Topic 1 - key concepts in chemistry

dalton model of an atom

John Dalton (1803)

• he published his own three-part atomic theory:

  1. all substances are made of atoms. atoms are small particles that cannot be created, divided or destroyed

  2. atoms of the same element are exactly alike, and atoms of different elements are different

  3. atoms join with other atoms to make new substances

     • much of dalton’s theory was correct, but some of it was later proven incorrect and revised as scientists learned more about atoms

J.J Thomson model of an atom

jj thomson (1897)

• used a cathode-ray tube to conduct an experiment

• this discovery identified an error in Dalton’s atomic theory. atoms can be divided into smaller parts

• because the beam moved away from the negatively charged plate and toward the positively charged plate, Thomson knew that the particles must have a negative charge

• Thomson proposed a model of an atom called the “plum pudding model”, in which negative electrons are scattered throughout soft blobs of positively charged material

Ernest Rutherford atomic model

Ernest Rutherford (1909)

• shot a beam of positively charged particles into a sheet of gold foil

• most of the particles did continue in a straight line (as expected from plum pudding model). however some of the particles were deflected to the sides a bit, and a few bounced straight back

• Rutherford developed a new model which said that most of the atom’s mass is found in the region in the centre called the nucleus

• in Rutherford’s model the atom is mostly empty space, and the electrons travel in random paths around the nucleus

describe the structure of an atom

a nucleus containing protons and neutrons, surrounded by electrons in shells

relative charge and relative mass of a proton

+1, 1

relative charge and relative mass of a neutron

0, 1

relative charge and relative mass of an electron

-1, 1/1836

explain why atoms contain equal number of protons and electrons

• atoms are neutral and the charges on a proton are +1 and on an electron -1

• therefore amount of protons = amount of electrons, so that the charges cancel

describe the size of the nucleus of an atom

the nucleus of an atom is very small compared to the overall size of the atom

where is most of the mass of an atom concentrated

in the nucleus

meaning of the term mass number of an atom

mass (nucleon) number = number of protons + neutrons

what is an isotope

isotopes are different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei

calculate the numbers of protons, neutrons and electrons in atoms given the atomic number and mass number

• atomic (proton) number = number of protons (= number of electrons if it’s an atom, because atoms are neutral)

• therefore, you can calculate number of neutrons by doing mass number - atomic number

explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers

• because isotopes have the same number of protons but different numbers of neutrons, they are still atoms of the same element, but they have different atomic masses

• the relative atomic mass is calculated using the abundance of different isotopes and because it is an average it can lead to the relative atomic mass not being a whole number (atomic number and mass number will always be whole number since they are not averages)

• since the mass of atoms is so small, we compare their masses to each other. a carbon atom having a mass number 12, (^12C) is taken as standard for this comparison and its relative atomic mass is 12

• it is written as Ar or R.A.M.

calculate the relative atomic mass of an element from the relative masses and abundances of its isotopes

the average mass, or R.A.M. of chlorine can be calculated using the following equation:

R.A.M. = (mass of isotope-A x % of isotope-A) + (mass of isotope-B x % of isotope-B) / 100

describe how Mendeleev arranged the elements, known at that time, in a periodic table

by using properties of these elements and their compounds

• he ordered his table in order of atomic mass, but not always strictly - ie. in some places he changed the order based on atomic weights

• left gaps for elements that he thought had not been discovered yet

describe how Mendeleev used his table to predict the existence and properties of some elements not discovered yet

• Mendeleev realised elements with similar properties belonged in the same groups in the periodic table so was able to leave gaps and place the discovered elements where they fit best

• elements with properties predicted by Mendeleev were later discovered and filled the gaps

explain that Mendeleev thought he has arranged elements in order of increasing relative atomic mass but this was not always true

because of the relative abundance of isotopes of some pairs of elements in the periodic table

• knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct, because some elements have a higher mass than others when isotopes are taken into account, but a lower one if you look at one specific isotope

explain the meaning of atomic number of an element

in terms of position in the periodic table and number of protons in the nucleus

• elements are arranged in order of atomic (proton) number (bottom number) and so that elements with similar properties are in columns, known as groups

• elements in the same group have the same amount of electrons in their outer shell, which gives them similar chemical properties

describe how elements are arranged in the periodic table

• elements are arranged in order of increasing atomic number, in rows called periods and elements

• elements with similar properties are placed in the same vertical columns called groups

metals in periodic table

metals = elements that react to form positive ions

• majority of elements are metals

• found to the left and towards the bottom of the periodic table, because they lose electron(s) in order to form these positive ions, forming an electronic structure that is stable, like that of a noble gas

non metals in the periodic table

non-metals = elements that do not form positive ions

• found towards the right and top of the periodic table, because they gain electron(s) in order to form these negative ions, forming an electronic structure that is stable, like that of a noble gas

predict the electronic configurations for the first 20 elements in the periodic table as diagrams and in the form

• the electronic configuration of an element tells you how many electrons are in each shell around an electron’s nucleus

• for example, sodium has 11 electrons: 2 in its most inner shell, then 8, then 1 in its outermost shell

  • you can represent sodium’s electronic configuration as: 2.8.1

  • or in a diagram

explain how the electronic configuration of an element is related to its position in the periodic table

• the group an electron is in tells you how many electrons are in its outermost shell aka group 1 elements have 1 electron in their outer shell

• the period of an electron tells you which number shell an element’s outermost electron is found in aka period 3 elements have their outermost electrons in shell 3

• remember all the shells up until the shell will be full (for the 1st shell means 2 electrons and for shells 2 and 3 this means 8 electrons)

explain how ionic bonds are formed (including the use of dot and cross diagrams)

formed by the transfer of electrons between atoms to produce cations and anions

• metals + non-metals: electrons in the outer shell of the metal atom are transferred

  • metal atoms lose electrons to become positively charged ions (cation)

  • non-metal atoms gain electrons to become negatively charged ions (anion)

• electron transfer during the formation of an ionic compound can be represented by dot and cross diagram

what is an ion

an ion is an atom or group of atoms with a positive or negative charge

• since an ion is formed from a metal losing an electron, ie. becoming a positive metal ion or from a non-metal gaining an electron, ie. becoming a negative ion

calculate the numbers of protons, neutrons and electrons in simple ions given the atomic number and mass number

• atomic number = proton number = number of protons

• mass number = nucleon number = number of protons + neutrons

• in an atom number of protons = number of electrons, but in an ion there is a different number of electrons to protons. to work out electrons in an ion:

  • work out how many electrons an atom of the element would have (same as proton number)

  • work out how many electrons have been lost or gained (using charge, -ve means electrons gained, +ve means electrons lost)

  • calculate number of electrons in atom plus electrons gained or minus electrons lost

explain the formation of ions in ionic compounds from their atoms, limited to compounds of elements in groups 1,2,6 and 7

• ions produced by metals in groups 1 and 2 and by non-metals in groups 6 and 7 have the electronic structure of a noble gas (group 0)

• this means group 1 metals will lose 1 electron and from +1 ions

• group 2 metals will lose 2 electrons and form +2 ions

• group 6 non-metals will gain 2 electrons and form 2- ions

• group 7 non-metals will gain 1 electron and form 1- ions

• remember a compound will have an overall charge of 0 so you need to balance out the + and - charges

explain the use of the endings -ide and -ate in the names of compounds

• these endings are used for the negatively charged ions in a compound

• -ide means the compound contains 2 elements (one is the non-metal -ve ion)

• -ate means the compound contains at least 3 elements, one of which is oxygen

deduce the formulae of ionic compounds (including oxides, hydroxides, halides, nitrates, carbonates and sulfates)

• oxide → involves O²- ion (eg. sodium oxide: Na₂ O)

• hydroxide → involves OH^1- ion (eg. sodium hydroxide: NaOH)

