Module 2 Chemistry Revision Card

Protons have a 1+ charge and a mass of 1 - number of protons is the atomic number

Electrons have a 1- charge and a mass of about 0.0005 (1/1836)

Neutrons have a charge of 0 and a mass of 1

Isotopes - atoms of the same element with a different number of neutrons and different masses

Mass number is the number of protons and number of neutrons, electrons and protons are equal

An ion is a charged atom where electrons are different to the number of protons. A cation is a postive ion and an anion is a negative ion

• Cations such as Mg2+ have lost 2 electrons so protons number is 12 but electron number will be 10 (12-2)
• Anions such as Cl- have gained an electron so proton number is still 17 but electron number is now 18 (17+1)

Relative isotopic mass - mass of an isotope relative to 1/12th of the mass of an atom of carbon 12 e.g. oxygen has a relative isotopic mass of 16 (assume same as mass number)

Relative atomic mass - weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12

• Takes into account percentage abundance of each isotope and relative isotopic mass of each isotope
• Relative atomic mass (Ar) can be determined using a mass spectrometer
• Sample is placed in the mass spectrometer and vaporised then ionised to form cations. Ions are accelerated and heavier ions move more slowly so isotopes are separated. Ions are detected on a mass spectrum as a mass-to-charge ratio
• To calculate this do 

Atoms of metal on the LHS of the periodic table lose electrons to form cations

Atoms of non-metals on the RHS of the periodic table gain electrons to form anions

Binary compounds contain two elements only e.g. sodium oxide

There are some polyatomic ions (more than one element) which have to be learnt

• NO3 - = nitrate
• CO3 2- = carbonate
• SO4 2- = sulfate
• OH- = hydroxide
• NH4 + = ammonium
• Zn2+ = zinc
• Ag+ = silver

Formula from ions can be written by comparing charges which should cancel when the ions react to form a compound

Some elements exist as diatomic molecules - two atoms bonded together e.g. N2, H2, O2 or as other small molecules e.g. P4 and S8

State symbols include (g) for gas, (aq) for aqueous, (l) for liquid and (s) for solid

When balancing equations numbers must go at the front so there are the same number of atoms of each element on each side. BE CAREFUL WHEN THERE ARE BRACKETS

Amount of substance is used to count the number of particles in a substance measured in mol

• One mole contains 6.02 x 10 ^23 particles = Avogadro's constant
• So number of moles x Avogadro's constant gives number of particles

Molar mass gives the mass in grams in each mole of the substance

• Mass per mole of substance measure in g mol ^-1
• Molar gas volume is gas volume per mole and is measured in dm3 mol-1
• Mole = mass divided by Mr

Molecular formula - number of atoms in each element in a molecule. Can be found from empirical formula to multiply to get given Mr

Empirical formula - simplest whole number ratio of atoms in each element in a compound

• Calculated by doing mass (or %) divided by Mr to get number of moles
• The divide by smallest number of moles to get ratio 
• May need to double etc to get whole number

Relative molecular mass (Mr) - compares mass of a molecule with mass of an atom of carbon-12

• Calculated by adding up relative atomic masses of elements in a molecule

Relative formula mass - compares mass of a formula unit (compound made from ions) with the mass of an atom of carbon-12

• Calculated by adding relative atomic masses of elements in the empirical formula

Hydrated salts

• Water molecules are part of the crystalline structure - this is the water of crystallisation
• The anhydrous salt is the salt with the water of crystallisation removed
• Shown by the salt then a large dot then the number of water molecules
• Number of water molecules can be calculated using empirical formulas
• Sample should be heated and weighed constantly until mass becomes constant so all water has been removed

Assumptions made with hydrated salts

• If the hydrated salt and anhydrous compound has different colours then you can tell when the water has been removed however, there could still be water underneath. Therefore, reheat repeatedly until mass becomes constant so all water removed
• Many salts decompose further when heated so it is difficult to judge if further decomposition has occured

Volume can be measured in cm3 and this can be converted to dm3 by dividing by 1000

• mole = volume (dm3) x concentration (moldm-3)

Standard solutions are solutions of known concentration, they can be prepared by dissolving exact mass of the solute in a solvent and making up the solution to an exact volume

• Mole can be calculated then using Mr the mass can be found
• Mass concentration can be written as g dm-3

Molar gas volume - volume per mole of gas molecules at a stated temperature and pressure

• Can occur at RTP which is 20C and 101kPa
• This means mole = volume in cm3 divided by 24000 or if in dm3 volume divided by 24

Ideal gas equation is used when not at RTP. Following assumptions are made about molecules in an ideal gas

• Random motion, elastic conditions, negligible size and no intermolecular forces

pv = nrt

• Where p is the pressure in Pa   (kPa -> Pa x 10^3)
• Where v is the volume in m^3     (cm3 to m3 is x10-6 or dm3 to m3 is x10-3)
• Where n is the number of moles
• Where r is the gas constant which is 8.314
• Where t is the temperature in K    (C to K is +273)

