Module 2
Chapter 2: Atoms, ions and compounds
Atomic structure
Nuclear model
● Has nucleus made of protons and neutrons, electrons are arranged around nucleus in shells
● Atomic number: number of protons in an element
● Mass number: number of protons + neutrons in an element
● Relative masses are used instead of actual masses
● Proton has virtually same mass as neutron
1
● Electron has 1836
th the mass of proton
● Therefore nearly all of an atom’s mass is in the nucleus
● Proton has equal positive charge as electron has negative charge
● Therefore atoms have the same number of protons as electrons
● Neutron has no charge - holds nucleus together despite of electrostatic repulsion of protons
● Therefore as nucleus gets bigger, more neutrons are needed (typically same or more than
number of protons)
Isotopes
● Atom of same element (has same number of protons and electrons) but different number of
neutrons and therefore mass
12 12
● All atoms/isotopes are written in 3 ways: “carbon-12”, 6
𝐶 or 𝐶
● Number of neutrons has no effect on chemical reactions but may slightly change physical
properties eg. more neutrons means higher mass so higher melting point
Ions
● Same number of protons but different number of electrons
● Cations: positive ion - more protons than electrons
● Anions: negative ion - more electrons than protons
35 −
● Written like 17
𝐶𝑙
Relative mass
● Strong nuclear force holding protons and neutrons together causes loss of some mass - mass
lost is “mass defect”
● Standard isotope needed to base all atomic masses - carbon-12 isotope is used
● Carbon-12 has exactly 12 atomic mass units (12u) so on this scale, 1u is a proton or neutron
1
● Relative isotopic mass (RIM): mass of an isotope relative to 12
th of mass of carbon-12 atom
1
● Relative atomic mass (RAM): weighted mean mass of an atom of an element relative to 12
th
of mass of carbon-12 atom - uses percentage abundance and RIM of each isotope
(% 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 1×𝑀𝑟 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 1)+(% 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 2×𝑀𝑟 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 2) ...
● 𝐴𝑟 = 100
● Percentage abundances found experimentally using mass spectrometer - process:
1. Sample placed in the mass spectrometer
2. Sample vaporised then ionised to form positive ions
3. Ions accelerated - heavier move slowly and harder to deflect than light so ions of
each isotope are separated
4. Ions detected on mass spectrum as mass-to-charge ratio m/z - more ions means
larger signal so higher peak
5. For ion with 1+ charge ratio is same is RIM, recorded on x-axis
,Formulae and equations
Forming ions
● Elements to left of group 4 lose electrons to form cations
● Elements to right of group 4 gain electrons to form anions
● Typically transition metals form ions with different charges - shown with roman numerals
Compounds
● Binary compound: contains 2 elements - in naming, metal comes first and ending of second
element changes to “-ide”
● Polyatomic ions: ion containing atoms of 1+ elements (necessary to know)
1+ 1- 2- 3-
Ammonium (NH4+) Hydroxide (OH-) Carbonate (CO32-) Phosphate (PO43-)
Nitrate (NO3-) Sulfate (SO42-)
Nitrite (NO2-) Sulfite (SO32-)
Hydrogencarbonate Dichromate(VI)
(HCO3-) (Cr2O72-)
Manganate(VII) /
permanganate (MnO4-)
● Diatomic molecules: 2 atoms bonded together to form molecule or compound
Equations
● Ionic equation: shows reacting ions
● Half equation: shows oxidation or reduction of ion
State symbols
● Gas: (g)
● Liquid: (l)
● Solid: (s)
● Aqueous: (aq)
3: Amount of substance
Amount and the mole
● 1 mole: 6.02 x 1023 particles
● Avogadro constant: 6.02 x 1023 particles/mol of carbon-12
● Example: 12g of carbon-12 has 1 mole - 6.02 x 1023 carbon atoms
● 1 mole is dependent on particle:
○ 1 mol of H: 1 mol of hydrogen atoms
