have similar physical and chemical properties

• zinc is not a transition metal

  • has a full d sublevel in both species

• physical properties

  • high electrical and thermal conductivity

  • high melting point

  • high tensile strength

  • malleable and ductile

• chemical properties

  • forms compounds with more than one oxidation state

  • form a variety of complex ions

  • form coloured compounds

  • acts as catalysts when either elements or compounds

• magnetic properties

  • only found in iron, nickel, and cobalt

    • due to presence of unpaired electron

    • every spinning electron can act as a magnet

Variable Oxidation States

• transition metals display a wide range of oxidation states

• all transition metals show both +2 and +3 oxidation states

• maximum oxidation states increase in steps of +1 and reaches a maximum at manganese then decreases in steps of -1

• oxidation states above +3 generally show covalent character

• compounds with higher oxidation states tend to be oxidizing agents

Coloured Compounds

• transition metal ions in solution have a high charge density

  • attracts water molecules which form coordination bonds with the positive ions

• complex ions are formed when a central ion is surrounded by molecules/ions that possess at least one lone pair of electrons (ligands)

  • number of coordination bonds from the ligands to the central ion is called the coordination number

• 

• 

• colours appear because of a spurt in the d-orbital’s energy levels

  • when light passes through, it excites electrons and increases the energy level for these electrons

  • the ions absorb some colours and reflects the ones opposite it

  • when the energy (light) is absorbed, the d-orbitals split into two levels

3.2 Functional Groups: Classification of Organic Compounds

• empirical formula → simplest whole-number ratio of atoms

• molecular formula → actual number of atoms

• full structural formula → shows every bond and atom at 90°/180°

• condensed structural formula → omits bonds, displays minimal information

• skeletal formula → shorthand representation of a structural formula

• aromatic compounds: molecules which contain a benzene ring

• catenation: ability of carbon to link to itself and form chains/rings

Functional Groups

• atoms/groups of atoms that are present in organic compounds and are responsible for a compound’s physical properties and chemical reactivity

• 

• *halogen atoms are regarded as substituents as they have taken the position of a hydrogen atom

  • IUPAC names example: chloroethane, 2-bromopropane, etc.

• **syllabus does not require knowledge of arenes as compounds but expects you to recognize the phenyl group when it is present in a structure

  • naming is also not required

Functional Groups: chemical reactivity

• reaction pathway → several reactions to produce a target compound

  • product of one reaction is reactant of the next

  • ex: ethane (CH2) → ethanoic acid (CH3COOH)

    • 

Amino acids: condensation reaction to form link

• amino acids contain 2 functional groups:

  • amine (-NH2) and carboxylic acid (-COOH)

• amino acids react together via condensation reaction

  • molecule of water eliminated, acid and amino groups form new bond

    • bond is a substituted amide link (peptide bond) forming a dipeptide

• 

• the dipeptide still has a functional group (-NH2, -COOH)

  • can perform a condensation reaction again, forming a tripeptide and eventually a chain of many linked amino acids (polypeptide)

Homologous Series

• organic compounds are classified into ‘families’ of compounds

• successive members of a homologous series always differ by CH2

  • ex: C2H6, C3H8, C4H10 → alkanes

• members of a homologous series can be represented by the same general formula

• members of a homologous series show a trend in physical properties

  • because they differ by CH2, carbon chains get progressively longer

    • so higher B.P. for example

• longer chians = increased London dispersion forces

Functional groups: physical properties

• most volatile → least volatile

  • alkane > halogenoalkane > aldehyde > ketone > alcohol > carboxylic acid

  • London dispersion forces → dipole-dipole interaction → hydrogen bonding

  • increasing strength of intermolecular attraction →

    • increasing boiling point →

• chain length and functional groups affect intermolecular forces

  • polar functional groups = dipole-dopole or hydrogen bonding

IUPAC Naming

1. Identify the longest straight chian of carbon atoms

   • 1=meth-, 2=eth-, 3=prop-, 4=but-, 5=pent-, 6=hex-, etc.

2. Identify the functional group

   • numbered → nuimber has to be the smallest value possible

3. Identify the side chains or substituent groups

   • halogenoalkane (-F, -Cl, -Br, -I), amine (-NH2)

Esters and Ethers

• esters → form when the alkyl group of an alcohol replaces the hydrogen of a carboxylic acid in a condensation reaction

  • R-COOH + R’OH → R-COO-R’ + H2O

  • the stem comes from the parent acid but the alkyl group of the alcohol is the prefix

    • ex: ethanol + ethanoic acid → ethylethanoate

• ethers → 2 alkyl chains linked by an oxygen atom

  • R-O-R’

