foundations in chemistry Flashcards
proton
neutron
electron
- give the relative masses and relative charges
1, +1
1, 0
1/1800, -1
atomic number
the number of protons in the nucleus
mass number
total number of protons and neutrons in the atom
number of neutrons =
mass number - atomic number
isotopes
atoms of the same element with the same number of protons but different numbers of neutrons
why do isotopes form
to gain a more stable ionic compound
isotopes have similar…
isotopes have varying…
similar chemical properties as they have the same electronic structure - same number of electrons in the outer shell
varying physical properties as they have different masses
relative isotopic mass
the mass of one isotope compared to one-twelfth of the mass of one atom of carbon-12
relative atomic mass
the weighted mean mass of one atom compared to one-twelfth of the mass of one atom of carbon-12
relative molecular mass
the average mass of a molecule compared to one-twelfth of the mass of one atom of carbon-12
how to calculate relative atomic mass
relative atomic mass = total(isotopic mass x % abundance)/100
ammonium formula
NH4+
phosphate formula
PO4 3-
carbonate formula
CO3 2-
sulfate formula
SO4 2-
nitrate formula
NO3-
hydroxide formula
OH-
what are spectator ions
- ions that are not changing state
- ions that are not changing oxidation number
ions definition
have a different number of electrons and protons
when are ions created
when electrons are transferred from one atoms to another. they then attract to each other and form compounds
mole definition
the amount of substance in grams that has the same number of particles as there are atoms in 12 grams of carbon-12
molar mass
the mass in grams of 1 mole of a substance
given the unit g/mol-1
how to calculate molar mass
add up all the mass numbers of each element in the compound
molar gas volume definition
units = dm3mol-1
the volume of 1 mole of a gas at a given temperature and pressure.
at 1atm and 25 degrees (room temp and pressure), the molar gas volume is 24dm3mol-1
avogadro’s constant definition
there are 6.02x10 23 atoms in 12 grams of carbon-12
what can avogadro’s constant be used for
atoms, molecules and ions
eg one mole of carbon atoms will contain 6.02x10^23 atoms
mole =
mass/mr
how many mg in a gram
how many grams in a kg
how many kg in a tonne
all 1000
empirical formula definition
the simplest ratio of atoms of each element in the compound
molecular formula definition
the actual number of atoms of each element in the compound
hydrated salt definition
contains water of crystallisation
heating in a crucible experiment
- what is it used for?
- method with example equation using calcium sulfate
- why should large amounts or small amounts of reagent not be used?
- why does the crucible need to be dry?
- why do we need a lid
- used for measuring mass loss in thermal decomposition reactions, and mass gain when reacting magnesium in oxygen
- CaSO4.xH20 = CaSO4 + H20
- weigh an empty clean dry crucible and lid
- add 2g of hydrated calcium sulphate to the crucible and weigh again
- heat strongly with a bunsen for a couple of minutes
- allow to cool
- weigh the crucible and contents again
- heat crucible again and reweigh until you reach a constant mass to ensure the reaction is complete
- large: to ensure the decomposition is not incomplete
- small: errors in weighing are too high
- wet crucible would cause mass loss be too large as water would be lost during heating
- improves the accuracy - prevents the loss of solid from the crucible but is also loose fitting enough to allow the gas to escape
concentration =
n/v (moles/volume)
conc = moldm-3
volume = dm3
cm3 to dm3
cm3 to m3
dm3 to m3
m3 to dm3
divide by 1000
divide by 1000000
divide by 1000
times by 1000
mass concentration equation and units
mass concentration = mass/volume
units:
mass conc = g dm-3
mass = grams
volume = dm3
method for making a solution
- weigh the sample bottle containing the required mass of solid on a 2dp balance
- transfer to a beaker and reweigh the sample bottle
- record the difference in mass
- add 100cm3 of distilled water to the beaker, use a glass rod to help dissolve the solid. heat gently if solid is not dissolving.