• halide → involves a -1 halidade ion (eg. sodium chloride NaCl)

• nitrate → involves NO₃^1-ion (eg. sodium nitrate: NaNO₃)

• carbonate → involvs CO₃^2- ion (eg. sodium carbonate: NaCO₃)

• sulfate -→ involves SO₄^2- ion (eg. sodium sulfate NaSO₄)

  always need to have balanced + and - charges

explain the structure of an ionic compound

• a giant structure of ions = ionic compound

• held together by strong electrostatic forces of attraction between oppositely charged ions

• the forces act in all directions in the lattice, and this is called ionic bonding

• the lattice has a regular arrangement of ions

explain how a covalent bond is formed

• covalent bonding occurs in most non-metallic elements and in compounds of non-metals

• when atoms share pairs of electrons, they form covalent bonds. these bonds between atoms are strong

what does covalent bonding result in

the formation of molecules

• covalently bonded substances may consist of small molecules eg. HCl, H₂, O₂

• some have very large molecules, such as polymers

• some have giant covalent structures (macromolecules) eg. diamond, silicon dioxide

recall the typical size (order of magnitude) of atoms and small molecules

• simple molecular substances consist of molecules in which the atoms are joined by strong covalent bonds

• therefore, atoms are smaller than small molecules

properties of ionic compounds

• ionic compounds are made up of a metal and a non-metal

• ionic compounds have regular structures (giant ionic lattices) in which there are strong electrostatic forces of attractions in all directions between oppositely charged ions

• they have high melting and boiling points, because a lot of energy is required to break the many strong bonds (electrostatic forces)

• when melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and carry current, and they do not conduct electricity as solids, because the ions are fixed and are not able to move, carrying charge with them

• often dissolve in water to from an aqueous solution

properties of simple molecular compounds

• substances that consist of small molecules are usually gases or liquids that have low boiling and melting points. they are made up of non-metal elements

• substances that consist of small molecules have weak intermolecular forces between the molecules. these are broken in boiling or melting, not the covalent bonds

  • the intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points

• substances that consist of small molecules dont conduct electricity, because small molecules do not have an overall electric charge. although, some breakdown in water to form ions which can conduct electricity

• many are insoluble in water, but some are soluble because they can form intermolecular forces with water which are stronger than those between water molecules or their own molecules already (eg. CO₂ and NH₃ are soluble)

giant covalent structures

• they are made up of non-metal elements

• substances that consist of giant covalent structures are solids with very high melting points

  • all of the atoms in these structures are linked to other atoms by strong covalent bonds

    • these bonds must be overcome to melt or boil these substances

• some giant covalent structures can conduct electricity, whereas others can’t

properties of metals

• metals consist of giant structures of atoms arranged in a regular pattern. they are always made up of just metallic elements

• the electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure

• the sharing of delocalised electrons gives rise to strong metallic bonds

• metals have giant strcutres with strong metallic bonding

  • therefore, most metals have high melting and boiling points

  • they can conduct heat and electricty because of the delocalised electrons in their structure

  • conduction depends on the ability for electrons to move throughout the metal

  • the layers of atoms in metals are able to slide over each other, so metals can be bent and shaped

  • insoluble in water - but some will react with it instead

what are graphite and diamond

allotropes of carbon and they are examples of giant covalent substances

structure of diamond

• in diamond, each carbon is joined to 4 other carbons covalently

  • it’s very hard, has a very high melting point and does not conduct electricity

structure of graphite

• in graphite, each carbon is covalently bonded to 3 other carbons, forming layers of hexagonal rings, which have no covalent bonds between the layers

  • the layers can slide over each other due to no covalent bonds between the layers, but weak intermolecular forces. meaning that graphite is soft and slippery

• one electron from each carbon atom is delocalised

  • this makes graphite similar to metals, because of its delocalised electrons

  • it can conduct electricity - unlike diamond

graphite and diamond uses

• graphite uses

  • electrodes - graphite can conduct electricity, unlike diamond

  • lubricant - weak intermolecular forces and no covalent bonds between the layers, therefore it is soft and slippery

