Module 2 Foundations in Chemistry Flashcards
The mass of an electron
1/1836 the mass of a proton
Isotopes
Atoms of the SAME element with the same number of protons but a DIFFERENT number of neutrons
Relative Isotopic Mass
The mass of an isotope relative to 1/12th the mass of an atom of carbon-12
Relative Atomic Mass
The weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of carbon-12
1 mole (mol)
The amount of substance containing 6.02 x 10^23 number of particles
Equation for the amount of substance
Moles (n) = mass (m)/Mr (M)
Number of particles (Av Constant)
Number of particles = Na (Avogardro’s constant) x moles
The empirical formula
The simplest whole number ratio of atoms of each element present in a compound
The complete combustion equation
CxHy + O2 —> 1/2yH2O + XCO2
Molecular formula
The actual number of atoms of each element in a molecule
Stoichiometry
Ratios former by the balancing number to determine the number of moles
Concentration
This gives an accurate indication of the amount, in moles, of a solute that is dissolved in a volume of the solvent
Standard solution
Solution of known concentration
How do you convert between g dm-3 to mol dm-3
Divide by the Mr
Which equation do we use to find the concentration of a solution when the volume contains a dm3 unit
Amount (n) = C x V
State the equation used when calculating the moles of a solution when our volume of solution value is in cm3
Amount (n) = C x V/1000
State what 1cm3 is equal to
1ml
State what 1dm3 is equal to
1 litre
How do we convert between cm3 and dm3
Divide by 1000
Solution
In a solution, a solute is dissolved into a solvent
Concentrated solutions contain…
A large amount of solute in 1dm3
Dilute solutions contain…
A small amount of solute in 1dm3
Formula for atom economy
Mr of desired product/Total Mr of all products x 100
Formula for % yield
Actual yield/Theoretical yield x 100
When writing a dissociation equation for a weak acid, which type of arrow is used
Reversible
Definition of Acids
Proton donors
Strong acids________________dissociate all H+ ions in ________________ _______________
Completely
Aqueous solution
Which proton is lost of the carboxyl group in the dissociation of a weak acid
The most acidic proton
Formula for ethanoic acid
CH3COOH
How do you write out the dissociation of a strong acid?
Separate into constituent ions
How do you write an ionic equation?
- Identify all ionic compounds excluding liquids and gases (acids ionise)
- Separate into their constituent ions
- Rewrite equation
- Remove spectator ions
- Rewrite equation
Definition of a base
Proton acceptors
4 types of bases?
Metal oxides, hydroxides, carbonates and ammonia
Alkalis have ____ in them
OH-
In neutralisation, what happens?
H+ ions react with a base to form salt and water
When is a salt produced
When the H+ ions react of an acid is replaced by a metal ion or ammonia
Acid + metal oxide ——->
Salt + water
Acid + metal hydroxide ——->
Salt + water
Acid + ammonia ——->
Ammonium salt
Acid + metal carbonate ——->
Salt + water + carbon dioxide
Definition of a titration
A technique to accurately measure the volume of one solution that reacts with another
Describe how to prepare a standard solution
See notes
What is a known and unknown solution
Known - when we know the concentration, unknown - when we don’t know the concentration
Which acids are weak
Carboxyl - defined by ‘oic’
What is an orbital
Region within nucleus of an atom which holds 2 electrons with opposite spins
State the 4 subshells
S,P,D,F
State how many electrons each sub-shell can hold
S=2
P=6
D=10
F=14
State the electron configuration of oxygen
1s2 2s2 2p4
Describe how to work out the shorthand electron configuration
Take the final electron configuration of atom, travel along periodic table to the noble gases, go up one element and use the selected noble gases symbol, and just place the final configuration next to it
What shape is an S orbital
Spherical
What shape is a P orbital
Dunbell
Why does 4s2 fill before 3d10
Order of increasing energy level
Define an alkali
Bases that can dissolve in water and release OH- ions
Give an example of a strong acid
Anything that isn’t a carboxyl acid
Why do we rinse the beaker 3 times when preparing a standard solution
So all the solution is transferred to the volumetric flask, otherwise, the concentration would diluet
What will happen if the meniscus goes above/below the graduation mark
Too much/little water added affecting concentration
What is the ‘mass by difference method’
Full weighing boat - empty weighing boat = final accurate mass
Describe how to set out reacting masses calculation
Table should include…
1. Molar ratio
2. Mass
3. Mr
4. Moles
Which equipment is involved in titrations
- Burette
- Conical flask
- Bulb pippete
When calculating mean titre, which values do we include?