- pour solution into a 250cm3 graduated flask via a funnel
- rinse beaker and funnel and add washings from the beaker and glass rod to the volumetric flask
- make up to the mark with distilled water using a dropping pipette for last few drops
- invert flask to ensure a uniform solution
REMEMBER TO FILL TO THE MENISCUS
method for diluting a solution
2. why use a volumetric pipette
3. why use a teat pipette
- pipette 25cm3 of original solution into a 250cm3 volumetric flask
- make up to the mark with distilled water, using a dropping pipette for last few drops
- invert to ensure a uniform solution
2. more accurate than a measuring cylinder as it has a smaller uncertainty
3. to ensure volume of solution is accurately measured and does not go over the line
formula for a new diluted concentration
new diluted concentration = original concentration x (original volume/new diluted volume)
ideal gas equation and units
PV = nRT
pressure = Pa
Volume = m3
Temp = K
n=moles
R = 8.31JK-1mol-1
how to convert from celsius to kelvin
add 273
what are the potential errors when using a gas syringe
- gas escapes before a bung is inserted
- syringe sticks
- gases like co2 or so2 are soluble in water so the true amount of gas is not measured
molar gas volume equation
v(dm3) = n x 24
number of particles/ions/molecules/atoms equation
moles x avogadro’s constant
density equation
and density units
density = mass/volume
density units = gcm-3
percentage yield equation
actual/theoretical x 100
how can %yield be lowered
- incomplete reactions
- side reactions
- losses during transfer of substances
- losses during purification stages
atom economy equation
mass of useful products/mass of all reactants x 100
USE BALANCING NUMBERS
if a process has a waste product, how can the economics of the process be improved
by selling the bi-product for other uses
acid definition
releases H+ ions in aqueous solutions, proton (H+) donor
common strong acids
common weak acid
HCL
H2SO4
HNO3
CH2COOH (ethanoic acid) - weak acid
base definition
proton (H+) acceptor
common bases
metal oxides
metal hydroxides
ammonia
alkali definition
soluble base that releases OH- ions in aqueous solutions
common alkalis
sodium hydroxide (NaOH)
potassium hydroxide (KOH)
ammonia (aq) - NH3
strong and weak acid definitions
strong = completely dissociate when dissolved in water
weak = slightly dissociate when dissolved in water, to form an equilibrium mixture
salt definition
formed when the H+ ion of an acid is replaced by a metal ion or an ammonium ion
acid + base =
acid + carbonate =
salt + water
salt + water + carbon dioxide
observations in a carbonate reaction
- effervescence due to CO2 gas
- solid carbonate will dissolve
titrations
1. method
2. how to work out average titre results
3. safety procedures
4. what does it mean if the jet space is not filled properly
5. why is a conical flask used and not a beaker
6. why do we only add a few drops of indicator
7. why can be distilled water be added during the titration
- rinse equipment (burette with acid, pipette with alkali, conical flask with distilled water)
- pipette 25cm3 of alkali into conical flask
- touch surface of alkali with pipette (to ensure the correct amount is added)
- add acid solution from burette
- make sure the jet space is filled with acid
- add a few drops of indicator and refer to colour change at end point
- use a white tile underneath the flask to help observe the colour change
- add acid to alkali whilst swirling the mixture and add acid dropwise at the end point
- note burette reading before and after addition of acid
- repeat until at least 2 concordant results
- make an average using only the concordant results
- acids and alkalis are corrosive: wear safety gloves and eye protection
- will lead to a larger than expected titre reading
- easier to swirl without spilling the contents
- indicators are weak acids: too much will affect the titration result
- so all the acid is washed into the reaction mixture, does not react with reagents or change the number of moles of acid added
titration indicators
- colour changes
- when to use
- phenolphthalein
pink (alkali) to colourless (acid)
end point = pink colour just disappears
use if NaOH is used - methyl orange
yellow (alkali) to red (acid)
end point = orange
use if HCl is used
apparatus uncertainties
- balance
- volumetric flask
- 25cm3 pipette
- burette
+- 0.001g
+-0.1cm3
+-0.1cm3
+-0.10cm3 (as two readings are taken)
% uncertainty equation
= uncertainty/ measurement made on apparatus x100
how to decrease apparatus uncertainties
- decrease the sensitivity uncertainty by using apparatus with greater resolution
- increase the size of the measurement made
how to reduce uncertainties in a titration
- replace measuring cylinders with pipettes or burettes which have lower apparatus uncertainty
- make the titre a larger volume by increasing the volume and conc of the substance in the conical flask or decreasing the conc of the substance in the burette: this reduces the % uncertainty in a burette reading
how to reduce uncertainties in measuring mass
- using a balance that measures to more dp
- using a larger mass
oxidation
the process of electron loss
involves an increase in oxidation number
reduction
process of electron gain
involves a decrease in oxidation number
oxidation numbers for…
group 1 metals
group 2 metals
Al
H
F
Cl, Br,I
O
+1
+2
+3
+1 (except in metal hydrides where it is -1 eg NaH)
-1
-1 (except in compounds with oxygen and fluorine)
-2 (except in peroxides where it is -1 and compounds with fluorine)
electron structure
1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4p6, 5s2, 4d10…
shapes of s sub levels and p sub levels
s sub levels = spherical
p sub levels = dumbbells
ionic bonding definition
electrostatic force of attraction between oppositely charged ions formed by electron transfer
what makes ionic bonding stronger and the melting points higher?
