module 2- foundations in chem Flashcards
Define isotopes.
Isotopes - atoms of the same element with different numbers of neutrons and different masses.
Define relative isotopic mass.
Relative isotopic mass - the mass of an atom of an isotope compared with 1/12 of the mass of an atom of C-12.
Define relative atomic mass.
Relative atomic mass - the weighted mean mass of an atom of an element compared to 1/12 of the mass of an atom of C-12.
What is relative mass and why do we use it to weigh atoms?
Relative mass - comparing the mass of atoms to others. Used because atoms are too small to see and have a mass much smaller than most other things.
What is the international standard to which atomic masses are compared?
International standard - carbon 12 isotope.
Why was carbon-12 chosen as the standard for atomic masses?
Can reliably be obtained in an isotopically pure state and is unreactive.
How do you calculate relative atomic mass?
Found experimentally using a mass spectrometer.
How do we determine relative atomic mass using a mass spectrometer?
Measures masses by measuring how quickly positive ions of an element sample travel; heavier ones move more slowly.
Define relative molecular mass.
Molecular mass - for simple molecules.
Define relative formula mass.
Formula mass - compounds with giant structures (ions, carbon, sulfur, etc.).
Define first ionisation energy.
First ionisation energy - energy required to remove one electron from each atom in one mole of gaseous atoms of an element to form one mole of gaseous 1+ ions.
What factors affect ionisation energy?
- Atomic radius
- Nuclear charge
- Electron shielding
Explain the trend in first ionisation energy down a group.
Atomic radius increases significantly, shielding increases, and these factors outweigh the increase in nuclear charge.
Explain the trend in first ionisation energies across periods 2 and 3.
Atomic charge increases, shielding stays equal, atomic radius decreases so general trend is that FIE increases across a period.
Why does Be have a lower FIE than B?
Electron removed from Be comes from 2s subshell whereas that of B comes from 2p subshell; 2p electrons are higher energy.
Why does O have a lower FIE than N?
In O, there is a pair of electrons in one of the 2p orbitals, while in N, all 2p orbitals contain 1 electron each.
Define atomic orbital.
Atomic orbital - region around the nucleus that can hold up to two electrons with opposite spins.
What are the shapes of s- and p-orbitals?
- s-orbital - spherical
- p-orbital - dumbbell shaped
What is the order of filling up orbitals?
1 electron goes in each orbital before any electrons pair up in any orbital.
What is the order of filling sub shells?
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p.
How does the order of filling orbitals differ from the order in which electrons are lost during ionisation?
Electrons are always lost from the orbital furthest from the nucleus, regardless of energy level.
What do roman numerals denote in formulae?
Oxidation number of the metal when written after metal cations; oxidation number of the element that is not oxygen when written after polyatomic anions.
Define oxidation.
Loss of electrons.
Define reduction.
Gain of electrons.
What are oxidising and reducing agents?
- Oxidising agent - a reagent that oxidises another species.
- Reducing agent - a reagent that reduces another species.
Summarise atomic structure in terms of protons, neutrons, and electrons.
Consist of a positive nucleus (protons and neutrons) and a cloud of negative electrons surrounding it.
What is the charge of group 1 metal ions?
Always form 1+ ions.
What is the charge of group 2 metal ions?
Always form 2+ ions.
What is the charge of group 3 metal ions?
Always form 3+ ions.
What is the charge of group 6 non-metal ions?
Always form 2- ions.
What is the charge of group 7 non-metal ions?
Always form 1- ions.
What is the charge of a nitride ion?
N3-.
What is the formula of a Silver ion?
Ag+.
What is the formula of a Copper(II) ion?
Cu2+.
What is the formula of an Ammonium ion?
NH4+.
What is the formula of an Iron(II) ion?
Fe2+.
What is the formula of an Iron(III) ion?
Fe3+.
What is the formula of a zinc ion?
Zn2+.
What is the formula of a Lead(II) ion?
Pb2+.
What are the formulae of the following polyatomic ions: carbonate, sulfate, nitrate, phosphate, hydrogencarbonate, hydroxide ions?