○ 1 mol of H2: 1 mol of hydrogen molecules
● Molar mass links moles with mass for any chemical substance eg. M(C) = 12.0g mol-1
𝑚𝑎𝑠𝑠
● 𝑚𝑜𝑙𝑒𝑠 = 𝑀𝑟
● Moles and amount are the same
Formulae
● Molecular formula: number of atoms of each element in molecule
● Empirical formula: simplest whole-number ratio of atoms of each element in compound -
important for substances that don’t exist as molecules (giant crystalline structures eg. ionic
compounds)
● Relative molecular mass: mass of molecule relative to mass of atom of carbon-12
● Relative formula mass: mass of formula unit relative to mass of atom of carbon-12
,Hydrated salts
● Many coloured crystals are hydrated (water molecules part of crystalline structure) - water is
known as “water of crystallisation”
● When hydrated salt is heat, bonds holding water within crystal are broken and driven off,
leaving white anhydrous salt
● Very difficult to remove all traces of water
● Example:
○ Hydrated copper(II) sulfate (blue) → anhydrous copper(II) sulfate (white) + water
𝐶𝑢𝑆𝑂4 • 5𝐻2𝑂 (𝑠) → 𝐶𝑢𝑆𝑂4 (𝑠) + 5𝐻2𝑂 (𝑙)
○ Water of crystallisation is shown separately using •
Assumptions
● Made in experimental formulas that mean real experiments may not work out the same way
● All of the water is lost:
○ Can be fairly sure when water has been mostly removed
○ Can’t see water left inside the crystals (only see surface of crystals)
○ If colour of hydrated and anhydrous is similar, can be difficult
○ Heat to constant mass - reheat crystals repeatedly until mass no longer changes
(suggests all water is removed)
● No further decomposition:
○ Many salts decompose further when heated or by too much
○ Can change to different colour or not change colour
Moles and volumes
● 1cm3 = 1ml
● 1dm3 = 1000cm3 = 1000ml = 1l
3 𝑚𝑜𝑙𝑒𝑠
● 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (𝑚𝑜𝑙/𝑑𝑚 ) = 3
𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 )
3 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔)
● 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (𝑔/𝑑𝑚 ) = 3
𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 )
● Standard solution: solution of known concentration - prepared by dissolving exact mass of
solute in solvent and making solution to an exact volume
Gases
● At same temperature and pressure, equal volumes of different gases have same number of
molecules - measuring volume indirectly measures number/amount (in mol) of gas molecules
● Molar gas volume: volume per mole of gas molecules at stated temperature and pressure
● Many experiments carried out at room temperature and pressure (RTP)
● RTP: 20oC, 101kPa (1atm), 1 mole of gas has volume approx. 24.0dm3/mol
3 3
● 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 ) = 𝑚𝑜𝑙𝑒𝑠 × 𝑚𝑜𝑙𝑎𝑟 𝑔𝑎𝑠 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 /𝑚𝑜𝑙)
● Ideal gas equation factors temperature and pressure as well:
3
𝑝𝑟𝑒𝑠𝑠𝑢𝑟𝑒 (𝑃𝑎) × 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 ) = 𝑚𝑜𝑙𝑒𝑠 × 𝑔𝑎𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 × 𝑡𝑒𝑚𝑝𝑒𝑟𝑎𝑡𝑢𝑟𝑒 (𝐾) (𝑝𝑉 = 𝑛𝑅𝑇)
● Gas constant: 8.314J/mol/K
Stoichiometry
● Ratio of reactants and products in moles
● Example:
2𝐻2 + 𝑂2 → 2𝐻2𝑂
2 𝑚𝑜𝑙 + 1 𝑚𝑜𝑙 → 2 𝑚𝑜𝑙
Reactions
Percentage yield
● How much product is actually made out of maximum possible amount
● Actual yield is usually less than theoretical yield:
, ○ Reaction doesn’t go to completion
○ Other reactions (side reactions) may have taken place
○ Purification of the product may result in loss of some product
𝑎𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
● 𝑝𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑦𝑖𝑒𝑙𝑑 = 𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
× 100
Limiting reagent
● Reactant used up completely first and stops reaction
● Has lowest moles of all reactants
Atom economy
● How many of the reactant atoms are part of the useful products
● Important for reactions to have high atom economy:
