Edexcel Chemistry ( AS )

Oxidation number of Hydrogen

- Always +1 except in metal hydrides (becomes -1) and in H₂; 0

Aromatic compounds

- Contain benzene rings; a ring of 6 carbon atoms with a delocalised ring of electrons

Exothermic Reactions

- Give out heat energy - ΔH is negative - Temperature often goes up

Elimination reactions of Halogenoalkanes

- If you react a halogenoalkane with a warm alkali dissolved in ethanol, an alkene is produced; mixture must be heated under reflux else volatile product will be lost - e.g. bromoethane with warm ethanolic KOH under reflux makes ethene + water + KBr - The hydroxide ions act as a base to remove a H⁺ ion from the halogenoalkane

Isoelectronic ions

- Ions of different elements with the same number of electrons

Chain isomers

- The carbon skeleton can be arranged differently; e.g. straight chain or branched in different ways - Have similar chemical properties but different physical properties such as boiling point

Covalent bond

- When atoms share electrons to fill their outer shells - Hold molecules together "The strong electrostatic attraction between the two positive nuclei and the shared electrons in the bond"

Effect of ionic charges on ionic bonding

- greater charge; stronger ionic bond - stronger bond; higher melting/boiling points - e.g. NaF (Na⁺ + F⁻) has melting point of 993°C, but CaO (Ca²⁺+ O²⁻) has a melting point of 2572°C

Alcohols

-OH suffix, e.g. ethanol; CH₃CH₂OH

Redox reaction

A reaction in which electrons are transferred - Loss of electrons called oxidation; oxidation number +1 for each lost - Gain of electrons called reduction; oxidation number -1 for each gained - Both can happen simultaneously, hence redox - Oxidising agents accept electrons and get reduced - Reducing agents donate electrons and get oxidised

Species

An atom, an ion, a radical or a molecule

Combustion

Any alcohol burned with an excess of oxygen will produce carbon dioxide and water

Relative isotopic mass

The mass of an atom of an isotope of an element compared with 1/12th if the mass of an atom of carbon-12

Ionisation

The removal of one or more electrons

Ionisation energy increases across a period (periodicity)

This is due to: - number of protons increasing; stronger nuclear attraction - All extra electrons are at roughly the same energy level, even if the outer electrons are in different orbital types - Generally little extra shielding or extra distance to lessen nuclear attraction

Migration of Ions

This is evidence for charged particles because, when electrolysing an ionic compound in solution, positive ions move to the cathode and negative ions move to the anode

Dative covalent bonding

This is where both electrons in a covalent bond come from the same atom. E.g. CO, NH₄⁺

Subshells

This table shows the subshells and how many electrons can be contained in each.

Heterogeneous Catalyst

- A catalyst that is in a different phase to the reactants; e.g. in Haber process, gas reactants passed over solid iron catalyst - Reaction happens on the surface of a heterogeneous catalyst; increasing surface area of catalyst increases number of molecules that can react at the same time

Identifying Compounds with Mass Spec

- A mass spectrum is produced by a mass spectrometer; the molecules in a sample are bombarded with electrons to remove an electron from the molecule to form a molecular ion, M⁺ (g) - Find the relative molecular mass of a compound by looking at the molecular ion peak; the peak with the highest m/z value

Electron pair repulsion theory

- A way of predicting molecular shape using electron pair repulsion. - Need to know shapes and angles for; methane, ammonia, water

Endothermic reactions

- Absorb heat energy - ΔH is positive - Temperature often falls

Hydrogen halides

- Acidic gases - Colourless - Can dissolve in water; e.g. Hydrochloric acid, hydrobromic acid, etc - Will turn damp blue litmus paper red - Can react with ammonia gas to produce white fumes e.g. hydrogen chloride makes ammonium chloride

Test for sulphate ions

- Add dilute HCl then barium chloride solution - If a white precipitate of barium sulphate forms; positive result

Test for halides

- Add dilute nitric acid to remove any interfering ions - Add silver nitrate solution (AgNO₃ (aq)) - A precipitate of silver halide is formed - Colour of precipitate identifies halide; AgF-no colour(soluble), AgCl-white, AgBr-cream, AgI-yellow - Precipitates can look similar; further testing can be done with ammonia solution - AgCl dissolves in dilute ammonia solution - AgBr dissolves in concentrated ammonia but not dilute - AgI does not dissolve in ammonia

Balancing half equations

- Add water to balance oxygens - Add H⁺ ions to balance hydrogens - Add electrons to balance charges

Maxwell-Boltzmann distribution

- Addition of a catalyst moves the activation energy line to the right; shows more molecules have at least the activation energy because the activation energy is lower

Remove traces of water

- After separation is used, the organic layer will end up containing trace amounts of water; must be dried - To dry it, add anhydrous salt e.g. magnesium sulphate; this binds to any water present - All the water has been removed when you can swirl the mixture and it looks like a snow globe - Filtering the mixture will remove the drying agent

Making alcohols

- Alkenes can be hydrated by steam at 300°C and pressure of 60-70 atm with a phosphoric(V) acid catalyst (H₃PO₄) - The reaction is often used to manufacture ethanol from ethene

Dehydration of alcohols

- Alkenes can be produced by eliminating water from alcohols in an elimination reaction - The alcohol is mixed with an acid catalyst such as concentrated phosphoric acid (H₃PO₄) - e.g. C₂H₅OH →(dehydration) CH₂=CH₂ + H₂O - Water molecule is made up of hydroxyl group and one hydrogen atom that was bonded to a carbon atom adjacent to the hydroxyl carbon - Often more than one product can be formed; watch out for E-Z isomers too

Oxidation number of Oxygen

- Always -2 except in peroxides (O₂²⁻); -1 and in O₂; 0

Test for Ammonium ions

- Ammonia is alkaline; will turn damp red litmus paper blue (dissolves into the water making it damp first - Use this to test for ammonium ions by adding sodium hydroxide to test substance and gently heat; if ammonia is given off there are ammonium ions in test substance

Bond strength affects melting and boiling points across a period

- As you go across a period, the type of bond formed between atoms of an element changes - For metals, melting and boiling points increase across a period because the metallic bonds get stronger due to increased charge density - Elements that form giant covalent lattices (C and Si) have strong covalent bonds between all their atoms; a lot of energy is needed to break all these bonds; they have extremely high melting points - Simple molecular structures (e.g. N₂, O₂, P₄, etc) have low melting and boiling points due to the weak London forces between their molecules. More electrons means stronger London forces - Noble gases have the lowest melting points in their periods due to being monoatomic; they have very weak London forces

CIP Rules

- Atoms with a larger atomic number are given a higher priority; use these to determine whether E/Z - Look at how the atoms of highest priority are assigned; if the higher priority atom on each side are both above the bond; Z isomer; if one is above and one below then E-isomer - Sometimes you may have to look further along the chain; e.g. an ethyl group takes priority over a methyl group

Harmful emissions from alkanes used as fuels

- Big source of pollution - Unburnt hydrocarbons, carbon monoxide, carbon particulates, oxides of sulfur and nitrogen from impurities

Homolytic Fission

- Bond breaks evenly and each bonding atom receives one electron from the bonded pair. - Two electrically charged radicals are formed - Radicals are particles with an unpaired electron - They are shown in mechanisms by a big dot next to the molecular formula - They are very reactive - E.g. X-Y → X∙+Y∙

Bond Enthalpy

- Bond enthalpy is the amount of energy required to break 1 mol of a type of bond in a molecule in the gas phase - Bond enthalpies are always positive, as all bonds require energy to break

Percentage Ionic Character

- Can be calculated using the Pauling scale - Only bonds between atoms of the same element can be purely covalent; identical electronegativites - Most compounds fall between the two extremes; often have ionic and covalent properties - Uses information in data book; Pauling scale and electronegativity difference/% ionic character table

Oxidation of primary alcohols

- Can be oxidised twice; first to form aldehydes and then to form carboxylic acids - To get an aldehyde, heat it with a controlled amount of oxidising agent in distillation apparatus, so the aldehyde (which has a lower boiling temp than the alcohol) is distilled off immediately - To produce carboxylic acid, alcohol must be heated with excess oxidising agent under reflux

Displacement reactions

- Can displace halide ions from solution - Displacement reaction = when one element replaces another in a compound e.g. Cl₂ (aq) + 2KBr (aq) → 2KCl (aq) + Br₂ (aq) - More reactive halogen displaces less reactive halide in solution - Are redox reactions; the thing thats displaced is oxidised, the displacing thing is reduced - If a reaction takes place, colour change will occur; if bromide displaced and bromine formed → orange, for iodide/iodine; brown - can add to organic solvent to make results clearer

Economic benefits of catalysts

- Can dramatically lower production costs - Give more product in shorter time - Make better products

Conditions for Enthalpy Changes

- Cannot directly measure the enthalpy of a system; does not matter; only enthalpy change really matters - Enthalpy changes found in data books are usually standard enthalpy changes; done at 100kPa and 298K - This is important because changes in enthalpy are affected by temperature and pressure; standard conditions means everyone knows exactly what the enthalpy change is describing

Effect of double bonds

- Carbon atoms in a C=C bond and the atoms bonded to these carbons all lie in the same plane; considered to be trigonal planar - In ethene, the whole molecule is planar, but in larger alkenes only the >C=C< unit is planar - The atoms cannot rotate around the carbons in a C=C bond like they can in a C-C bond; double bonds are rigid; do not bend much and do not rotate - Things can still rotate about any ingle bonds in the molecule - Restricted rotation causes alkenes to form stereoisomers