• diamond uses

  • cutting tools - very hard, due to its rigid structure

explain the properties of fullerenes including C₆₀ and graphene in terms of their structures and bonding

• graphene

  • single layer of graphite

  • has properties that make it useful in electronics and composites

• carbon can also form fullerenes with different numbers of carbon atoms

  • molecules of carbon atoms with hollow shapes

  • they are based on hexagonal rings of carbon atoms, but they may also contain rings with five or seven carbon atoms

  • the first fullerene to be discovered was Buckminsterfullerene (C₆₀), which has a spherical shape

• carbon nanotubes

  • cylindrical fullerenes with very high length to diameter ratios

  • their propertieis make them useful for nanotechnology, electronics and materials

• examples of uses

  • they can be used as lubricants, to deliver drugs in the body and catalysts

  • nanotubes can be used for reinforcing materials, for esxample tennis rackets

describe, using poly(ethene) as the example, that simple polymers consist of large molecules containing chains of carbon atoms

• polymers have very large molecules

• atoms in the polymer molecules are linked to other atoms by strong covalent bonds

• intermolecular forces between polymer molecules are relatively strong and so these substances are solids at room temperature

explain the properties of metals

• malleable - the layers of atoms in metals are able to slide over each other

• can conduct electricity - delocalised electrons can move

describe the limitations of particular representations and models, to include dot and cross, ball and stick models and two and three dimensional representations

• main limitation is that it applies really well only to the small class of solids composed of group 1 and 2 elements with highly electronegative elements such as the halogens

• in covalent molecular, the dot-cross diagrams don’t express the relative attraction of shared electrons due to electronegativity

• 2d diagrams don’t show 3d arrnagement of atoms, and 3d diagrams don’t show the share or transfer of electrons

describe most metals

as shiny solids which have high melting points, high density and are good conductors of electricity whereas most non-metals have low boiling points and are poor conductors of electricity

calculate relative formula mass given relative atomic mass

• relative formula mass (Mr) of a compound: sum of the relative atomic masses of the atoms in the numbers shown in the formula

calculate the formulae of simple compounds from reacting masses

• to calculate formula from reacting masses

  a. work out moles of each using: moles = mass/molar mass

  b. work out ratio of moles (dividing the moles by the smaller one)

  c. times the ratio so that you get the smallest whole numbers possible

  d. find the formula by timesing each element by their number in the ratio (remember to use little numbers not a big number at the front)

• this is an empirical formula because it shows the simplest ratio of the number of atoms of different types of elements in a compound

deduce: the empirical formula of a compound from the formula of its molecule, and the molecular formula of a compound from its empirical formula and its relative molecular mass

• empirical formula from the formula of molecule:

  • if you have a common multiple eg Fe₂ O₄, the empiraical formula is the simplest whole number ratio, which would be FeO₂

  • if there is no common multiple, you already have the empirical formula

• molecular formula from empirical formula and relative molecular mass

  • find relative molecular mass of the empirical formula

  • divide relative molecular mass of compound by that of the empirical formula

  • multiply the number of each type of atom in the empirical formula by this number

  • eg. if the answer was 2 and the empirical formula was Fe₂ O₃ then the molceular formula would be the empirical formula x 2 = Fe₄O₆

describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide

• weigh some pure magnesium

• heat magnesium to burning in a crucible to from magnesium oxide as the magnesium will react with the oxygen in the air

• weigh the mass of the magnesium oxide

• known quantities: mass of magnesium used & mass of magnesium oxide produced

• required calculations:

  • mass oxygen = mass magnesium oxide - mass magnesium

  • moles magnesium = mass magnesium/molar mass magnesium

  • calculate ratio of moles of magnesium to moles of oxygen

  • use ratio to form empirical formula

explain the law of conservation of mass applied to: a closed system including a precipitation reaction in a closed flask and a non-enclosed system including a reaction in an open flask that takes in or gives out a gas