○ Produces large proportion of desired products and few waste products
○ Important for sustainability - makes best use of resources
𝑠𝑢𝑚 𝑜𝑓 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑑𝑒𝑠𝑖𝑟𝑒𝑑 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
● 𝑎𝑡𝑜𝑚 𝑒𝑐𝑜𝑛𝑜𝑚𝑦 = 𝑠𝑢𝑚 𝑜𝑓 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑎𝑙𝑙 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
× 100
Sustainability
● Percentage yield and atom economy aren’t only things that matter
● Accessibility of raw materials
● Waste products’ usability and environmental impact eg. greenhouse gases
● Energy needed for reaction
Conversions
● cm3 x 10-6 = m3
● dm3 x 10-3 = m3
● oC + 273 = K
● kPa x 103 = Pa
4: Acids and redox
Solutions
Acids
● When dissolved in water, releases H+ ions into solution
● pH less than 7 - in logarithmic scale
Bases and alkalis
● Base: neutralises acid to form salt
● Alkali: base that dissolves in water, releases OH- ions into solution
Neutralisation
● Acid’s H+ ions react with alkali’s OH- ions to form salt and neutral water
● Hydrogen in acid is replaced by metal or ammonium ions to form salt
● Salt: chemical compound of positively and negatively charged ions
● Salt naming: alkali acid
● Ionic equation will show neutralisation of H+ ions by OH- ions to form neutral H2O
● Acid + metal oxide (base) (s) → salt + water
● Acid + metal hydroxide (alkali) (aq) → salt + water
● Acid + carbonate (s) → salt + water + carbon dioxide
● Acid + metal (s) → salt + hydrogen (not neutralisation - no water formed)
Titrations
Purpose
● Accurately measure volume of one solution that reacts exactly with another solution
● Can be used to:
○ Find concentration of solution
○ Identify unknown chemicals
Chapter 2: Atoms, ions and compounds
Atomic structure
Nuclear model
● Has nucleus made of protons and neutrons, electrons are arranged around nucleus in shells
● Atomic number: number of protons in an element
● Mass number: number of protons + neutrons in an element
● Relative masses are used instead of actual masses
● Proton has virtually same mass as neutron
1
● Electron has 1836
th the mass of proton
● Therefore nearly all of an atom’s mass is in the nucleus
● Proton has equal positive charge as electron has negative charge
● Therefore atoms have the same number of protons as electrons
● Neutron has no charge - holds nucleus together despite of electrostatic repulsion of protons
● Therefore as nucleus gets bigger, more neutrons are needed (typically same or more than
number of protons)
Isotopes
● Atom of same element (has same number of protons and electrons) but different number of
neutrons and therefore mass
12 12
● All atoms/isotopes are written in 3 ways: “carbon-12”, 6
𝐶 or 𝐶
● Number of neutrons has no effect on chemical reactions but may slightly change physical
properties eg. more neutrons means higher mass so higher melting point
Ions
● Same number of protons but different number of electrons
● Cations: positive ion - more protons than electrons
● Anions: negative ion - more electrons than protons
35 −
● Written like 17
𝐶𝑙
Relative mass
● Strong nuclear force holding protons and neutrons together causes loss of some mass - mass
lost is “mass defect”
● Standard isotope needed to base all atomic masses - carbon-12 isotope is used
● Carbon-12 has exactly 12 atomic mass units (12u) so on this scale, 1u is a proton or neutron
1
● Relative isotopic mass (RIM): mass of an isotope relative to 12
th of mass of carbon-12 atom
1
● Relative atomic mass (RAM): weighted mean mass of an atom of an element relative to 12
th
of mass of carbon-12 atom - uses percentage abundance and RIM of each isotope
(% 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 1×𝑀𝑟 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 1)+(% 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 2×𝑀𝑟 𝑜𝑓 𝑖𝑠𝑜𝑡𝑜𝑝𝑒 2) ...