Hydrogen bonding in Water

- Causes it to have a relatively high boiling point for such a small molecule - Explains why ice floats; the water molecule are arranged so that there is the maximum number of hydrogen bonds; lattice structure formed wastes a lot of space - As ice melts; H-bonds are broken and the lattice breaks down, allowing molecules to fill the space - This means ice is much less dense than water; hence it floats

Group 1 Nitrates

- Decompose to form the metal nitrite and oxygen e.g. 2KNO₃→2KNO₂ + O₂ - Except LiNO₃, which makes Li₂O, NO₂ and O₂

Group 2 Carbonates

- Decompose to form the oxide and carbon dioxide e.g. CaCO₃ → CaO + CO₂

Group 2 Nitrates

- Decompose to form the oxide, nitrogen dioxide and oxygen e.g. 2Ca(NO₃)₂ → 2CaO + 4NO₂ + O₂

Ionisation Energy

- Decreases down group - Each element in group 2 has an extra electron shell compared to the one above it - Extra inner shell shields outer electrons from nuclear attraction - Extra shell also means outer electron is further away from nucleus; greatly reduces pull from nucleus - Explains trend in reactivity; react by losing their 2 outermost electrons; so higher ionisation energy; less likely to lose electrons; less reactive; reactivity increases down group

Solubility

- Depends on compound anion - Generally compounds of group 2 element that contain singly charged anions (e.g. OH⁻) increase in solubility down group - Compounds containing doubly charged anions decrease in solubility down group - All sulphates but barium sulphate are soluble in water

Primary, secondary or tertiary

- Depends on how many alky groups are attached to the carbon the OH group joins to - Primary; 1 R group on carbon atom; secondary; 2 R groups; tertiary; 3 R groups

Chlorine (and iodine, bromine) in water

- Disproportionation - Produces hydrochloric acid and hypochlorous acid; Cl₂ (g) + H₂O (l) ⇌ HCl (aq) + HClO (aq); then hypochlorous acid ionises to make chlorate(I) ions; kill bacteria - Same reaction happens for bromine or iodine with water

Reaction with Hot Alkalis

- Disproportionation - e.g. 3X₂ + 6NaOH → NaXO₃ + 5NaX + 3H₂O

Reaction with Cold Alkalis

- Disproportionation reaction - E.g. X₂ + 2NaOH → NaOX + NaX + H₂O - Chlorine + NaOH make bleach

Non-polar substances

- Dissolve best in non-polar solvents - Non-polar substances e.g. ethene have London forces between their molecules; can form London forces with similar solvents e.g. hexane; dissolves

Alcohols

- Dissolve in polar solvents - Due to -OH group forming H-bonds with water - More carbon atoms; less soluble

Ionic substances

- Dissolve in polar solvents e.g. water - When an ionic substance is mixed with water, the ions in it are attracted to the oppositely charged end of the water molecule - Ions are pulled away from ionic lattice by water molecules, which then surround the ions; called hydration - Some ionic molecules do not dissolve; binds between their ions are stronger than those that would be formed with water e.g. Al₂O₃

Modern period table organises elements by Proton Number

- Dmitri Mendeleev created base for modern periodic table in 1869

Oxidation of tertiary alcohols

- Do not react at all with oxidising agents - Only way to oxidise them is to burn them

Roman numerals in written name

- E.g. Copper(I) oxide; copper has oxidation number of +1 - Ions ending in -ate e.g. sulfate(VI); sulphur has an oxidation number of +6; SO₄²⁻ ion

Atomic emission spectra

- Each set of lines represents electrons moving to a different energy level - One set of lines is produced when electrons fall to the n=1 (ground state) level, another when they fall to n=2, etc. - When they drop to n=1, the series of lines is produced in the ultraviolet part of the electromagnetic spectrum - n=2 produces lines in the visible part of the spectrum - n=3 produces in the infrared part of the spectrum

Electrophiles

- Electron pair acceptors - Often positively charged ions (e.g. H⁺) or δ+ areas - Electron poor; attracted to electron rich areas - Like to react with negative ions, atoms with lone pairs and the electron rich area around C=C bond - Alkene molecules undergo electrophilic addition; in a molecule with a polar bond, such as HBr, the Hδ⁺ acts as an electrophile and is strongly attracted to the C=C bond (which polarises the H-Br bond even more until it breaks)

Nucleophiles

- Electron pair donors - Often negatively charged ions (e.g. halide ions) or species that contain a lone pair of electrons (e.g. the oxygen atom in water) - Electron rich, so attracted to areas that are electron poor; like to react with positive ions or molecules with polar bonds (due to the δ⁺ area) - Attracted to the Cδ+ atom in a polar carbon-halogen bond; the C-X bond breaks and the nucleophile takes the halogen's place (this is nucleophilic substitution)

Electronic configuration

- Electrons fill up the lowest energy subshells first - Electrons fill orbitals singly before they start pairing up - Exceptions: Chromium and Copper - donate a 4s electron to the 3d subshell because they are more stable with a full or half-full d-subshell

Atomic emission spectra - electron excitement

- Electrons release energy in fixed amounts - In their ground state, atoms have their electrons in their lowest possible energy levels - If an atom's electrons take in energy from their surroundings, they can move up energy levels, getting further from the nucleus. These are known as excited electrons - Excited electrons release energy by dropping from a higher energy level down to a lower one. The energy levels all have fixed values - they are discrete - An emission spectrum shows the frequency of light emitted when electrons drop down from a higher energy level to a lower one. These frequencies appear as coloured lines on a dark background - Each element has a different electron arrangement, so the frequencies of radiation absorbed and released are emitted. This causes the spectrum for each element to be unique

Drop in ionisation energy between groups 5 and 6 is due to electron repulsion

- Elements with singly filled or full subshells are more stable than those with partially filled shells, hence they have higher first ionisation energies - e.g Sulphur 1s²2s²2p⁶3s²3p⁴ (1IE: 1000kJmol⁻¹ and Phosphorus 1s²2s²2p⁶3s²3p³ (1IE: 1012kJmol⁻¹) - shielding is identical in both, and electron being removed from same orbital - In P, electron is being removed from a singly-occupied orbital, in S its being removed from an orbital containing 2 electrons - The repulsion between two electrons in an orbital means that electrons are easier to remove from a shared orbital

Emission Spectra support the idea of Quantum shells

- Emission spectra show clear lines for different energy levels - supports idea that energy levels are discrete; electrons jump between levels with no in-between stage.

Making and breaking bonds

- Energy is needed to break bonds, so bond breaking is endothermic (∆H is positive) - Energy is released when bonds are formed, so bond formation is exothermic (∆H is negative) - The enthalpy change for a reaction is the overall effect of these two changes; if more bonds are broken than made, ∆H is positive. If its less, ∆H is negative

Enthalpy Level diagrams

- Enthalpy/energy level diagrams show the relative energies of the reactants and products in a reaction; the difference in the enthalpies is the enthalpy change of the reaction - The less enthalpy a substance has, the more stable it is.

Differentiating between similar molecules

- Even if two different compounds contain the same atoms, you can still tell them apart with mass spec because they will produce slightly different sets of fragments - E.g. propanal and propanone; same Mr but different structures; propanal will have a C₂H₅⁺ fragment but propanone will not - Every compound produces a different mass spectrum; larger computer databases can be used to identify a compound from its spectrum

Graphite can conduct electricity

- Exception to the rule - The carbon atoms form sheets with each carbon atom sharing 3 of its outer shell electrons with 3 other carbon atoms; leaving fourth free to move between sheets

Properties of giant structures provide evidence for covalent bonds

- Extremely high melting points; many very strong bonds need to be broken, which takes a lot of energy - Often extremely hard; due to lots of strong bonds throughout the lattice arrangement - Good thermal conductors; vibrations travel easily through stiff lattices - Insoluble; atoms more attracted to others in the lattice than solvent molecules - insolubility in polar solvents like water shows lack of ions - Cannot conduct electricity; no charged ions (in most giant covalent structures) or free electrons

KBr with H₂SO₄

- First reaction gives misty fumes of HBr - Br⁻ ions are stronger reducing agents; react with H₂SO₄ in a redox reaction to form sulphur dioxide and bromine gas; 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O

KI with H₂SO₄

- First reaction produces HI - Second reaction reduces H₂SO₄ same as above - Third reaction reduces SO₂ to H₂S; 6HI + SO₂ → H₂S + 3I₂ + 2H₂O

Reaction with water, oxygen and chlorine

- Form ions with 2+ charge - React with water to form hydroxides e.g. Ca(OH)₂ - Burn in oxygen to form oxides e.g CaO - React with chlorine to form chlorides

Addition Polymers

- Formed when alkenes join up - The double bonds in alkenes can open up and join together make long chains called polymers - The individual, small alkenes are called monomers - E.g. poly(ethene) is made by addition polymerisation of ethene - to find the monomer used to form an addition polymer, take the repeat unit and add a double bond between the carbon atoms and and remove the single bond from each end - To find a repeat unit from a monomer, do the reverse