Each 1K rise in temperature is the same as a 1C rise in temperature

Balancing number in equations gives the ratio of moles of each substance = stoichiometry

• Can be used to calculate moles of other substances by comparing ratios

Percentage yield 

• Maximum possible amount of product = theoretical yield but this is difficult to achieve as reaction may not have gone to completion, other reactions may have occured or purification could have lost some of the product
• Actual yield is usually less than theoretical yield
• Percentage yield = actual yield divided by theoretical yield x100

Limiting reagent

• Reactant not in excess will be used up first and is the limiting reagent
• To find out which reactant is in excess work out moles 
• If 2:1 then the element with 2 will be used up first because more is required so this is the limiting reagent and calculations need to be based on number of moles of this

Atom economy - measure of how well atoms have been utilised. Reactions with high atom economies

• Produce large proportion of desired products and few waste products
• Are important for sustainability 

Atom economy = (sum of molar masses of desired products) divided by (sum molar masses of all products) x100

Efficiency of reactions depends on other factors

• Cost of getting raw materials and producing reactant
• Percentage yield obtained (reactants converted to products)
• Harmful products such as carbon dioxide could be produced

Strong acids completely dissociate their H+ ions, weak acids only partially dissociate so are in equilibrium.

Metal oxides, metal hydroxides, metal carbonates and ammonia are classed as bases and they neutralise an acid to form a salt.

An alkali is a base that dissolves in water releasing hydroxide ions into the solution.

When H+ ions react with a base to form a salt and water, the H+ ions from the acid are replaced by metal or ammonium ions from the base

• With alkalis, the reactants are in solution but acid + alkali ----> salt + water
• Ionic equation is just H+  + OH- ----> H20
• Carbonates also neutralise acids to form a salt and water

Examples of acids - HCl, H2SO4, HNO3, CH3COOH

Examples of bases - NaOH, KOH, NH3

A titration is used to accurately measure the volume of one solution that reacts exactly with another solution. They can be used for

• Finding the concentration of a solution
• Identification of unknown chemicals
• Finding the purity of a substance

Standard solutions are solutions of known concentration that are made in volumetric flasks. A 100cm3 volumetric flask has a tolerance level of +-0.20cm3 and a 250cm3 flask had a tolerance level of +- 0.30cm3. To prepare a standard solution

• Solid weighed accurately then dissolved in a beaker using less distilled water than needed to fill the volumetric flask
• Solution is transferred to a volumetric flask and the beaker rinsed with distilled water to remove any traces of solid
• Flask is filled to the graduation line by adding distilled water until the bottom of the meniscus lines up with the mark
• The volumetric flask is then inverted several times to mix the solution

Acid is titred against a base using a pipette and burette which have the tolerance levels of:

• 10cm3 pipette +-0.04cm3
• 25cm3 pipette +- 0.06cm3
• 50cm3 burette +-0.10cm3
• Burette readings are measured to the nearest two decimal places +-0.05cm3

Procedure

• Add a measured volume of one solution to a conical flask using a pipette
• Add the other solution to a burette and record the initial burette reading
• Add a few drops of indicator to the solution in the conical flask
• Run the solution in the burette into the solution in the conical flask and swirl, colour change shows end point of the titration
• Record final burette reading and carry out trial titration to find approximate titre
• Repeat titration until two titres are concordant - values within 0.1cm3

Mean titre is calculated by only using concordant accurate titres. From a titration you known the concentration and volume of one solution and reacting volume of another.

Oxidation numbers are always 0 for elements

Oxidation numbers have a sign which is always placed BEFORE the number. There are some special cases to learn

• H in metal hydrides such as NaH and CaH2 is -1 rather than +1
• O in peroxides is always -1 in H2O2
• O is bonded to F it has a charge of + 2 e.g. in Fe2O

Sum of oxidation numbers = total charge. Sometimes will have to be put equal to charge and rearranged to find oxidation number

Roman numerals can be used in naming for example iron(II) is Fe2, nitrite is NO2- and nitrate (V) is NO3-

Reduction - gain of electrons, decrease in oxidation number

Oxidation - loss of electrons, increase in oxidation number

The number of electrons that can fill each shell (n=1 to n=4) is 2,8,18,32

• Shells are classed as energy levels where energy increases as shell number increases
• Shell number or energy level number is called principle quantum number, n 

Shells are made of atomic orbitals which is a region around the nucleus which can hold up to two electrons with opposite spins

• s-orbitals are shaped like a sphere and can hold two electrons, greater shell number, greater radius of s-orbital
• p-orbtials are shaped like a dumb-bell, there are three orbitals - px, py and pz each hold two electrons. Each shell from n=2 contains three p-orbitals so 8 (+2 from s)
• d orbitals - n=3 contains five d-orbitals, each with two electrons so contains 18 in total (10+8)
• f-orbitals - n=4 contains seven f-orbitals, each with two electrons (14+18) so 32 total

Sub-shells:

Orbitals fill in order of increasing energy using the sequence on the previous side e.g. n=2 is 2s, 2p then n=3 is 3s, 3p, 3d

• d orbitals are at a higher energy level than s orbitals so s orbitals will fill and be emptied first

Electrons are negatively charged so repel each other so electron pairs are drawn with opposite spins to help counteract repulsion.