● 𝐴𝑟 = 100
● Percentage abundances found experimentally using mass spectrometer - process:
1. Sample placed in the mass spectrometer
2. Sample vaporised then ionised to form positive ions
3. Ions accelerated - heavier move slowly and harder to deflect than light so ions of
each isotope are separated
4. Ions detected on mass spectrum as mass-to-charge ratio m/z - more ions means
larger signal so higher peak
5. For ion with 1+ charge ratio is same is RIM, recorded on x-axis
,Formulae and equations
Forming ions
● Elements to left of group 4 lose electrons to form cations
● Elements to right of group 4 gain electrons to form anions
● Typically transition metals form ions with different charges - shown with roman numerals
Compounds
● Binary compound: contains 2 elements - in naming, metal comes first and ending of second
element changes to “-ide”
● Polyatomic ions: ion containing atoms of 1+ elements (necessary to know)
1+ 1- 2- 3-
Ammonium (NH4+) Hydroxide (OH-) Carbonate (CO32-) Phosphate (PO43-)
Nitrate (NO3-) Sulfate (SO42-)
Nitrite (NO2-) Sulfite (SO32-)
Hydrogencarbonate Dichromate(VI)
(HCO3-) (Cr2O72-)
Manganate(VII) /
permanganate (MnO4-)
● Diatomic molecules: 2 atoms bonded together to form molecule or compound
Equations
● Ionic equation: shows reacting ions
● Half equation: shows oxidation or reduction of ion
State symbols
● Gas: (g)
● Liquid: (l)
● Solid: (s)
● Aqueous: (aq)
3: Amount of substance
Amount and the mole
● 1 mole: 6.02 x 1023 particles
● Avogadro constant: 6.02 x 1023 particles/mol of carbon-12
● Example: 12g of carbon-12 has 1 mole - 6.02 x 1023 carbon atoms
● 1 mole is dependent on particle:
○ 1 mol of H: 1 mol of hydrogen atoms
○ 1 mol of H2: 1 mol of hydrogen molecules
● Molar mass links moles with mass for any chemical substance eg. M(C) = 12.0g mol-1
𝑚𝑎𝑠𝑠
● 𝑚𝑜𝑙𝑒𝑠 = 𝑀𝑟
● Moles and amount are the same
Formulae
● Molecular formula: number of atoms of each element in molecule
● Empirical formula: simplest whole-number ratio of atoms of each element in compound -
important for substances that don’t exist as molecules (giant crystalline structures eg. ionic
compounds)
● Relative molecular mass: mass of molecule relative to mass of atom of carbon-12
● Relative formula mass: mass of formula unit relative to mass of atom of carbon-12
,Hydrated salts
● Many coloured crystals are hydrated (water molecules part of crystalline structure) - water is
known as “water of crystallisation”
● When hydrated salt is heat, bonds holding water within crystal are broken and driven off,
leaving white anhydrous salt
● Very difficult to remove all traces of water
● Example:
○ Hydrated copper(II) sulfate (blue) → anhydrous copper(II) sulfate (white) + water
𝐶𝑢𝑆𝑂4 • 5𝐻2𝑂 (𝑠) → 𝐶𝑢𝑆𝑂4 (𝑠) + 5𝐻2𝑂 (𝑙)
○ Water of crystallisation is shown separately using •
Assumptions
● Made in experimental formulas that mean real experiments may not work out the same way
● All of the water is lost:
○ Can be fairly sure when water has been mostly removed
○ Can’t see water left inside the crystals (only see surface of crystals)
○ If colour of hydrated and anhydrous is similar, can be difficult
○ Heat to constant mass - reheat crystals repeatedly until mass no longer changes
(suggests all water is removed)
● No further decomposition:
○ Many salts decompose further when heated or by too much
○ Can change to different colour or not change colour
Moles and volumes
● 1cm3 = 1ml
● 1dm3 = 1000cm3 = 1000ml = 1l
3 𝑚𝑜𝑙𝑒𝑠
● 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (𝑚𝑜𝑙/𝑑𝑚 ) = 3
𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 )