Halogenoalkanes wth ammonia

- Forms amines; organic compounds based on ammonia, but one or more of the hydrogen atoms are replaced by alkyl groups - If you warm a halogenoalkane with excess ethanolic ammonia (ammonia dissolved in ethanol), the ammonia swaps places with the halogen to form a primary amine; a nucleophilic substitution - 2 step mechanism; first step; ammonia acts as a nucleophile; second step is a reversible reaction between amine and ammonium

Properties of metallic structures

- Generally high melting points due to strong bonds; number of delocalised electrons per atom affects the melting point; more=higher, also affected by size of metal ion and size of the whole structure - No bonds holding specific ions together; malleable (can be shaped) and ductile (can be drawn into wire) - Delocalised electrons can pass kinetic energy to each other; good thermal conductors - Delocalised electrons free to move and carry a current; good electrical conductors however impurities can strongly reduce this - Insoluble except into liquid metals because of strength of bonds

Distillation

- Gently heating mixture in distillation apparatus; the substances will evaporate out of the mixture in order of increasing boiling point - The thermometer shows the boiling point of the substance that is evaporating at any given time - If you know the boiling point of your pure product, you can use the thermometer to tell you when its evaporating and therefore when its condensing - If the product as a lower boiling point than the reactants, then the reaction mixture can be heated so that the product evaporates as it is formed - If starting material has a higher boiling point than the product, as long as the temp. is carefully controlled, it won't evaporate out of the reaction mixture - If a product and its impurities have different boiling points, they can be separated using distillation

Can use properties to predict structure

- Given a list of properties of a strucure, know what type of structure it is

Reaction with Group 1 and 2 Metals

- Halogen is reduced; metal oxidised - Form halide salts

Reactivities of Halogenoalkanes

- Halogenoalkane + water→ alcohol + halide + hydride - Add silver nitrate solution to mixture too to react with halide ions and form a silver halide precipitate - The quicker the precipitate forms, the faster the rate of hydrolysis for that halogenoalkane - Rate of hydrolysis depends on carbon-halogen bond enthalpy; weaker C-X bonds break more easily; react faster - Larger halogen; longer bond length; lower bond enthalpy - size of halogen increases down group 7; fluoroalkanes hydrolyse slowest, iodoalkanes hydrolyse much faster

Hydrolysis of Halogenoalkanes

- Halogenoalkanes can be hydrolysed to alcohols in nucleophilic substitution reactions - Water can be used to do this - General equation; R-X + H₂O → R-OH + H⁺ + X⁻ - e.g. bromoethane + water; CH₃CH₂Br + H₂O + C₂H₅OH+H⁺+Br⁻

Halogenoalkanes with aqueous KOH

- Halogenoalkanes react with hydroxide ions by nucleophilic substitution to form alcohols; warm potassium hydroxide and done under reflux, else will not work - R-X + KOH → ROH + KX (R stands for an alkyl group, X stands for a halogen - bond breaks heterolytically

Exceptions (Polar molecules that do not dissolve in water

- Halogenoalkanes; dipoles are two weak to form h-bonds with water - H-bonding between water molecules is stronger than the bonds that would be formed with halogenoalkanes; they don't dissolve - However can form perm-perm dipole bonds so will happily dissolve in polar solvents that also form perm-perm dipole bonds

Nucleophilic Substitution of Halogenoalkanes

- Halogens are generally more electro negative than carbon; C-X bond is polar - δ+ carbon doesn't have enough electrons; it can be attacked by a nucleophile - OH⁻, NH₃ and CN⁻ are nucleophiles that react readily with halogenoalkanes; water is also a weak nucleophile - One-step mechanism - 3 examples are needed for the exam; with aqueous KOH, cyanide ions, and ammonia

Radicals

- Have an unpaired electron; e.g. chlorine atoms produced when UV light splits a Cl₂ molecule . - The presence of an unpaired electron makes them very reactive; they react with anything; positive, negative or neutral

Structural Isomers

- Have same number of each element, but different structural arrangements - Chain isomers - Positional isomers - Functional group isomers

Cracking

- Heavy fractions can be cracked (C-C bonds broken) to make smaller molecules; e.g. decane into ethene and octane - Two types; thermal cracking and catalytic cracking

Bond fission

- Heterolytic Fission - Homolytic Fission

How theory of Ionic Bonding fits Physical Evidence

- High melting points; shows ions are held together by strong attractions - Often soluble in water, but not in non-polar solvents; this shows that the particles are charged - Do not conduct electricity when solid, but do when dissolved or molten - Brittle; when pushed or pulled, oppositely charged ions are moved closer together and the repulsion causes the solid to break. Supports lattice model

Cis-trans isomerism

- If the carbon atoms have at least one group in common they can be labelled as cis or trans isomers (as well as E o Z) - cis means groups are on same side of double bond - trans means groups are on opposite sides - The cis isomer is not necessarily the Z isomer and the trans isomer is not necessarily the E isomer; cis-trans isomerism only looks at the identical groups

Separation

- If the product is insoluble in water then separation can be used to remove any impurities that do dissolve in water such as salts or soluble organic compounds - Once the reaction has completed, pour the mixture into a separating funnel, then add water - Shake the funnel and allow it to settle. The organic layer and aqueous layer are immiscible; open the tap and run each layer off into a separate container

Heterogeneous equilibria and Kc

- In a heterogeneous system, not everything is in the same state - Writing expressions for Kc is more difficult - Not everything is included - Pure solids and liquids are not included; this is because their concentrations stay constant throughout

Homogeneous equilibria and Kc

- In a homogeneous system, everything is in the same physical state - Kc can be calculated using the concentrations of products and reactants at equilibrium - For homogeneous equilibria, all the products and reactants are included in the expression for Kc aA + bB ⇌ dD + eE, Kc = ([D]^d[E]^e)/([A]^a[B]^b) lowercase are the moles of each substance, the square brackets indicate mean concentration in mol dm⁻³

Effect of electronegativity on Covalent bonds

- In covalent bonds, the bonding electrons sit in orbitals between two nuclei. If both are of similar electronegativity, electrons will sit roughly midway; bond will be non-polar - Homonuclear, diatomic gases such as H₂ are non-polar because identical electronegativities - Some have similar electronegativities e.g. C-H bonds are essentially non-polar - Different electronegativities; bonding electrons pulled towards more electronegative atom; electrons spread unevenly; charge across bond; each atom as a partial charge; δ⁺ or δ⁻; bond is said to be polar - In a polar bond; difference in electronegativity causes a dipole. A dipole is a slight difference in charge between the two atoms caused by a shift in electron density in bond - greater the difference in electronegativity; the more polar the bond

Bond length

- In covalent molecules, the positive nuclei are attracted to the electron density (where shared electrons are) - However there is also repulsion; the two positively charged nuclei repel each other, as do the electrons. - To maintain the covalent band, there must be a balance between these forces - The distance between the two nuclei is the distance where the attractive and repulsive forces balance each other - This distance is the bond length

Homogeneous catalyst

- In the same physical state as the reactants - Usually an aqueous catalyst for 2 aqueous solutions - Reacts with reactants to form an intermediate species, which then react to form the products and reform the catalyst

Trends in ionic radii

- Increases down a group - Ionic radius of a set of isoelectronic ions decreases as the atomic number increases

Thermal stability of carbonates and nitrates

- Increases down group; carbonate and nitrate ions are large negative ions and can be made unstable by the presence of a cation - Cation polarises anion, distorting it; greater distortion→less stable compound - Large cations cause less distortion than small cations as they have lower charge density - Group 2 compounds are less thermally stable than group 1 compounds due to increased charged density of cations

Catalysts

- Increases the rate of reaction by providing an alternative reaction pathway with a lower activation energy; a greater proportion of collisions result in reaction - Catalyst remains chemically unchanged - Only a small amount of catalyst is needed - Catalysts are usually quite specific; different catalysts are needed for different reactions

Increasing temperature

- Increasing temperature will, on average, increase the kinetic energy of the molecules and they will move faster - A greater proportion of molecules will have at least the activation energy; this changes the shape of the Maxwell-Boltzmann distribution curve; moves lower and to the right - Because the molecules are moving faster, they will collide more often. This also increases the rate of reaction

Giant Ionic lattices

- Ionic crystals (e.g. NaCl) are giant lattices of ions; called giant because its made up of a repeated basic unit - Forms because each ion is electrostatically attracted in all directions to ions of opposite charge - In NaCl, the Na⁺ and Cl⁻ ions are packed together alternately in a lattice

Half-Equations

- Ionic half-equations show oxidation or reduction - E.g. half-equation for oxidation of sodium: Na → Na⁺ + e⁻ - Half equations can be combined to make full equations for redox reactions e.g O₂ + 4e⁻ → 2O²⁻ and Mg → Mg²⁺ + 2e⁻ - ensure both have same number of electrons; so double everything in second equation to make 2Mg → 2Mg²⁺ + 4e⁻ - Combine to get 2Mg + O₂ → 2MgO

Ions

- Ions are charged atoms; positive ions have more protons than electrons, and vice versa for negative ions - Ions have different numbers of electrons to their parent elements' - e.g. Li⁺ has only 2 electrons, whereas Li has 3 - e.g. F has 9 electrons, F⁻ has 10

Isotopes

- Isotopes of an element are atoms with the same number of protons but a different number of neutrons. - E.g. ³⁵Cl has 18 neutrons and ³⁷Cl has 20 - Number and arrangement of electrons dictate the chemical properties of an element, so all isotopes of an element have the same chemical properties - However isotopes of the same element can have different physical properties such as density and diffusion rates