  • the longer chain will be the stem and retains its alkane name

  • the shorter chain is regarded as a substituent and is given the prefix alkoxy

    • ex: methoxypropane, ethoxyethane

• prefix - stem - suffix

  • prefix - position, number, and name of substituents

  • stem - number of carbon atoms in longest chain

  • suffix - class of compound determined by functional group

Structural Isomers

• same molecular formula but different arrangements of the atoms

• each isomer is a distinct compound

  • unique physical and chemical properties

• the more branching that is present in an isomer, the lower its boiling point

  • reduced surface contact weakens London dispersion forces

• primary carbon → attached to functional group and at least 2 hydrogen atoms

• seconday carbon → attached to functional group, one hydrogen atom, and 2 alkyl groups

• tertiary carbon → attached to functional group and 3 alkyl groups

Stereoisomers

• atoms are attached in the same order but differing in spatial or 3D arrangements → requires 3D representation

  • Isomerism - structural + stereo

    • Stereo - configurational + conformational

      • Configurational - cis-trans + optical

Conformational Isomers (not needed)

• spontaneously interconnect through bond rotations and so cannot be isolated seperately (usually)

  • some conformers of a compound may be more stable than others so are favoured → influences reactivity of the compound

Configurational isomers

• permanent difference in geometry

  • cannot be interconverted and exist as seperate compounds with some distinct properties

Cis-trans isomers

• double-bonded molecules

  • consists of one sigma and one pi bond (pi bond forming by sideways overlap of two p orbitals)

  • free rotation around this double bond is not possible

    • would push p orbitals out of position and pi bond breaks

• 

  • when the molecule contains two or more different groups attached to the double bond, these can be arranged to give 2 different isomers

    • cis → same side, trans → opposite side s

• cyclic molecules

  • cycloalkanes contain a ring of carbon atoms that restricts rotation

    • bond angles are strained from the tetrahedral angles in parent alkane

Optical isomers

• chiral - carbon atom attaches to 4 different atoms/groups

  • also known as asymmetric or stereocentre

• the four groups arranged tetrahedrally with bond angles of 109.5° can be arranged in 2 different 3D configurations which are mirror images

  • known as enantiomers → chiral molecules

    • have opposite configurations at each chiral center

• diastereoisomers → have opposite configurations at one or more (but not all) chiral centers

  • not mirror images of each other

• 

Properties of enantiomers

• optical activity → interaction with light

  • when a beam of plane-polarized light passes through a solution of optical isomers, they rotate the plane of polarization

  • optically active → seperate solutions of enantiomers (at the same concentration) rotate plane-polarized light in equal amounts but opposite directions

    • racemic mixture → chiral compound with equal concentration of 2 optical isomers

  • two optical isomers’ rotations cancel out, so racemic mixtures are optically inactive

• reactivity with other chiral molecules

  • when a racemic mixture is reacted with a single enantiomer of another chiral compound, the two components of the mixture (+ and -) react to produce different products

    • products have distinct chemical and physical properties so can be seperated

    • resolution → two enantiomers seperated from a racemic mixture

Mass Spectrometry

• used to find mass of individual atoms and finding relative abundances of different isotopes → also finds relative molecular mass of a compound

Fragmentation patterns

• ionization process → shooting an electron from electron gun then hitting the incident species and removing an electron

  • X(g) + e- → X+(g) + 2e-

    • X is a molecule

  • collision can be so energetic the molecule breaks into different fragments

• fragmentation pattern is used as evidence to find the structure of a compound

  • peak (largest mass/charge) is molecular ion that passed without fragmenting

Infrared Spectroscopy

• frequency of radiation is often measured as number of waves per centimeter (wavenumber)

• radio waves can be absorbed by certain nuclei, reversing their nuclear spin (environment)

  • used in nuclear magnetic resonance (NMR)

• microwaves cause molecules to increase their rotational energy (bond lengths)

• infrared radiation is absorbed by certain bonds causing them to stretch/bend (bonds)

• visible/ultraviolet light can produce electronic transitions (electronic energy levels)

• x-rays are produced when electrons make transitions between inner energy levels

  • produce diffraction patterns (molecular/crystal structure)

Natural frequency of a chemical bond

• chemical bonds are like springs/rulers

  • each bond vibrates and bends at a natural frequency

    • depends on strengths and atom masses

  • light atoms vibrate at higher frequencies (less weight)

  • multiple (stronger) bonds vibrate at higher frequencies

• simple diatomic molecules can only vibrate when the bond stretches

Exciting bonds

• energy needed to excite bonds occur in the infrared (IR) region

• only polar diatomic molecule bonds will interact with IR radiation

  • presence of positive and negative charge allows the electric field component of the IR radiation to excite the vibrational energy