3 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔)
● 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (𝑔/𝑑𝑚 ) = 3
𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 )
● Standard solution: solution of known concentration - prepared by dissolving exact mass of
solute in solvent and making solution to an exact volume
Gases
● At same temperature and pressure, equal volumes of different gases have same number of
molecules - measuring volume indirectly measures number/amount (in mol) of gas molecules
● Molar gas volume: volume per mole of gas molecules at stated temperature and pressure
● Many experiments carried out at room temperature and pressure (RTP)
● RTP: 20oC, 101kPa (1atm), 1 mole of gas has volume approx. 24.0dm3/mol
3 3
● 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 ) = 𝑚𝑜𝑙𝑒𝑠 × 𝑚𝑜𝑙𝑎𝑟 𝑔𝑎𝑠 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 /𝑚𝑜𝑙)
● Ideal gas equation factors temperature and pressure as well:
3
𝑝𝑟𝑒𝑠𝑠𝑢𝑟𝑒 (𝑃𝑎) × 𝑣𝑜𝑙𝑢𝑚𝑒 (𝑑𝑚 ) = 𝑚𝑜𝑙𝑒𝑠 × 𝑔𝑎𝑠 𝑐𝑜𝑛𝑠𝑡𝑎𝑛𝑡 × 𝑡𝑒𝑚𝑝𝑒𝑟𝑎𝑡𝑢𝑟𝑒 (𝐾) (𝑝𝑉 = 𝑛𝑅𝑇)
● Gas constant: 8.314J/mol/K
Stoichiometry
● Ratio of reactants and products in moles
● Example:
2𝐻2 + 𝑂2 → 2𝐻2𝑂
2 𝑚𝑜𝑙 + 1 𝑚𝑜𝑙 → 2 𝑚𝑜𝑙
Reactions
Percentage yield
● How much product is actually made out of maximum possible amount
● Actual yield is usually less than theoretical yield:
, ○ Reaction doesn’t go to completion
○ Other reactions (side reactions) may have taken place
○ Purification of the product may result in loss of some product
𝑎𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
● 𝑝𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑦𝑖𝑒𝑙𝑑 = 𝑡ℎ𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑
× 100
Limiting reagent
● Reactant used up completely first and stops reaction
● Has lowest moles of all reactants
Atom economy
● How many of the reactant atoms are part of the useful products
● Important for reactions to have high atom economy:
○ Produces large proportion of desired products and few waste products
○ Important for sustainability - makes best use of resources
𝑠𝑢𝑚 𝑜𝑓 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑑𝑒𝑠𝑖𝑟𝑒𝑑 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
● 𝑎𝑡𝑜𝑚 𝑒𝑐𝑜𝑛𝑜𝑚𝑦 = 𝑠𝑢𝑚 𝑜𝑓 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠𝑒𝑠 𝑜𝑓 𝑎𝑙𝑙 𝑝𝑟𝑜𝑑𝑢𝑐𝑡𝑠
× 100
Sustainability
● Percentage yield and atom economy aren’t only things that matter
● Accessibility of raw materials
● Waste products’ usability and environmental impact eg. greenhouse gases
● Energy needed for reaction
Conversions
● cm3 x 10-6 = m3
● dm3 x 10-3 = m3
● oC + 273 = K
● kPa x 103 = Pa
4: Acids and redox
Solutions
Acids
● When dissolved in water, releases H+ ions into solution
● pH less than 7 - in logarithmic scale
Bases and alkalis
● Base: neutralises acid to form salt
● Alkali: base that dissolves in water, releases OH- ions into solution
Neutralisation
● Acid’s H+ ions react with alkali’s OH- ions to form salt and neutral water
● Hydrogen in acid is replaced by metal or ammonium ions to form salt
● Salt: chemical compound of positively and negatively charged ions
● Salt naming: alkali acid
● Ionic equation will show neutralisation of H+ ions by OH- ions to form neutral H2O
● Acid + metal oxide (base) (s) → salt + water
● Acid + metal hydroxide (alkali) (aq) → salt + water
● Acid + carbonate (s) → salt + water + carbon dioxide
● Acid + metal (s) → salt + hydrogen (not neutralisation - no water formed)
Titrations
Purpose
● Accurately measure volume of one solution that reacts exactly with another solution
● Can be used to:
○ Find concentration of solution
○ Identify unknown chemicals