Orbitals are occupied singly at first and then filled up with the second electron

Have to be able to write electron configurations for atoms up to ATOMIC NUMBER 36

• Electron configurations show how sub-shells are occupied by electrons, and number of electrons = atomic number
• REMEMBER fill 4s first then the 3d orbital
• Shorthand electron configurations can be written using the previous noble gas in the table then the outer electron sub-shells

• Positive ions (cations) form when atoms lose electrons
• Negative ions (anions) form when atoms gain electrons

When doing electron configurations for ions, have to take the charge into account e.g. if Mg2+ then the electron configuration would be done for 12-2 so 10 electrons

The periodic table can be divided into blocks corresponding to highest energy sub-shells

Ionic bonding - electrostatic attraction between positive ions and negative ions, holds cations and anions together in ionic compounds

Ionic compounds contain metals and non-metals and can be drawn using dot an cross diagrams

• Each molecule is drawn separately and has its own charge

Giant ionic lattices are formed from oppositely charged ions being strongly attracted in all directions. Properties of ionic compounds

• Melting and boiling point - solids at room temperature as insufficient evidence to overcome strong electrostatic forces of attraction. High melting and boiling point to be able to overcome forces. Melting point is higher for lattices containing ions with a greater ionic charge because there is stronger attraction between ions. Size also important
• Most ionic compounds dissolve in polar solvents which break down the lattice and surround each ion in solution. In compounds with ions of large charges the ionic attraction may be too strong for water so the compound is not very soluble. Solubility depends on strength of attractions within the giant ionic lattice and attraction between ions and water

• Solubility requires the ionic lattice to be broken and for water molecules to surround ions
• In a solid state, an ionic compound does not conduct electricity because ions are fixed in position so there are no mobile charge carries. When molten or dissolved, ions are free to move and carry the charge

High melting and boiling point

Dissolve in polar solvents

Conduct when molten or dissolved

Covalent bonding - strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atom. Occurs between two non-metals

• Atomic orbitals overlap to give a shared pair of electrons which is attracted to the nuclei of both bonding atoms

In covalent bonds the attraction is localised between the electron and nuclei of the bonded atom causing a molecule to form which consists of two or more atoms. Dot and cross diagrams can also be used for covalent bonding except electrons are shared so they are not drawn separately 

• Displayed formula can also be used which shows positing of atoms, their bonding and any lone pairs
• Number of covalent bonds depends on number of electrons in outer shell 
• BF3 is unique - only has 6 electrons in outer shell once 3F joined
• Different numbers of unpaired electrons lead to different covalent compounds in sulfur
• Expansion of the octet - occurs in sulfur hexafluoride where a d-shell becomes available for bonding

A multiple covalent bond exists when two atoms share more than one pair of electrons

• Double covalent bond - electrostatic attraction is between two shared pairs of electrons and the nuclei of the bonding atom e.g. CO2
• Triple covalent bond - electrostatic attraction is between three shared pairs of electrons and the nuclei of the bonding atoms e.g. N2, HCN

Dative covalent or coordinate bond - covalent bond where the shared pair of electrons has been supplied by one of the bonding atoms only

• Shared electron pair was originally a lone pair
• E.g. NH4 + where an NH3 donates a lone pair of electrons to a H+ ion which has no electrons as it has lost one. Both drawn with same symbol and charge on outside of molecule after square brackets

Average bond enthalpy - measurement of covalent bond strength. Larger value, stronger covalent bond

Electron-pair repulsion theory

• Electron pairs surrounding a central atom determine the shape of the molecule or ion
• Electrons repel as far as possible

A lone pair of electron is slightly closer to the central atom so occupies more space, meaning it has more repulsion. Lone pairs reduce the angle by 2.5 for each lone pair and has more repulsion than a bonding pair

Wedges are used to show a 3D shape of a molecule

Attraction of a bonded atom for the pair of electrons in a covalent bond = electronegativity

• Pauling scale is used to compare electronegativities
• Electronegativities increase along and up the periodic table so F is the most electronegative (4) and potassium the least
• Noble gases not included as they don't tend to form compounds

