Edexcel A level chemistry paper 1 topics questions and answers, Exams of Nursing

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Atoms - Made up of 3 subatomic particles - electrons, neutrons and protons Electrons - - Subatomic particle

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

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• Charge of +
• 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 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

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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- 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 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- Relative molecular/formula mass - The average mass of a molecule compared to the mass of an atom of carbon- 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. Multiply each relative isotopic mass by its relative isotopic abundance, and add up the results
2. Divide by the sum of the isotopic abundances

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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 Predicting mass spectra for diatomic molecules (E.g. Cl₂) - 1. Express each % as a decimal (e.g. 75%→0.75 and 25%→0.25)
3. 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.
4. Look for any values in the table that are the same and add up their abundances
5. 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
6. 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

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

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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 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 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

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• 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 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 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.

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Ionisation - The removal of one or more electrons 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 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. 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.

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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. 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⁻ 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

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Modern period table organises elements by Proton Number - - Dmitri Mendeleev created base for modern periodic table in 1869 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 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.

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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 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 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

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• 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 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 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

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sharp drop (due to simple molecular structures), then slight downward trend across the rest of the period. Ionic bond - An ionic bond is the strong electrostatic attraction between two oppositely charged ions. 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 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 Trends in ionic radii - - Increases down a group

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• Ionic radius of a set of isoelectronic ions decreases as the atomic number increases Isoelectronic ions - - Ions of different elements with the same number of electrons 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 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

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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 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" 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 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 bond, there must be a balance between these forces

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• The distance between the two nuclei is the distance where the attractive and repulsive forces balance each other
• This distance is the bond length Dative covalent bonding - This is where both electrons in a covalent bond come from the same atom. E.g. CO, NH₄⁺ Molecular shape - Depends on electron pairs around the central atom
• Number of electron pairs
• Type of electron pairs 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

Bonding pair type / bond angle size - Largest angle → Smallest angle:

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• Lone pair/lone pair angles
• Lone pair/bonding pair angles
• Bonding pair/bonding pair angles Electron pair repulsion theory - - A way of predicting molecular shape using electron pair repulsion.
• Need to know shapes and angles for; methane, ammonia, water 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:

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• Count a double bond as 2 bonds
• Bonding centres are the atoms bonded to the central atom 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° 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 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

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• Two lone pairs; distorted T; 87.5° angles e.g. ClF₃ 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 Lone pairs - Pairs of electrons not shared between atoms 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 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

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

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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 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

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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 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 both atoms have 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 has a partial charge; δ⁺ or δ⁻; bond is said to be polar

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• 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 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 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

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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 Intermolecular forces - These are the weak electrostatic attractions between molecules, including:

• London forces (instantaneous-induced dipole attractions)
• Permanent dipole-permanent dipole attractions
• Hydrogen bonding (strongest type of intermolecular force) 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

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• 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 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 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