KF or KCl with H₂SO₄

- KF or KCl with H₂SO₄ produces KHSO₄ + HF or HCl; will see misty fumes as the gas comes into contact with moisture in air - Fluoride and chloride ions are too weak oxidising agents so the reaction stops there. Not redox; oxidation numbers stay the same

Sulfur dioxide and oxides of nitrogen

- Lead to acid rain - Sulfur dioxide produced dissolves in moisture in the atmosphere, forming sulfuric acid - Oxides of nitrogen produced when high temp. and pressure in engine cause oxygen and nitrogen in the air to react together; dissolve in atmosphere to form nitric acid - Acid rain destroys vegetation, trees, corrodes buildings and statues and kills fish in lakes

Hydrogen bonding in alcoholds

- Less volatile than similar alkanes - All contain a polar hydroxyl group; allows them to form h-bonds - Low volatility (high boiling points) compared to non-polar compounds of similar sizes

Biofuels

- Made from living matter over a short period of time - bioethanol; ethanol made by fermentation of sugar from crops such as maize - biodiesel; made by refining renewable fats and oils e.g. vegetable oil - biogas; produced by breakdown of organic waste matter - These fuels produce CO₂ when burnt; but its the CO₂ that was absorbed when the plants photosynthesised while growing; considered carbon neutral - Biodiesel and biogas are made from waste which would otherwise go to landfill - One problem is that petrol engines would need to be modified use fuels with high ethanol concentrations - The land used to grow crops for fuel cant be used for food

Electron Shells

- Made up of subshells and orbitals - Electrons move around the nucleus in quantum shells (aka energy levels) - Shells further from the nucleus have a greater energy level than those closer to the nucleus - Shells contain different types of subshell, each of which have different numbers of orbitals which can each hold 2 electrons

Reversible reactions

- Many reactions are reversible, shown by a ⇌ rather than a mono-directional arrow - As the reactants are used up, the forward reaction slows down; as more product is formed, the reverse reaction speeds up - When the forward and backward reactions are happening at the same rate, the amounts of product and reactant stop changing and a dynamic equilibrium has been reached

Mechanisms

- Mechanisms break down reactions into individual stages to show how the substances react together - Curly arrows are used to show movement of electron pairs when bonds are made or broken

Physical properties of a solid

- Melting and boiling points depend on intermolecular forces - Substance will only conduct electricity if it contains charged particles that are free to move - How soluble a substance is in water depends on the type of particle is contains

Giant metallic structures

- Metal elements exist as giant metallic lattice structures - The electrons in the outermost shell of the metal atoms are delocalised; free to move; leaving positive metal ions (e.g. Na⁺, Mg²⁺, Al³⁺) - The positive metal ions are electrostatically attracted to the delocalised negative electrons; form a lattice of closely packed positive ions in a sea of delocalised electrons - This is known as metallic bonding; very strong bond - Overall structure; layers of ions separated by layers of electrons

Crude oil

- Mixture of many hydrocarbons; mostly alkanes: range from small alkanes like pentane to massive ones of more than 50 carbons - Not useful in this form; separated out by fractional distillation

Reforming alkanes into cycloalkanes and aromatic hydrocarbons

- Most cars run on petrol or diesel, both of which contain a mixture of alkanes, other hydrocarbons, impurities and additives - Some are straight chain alkanes - Knocking is where alkanes explode of their own accord when the fuel/air is compressed; straight chain alkanes are the most likely to undergo knocking; branched chain and cyclic hydrocarbons are less likely to make knocking happen; make combustion more efficient - Can convert straight chain alkanes into branched and cyclic alkanes by reforming; this uses a catalyst (e.g. platinum on aluminium oxide)

Molecule shapes: 6 electron pairs

- No lone pairs; octohedral; bond angle 90° e.g. SF₆ - One lone pair; square pyramidal; 90° and 81.9° angles e.g. IF₅ - Two lone pairs; square planar; 90° angles e.g. XeF₄

Molecule shapes: 4 electron pairs

- No lone pairs; tetrahedral; bond angle 109.5° e.g. NH₄⁺ - 1 lone pair; trigonal pyramidal; bond angle 107° e.g. PF₃ - 2 lone pairs; non-linear/bent; bond angle 104.5° e.g. H₂O

Molecule shapes: 5 electron pairs

- No lone pairs; trigonal bipyramidal; 90° angles e.g. PCl₅ - One lone pair; seesaw; a 102° angle and 2 87° angles e.g. SF₄ - Two lone pairs; distorted T; 87.5° angles e.g. ClF₃

Molecule shapes: 3 electron pairs

- No lone pairs; trigonal planar; bond angle of 120° e.g. BCl₃ - 1 lone pair; non-linear/bent; bond angle 119° e.g. SO₂

Halogens

- Non-metals of group 7 - Highly reactive - Exist as diatomic molecules - Low solubility in water - Dissolve easily in organic compounds e.g. hexane

Renewability of fossil fuels

- Non-renewable - Main fossil fuels are running out - Will become more expensive as they become more scarce

Effect of London forces on melting and boiling points

- Not all London forces are same strength; larger molecules; more electrons; stronger dipoles; stronger London forces - Molecules with greater surface area; bigger exposed electron clouds; stronger London forces - When boiling a liquid, intermolecular forces must be overcome; requires energy; more energy required when intermolecular forces are stronger; stronger London forces; higher boiling points - Melting solids also involves overcoming intermolecular forces; stronger London forces; higher melting points - Alkanes are good demonstrators of this

Molecule Energy

- Not all molecules in a system have the same energy; some have low energy, some have high and some have moderate - A Maxwell-Boltzmann distribution shows the number of molecules plotted against their kinetic energy - The proportion of molecules with at least the activation energy determines the speed of the reaction; only these molecules can react

Catalysts

- Not included in expressions for Kc - They do not affect the equilibrium concentrations of products or reactants; simply speed up rate at which equilibrium is reached

Factors affecting ionisation energy

- Nuclear charge: more protons; more positively charged nucleus; stronger attraction for electrons - Electron shell: attraction falls off rapidly with distance; an electron in a shell close to the nucleus is much more strongly attracted than one in a shell further away - Shielding: as the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge and are repelled by the negatively charged electrons between them and the nucleus. This lessening of the pull of the nucleus by inner shells of electrons is called shielding.

Graphene

- One layer thick graphite - Carbon sheet arranged in hexagons; one atom thick; two-dimensional - Conducts electricity - Incredibly strong - Transparent - Very light

Orbitals

- Orbitals within the same subshell have the same energy - s-orbitals are spherical - p-orbitals are dumbbell-shaped. There are 3 p-orbitals and they are at right angles to each other

Refluxing

- Organic reactions are slow and the substances are usually flammable and volatile - Refluxing prevents these substances from evaporating or catching fire - The mixture is heated in a flask with a vertical Liebig condenser; this continuously boils, evaporates and condenses the vapours and recycles them into the flask, giving them time to react - The heating is usually electrical; hot plates, heating mantles or electrically controlled water baths; avoids naked flames which might ignite the compounds

The oxides and hydroxides formed are bases

- Oxides of group 2 metals react readily with water; form metal hydroxides which can dissolve; strongly alkaline due to OH⁻ ions - Exception: Beryllium oxide; does not react with water and Beryllium hydroxide is insoluble - Exception: Magnesium oxide; reacts slowly and hydroxide is not very soluble - Form more strongly alkaline solutions going down group - Both the oxides and hydroxides will neutralise dilute acids, forming the corresponding salts e.g. MO (s) + 2HCl (aq) → MCl₂ (aq) + H₂O (l) M = any group 2 metal

Alkanes

- Oxidising (burning) alkanes with oxygen produces carbon dioxide and water; combustion reaction; e.g. combustion of propane; C₃H₈(g)+5O₂(g)→3CO₂(g)+4H₂O(g) - When there is not enough oxygen, incomplete combustion takes place; produces carbon monoxide, carbon and water e.g. C₂H₆(g)+2O₂(g)→C(s)+CO(g)+3H₂O(g) - Combustion reactions happen between gases, liquid alkanes must be vaporised first; smaller alkanes vaporise more easily; they are more volatile; burn more easily too - Combustion reactions are exothermic - Larger alkanes require much more energy per mole because they have more bonds - Alkanes are excellent fuels because they release a lot of energy when burned

Le Chatelier's Principle

- Predicts what will happen if conditions are changed - If you change pressure, concentration or temperature of a reversible reaction, the position of equilibrium is usually altered; different amounts of reactants and products at equilibrium The principle states; If there is a change in concentration, pressure or temperature, the equilibrium will move to counteract the change

Adding Hydrogen

- Produces alkanes - E.g. ethene will react with hydrogen gas to produce ethane. requires a nickel catalyst and a temp of 150°C - E.g. margarine is made by hydrogenating unsaturated vegetable oils; by removing some double bonds, the melting point increases making it a solid at room temp.