- Carbonate - CO3^2-
- Sulfate - SO4^2-
- Nitrate - NO3^-
- Phosphate - PO4^3-
- Hydrogen carbonate - HCO3^-
- Hydroxide - OH^-
What is an acid?
Acid - proton (H+) donor.
How do acids behave in water solution?
Acids dissociate in water to release H+ ions.
What are strong acids? Give examples.
Strong acid - Acid that dissociates fully in solution. Examples: HNO3, H2SO4, H3PO4, HBr, HClO4.
What are weak acids? Give examples.
Weak acid - Acid that dissociates partially in solution. Examples: carboxylic acids.
Explain the difference between strong/weak acids and concentrated/dilute acids.
Strong and weak refer to how an acid dissociates in solution; concentrated and dilute refer to how many particles are in a given volume.
What is a base?
Base - proton (H+) acceptor.
What is an alkali?
Alkali - soluble base that dissolves in water to produce OH- ions.
Give examples of common alkalis and how they produce OH- ions.
- Metal hydroxides - Group I and II metal hydroxides.
- Metal oxides - Group I and II metal oxides react with water.
- Ammonia - reacts with water to produce a solution containing hydroxide ions.
What is neutralisation? Provide the ionic equation.
Neutralisation - reaction where an acid and a base react to form salt and water. Ionic equation: H+(aq) + OH-(aq) –> H2O(l).
What are salts?
Salts - ionic compounds formed when the H+ ions in an acid are replaced by positive metal or ammonium ions.
What are the 3 ways of making salts?
- Making a soluble salt from an insoluble base - filtration or crystallisation.
- Making a soluble salt from a soluble base - titration.
- Making an insoluble salt from a soluble base - precipitation.
What reagent is used to test for halide ions and why?
Aqueous silver nitrate is used because silver halides are all insoluble and form precipitates of different colors.
What is the method for testing for halide ions?
Add aqueous silver nitrate to the unknown solution; if halide ions are present, a precipitate will form.
How do you test for carbonates?
Add dilute nitric acid to the test substance; effervescence indicates carbonate ions are present.
How do you carry out a test for sulfate ions?
Add aqueous solution containing Ba ions to the unknown solution; if sulfates are present, a white precipitate forms.
What reagent is used for sulfate ion test and why?
Barium ions are used because most sulfates are soluble except BaSO4.
How do you carry out a test for ammonium ions?
Add NaOH solution to the unknown solution, heat, and hold damp red litmus paper over the solution; if it turns blue, ammonium ions are present.
Explain the ammonium ion test.
Ammonium salts react with NaOH to form ammonia gas; damp litmus paper turns blue due to OH- ions produced.
What is the meaning of amount of substance?
Amount of substance - quantity whose unit is the mole, used for counting species such as atoms, ions, or molecules.
Define the mole.
Mole - amount of any substance containing as many elementary particles as there are carbon atoms in exactly 12g of carbon-12.
Define the Avogadro constant.
Avogadro constant - number of atoms per mole of the carbon-12 isotope (6.02 * 10^23).
Define molar mass.
Molar mass - the mass per mole of a substance in gmol-1.
What is the relationship between amount of substance, mass of substance, and molar mass?
Amount of substance in mol, mass in g, molar mass in gmol-1.
Define empirical formula.
Empirical formula - formula that shows the simplest whole number ratio of atoms of each element present in a compound.
Define molecular formula.
Molecular formula - chemical formula that indicates the kind of atoms and the number of each kind in a molecule or compound.
Define anhydrous.
Anhydrous - containing no water molecules.
Define hydrated.
Hydrated - a crystalline compound containing water molecules.
Define water of crystallisation.
Water of crystallisation - water molecules that are bonded into a crystalline structure of a compound.
What are the types of uncertainties?
- Procedural - to do with the experimental method (non-quantifiable).
- Measurement - to do with measurements (quantifiable).
What is the definition of concentration?