Concordant - within 0.1 of each other
Phenolphthalein turns what in alkali?
Pink
Phenolphthalein turns what in acid?
Clear
Methyl orange turns what in acid?
Orange
Methyl orange turns what in alkali?
Yellow
True or false
The rough titre is included in the mean tire calculation
False
Oxidation number
The number of electrons involved in bonding to a different element
What is the oxidation number equal to?
The charge
When calculating the oxidation number, which element do we always begin with?
The most electronegative
Uncombined elements have an oxidation number of?
0
Oxidation number of oxygen?
2-
Oxidation number of hydrogen?
1+
Oxidation number of hydride?
1-
Explain what the roman numerals in a compound mean
The positive charge of the ion and its oxidation state
What is an oxyanion?
A negative ion containing an oxide
Oxyanions always end in which suffix?
‘Ate’
Using oxidation states, state the formula of chlorate (VII)
- Oxygen and chlorine present, so ClO
- Cl has a charge of 7+
- Use the minimum amount of oxidation numbers for oxygen to make it negative
- In this case, it will be 4
- So, we used 4 oxygens, meaning a small four will go next to the oxygen and it will have a negative charge
ClO4-
Oxidation is when electrons are
Lost and the oxidation number increases
Reduction is when electrons are
Gained, and the oxidation number decreases
How can you work out oxidising or reducing agent?
Whatever is oxidised in RHS of equation, the same atom on the LHS will be the reducing agent
(Same for oxidising agent)
Hydrated
When water of crystallisation is present in a compound
Anhydrous
When all water of crystallisation has been removed from a compound
Water of crystallisation
When water is present in a compound giving a crystalline appearance (.xH2O)
Redox reaction
When oxidation and reduction take place simultaneously (OILRIG)
Oxidation is when the oxidation number __________, and electrons are _________
- Increases
- Lost
Reduction is where the oxidation number _______________ and electrons are ________________
- Decreases
- Gained
Explain how to identify the oxidising and reducing agents
Using oxidation numbers, whatever is oxidised on the RHS, the same atom on the LHS will be the reducing agents.
Converse for oxidising agent
How to tell if a compound is oxidised?
Oxidation number increases
When writing a half equation for oxidation, the electron number is added on to which side of the equation?
Right hand side
When writing a half equation for reduction, the electron number is added on to which side of the equation?
Left hand side
Disproportionation reaction
The same element is is oxidised and reduced
What constructing difficult half equations, which 4 steps are used?
- Balance the element being oxidised or reduced (same number of the element on each side)
- Write out oxidation numbers, if reduction, electrons go on RHS, if oxidation, LHS
- Balance oxidation numbers using electrons e.g., 7+ to 2+, 5e- needed
- Balance the charges of each side using H+ or OH-
- Use water to balance the H and O on the side where they aren’t present
When constructing half equations in acidic conditions, how should the charges be balanced?
Using H+ ions
When constructing half equations in alkaline conditions, how should the charges be balanced?
Using OH- ions
All atoms aim to be like….
The noble gases
The octet rule
Atoms of elements chemically combine with each other so each atom has 8 outer electrons. This is achieved by ionic/covalent bonding
Ionic bonding is between….
Metals and non metals
Definition of ionic bonding
Electrostatic force of attraction between oppositely charged ions
The net charge of any ionic compound is…
0
Structure of ionic compound
- Giant ionic lattice
- 6:6 coordination
- Regular repeated arrangement of ions
Melting/boiling point of ionic compound
High, strong electrostatic forces of attraction between oppositely charged ions in all directions of the lattice - take lots of thermal energy to break
Electrical conductivity of ionic compounds
- Non-electrical conductors when solid, ions fixed in place, can’t move and carry charge
- Electrical conductors when molten/aqueous, good electrical conductors, ions free to move and carry charge
Solubility of ionic compounds
Soluble in water and polar solvents
- Giant ionic lattice breaks down
- Water molecules surround each ion in solution
Covalent bonding is between….