Thermal Cracking

- Produces lots of alkenes - Takes place at high temp and high pressure (up to 1000°C and 70 atm) - These alkenes are useful for making valuable products like plastics, e.g. poly(ethene)

Types of mechanism

- Radical substitution of halogens in alkanes to make halogenoalkanes - Electrophilic addition of halogens and hydrogen halides to alkenes to make halogenoalkanes - Nucleophilic substitution of primary halogenoalkanes with aqueous potassium hydroxide to alcohols and with ammonia to make amines

Reaction Profile diagrams

- Reaction profile diagrams show how the enthalpy changes during a reaction - The activation energy, Ea, is the minimum amount of energy needed to begin breaking reactant bonds and start a chemical reaction

Halogenoalkanes with cyanide ions

- Refluxing a halogenoalkane with a potassium cyanide in ethanol causes the cyanide ions to react with the halogenoalkane by nucleophilic substitution to form a nitrile - e.g. R-X + CN⁻→(ethanol reflux)→ R-C≡N + X⁻

Oxidation of secondary alcohols

- Refluxing a secondary alcohol with acidified dichromate(VI) will produce a ketone - Ketones cannot be oxidised easily, so continued refluxing will not oxidise them further

Catalytic converters

- Remove pollutants from exhaust fumes using a platinum catalyst to change them into harmless gases e.g. water vapour and nitrogen or less harmful gases like CO₂ - E.g. 2NO + CO → N₂ + CO₂

Problem with radical substitution

- Results in a mixture of products - E.g. making chloromethane with too much chlorine in mixture will result in some remaining H's on the CH₃Cl to be replaced with Cl's making dichloromethane - Can repeat again to make trichloromethane - And again to make tetrachloromethane - Reduce the chance of this by using an excess of methane; more chance for radical to collide with methane molecule rather than chloromethane - Can also take place at any point on carbon chain, so a mixture of structural isomers can be formed e.g. propane + chlorine will make 1-chloropropane and 2-chloropropane

Enthalpy changes from enthalpies of combustion

- Route 1 = route 2 - Arrow on cycle point down (towards the combustion products)

Alkanes

- Saturated hydrocarbons - Each carbon has four single bonds with other atoms

Oxidation with acidified potassium manganate(VII)

- Shaking an alkene with acidified potassium manganate(VII) decolourises the purple solution; the alkene is oxidised and a colourless diol is formed (an alcohol with 2 OH groups)

Checking if a reaction has happened with IR Spec

- Shows if the functional group has changed during a reaction - e.g. if an alcohol is oxidised to an aldehyde, the O-H will disappear from the spectrum, and a C=O absorption will appear

Effect of ionic radii on ionic bonding

- Smaller ions can pack closer together than larger ions - Electrostatic attraction gets weaker with distance - This means small, closely packed ions have stronger ionic bonding than larger ions. - Therefore ionic compounds with smaller ions have higher boiling and melting points

Giant covalent structures

- Sometimes covalent bonds lead to the formation of huge lattices containing billions of atoms - Electrostatic attractions between atoms in giant covalent lattices are much stronger than in simple covalent molecules - Carbon and silicon often form these giant structures because they can each form four strong covalent bonds

Catalysts

- Speed up reaction - Lower activation energy required by providing an alternative route - This increases the number of particles with at least the activation energy

Types of ΔH

- Standard enthalpy change of reaction - Standard enthalpy change of formation - Standard enthalpy change of combustion - Standard enthalpy change of neutralisation

E/Z Isomerism

- Stereoisomers have the same structural formula but a different arrangement in space due to lack of rotation in C=C bond - Stereoisomers occur when the two double bonded carbons have two different atoms or groups attached to them - E-isomer and Z-isomer - Z-isomer has the same groups either both above or both below the double bond, while E-isomers have the same groups positioned across the double bond - Z stands for zusammen; german for together - E stands for entgagen; german for opposite

Hydrogen bonding

- Strongest intermolecular force - Only happens when hydrogen is covalently bonded to fluorine, nitrogen or oxygen - Fluorine, nitrogen and oxygen are very electronegative; so draw bonding electrons away from the hydrogen - Bond is polarised; hydrogen has very high charge density because its so small; hydrogen atoms form weak bonds with lone pairs of electrons on the fluorine, nitrogen or oxygen atoms of other molecules - Water, ammonia and hydrogen fluoride all have h-bonding - Organic molecules that form hydrogen bonds often contain -OH or -NH groups e.g. alcohols and amines - Look out for those groups when asked to predict if a substance forms hydrogen bonds

Protons

- Subatomic particle - Charge of +1 - Relative mass of 1 - Contained in nucleus - Dictates the type of element that an atom is; all atoms of the same element have the same number of protons

Electrons

- Subatomic particle - Charge of -1 - Arranged in orbitals - Relative mass of 0.0005 - negligible

Neutrons

- Subatomic particle - No charge - Relative mass of 1 - Contained in nucleus - Dictates the isotope of an element that an atom is; not all atoms of the same element have the same number of neutrons

Effect of Hydrogen bonds on substance behaviour

- Substances that form hydrogen bonds have high melting and boiling points - Group 7 and 6 hydrides; trend = HF (G7) and H₂O (G6) very high, drops to HCl/H₂S due to decreased permanent-permanent dipole attractions, increases again to HBr (H₂Se) and HI (H₂Te) due to increased London forces overriding decreased permanent dipole attractions - Substances that can form hydrogen bonds are soluble in water

Oxidation number of molecular ions

- Sum of oxidation number = overall charge - Each constituent atom has its own oxidation number

Electron Pairs repel each other

- The amount of repulsion depends on the type of electron pair - Lone pairs repel more than bonding pairs - Greatest angles are between lone pairs, and bond angles between bonding pairs are often reduced because they're pushed together by lone pair repulsion -

Fragmentation of molecular ion

- The bombarding electrons cause some of the molecular ions to break into fragments - The fragments that are ions show up on the mass spec, making a fragmentation pattern; these can be used to identify molecules and their structure - To work out structural formula, work out which ion could have made each peak from its m/z value (assume m/z value of peak matches mass of ion that made it) 1. Identify fragments 2. Piece together to form molecule with correct Mr

Heterolytic Fission

- The bond breaks unevenly, with one of the bonded atoms receiving both of the electrons from the bonded pair - Two different substances can be formed; a positively charged cation or negatively charged anion

Dynamic equilibrium

- The concentrations of products and reactants remain constant - Can only happen in a closed system

Functional group isomers

- The same atoms can be arranged into different functional groups - These have very different physical and chemical properties

Positional isomers

- The skeleton and functional group could be the same, just with the functional group attached to a different carbon atom - Different physical properties, can have different chemical properties

Hess's Law

- The total enthalpy change is independent of the route taken - The total enthalpy change of a reaction is always the same, no matter the route taken; Route 1 = Route 2

Group 1 Carbonates

- Thermally stable; cannot be decomposed with the heat of a Bunsen - Except Li₂CO₃; decomposes to Li₂O and CO₂

Aldehydes and Ketones

- They are carbonyl compounds; they have the functional group C=O - General formula is CnH2nO - Aldehydes have a hydrogen and an alkyl group attached to the carbonyl carbon atom - Ketones have 2 alkyl groups attached to the carbonyl carbon atom - Test for whether a compound is an aldehyde or ketone using Benedict's solution; if heated with an aldehyde the colour changes from blue to a brick-red precipitate; if heated with a ketone, nothing happens

Carbon monoxide

- Toxic; better at binding to haemoglobin in blood than oxygen; less oxygen carried around body; oxygen deprivation; can be fatal

Isomers

- Two molecules that have the same molecular formula but different arrangements of atoms - Two types; structural isomers and stereoisomers

Adding hydrogen Halides

- Undergoes similar process as addition of halogens; the hydrogen end become slightly positive in this reaction, halogen end slightly negative - Adding hydrogen halides to unsymmetrical alkenes forms two products; the amount of each product formed depends on the stability of the carbocation formed in the middle of the reaction - The major product from the addition of a hydrogen halide to an unsymmetrical alkene is the one where hydrogen adds to the carbon with the most hydrogens already attached

Using oxidation numbers to write chemical formulae

- Unless told otherwise, assume overall charge of compound is 0 - Work out ratio of cations:anions that gives charge of 0 e.g. barium(II) nitrate; barium has +2, NO₃⁻ has -1, so 1:2 ratio to make zero; Ba(NO₃)₂

Alkenes

- Unsaturated hydrocarbons - Have at least one C=C double bond - Unsaturated because they could make more bonds with other atoms in addition reactions

Determine Purity using Boiling point

- Use distillation apparatus to determine boiling point by gently heating the liquid until it evaporates - Then look up boiling point in data books and compare to measurement; if impurities are present the measured boiling point will be higher than the recorded value

Making polymers sustainably

- Use reactant molecules that are safe for humans and the environment - Use as few other materials, such as solvents, as possible; or choose ones which are not harmful to environment - Renewable raw materials should be used when possible - Energy use should be minimum; e.g. use catalysts - Limit waste products made; especially hazardous ones - Make sure lifespan of polymer is appropriate for its use

Catalytic Cracking

- Uses a zeolite catalyst (hydrated aluminosilicate) at slight pressure and high temperature (around 450°C) - Produces aromatic compounds and motor fuels - Cuts costs because catalyst allows it do be done at low pressure and lower temperature and speeds up the reaction

Test for Carbonate and Hydrogencarbonate Ions

- When added to dilute hydrochloric acid, fizz and give off carbon dioxide - Test for carbon dioxide with limewater; turns cloudy with CO₂; if this happens; positive result

Oxidation numbers give the movement of electrons

- When atoms react with or bond to other atoms, they can gain or lose electrons. The oxidation number shows how many electrons an atom has gained or lost to form an ion or part of a compound