Concentration - the amount of solute dissolved in a given quantity of solvent.
What is meant by a standard solution?
Standard solution - solution where we know its concentration.
Write down the formula that relates concentration, number of moles, and volume.
Concentration (mol/dm^3) = Number of moles / Volume (dm^3).
What is meant by weak, concentrated, strong, and dilute?
Weak and strong refer to how an acid dissociates; concentrated and dilute refer to the number of particles in a given volume.
Define concordant results in a titration.
Concordant results - results which are within ±0.1cm3 of each other.
Which results should be used to generate the mean titration value for calculations?
Mean titre - include concordant results, NOT including rough titre.
What should you do if there are 3 or more concordant results in accurate titrations?
By convention, all 3 should be included in the mean.
Define molar gas volume and state what it is at room temperature & pressure.
Molar gas volume - volume per mol of gas molecules at a stated temp and pressure; at RTP it is 24 dm^3.
What is the equation linking molar gas volume and number of moles?
Volume = Number of moles × Molar gas volume.
State the ideal gas equation & units.
PV = nRT, where P is pressure in Pa, V is volume in m^3, n is number of moles, R is the gas constant (8.314 J/(mol·K)), and T is temperature in Kelvin.
How do you convert from degrees Celsius to Kelvin?
K = °C + 273.15.
How do you convert from atm to Pa?
1 atm = approximately 100,000 Pa.
What are the three requirements of a primary standard?
- Must be readily available in a very pure state
- Must be stable
- Must not absorb water vapour
What is an ideal gas?
Ideal gas - a hypothetical gas whose molecules occupy negligible space and have no interactions, obeying the gas laws exactly.
When is the ideal gas law the most accurate?
- High temperatures
- Very small gaseous particles
- Low pressures
Give a method for making a standard solution for a titration.
Dissolve a known mass of solute in a volumetric flask and dilute to the mark with solvent.
What are the four characteristics of an ideal gas?
Be readily available in a very pure state, Must be stable, Must not absorb water vapour, Should have a high Mr
Define ideal gas
A hypothetical gas whose molecules occupy negligible space and have no interactions, and which consequently obeys the gas laws exactly
When is the ideal gas law the most accurate?
At high temperatures, with very small gaseous particles, and at low pressures
List the steps for making a standard solution for a titration
- Weigh and record the mass of the bottle
- Tip solid into 100cm3 beaker and reweigh empty bottle
- Rinse volumetric flask and its stopper with deionised water
- Add around 50cm3 deionised water to the beaker and stir until dissolved
- Pour solution through a glass funnel into a 100cm3 volumetric flask
- Rinse beaker and glass rod with deionised water and pour washings into the flask
- Top up volumetric flask with deionised water to the line
- Insert stopper and invert flask to mix the solution
Why is the bottle reweighed after tipping solid into a beaker?
To account for particles of solid which were not transferred
Why do we use deionised water rather than tap water in making standard solution?
Tap water contains dissolved substances which could affect the titration
Why do we wash out the beaker and transfer the washings?
To avoid errors due to some of the solution remaining in the beaker
Why do we add deionised water dropwise when filling the flask to the calibration line?
To avoid overfilling the flask
How do we know when the correct amount of water was added to the volumetric flask?
The bottom of the meniscus should be on the calibration line when viewed at eye level
Why do we invert the flask several times?
To get a uniform/homogenous solution
What should be done if the volumetric flask was overfilled?
Discard solution and start again - concentration will be unknown
Why should we take the bottle off the balance before adding the solid?
To prevent the solid damaging the balance pan if spilt
Why should a precise balance be used when making a standard solution?