  • change in vibrational energy produces change in dipole moment

    • intensity of absorption depends on polarity

Stretching and Bending

• in a polyatomic molecule (like water), it is more correct to consider the molecule stretching and bending as a whole, rather than considering individual bonds

  • ex: water can vibrate at 3 fundamental frequencies

    • 

    • symmetric stretch, asymmetric stretch, symmetric bend

  • each of the modes of vibration results in a change of dipole in the molecule

    • can be detected with IR spectroscopy

• for a symmetrical linear molecule (like carbon dioxide), there are 4 modes of vibration

  • symmetric stretch is IR inactive: no change in dipole moment

    • dipoles of both C=O bonds are equal and opposite throughout the interaction

    • 

Greenhouse Gases

• greenhouse effect: solar radiation passes through the atmosphere and warms the surface of the Earth. The surface radiates some of this energy as longer wavelength infrared radiation which is absorbed by greenhouse molecules which makes the air warmer, causing it to radiate heat. Some of this radiation is re-radiated back to the Earth’s surface and some back to space

• the ability of a molecule to absorb infrared radiation depends on the change in dipole moment that occurs as it vibrates

  • some greenhouse gases are much more effective than others in absorbing IR radiation

• global warming potential: amount of infrared radiation that one tonne of a gas would absorb compared to the amount that would be absorbed by one tonne of carbon dioxide

  • depends on effectiveness and atmospheric lifetime of the gas

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Wavenumbers

• absorption of certain wavenumbers of IR radiation helps to identify bonds in a molecule

• 

• some bonds can be identified by shapes of their signal

  • ex: O-H bond is broad, C=O is sharp

• hydrogen bonding broadens IR absorption so can be detected

  • ex: O-H in carboxylic acids have broader absorption

• molecules with several bonds can vibrate in many different ways and with different frequencies

  • complex pattern can be used as a ‘fingerprint’ to be matched

  • comparison of spectrum of a sample with a pure compound can be used as a test of purity

Nuclear Magnetic Resonance Spectroscopy (NMR)

• nuclei of atoms with an odd number of nucleons (H, C, F) have a property called nuclear spin and behave like tiny bar magnets

  • when placed in an external magnetic field, these nuclei can exist in two distinct energy levels

    • depending on whether magnetic field is aligned with/opposed to the external magnetic field

    • energy gap between the energy levels is very small and only requires absorption of low-energy radio waves to close the gap between energy levels

• 

• as electrons shield nucleus from full effects of external magnetic field, differences in electron distribution produce different energy seperations between the two spin energy levels

  • nuclei in different chemical environments produce different signals

    • proton = hydroegn because hydrogen has 1 proton

Magnetic Resonance Imaging (MRI)

• application of NMR spectroscopy

  • uses H’s magnetic moment

• with a powerful magnet, radio waves are used to generate an electronic signal that can be decoded to produce images

• useful in diagnosis of living tissue due to hydrogen in water

H NMR Spectroscopy

• NMR provides:

  1. number of signals in the spectrum

  2. position/chemical shifts of each signal

  3. size/integrated area of each signal

  4. splitting pattern observed for each signal

     • gives information on chemical environments and therefore structure

Chemical environments

• hydrogen nuclei (protons) that have the same chemical environment are said to be equivalent as they give the same signal in NMR

  • number of singals observed therefore depends on number of chemical environments

Chemical shifts

• position where a signal appears in NMR spectrum is measured in terms of chemical shift which has units of parts per million (ppm)

• the closer a hydrogen atom is to an electronegative atom, the more pronounced the electron-withdrawing effect and the higher chemical shift observed

  • the high electronegativity effectively pulls electrons away from the hydrogen atoms thus deshielding the hydrogens’ nuclei

  • nuclei are now more susceptible to effects due to external magnetic field

• hydrogen nuclei in particular environments have characteristic chemical shifts

  • found in section 21 of data booklet

Splitting patterns

• individual signals in NMR do not consist of a single peak

  • signals are split/resolved into distinctive patterns

• splitting occurs as the effective magnetic field experienced by particular nuclei is modified by the magnetic field produced by neighbouring protons

  • spin-spin coupling

• the number and intensity of lines produced are easily predicted

  • based on the number of neighbouring hydrogens involved in coupling

• number of lines: n+1 → n=number of hydrogen atoms on the neighbouring carbons

• intensity: pascal’s triangle

• pattern will continue for each additional proton on neighbouring carbons

• protons bonded to the same atom do not interact as they are equivalent and behave as a group

• protons on carbon atoms not adjacent to each other do not generally interact as they are too far apart for their magnetic fields to interact

• alcohol protons (OH) typically do nto engage in spin-spin coupling

  • signals for OH protons are not split and appear as singlets

  • OH protons are not counted when applying the n+1 rule