Enthalpy changes

- When chemical reactions occur, some bonds are broken and some are made. More often than not, this causes a change in energy. This is known as enthalpy change, ΔH. This is the heat energy change in a reaction at a constant pressure. The units of ΔH are kJmol⁻¹ - ΔH° is written to show that the measurements were taken under standard conditions and that the elements were in their standard states. - Standard conditions are 100kPa and a specified temperature, usually 298K

If all groups are different

- When the carbons on either side of the double bond have the same groups attached, then its easy to work out which is the E-isomer and Z-isomer - Becomes problematic if carbons all have totally different groups attached - Using CIP rules, can easily determine which is which

Nucleus

- Where most of the mass of the atom is contained - Made up of protons and neutrons - Diameter is much smaller than that of whole atom

Polar bonds and molecular polarity

- Whether a molecule is polar or not depends on its shape and the polarity of its bonds - Polar molecule has an overall dipole; which is just a dipole caused by the presence of a permanent charge across the molecule 1) Simple molecule, e.g H→Cl, polar bond gives whole molecule a permanent dipole - polar molecule 2) More complex molecules may have several polar bonds. If these are arranged in opposite directions; cancel each other out; overall non-polar molecule e.g. O←C→O 3) If all the polar bonds point in roughly the same direction, the molecule will be polar e.g. CHCl₃ or water → shows direction electrons are being pulled

Ionisation Energies show Shell structure

- Within each shell, successive ionisation energies increase; less repulsion from other electrons each time, so there is stronger attraction to the nucleus - Big jumps occur between shells; an electron is being removed from a shell closer to the nucleus - This type of graph can tell you which group an element belongs to; count how many electrons are removed before the first big jump

Successive ionisation energies

- You can remove all the electrons from an atom, leaving just the nucleus - Each time an electron is removed, there's a successive ionisation energy which is greater than the previous ionisation energy - n'th ionisation energy can be written as: X(ⁿ⁻¹)⁺(g) →Xⁿ⁺(g) + e⁻

Mean bond enthalpies

- e.g. in water... there are two O-H bonds, and you'd think it would take the same amount of energy to break both. However the first bond broken needs +492kJmol⁻¹, whereas the second needs only +428kJmol⁻¹. This is due to the extra electron repulsion. (492+428)/2=460; mean bond enthalpy of O-H in water is 460kJmol⁻¹ - Data book reads bond enthalpy for O-H is +463kJmol⁻¹; its an average of a large range of molecules containing O-H bonds - The energy needed to break one mole of bonds in the gas phase, averaged over many different compounds

Electronic configuration decides the chemical properties of an element (periodicity)

- s-block elements (Groups 1 and 2) have 1 or 2 outer shell electrons, which are easily lost to form positive ions with the electron configuration of an inert gas e.g. Na: 1s² 2s² 2p⁶ 3s¹ → Na⁺: 1s² 2s² 2p⁶ (electron configuration of Neon) - p-block elements (Groups 5,6,7) can gain 1,2 or 3 electrons to form negative ions with an inert gas configuration e.g. O: 1s² 2s² 2p⁴ → O²⁻: 1s² 2s² 2p⁶ - groups 4-7 can also share electrons when they form covalent bonds - Group 0 (inert gases) have completely filled s and p subshells and don't need to gain, lose or share electrons; their full subshells are what make them inert - d-block elements (transition metals) tend to lose s and d electrons to form positive ions

Periodic table electron configuration blocks

- s-block elements have an outer shell electronic configuration of s¹ or s² - p-block elements have an outer shell electronic configuration of s²p¹ to s²p⁶

Permanent-permanent dipole attractions

- δ⁺ and δ⁻ charges on polar molecules cause weak electrostatic attractions between molecules - δ⁻ end of one molecule is attracted to δ⁺ end of another - Happen alongside London forces; polar molecules have both acting on them; generally have higher melting/boiling points than non-polar molecules with similar London forces

Enthalpy change from enthalpies of formation

- ∆fH° for all reactants must be known for all reactants and products that are compounds; elements have a ∆fH° of zero - ∆fH° + sum of ∆fH°(reactants) = sum of ∆fH° products - therefore; ∆fH° = sum of ∆fH°(products) - sum of ∆fH°(reactants) - On a Hess cycle, the arrows point up

Aldehydes

-al suffix; e.g. ethanal; CH₃CHO

Carboxylic acids

-oic acid suffix; e.g. ethanoic acid; CH₃COOH

Ketones

-one suffix; e.g. propanone; CH₃COCH₃

Oxidation number of uncombined elements

0 - Have not accepted nor donated any electrons - Molecules made up of all the same element are also 0

Using electron pairs to predict the shapes of molecules

1) Find central atom (one all the others are bonded to) 2) Work out number of electrons in outer shell of centre atom 3) The molecular formula tells you how many atoms the central atom is bonded to (this allows you to work out how many electrons are shared with the central atom) 4) Add up all electrons and divide by 2 for number of electron pairs on central atom (if its an ion, remember to account for charge) 5) Compare number of electron pairs and number of lone pairs and bonding centres around central atom to work out shape Note: - Count a double bond as 2 bonds - Bonding centres are the atoms bonded to the central atom

Uses of different length hydrocarbons and their boiling points

1-4 carbons; camping gas 5-12 carbons; petrol 7-14 carbons; Naptha 11-15 carbons; Kerosene (paraffin) 15-19 carbons; Diesel 20-30 carbons; Lubricating oil 30-40 carbons; Fuel oil; ships, power stations 40-50 carbons; wax, grease; candles, lube 50+ carbons; Bitumen; roofing, road surfacing

Identify Organic molecules using IR Spec

1. An IR spectrometer produces a graph that shows the frequencies of radiation the molecules are absorbing; can be used to identify the functional groups in a molecule; use the data table information to match up the peaks on the spectrum with the functional groups that made them 2. The peaks show where radiation is being absorbed; the peaks on an IR spectra point downwards 3. Transmittance is always plotted on the y-axis and wavenumber on x-axis. Wavenumber is the measure used for frequency; 1/wavelength in cm

Fractional Distillation

1. Crude oil is vaporised at 350°C 2. Vaporised crude oil goes into a fractioning column and rises up through the trays; largest hydrocarbons do not vaporise at all due to very high boiling points; form a gooey resin instead 3. As crude oil vapour moves up the fractioning column it gets cooler; each length alkane has a different boiling point so they condense at different temperatures; fractions are drawn off at different levels in the column 4. Hydrocarbons with the shortest chain lengths/lowest boiling points do not condense; are drawn off as gases at the top of the column

Predicting mass spectra for diatomic molecules (E.g. Cl₂)

1. Express each % as a decimal (e.g. 75%→0.75 and 25%→0.25) 2. Make a table showing all the different Cl₂ molecules. For each, multiply the abundances of each isotope to get the relative abundance of each molecule. 3. Look for any values in the table that are the same and add up their abundances 4. Divide all the relative abundances by the smallest relative abundance to get the smallest whole number ratio. And by working out the relative molecular mass of each molecule, you can predict the mass spectra 5. Plot the mass spectra with the relative abundances you worked out on the y-axis and the relative molecular masses (m/z) on the x-axis

Infrared Radiation

1. In infrared (IR) spectroscopy, a beam of IR radiation is passed through a sample of a chemical 2. The IR radiation is absorbed by the covalent bonds in the molecules, increasing their vibrational energy 3. Bonds between different atoms absorb different frequencies of IR radiation; bonds in different places in a molecule absorb different frequencies too - OH group in an alcohol and OH group in a carboxylic acid absorb different frequencies

Identifying compounds using mass spectrometry

1. Molecules in a sample are bombarded with electrons to remove an electron and form a molecular ion, M⁺ 2. The molecular mass is shown by the molecular ion peak - the peak with the highest m/z value, not including any M+1 peaks caused by presence of carbon-13

Solid heterogenous catalyst mechanism

1. Reactant molecules arrive at surface and bond with the solid catalyst. This is called adsorption. 2. The bonds between the reactants atoms are weakened and break up. This forms radicals; these then get together and form new molecules 3. The new molecules then detach from the catalyst; desorption

In Industry

1. The conditions chosen in industry are a compromise 2. In exam, may be asked to look at a process and work out which conditions should be used to give a good balance of high rate and high yield 3. e.g. ethanol formed from ethene and steam; reversible exothermic reaction; C₂H₄ + H₂O ⇌ C₂H₅OH, ∆H = -46 kJ mol⁻¹, carried out at 60-70 atm and a 300°C temp and a phosphoric(V) acid catalyst - Exothermic reaction; lower temps. favour forward reaction; better yield at lower temp. However lower temp; lower rate of reaction; 300°C is a compromise between max yield and high rate - Higher pressures favour forward reaction, however high pressures are expensive to maintain

Molecule shapes: 2 electron pairs

2 electron pairs around a central atom make a linear molecule, e.g. CO₂ or BeCl₂ Therefore bond angle is 180°

Curly Arrows

A curly arrow shows where a pair of electrons goes during a reaction

Halogenoalkane

A halogenoalkane is an alkane with one or more halogens in place of hydrogens

Particles must collide to react

A reaction won't take place between two particles unless; - They collide with the right direction; they need to be facing each other the right way - They collide with at least a certain minimum amount of kinetic energy (Collision Theory^) - This minimum amount of energy is known as the activation energy; the particles need this much energy to break the bonds and start the reaction - Reactions with low activation energies happen easily; however reactions with high activation energy need to be heated

General Formula

Algebraic formula that can describe any member of a family of compounds. E.g. Alcohols; CⁿH²ⁿ⁺¹OH

Halogenoalkanes

Alkanes with one or more halogens in place of a hydrogen. Prefix indicates which halogen; chloro-/bromo-/iodo- e.g. chloroethane; CH₃CH₂Cl

Ionic bond

An ionic bond is the strong electrostatic attraction between two oppositely charged ions.