Non metals
Define covalent bonding
Strong electrostatic forces of attraction between shared pair of electrons and the nuclei of bonded atoms
A covalent bond is an ___________ of atomic ___________
- Overlap
- Orbitals
What is meant by the term ‘localised’ in covalent bonding?
Attraction is localised, acting solely between shared pair of electrons ad nuclei of 2 bonded atoms resulting in a molecule
Define ‘lone pair’ in relation to electrons
Atoms not bonded covalently
Draw out a triple bond
3 dots, 3 crosses
Which bond is present in ethene?
Double bond (alkenes have double bonds)
True or false
Sulfur must have a full octet
False
Dative covalent bonding definition
Where only one of the atoms applies both of the electrons shared in covalent bonding
When adding a dative covalent bond in a displayed formula, how is this represented?
Through an arrow symbol
Melting and boiling point of simple covalent compound
- Low
- Little thermal energy needed to overcome weak intermolecular forces between molecules
Electrical conductivity of simple covalent compound
- Non electrical conductors
- No mobile electrons or ions
Solubility of simple covalent compounds
- Soluble in non-polar solvents e.g., hexane
- Forms induced dipole-dipole forces with these solvents
List examples of giant covalent compounds
- Diamond
- Graphite
- Buckminster fullerene
- Silicon
- Silicon dioxide
Melting point and boiling point of diamond
- High
- Strong covalent bonds between carbon atoms
- Lots of thermal energy needed to break covalent bonds
Melting point and boiling point of graphite
- High
- Strong covalent bonds between carbon atoms
- Lots of thermal energy needed to break covalent bonds
Electrical conductivity of diamond
- Non conductors
- No mobile electrons or ions
Electrical conductivity of graphite
- Electrical conductor
- Delocalised electrons can carry and move charge
Solubility of diamond
Insoluble in water
When adding a dative covalent bond onto a dot and cross diagram, where should this be added?
To a lone pair of electrons
Definition of metallic bonding
Strong electrostatic force of attraction between positively charged ions and delocalised electrons
Structure of metallic compound
Giant metallic lattice
Draw the structure of a metallic compound
Should include positive metals ions with an electron surrounding each one
Melting and boiling point of metallic compound
- High
- Strong electrostatic forces of attraction between positively charged ions and delocalised electrons
- Lots of thermal energy to break
Electrical conductivity of metallic compounds
- Good conductors
- Delocalised electrons can move and carry charge
Solubility of metallic compounds
Insoluble in water
Different isotopes have similar chemical properties. Explain why (1)
Same number of outer electrons
Reject same number of electrons/protons or different number of neutrons
Solid aluminium fluoride has a giant ionic lattice structure.
Describe what is meant by the term ionic lattice, in terms of type and arrangement of particles present (2)
Repeating pattern (1)
of oppositely charged ions (1)
ALLOW ‘regular’ OR ‘alternating’ OR ‘uniform (arrangement)’ for ‘repeating pattern’
ALLOW positive and negative ions OR aluminium ions and fluoride ions
ALLOW oppositely charged ions from a labelled diagram
If a molecule has a larger average bond enthalpy, what does this mean?
It has a stronger bond strength
Explain the electron pair repulsion theory (2)
- Bonded pairs repel each other to get as far apart as possible
- Lone pairs repel more strongly/further/greater than bonding pairs
Explain the shape of molecule in a methane atom (3)
- Tetrahedral shape with a bond angle of 109.5 degrees
- The 4 bonding pairs repel each other (as far as possible) equally
How much does each lone pairs reduce the bond angle by?
2.5 degrees
2 bonding pairs, 0 lone pairs
Linear, 180 degrees
3 bonding pairs, 0 lone pairs
Trigonal planar, 120 degrees
4 bonding pairs, 0 lone pairs
Tetrahedral, 109.5 degrees
6 bonding pairs, 0 lone pairs
Octahedral, 90 degrees
3 bonding pairs, 1 lone pair
Trigonal pyramidal, 107 degrees
2 bonding pairs, 2 lone pairs
Bent, 104.5 degrees
State the definition of electronegativity (2)
- The ability of an atom to attract electrons
- In a covalent bond
Which elements have the strongest electronegativity?