Allows concentration to be known to a high accuracy and precision
Describe the steps to prepare a burette for a titration
- Rinse burette with standard HCl solution
- Drain into waste beaker and fill with acid
- Remove air pocket below the tap
- Refill any displaced HCl
- Ensure the funnel is removed from the top of burette
Describe a correct method to transfer an unknown solution from a volumetric flask to the reaction vessel
- Rinse conical flask with deionised water
- Half fill a clean 100cm3 beaker with the unknown solution
- Rinse volumetric pipette with the unknown solution
- Measure accurately required volume by allowing meniscus to sit on the calibration line
- Allow adequate time for pipette to drain, do not force last drop out
Describe a method to carry out the titration
- Add 3-4 drops methyl orange indicator to flask
- Take initial reading of burette
- Carry out rough titration and record volume of acid added
- Repeat titration more accurately, adding 1 drop at a time as endpoint is approached
- Carry out further titrations until 2 results agree to ±0.1cm3
Describe a method for finding the molar volume of a gas
- Check for gas syringe leaks
- Find mass of magnesium
- Place magnesium strip into a flask with HCl
- Allow reaction to start and record final volume of gas
Describe a method for determining the formula of a hydrated salt
- Weigh crucible and record mass
- Place hydrated salt in crucible and weigh
- Heat until it turns anhydrous and reweigh
- Calculate the number of water molecules per unit of salt
What is ionic bonding?
Electrostatic attraction between positive and negative ions
Explain how ionic bonding occurs
Metal forms a positive ion by losing electrons; non-metal forms a negative ion by gaining electrons
Describe the structure of ionic compounds
Each ion is surrounded by oppositely charged ions, forming a giant ionic lattice
State and explain the melting point and boiling point of ionic compounds
- Solids at RTP due to strong electrostatic forces of attraction
- High melting and boiling points due to large energy needed to overcome ionic bonds
What factors affect the melting point and boiling point of ionic compounds?
- Ionic charge
- Size of ions
State and explain the solubility of ionic compounds
- Many are soluble in polar solvents
- Insoluble if ionic attraction is too strong for solvent to break down the lattice
What affects the solubility of ionic compounds?
Ionic charge - higher charge means stronger attractions in giant ionic lattice
State and explain the electrical conductivity of ionic compounds
- Not conductive when solid due to fixed position of ions
- Conductive when liquid or aqueous as ions are free to move
What is covalent bonding?
Strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
Where does covalent bonding occur?
- Non-metallic elements
- Compounds of non-metallic elements
- Polyatomic ions
What is a dative covalent bond?
A covalent bond where the shared pair of electrons has been supplied by one of the bonding atoms only
How many covalent bonds do carbon, nitrogen, oxygen, and hydrogen form?
Carbon - 4, Nitrogen - 3, Oxygen - 2, Hydrogen - 1
How many covalent bonds does boron form?
Three covalent bonds in BF3 and does not follow octet rule
What is average bond enthalpy?
Measurement of covalent bond strength
State VSEPR theory
Electron pairs repel each other and adopt positions separated by the largest possible angle
What is a lone pair?
Pair of electrons in the outer shell NOT used in bonding
How do lone pairs affect bond angles?
Lone pairs reduce bond angles by about 2.5˚ per lone pair
What is electronegativity?
Ability of an atom to attract the bonding electrons in a covalent bond
What is the most electronegative atom?
Fluorine (F) - 4.0
What is a non-polar bond?
Bonded electron pair is shared equally between bonded atoms
What is a polar bond?
Bonded electron pair is shared unequally between bonded atoms
What happens when electronegativity difference between bonded atoms is large?
One bonded atom will have much greater attraction to shared pair of electrons than the other
What is a dipole?
Separation of opposite charges
What are the conditions for a molecule to be polar?
- Must contain polar bonds
- Charges must not be symmetrical
What are the types of intermolecular forces?
- Induced dipole-dipole interactions (London forces)
- Permanent dipole-dipole interactions
- Hydrogen bonding
Explain how induced dipole-dipole interactions happen.
Electrons in atoms/molecules are constantly moving, leading to temporary dipoles that induce further dipoles in neighboring atoms.
What are permanent dipole-dipole interactions?
Intermolecular forces that arise from interactions between permanent dipoles in different polar molecules
What is the main factor that affects the strength of permanent dipole-dipole interactions?
Electronegativity difference
What are permanent dipole-dipole interactions?