Levels of oxidation

An oxidising agent such as acidified dichromate (e.g. K₂Cr₂O₇/H₂SO₄) can be used to mildly oxidise an alcohol - Primary alcohols are oxidised to aldehydes and then carboxylic acids - Secondary alcohols are oxidised to ketones only - Tertiary alcohols cannot be oxidised

Reduction reaction

Any reaction in which a species gains electrons

Oxidation reaction

Any reaction in which a species loses electrons

First ionisation energies decrease down a group

As you go down the group in the periodic table, ionisation energies generally fall, i.e. it gets easier to remove outer electrons. This happens because elements further down the group have extra shells, so the atomic radius is larger, so the outer electrons are further away, which greatly reduces attraction.

Mass Spectrometry

Can be used to work out the relative atomic mass. Particles measured with a mass spectrometer must be charged, so they are often bombarded with electrons in order to remove one, giving a charge of +1 1. Multiply each relative isotopic mass by its relative isotopic abundance, and add up the results 2. Divide by the sum of the isotopic abundances

Carbocation stabilty

Carbocations with more alkyl groups are more stable because alkyl groups feed electrons towards the positive charge; the more stable carbocation is more likely to form

London forces (instantaneous-induced dipole attractions)

Cause all atoms and molecules to be attracted to each other - Electrons in charge clouds are always moving really quickly. At any given time there may be more electrons at one side of an atom or molecule than the other. At this moment the atom/molecule would have an instantaneous/temporary dipole - This can induce another temporary dipole in the opposite direction on a neighbouring atom. The two dipoles are then attracted to one another - Second dipole can then induce yet another dipole on a third atom; domino effect - Electrons are constantly moving; dipoles being created and destroyed all the time; overall effect despite this constant change is for atoms to be attracted to each other

Flame tests

Colours Li: Red Na: Orange/yellow K: Lilac Rb: Red Cs: Blue Ca: Brick red Sr: Crimson Ba: Green Test: Mix small amount of compound with HCl, use a piece of nichrome wire to put mixture into hot flame. Note colour produced

Rules of principle

Concentration; - Increasing concentration of reactant; equilibrium tries to get rid of the extra reactant by making more product; equilibrium shifts right - Increasing concentration of the product; equilibrium tries to remove the extra product; increases rate of reverse reaction; equilibrium shifts left - Decreasing concentrations has the opposite effect Pressure (gases only); - Increasing pressure shifts the equilibrium to the side with fewer gas molecules; this reduces the pressure - Decreasing pressure shifts the equilibrium to the side with more gas molecules; this raises the pressure - e.g. 2SO₂ + O₂ ⇌ 2SO₃; 3 molecules on left, 2 on right so increasing pressure shifts equilibrium to the right Temperature; - Increasing temperature add heat; this shifts the equilibrium in the endothermic direction; absorbs the extra heat - Decreasing temperature removes heat; equilibrium shifts in exothermic direction to replace the heat - If forward reaction is endothermic, reverse reaction will be exothermic and vice versa§

Biodegradable polymers

Decompose in the correct conditions - Decompose when organisms digest them - Can be made from renewable raw materials such as starch or from oil fractions such as isoprene - Renewable raw material is advantageous; won't run out, carbon neutral, save energy over oil-based plastics - However are more expensive than non-biodegradable

Molecular shape

Depends on electron pairs around the central atom - Number of electron pairs - Type of electron pairs

Properties: Simple covalent structures

E.g. CO₂ - Low melting and boiling points - Usually liquid or gas, sometimes solid (E.g. I₂) - Does not conduct electricity as a solid - Does not conduct electricity as a liquid - Can dissolve in water if it can form H-bonds

Properties: Metallic

E.g. Fe - High melting and boiling points - Solid at room temp. and pressure - Does conduct electricity as both solid and liquid due to delocalised electrons - Insoluble in water

Molecular lattices due to London Forces

E.g. Iodine molecules - Iodine atoms held together by strong covalent bonds to form molecules of I₂ - Then held in a molecular lattice arrangement by weak London forces - Structure is known as a simple molecular structure

Properties: Ionic structures

E.g. NaCl - High melting and boiling point - Solid at room temperature and pressure - Does not conduct electricity as a solid - Conducts electricity as a liquid - Is soluble in water

Properties: Giant covalent

E.g. diamond - High melting and boiling points - Solid at room temp. and pressure - Don't conduct electricity as solid (except graphite/graphene) - Do not melt; sublime instead - Insoluble in water

Enthalpy changes using mean bond enthalpies

Enthalpy change of reaction = Sum of bond enthalpies of reactants - sum of bond enthalpies of products

Bonds

For a substance to dissolve in another: - Bonds in the substance must break - Bonds in the solvent must break - New bonds form between solvent and substance - Usually substances only dissolve if the new bonds formed are around as strong or stronger than the strength of those broken

Adding halogens

Forms dihalogenoalkanes 1. The halogens are added across the double bond; one to each carbon in the double bond; its electrophilic addition 2. (mechanism shown in attached picture) 3. When an alkene is shaken with brown bromine water, the solution decolourises; this is due to the addition reaction taking place forms a colourless dibromoalkane; bromine water is used to test for C=C bonds

Equilibrium constant, Kc

Gives an idea of how far left or right the equilibrium is

Primary, secondary or tertiary

Halogenoalkanes with just one halogen atom can be primary, secondary or tertiary; on the carbon with the halogen attached: 1. A primary halogenoalkane has two hydrogen atoms and one alkyl group 2. A secondary halogenoalkane has one hydrogen atom and 2 alkyl groups 3. A tertiary halogenoalkane has no hydrogen atoms but 3 alky groups They have different reactivities too; Most reactive: tertiary > secondary > primary :least reactive

Alkenes

Have one or more double bonds between carbons; suffix: -ene, e.g. propene; CH₃CH=CH₂

Nomenclature

IUPAC System for naming organic compounds 1. Count the number of carbons in the longest continuous chain; gives the stem e.g. 5 = pent-, 3 = eth-, etc 2. Main functional group usually indicates which homologous series the molecule is in; gives a prefix or suffix 3. Number the longest carbon chain so that the main functional group has the lowest number. If there's more than one longest chain, pick the one with the most side chains 4. Any side chains or less important functional groups are added as prefixes to the start of the name; put in alphabetical order, after the number of the carbon each is joined to. 5. If there's more than one identical side chain or functional group, use di-, tri-, tetra- etc. before that part of the name, but ignore this when determining alphabetical order

Effect of pressure

Increase in pressure speeds up reactions - If any reactants are gases, increasing pressure has same effect as increasing concentration in a solution

Reducing power of halides

Increases down group - Ions get bigger - More shielding

Effect of concentration

Increasing concentration speeds up reactions - Increasing the concentration of reactants in a solution increases the number of particles in a given volume of the solution; particles will collide more frequently; more chances to react

Addition reactions

Joining 2 or more molecules together to form a larger molecule

Polymerisation

Joining together lots of simple molecules to form a giant molecule

Bonding pair type / bond angle size

Largest angle → Smallest angle: - Lone pair/lone pair angles - Lone pair/bonding pair angles - Bonding pair/bonding pair angles

Trends in group 7

Less reactive down group: - Increased atomic radii and shielding make it harder to attract the extra ion they need to react Melting and boiling points increase down group - Increased London forces due to increased number of electrons

Reaction rate from graph

Linear graph; Calculate gradient Curved graph; draw tangent at specified point on curve (e.g. for initial rate, draw tangent at time=0) then calculate gradient of tangent

Atoms

Made up of 3 subatomic particles - electrons, neutrons and protons

Test thermal stability of nitrates

Measuring: - How long it takes for a certain amount of oxygen to be produced - How long it takes until an amount of brown gas (NO₂) has been produced - done in fume cupboard; toxic

Test thermal stability of carbonates

Measuring: - How long it takes for an amount of carbon dioxide to be produced

Calculating Isotopic Masses from Relative atomic mass

Need: Relative mass of element and all but one of the abundances and isotopic masses of its isotopes 1. Find abundance of last isotope; percentage abundances so do 100-(sum of known% abundances) 2. Put into equation for finding the relative atomic mass and rearrange for the unknown value

Calculating mean bond enthalpies from reaction enthalpies

Need: enthalpy change of reaction and all but one of the bond enthalpies of reactants; rearrange prior equation

Oxidation number of neutral compounds

Overall charge is 0, but each constituent atom has its own oxidation number

Lone pairs

Pairs of electrons not shared between atoms

Halogen reactions with Alkanes

Photochemical reactions; - Started by UV light - Hydrogen atom is substituted with a halogen Mechanism; chlorine + methane → chloromethane - Initiation reaction; radicals are produced; sunlight provides enough energy to break the Cl-Cl bond; homolytic fission results in radicals - Propagation reactions; radicals used up and created in chain reaction; Cl•+CH₄→•CH₃+HCl, •CH₃+Cl₂→CH₃Cl+Cl•; continues until all Cl₂ and/or CH₄ molecules are used up - Termination reactions; radicals are eliminated - joined together to form stable molecules; many possible termination reactions, e.g. Cl•+•CH₃→CH₃Cl, •CH₃+•CH₃→C₂H₆ (trace product)