• law of conservation on mass: no atoms are lost or made during a chemical reaction so the mass of the produces = the mass of the reactants

  • therefore, chemical reactions can be represented by symbol equations, which are balanced in terms of the numbers of atoms of each element involved on both sides of the equation

• with a precipitation reaction - precipitate that forms is insoluble and is a solid, as all the reactants and products remain in the sealed reaction container then it is easy to show that the total mass in unchanged

• does not hold for a reaction in an open flask that takes in or gives out a gas, since mass will change from what was at the start of the reaction as some mass is lost when the gas is given off

calculate masses of reactants and products from balanced equations, given the mass of one substance

• find moles of that one substance: moles = mass/molar mass

• use balancing numbers to find the moles of desired reactant or product (eg. if you had the equation: 2NaOH + Mg → Mg(OH)₂ + 2Na, if you had 2 moles of Mg, you would form 2 ×2 = 4 moles of Na)

• mass = moles x molar mass (of the reactant/product) to find mass

calculate the concentration of solutions in gdm^-3

• concentration of a solution can be measured in mass per given volume of solution eg. grams per dm³ (g/dm³)

• to calculate concentration of a solution use the equation

  concentration (gdm^-3) = mass of solute (g)/volume (dm³)

• to calculate mass of solute in a given volume of a known concentration use the equation: mass = conc x vol

how is one mole of particle of a substance defined

• the number of atoms, molecules or ions in one moe of a given subsance is the Avogadra constant: 6.02 × 10²³ per mole

• the mass of one mole of particles is the ‘relative particle mass’ in grams

calculate the number of: moles of particles of a substance in a given mass of that substance and vice versa, particles of a substance in a given number of moles of that substance and vice versa and particles of a substance in a given mass of that substance and vice versa

• chemical amounts are measured in moles. the symbol or the unit of mole is mol.

• the mass of one mole of a substance in grams is numerically equal to its relative formula mass

  • for example, the Ar of Iron is 56, so one mole of iron weighs 56g

  • the Mr of nitrogen gas (N₂) is 28 (2 × 14), so one mole is 28g

• one mole of a substance contains the same number of the stated particles, atoms, molecules or ions as one mole of nay other substance

• you can convert between moles and grams by using this equation: moles = mass (g)/relative atomic mass

  • eg how many moles are there in 42g of carbon?

    • moles = mass/Mr = 42/12 = 3.5 moles

• the number of particles, atoms, molecules or ions in a mole of a given substance is the Avogadro constant: 6.02 × 10²³ per mole

  • this means the number of particles in a given number of moles of a substance = moles x avogadro’s constant

explain why, in a reaction, the mass of product formed is controlled by the mass of the reactant which is not in excess

• in a chemical reaction with 2 or more reactants you will often use one in excess to ensure that all of the other reactant is used

  • the reactant that is used up / not in excess is called the limiting reactant since it limits the amount of products

• if a limiting reagent is used, the amount reactant in excess that actually reacts is limited to the exact amount that reacts with the amount of limiting reagent you have, so you need to use the moles/mass of the limiting reagent for any calculations

deduce the stoichiometry of a reaction from the masses of the reactants and the products

• stoichiometry refers to the balancing numbers in front of compounds/elements in reaction equations

• balancing numbers in a symbol equation can be calculated from the masses of reactant and products:

• eg. for the reaction: Cu + O₂ → CuO (not balanced), 127g Cu reacts, 32g o oxygen react an d 159g of CuO are formed. work out the balanced equation using the masses given:

  • moles: (moles = mass/Mr)

    Cu: moles = 127/63.5 = 2

    O₂: moles = 32/(16 × 2) = 32/32 = 1

  • therefore you have a ratio of 2:1:2 for Cu:O₂:CuO, making the overall balanced equation 2Cu + O₂ → 2CuO