F, O, N, Cl
When one atom is more electronegative than the other, it forms a ____________________ ___________
Permanent dipole
Why do electrons sit closer to the more electronegative atom?
The electron cloud is more dense than
OR
The electron density is higher
What does the delta symbol (δ) represent
A slight charge
Definition of a permanent dipole
Small difference in charge across a bond as a result of differing electronegativities
What is the name of a molecule where a permanent dipole is present?
Polar (molecule)
Definition of non-polar molecules
Two atoms with the same electronegativities
Large difference in electronegativity
Ionic bonding
Medium to small difference in electronegativity
Polar covalent bond
No difference in electronegativity
Non-polar covalent bond
Explain Pauling’s electronegativity values (2)
- (Across periodic table) as nuclear charge
- (and) atomic radius increases, electronegativity increases
True or false: Symmetrical molecules are non-polar
True
In symmetrical molecules, all the dipoles __________ ____
Cancel out
True or false:In unsymmetrical molecules, the dipoles cancel out
False
Why is a molecule containing lone pairs not symmetrical?
Because the dipoles aren’t exactly the same
Explain the solubility of the polar molecule NaCl (3)
- Water molecules attract the Na+ and Cl- ions
- Ionic lattice breaks down as it dissolves
- (In resulting solution) water molecules surround the Na+ and Cl- ions
Where do IM forces exist?
Between covalently bonded atoms
Explain how permanent dipole-dipole forces arise
- Occurs between molecules containing a permeant dipole
- A positive ion from one molecule will align with a negative ion from another molecule
- This forms an IM force
Which IM force is strongest?
Hydrogen bonds
Where do hydrogen bonds exist?
Between molecules with H directly bonded to F, O or N and the H atom of —NH, —OH and HF
Explain how induced dipole-dipole forces (London forces) arise
- Electrons constantly moving and constantly changing electron density
- Electrons aren’t always symmetrical around molecule, sometimes imbalanced forming instantaneous dipole
- This induces a dipole in the neighbouring molecule forming an attraction
Where do London forces occur?
Between non-polar molecules
Which IM force is the weakest?
Induced dipole-dipole forces (London forces)
Explain the strength of London forces (3)
- The larger the instantaneous dipole and induced dipole
- The greater the induced dipole-dipole interactions
- The stronger the attractive forces between the molecule
List the 2 factors affecting electronegativity
- Increases across a period as the number of protons increases and atomic radius decreases (because electrons in the same shell are pulled in more)
- Decreases down a group because the distance between the nucleus and outer electrons increases and the shielding of inner shell increases
When going down group 7, why do boiling points increase?
- Increasing number of electron (because they’re bigger molecules)
- This increases the size of the induced dipole-dipole interactions between the molecules
Explain why BF3 is a non-polar molecule (3)
- Symmetrical
- All dipoles cancel out
- Fluorine is the most electronegative
All non-polar molecule will always contain ______________ forces
London
Anytime a lone pair of electrons is present, the molecule is always _______________
Polar
Which two atoms are always δ+
Carbon and hydrogen
Despite having the same type of intermolecular forces of attraction, ammonia has a lower boiling point that water.
Explain why (2)
- Hydrogen bonds between water molecules are stronger (than ammonia);
*Allow H-bonds*
- Because oxygen is more electronegative (than nitrogen) and has two lone pairs (whereas ammonia only has one);
Ammonia has hydrogen bonding, bromine has London forces.
Explain why bromine’s boiling point is higher than ammonias (1)
London forces/induced dipole-dipole/dispersion forces/Van der Waals forces between Br molecules are stronger than the hydrogen bonds between the ammonia molecule.
DO NOT NEED TO EXPLAIN WHY
When describing the type of IM forces in a molecule, which 2 factors must be included in the statement?
- Type of molecule e.g., polar
- Bonding occurs between molecules
Why does water have anomalous properties?