Polar solvents and Non-polar solvents

Polar: - made up of polar molecules such as water - not all polar solvents can form h-bonds, e.g. propanone (acetone) only forms London forces and permanent-permanent dipole bonds Non-polar - E.g. hexane - Bond to each other by London forces Many substances are soluble in one type of solvent but not the other

Substitution reactions to form halogenoalkanes

Reacting alcohols with PCl₅ or HCl to produce chloroalkanes - With PCl₅; ROH + PCl₅ → RCl + HCl + POCl₃ - With HCl; ROH + HCl → RCl + H₂O - Rate of reaction is fastest for tertiary alcohols and slowest for primary alcohols OH⁻ can be swapped with bromine to make a bromoalkane - Alcohols will react with compounds containing bromide ions in a substitution reaction - The hydroxyl group is replaced by the bromide, so the alcohol transforms into a bromoalkane - The reaction also requires an acid catalyst, e.g. 50% concentrated H₂SO₄ Iodoalkanes can be made using red phosphorus and iodine - Reacting an alcohol with PI₃ makes an iodoalkane - The PI₃ is normally made in-situ - General equation; 3ROH + PI₃ → 3RI + H₃PO₃

Trends across Periods 2 and 3 for melting and boiling points

Relatively strong increase across the period (metals→giant lattices) until group 5, where there is a sharp drop (due to simple molecular structures), then slight downward trend across the rest of the period.

Oxidation number of monatomic ion

Same as its charge, e.g. Na⁺ or Mg²⁺

Displayed formula

Shows how the atoms are arranged and the bonds between them

Structural Formula

Shows the arrangement of atoms carbon by carbon, with the attached hydrogens and functional groups e.g. methylpropane; CH₃CH(CH₃)CH₃

Skeletal formula

Shows the bonds of the carbon skeleton only with any functional groups; the hydrogen and carbon atoms are not shown. E.g. pentane

Alkanes

Simplest hydrocarbons; no double bonds or side chains suffix = -ane; example: propane - CH₃CH₂CH₃ Branched alkanes; alkyl groups attached to main chain (suffix: -yl) e.g. methylpropane; CH₃CH(CH₃)CH₃

Empirical formula

Simplest whole number ratio of atoms of each element in a compound E.g. Ethane, C₂H₆ would have empirical formula of CH₃

Hydrolysis Reaction

Splitting a molecule into two new molecules by adding H⁺ and OH⁻ derived from water

Intermolecular forces in organic molecules

Strength depends on their shape; - Alkanes have covalent bonds inside the molecules - London forces between molecules hold them together - Longer the carbon chain, stronger the London forces; more molecular surface contact and more electrons to interact - As molecules get longer, it gets harder to separate them because more energy is needed to overcome the London forces - Branched-chain alkanes cannot pack closely together and their molecular surface contact is small compared to straight-chain alkanes of similar molecular mass

Like dissolves like - usually

Substances dissolve best in substances with similar molecular forces

Electronegativity

The ability of an atom to attract bonding electrons in a covalent bond - Usually measured on Pauling scale; higher value, more electronegative e.g. Fluorine = 4.0 (most electronegative element) - Least electronegative elements have values around 0.7 - More electronegative atoms have higher nuclear charges and smaller atomic radii - Increases across periods and up groups - Pauling values in data book

Molecular formula

The actual number of each element in a molecule e.g. Ethane; C₂H₆

Electrophilic Addition

The alkene's double bond opens up and atoms are added to the carbon atoms; 1. Electrophilic addition happens because the double bond has plenty of electrons and is easily attacked by electrophiles 2. Electrophiles are electron pair acceptors; usually slightly short on electrons so they are attracted to electron-dense areas 3. Electrophiles are usually positive ions and polar molecules

Relative molecular/formula mass

The average mass of a molecule compared to the mass of an atom of carbon-12

Reaction rate

The change in the amount of reactant or product per unit time

Standard enthalpy change of formation, ΔfH°

The enthalpy change when 1 mole of a compound is formed from its elements in their standard states, under standard conditions

Standard enthalpy change of combustion, ΔcH°

The enthalpy change when 1 mole of a substance is completely burned in oxygen, under standard conditions.

Standard enthalpy change of neutralisation, ΔneutH°

The enthalpy change when an acid and an alkali react together under standard conditions to form 1 mole of water

Standard enthalpy change of reaction, ΔᵣH°

The enthalpy change when the reaction occurs in the molar quantities shown in the chemical equation, under standard conditions

First Ionisation Energy

The first ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous ions with a charge of +1. It is an endothermic process

Bond Enthalpy in relation to bond length

The higher the electron density between the nuclei (the more electrons in the bond), the stronger the attraction between atoms; the higher the bond enthalpy and the shorter the bond length I.e. Shorter bonds; higher enthalpy

Markownikoff's rule

The major (most produced) product from the addition of a hydrogen halide (HX) to an unsymmetrical alkene is the one where the hydrogen adds to the carbon with the most hydrogens already attached

Relative atomic mass

The relative atomic mass is the weighted mean mass of an atom of an element compared to 1/12th of the mass of an atom of carbon-12 Can be worked out from Isotopic Abundances - Multiply the isotopic mass of each isotope by its % abundance, add them up then divide the total by 100

Covalent bonds; Sigma and Pi bonds

The way that atomic orbitals overlap causes different types of bonds to form 1. Single covalent bonds in organic molecules are sigma bonds; formed when two orbitals overlap in a straight line, in the space between two atoms; gives highest possible electron density between the two positive nuclei 2. This high electron density means there is a strong electrostatic attraction between the nuclei; high bond enthalpy; strongest type of covalent bond 3. Double bonds are made up of a sigma bond and a pi bond; a pi bond is formed when two lobes of two orbitals overlap sideways; one above and one below the molecular axis; e.g. p-orbitals 4. In a pi bond, the electron density is spread out above and below the nuclei; the electrostatic attraction between the nuclei and shared pair is weaker than in sigma bonds; pi bonds have relatively low bond enthalpy 5. This means that double bonds are less than twice as strong as single bonds 6. In alkenes, the C-C and C-H bonds are all sigma bonds. The C=C bonds contain both a sigma and a pi bond

Intermolecular forces

These are the weak electrostatic attractions between molecules, including: - London forces (instantaneous-induced dipole bonds) - Permanent dipole-permanent dipole bonds - Hydrogen bonding (strongest type of intermolecular force)

Drop in ionisation energy between groups 2 and 3 shows subshell structure

This is because group 3 elements have their outermost electron in a p orbital rather than an s orbital; the outermost electron therefore is further away from the nucleus, this, and the shielding provided by the s orbital below the p orbital, is enough to override the effect of the increased nuclear attraction

Atomic Radius decreases across a period (periodicity)

This is because: - As the number of protons increases, the positive charge on the nucleus also increases, causing electrons to be pulled closer to the nucleus. - The extra electrons that the elements gain across a period are added to the outer energy level so they don't provide any extra shielding.

Metals forming compounds

Usually donate electrons to form positive ions, oxidation numbers increase

Non-metals forming compounds

Usually gain electrons to form negative ions; oxidation numbers decrease

Methods for disposing of polymers

Waste plastics can be buried - Landfill can be used to dispose of waste plastics when the plastic is difficult to separate from other waste, not sufficient in quantities to make separation financially worthwhile or too technically difficult to recycle - Landfill needs to be reduced as much as possible because the amount of waste generated is becoming a problem Waste plastics can be reused - Many plastics are made from non-renewable oil-fractions; so it makes sense to reuse whenever possible - More than one way to reuse plastics; recycling (e.g. poly(propene)) by melting and remoulding or cracked into monomers for use to make more plastics or other chemicals Waste plastics can be burned - If recycling is not possible, waste plastics can be burned and the heat used to generate electricity - this process needs to be carefully controlled to reduce toxic gas production; e.g. polymers containing chlorine produce HCl when burned; must be removed - Waste gases from combustion are passed through scrubbers which neutralise gases such as HCl by reacting them with bases - Plastics can be sorted before they are burned to separate out any materials that will produce toxic gases

Elimination reactions

When a small group of atoms breaks away from a larger molecule

Disproportionation reaction

When an element in a single species is simultaneously oxidised and reduced in the same reaction e.g. Cl₂ (oxidation number 0) + 2OH⁻ → ClO⁻ (Cl has oxidation number +1) + Cl⁻ (-1) + H₂O

Substitution Reactions

When one species is replaced by another

Cycloalkanes

cycle-...-ane suffix/prefix; e.g. cyclohexane; C₆H₁₂

Calculate enthalpy changes

q = mcΔT, where; q = heat lost or gained (Joules) m = mass of water in calorimeter, or solution in the insulated container (g) c = specific heat capacity of water (4.18Jg⁻¹K⁻¹) ¹ ΔT = change in temperature of the water or solution (°C or K)

Work out rate from experimental data

rate of reaction = (amount of reactant used or amount of product formed)/time taken Rate of reaction is proportional to 1/time