Intermolecular forces that arise from interactions between permanent dipoles in different polar molecules
Slightly +ve and slightly -ve parts of different molecules are attracted to one another.
What type of intermolecular forces do F2 and HCl have?
F2 has London forces; HCl has permanent dipole-dipole interactions
HCl is predicted to have a higher boiling point due to stronger intermolecular forces.
What is the main factor that affects the strength of permanent dipole-dipole interactions?
Electronegativity difference
A larger difference leads to a larger dipole, making interactions stronger.
What are hydrogen bonds?
Strong permanent dipole-dipole attraction between a δ+ hydrogen on one molecule and a lone pair on a δ- oxygen, fluorine or nitrogen on a different molecule.
What is the order of strength of intermolecular forces from weakest to strongest?
London forces, Permanent dipole-dipole interactions, Hydrogen bonds.
What are the anomalous properties of water?
Solid is less dense than the liquid, relatively high melting point/boiling point, relatively high surface tension and viscosity.
Why is ice less dense than water?
Hydrogen bonds hold water molecules slightly apart, forming an open structure.
Explain why water has a relatively high melting point and boiling point.
Hydrogen bonds require a large quantity of energy to break, resulting in higher melting and boiling points than expected.
What is a simple molecular structure?
3D structure of molecules held together by weak intermolecular forces (hydrogen bonds, van der Waals forces) with strong covalent bonds between atoms.
What is the melting point and boiling point of simple molecular substances?
Low melting and boiling points due to weak intermolecular forces being broken.
What is the solubility of non-polar simple molecular structures?
Soluble in non-polar solvents, insoluble in polar solvents.
What does the solubility of polar simple molecular substances depend on?
Strength of the dipole formed between solute and solvent.
What is the electrical conductivity of simple molecular substances?
Non-conductors due to no charged particles that can move.
What is a giant metallic lattice?
Regular 3D structure of cations and delocalised electrons held together by strong electrostatic attractions.
What is metallic bonding?
Strong electrostatic attraction between cations and delocalised electrons.
What is a giant covalent lattice?
3D structure of atoms held together by strong covalent bonds.
What are the elements that form giant covalent lattices?
Carbon (diamond, graphite), Silicon, Boron.
What is the melting point and boiling point of giant covalent structures?
High melting and boiling points due to strong covalent bonds requiring a large quantity of energy to break.
What is the solubility of giant covalent lattices?
Insoluble in almost all solvents due to strong covalent bonds.
What is the general electrical conductivity of giant covalent lattices?
In diamond and silicon, non-conductive; graphite and graphene can conduct electricity due to delocalised electrons.
Why is diamond hard and graphite soft?
Diamond has a tetrahedral structure with strong covalent bonds; graphite has weak intermolecular forces between layers.
Describe the structure of diamond.
Tetrahedral arrangement of atoms with 109.5-degree bond angles; each atom forms 4 covalent bonds.
Describe the structure of graphite.
Each carbon atom is joined to three others at 120-degree angles; layers with hexagonal arrangements have weak London forces between them.
Describe the structure of graphene.
Single layer of hexagonally arranged carbon atoms linked by strong covalent bonds.
What is the solubility of giant metallic lattices?
Insoluble; interactions with polar solvents lead to chemical reactions rather than dissolving.
What is the melting point and boiling point of giant metallic lattices?
High melting and boiling points due to strong metallic bonds.
What does metallic bond strength depend on?
Charge of metal ion; higher charge leads to more delocalised electrons and stronger electrostatic attraction.
Why are metals malleable and ductile?
Layers of ions can easily slide over each other.
How do alloys make metals harder?
Different sizes of metal ions hinder the sliding of layers.
What is the electrical conductivity of giant metallic lattices?
Conductive when solid due to delocalised electrons; conductive when liquid due to both cations and delocalised electrons moving freely.
State number of bonded pairs, lone pairs, shape, bond angle and draw the 3D diagram of BF3.
Draw 3D diagram of NO3-
Explain why diamond is hard and graphite is soft.