Electronegativity values can be used to estimate the type of bonding

• Covalent - electronegativity difference of 0
• Polar covalent - electronegativity difference of 0-1.8
• Ionic - electronegativity difference of >1.8

Bond polarity

1) Non-polar bond - electron pair is shared equally between the bonded atoms

• Bonded atoms are the same or have similar/same electronegativities
• Pure covalent bond forms between the same element e.g. H2

2) Polar bonds - bonded electron pair is shared unequally between bonded atoms

• Bonding atoms are different and have different electronegativity values forming polar covalent bonds e.g. HCl
• Chlorine is more electronegative than hydrogen so has a greater attraction for the bonded pair of electrons so gets a delta negative charge and the hydrogen delta positive
• Separation of opposite charges = dipole, HCl is a polar molecule
• A dipole in a polar covalent bond does not change and is a permanent dipole 

Shapes in polar molecules can mean dipole either cancel each other out or reinforce each other to produce a larger dipole

• H2O is a polar molecule with the H and O having a permanent dipole
• The two dipoles act in different directions and produce oxygen with a delta negative charge and hydrogen with a delta positive charge

A carbon dioxide molecule is non-polar, the two C=O bonds have permanent dipoles. These act in opposite directions and exactly cancel each other out. Therefore, the overall dipole is zero.

Ionic lattices such as sodium chloride can be dissolved by water

• Water molecules attract Na+ and Cl-
• The ionic lattice breaks down as it dissolves forming a solution
• The Na+ ions are attracted towards the oxygen due to its negative charge
• The Cl - ions are attracted towards the hydrogen due to its positive charge

Intermolecular forces are weak interactions between dipoles of different molecules. These can either be

• Induced dipole-dipole interactions (London Forces)
• Permanent dipole-dipole interactions
• Hydrogen bonding

These tend to determine melting and boiling point whereas covalent bonding determines identity and chemical reactions of molecules. Strength:

London < permanent dipole < hydrogen bonding < single covalent

Induced dipole-dipole interactions (London forces)

• Weak intermolecular forces that exist between all molecules polar or non-polar
• Movement of electrons produces a changing dipole in a molecule which forms an instantaneous dipole
• This instantaneous dipole induces a dipole on neighbouring molecules
• Induced dipole induces further dipoles on neighbouring molecules which attract others

Induced dipoles result from interactions of electrons between molecules therefore, the more electrons the larger the instantaneous and induced dipoles and the stronger forces between molecules. More energy is needed to overcome the intermolecular forces so boiling point increases

Permanent dipole-dipole interactions - act between polar molecules

• Permanent dipoles form between H and Cl in two different molecules of HCl
• This also contains London forces so boiling point is higher than a non-polar molecule of F2 because extra energy is needed to break permanent dipoles
• A boiling point of -85 is higher than -220 and requires more energy

Simple molecular substances are made of small units containing a definite number of atoms and molecular formula e.g. H2. In a solid state, simple molecules form a regular structure = simple molecular lattice

• Molecules are held in place by weak intermolecular forces
• Atoms within each molecule are bonded together by strong covalent bond

1) Low melting and boiling point 

• Simple molecular substances are covalently bonded and can exist as solids, liquids or gases at room temperature
• All simple molecular substances can be solidified into a molecular lattice by reducing the temperature 
• They have low melting and boiling points
• When broken during melting only weak intermolecular forces break, covalent bonds do not break

2) Solubility

• Non-polar simple molecular substances are soluble in non-polar solvents
• They are insoluble in polar solvents because there is little interaction between molecules in the lattice and solvent molecules. Intermolecular bonding within the solvent is too strong to be broken
• Polar covalent substances may dissolve in polar solvents as they can attract each other
• Solubility depends on strength of the dipole but this can be difficult for molecules such as ethanol with polar and non-polar parts
• Hydrophilic parts are polar and hydrophobic parts are non-polar so won't interact with water

3) Electrical conductivity

• No mobile charged particles in simple molecular structures so they cannot conduct

Hydrogen bonding is found between molecules containing

• Electronegative atoms with a lone pair of electrons
• H-O, H-N or H-F

Hydrogen bonding acts between a lone pair of electrons on the electronegative atom in one molecule and a hydrogen atom in a different molecule - strongest type of intermolecular force. Shown by a dashed line, also needs dipoles and lone pairs

1) Ice is less dense than liquid water

• Hydrogen bonds hold water molecules apart in an open lattice structure so they are further apart than in water making ice less dense 
• Holes decrease the density of water, when melting the ice lattice collapses and molecules move closer together

2) Water has a high melting and boiling point

• Water has London forces between molecules but hydrogen bonds are extra forces which require more energy to break
• When ice lattice breaks, the arrangement of hydrogen bonds in broken and when the water boils the hydrogen bonds break completely
• Without hydrogen bonds water would exist as a gas at RTP

Water also has a high surface tension and viscosity so droplets form