Due to hydrogen bonds between molecules
List the 2 anomalous properties of water
- Ice is less dense than water
- Water has a relatively high melting and boiling point
Explain why ice is less dense than water (2)
- Water molecules held further apart by hydrogen bonds between
- Giving open lattice structure
Explain why water has a relatively high melting and boiling point (2)
- H-bonds relatively strong (stronger than other IM forces)
- More thermal energy needed to break
Suggest why ice has a higher melting and boiling point than solid ammonia (4)
- H-bonds between water molecules stronger than H-bonds between ammonia molecules
- (Because) O is more electronegative than N
- (So) the dipole across the O—H bond is stronger than across N—H bond (because difference in electronegativity is greater in O—H, so, stronger dipole)
- (Also) there are 2 hydrogen bonds between water molecules and only 1 between ammonia molecules
List the 2 factors that make H-bonds stronger
- Amount of H-bonds
- Greater difference in electronegativity leading to a stronger dipole
Why does water have 2 H-bond whereas ammonia only has 1?
Water has 2 lone pairs whereas ammonia only has 1 (H-bonds form from lone pairs)
Explain the boiling points of the 3 substances listed below (5)
- H2O = 100
- CHCl = 62
- Cl2 = -34
- Mentions that all are simple covalent structures
- Water is polar, H-bond between water molecules
- Trichloromethane is polar, permanent dipole-dipole interactions between molecules
- Chlorine is non-polar, London forces between molecules
- H-bonds strongest, most thermal energy needed, London forces weakest, least thermal energy to overcome, permenant dipole-dipole have intermediate strength
Phosphorus, P4 has a higher melting point than chlorine, Cl2.
Explain this difference (3)
- Phosphorus has more electrons
- Stronger London forces
OR
Stronger induced dipole(-dipole) interactions
- More energy required to break the intermolecular forces / bonds OR London forces
Remember - London forces are affected by number of electrons, not electronegativity like H-bonds
Explain the boiling points of the molecules below (5)
- Ammonia = -33
- Fluorine = -188
- Bromine = 59
- Ammonia has H-bonding
- Fluorine and bromine have London forces
- Forces of attraction are between molecules (or IMF)
Statement may read: Ammonia has H-bonding between molecules - The London forces in bromine are greater than in fluorine because bromine has more electrons than fluorine
- Bromine has a higher boiling point than ammonia because the London forces are stronger than the hydrogen bonds and hydrogen bonding in ammonia is stronger than the London forces in fluorine
RTP
Room temperature and pressure
When calculating the moles of a gas at RTP, state the equation when the volume is given in cm3
n (mol) = volume (dm3) ÷ molar gas volume (24000 cm3 mol-1)
When calculating the moles of a gas at RTP, state the equation when the volume is given in dm3
n (mol) = volume (dm3) ÷ molar gas volume (24 dm3 mol-1)
When calculating a volume and the question doesn’t ask for specific units, can any moles of a base equation be used, as long as the correct units follow?
Yes
List the 2 methods of collecting gases experimentally
- PAG1 method
- Gas syringe and conical flask
In which instance is the ideal gas equation used
When reactions don’t occur at RTP
List the conditions when molecules are deemed to make up an idea gas
- Random motion
- Elastic collision
- Neglible size
- No intermolecular forces
State the ideal gas equation
pv = nRT
How do you convert from kPa to Pa
Multiply by 1000
How do you convert from cm3 to dm3?
Divide by 1000
How do you convert from dm3 to m3?
Divide by 1000
How do you convert from cm3 to m3?
Divide by 1000 twice, or divide by 10^6
How do you to and from degrees Celsius and kelvin
- Degrees Celsius to Kelvin = add 273
- Kelvin to degrees Celsius = subtract 273
At room temperature K2S is a solid, but SF2 is a gas.
Use ideas about structure and bonding to explain this difference (2)
- K2S has strong ionic bond or giant ionic lattice
- SF2 has London forces which are weak between molecules
Shape and bond angle of SF6 molecule (2)
Octrahedral
90 degrees
Suggest why SF6 is unreactive (1)
No lone pairs
Why does a student carry out a rough titre? (1)
To estimate the tire
A 25.0 cm3 pipette was used to measure out the 25 cm3 glutaric acid solution for each titration.
Before use, one student washed the pipette with water instead of glutaric acid solution.
State the effect on the titre.
Explain your answer (2)
Effect: Lower titre
Explanation: Water dilutes/lowers concentration