- when the ions are smaller and have higher charges
what are the physical properties of ionic compounds?
give an example of the structure
- high melting points and bp: strong electrostatic forces between oppositely charged ions
- non conductor of electricity when solid: ions are held in lattice and so cannot move
- conductor of electricity when molten or in solution: ions are free to move and charge can be carried
- soluble in aq solutions
FORMS GIANT IONIC LATTICES
covalent bonding definition
the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
which period allows ‘expansion of the octet’
period 3
dative covalent bonding definition
- common examples (make sure you can draw them)
forms when the shared pair of electrons in the covalent bond comes from only one of the bonding atoms
- NH4+, H3O+, NH3BF3
larger average bond enthalpy =
stronger covalent bond
covalent bonding properties and types of structure formed
covalent bonding forms SIMPLE MOLECULAR structures
low mp and bp: due to weak intermolecular forces between molecules
solubility: poor
conductivity: poor as no ions to conduct and electrons are fixed in place
electronegativity definition
the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself
most electronegative atoms are…
F,O,N and Cl
factors affecting electronegativity…
- increases across a period as number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more
- decreases down a group as the distance between nucleus and outer electrons increases the shielding of inner shell electrons increases
what does a small electronegativity difference indicate?
the compound is covalent
how does a permanent dipole (polar covalent) bond form
when the elements in the bond have different electronegativities
whats a polar covalent bond
- has an unequal distribution of electrons in the bond, producing a charge separation (dipole)
are symmetrical molecules polar or non-polar and why?
- non-polar
- the individual dipoles on the bonds cancel out due to the symmetrical shape of the molecule: there is no net dipole moment
will compounds with a large electronegativity difference be ionic or covalent
IONIC
where do induced dipole-dipole interactions occur
- between all simple covalent molecules
- between the separate atoms in noble gases
- NOT IN IONIC SUBSTANCES
how do induced dipole-dipole interactions form
- in a molecule, the electrons are constantly moving
- as this happens, electron density fluctuates and parts of the molecule become more or less negative
- from this, temporary dipoles form
- these temporary dipoles cause dipoles to form in neighbouring molecules. these are induced dipoles.
factors affecting the size of induced dipole-dipole interactions
- the more electrons, the higher the chance that temporary dipoles will form. this makes the induced dipole-dipole interactions stronger and so boiling points will increase.
explain permanent dipole dipole forces
- occurs between polar molecules
-they are stronger than induced dipole–dipole interactions and so the compounds have higher boiling points - Polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)
- Polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms.
where does hydrogen bonding occur
- in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons
- it occurs in addition to induced dipole dipole interactions
explain the molecular structure of ice
Water can form two hydrogen
bonds per molecule, because
oxygen is very electronegative,
and it has two lone pairs of
electrons
The molecules are held
further apart than in liquid
water and this explains the
lower density of ice
explain the molecular structure of iodine
- covalent bonds between the iodine atoms in the I2 molecule
- crystals contain a regular arrangement of I2 molecules held together by weak induced dipole-dipole interactions
explain why the chemical properties of 35Cl and 37Cl are similar
- they are isotopes
- they have the same e- configuration/ same number of electrons in their outer shell
state the block in the periodic table in which silicon is placed and explain your answer
- p block
- silicon’s outermost electron is in a p orbital
suggest why the relative atomic mass on the periodic table may be different to the relative atomic mass of a sample analysed by mass spec
- other isotopes could be present
- some isotopes may be absent
- different abundances of isotopes
state the formula of two common silicon ions
Si4+
Si4-
name an instrument which would be used to measure the relative abundances of isotopes of an element
mass spectrometer
explain the graph which would be produced when measuring the relative abundances of isotopes in an element
describe how the graph and data obtained can be used to calculate the relative atomic mass of the element in question
- mass spectrum produced gives a percentage abundance of each isotope present
- against an m/z value for each isotope
to calculate the relative atomic mass:
- multiply each m/z value by its abundance
- add together and divide by total abundance
predict the size of the atomic radius of 35Cl relative to the atomic radius of 37Cl
- same atomic radius
- each isotope has the same number of protons and electrons
what makes it easier for an element to be reduced/gain electrons
- high electronegativity
- small atomic radius
- amount of electrons needed to gain full outer shell
explain why a reaction would have the lowest atom economy
- it has the greatest amount of waste product
fermentation reaction
- no other reactants
- decomposes to form alcohol and carbon dioxide
state an advantage of developing a chemical process with a high atom economy
- less waste produced
- reaction is more sustainable
limestone chemical formula
CaCo3
why may water of crystallisation experiment be inaccurate
- how to improve the experiment
- not heated to a constant mass
- heat for longer, use a smaller mass to make sure all the water has been removed
describe the effect on the value of x obtained if the salt was not heated strongly enough to a constant value
- some water remains with the salt
- the moles of water of crystallisation calculated are too low and the value of x would be too small
how to improve atom economy
- find a use for the waste product
state three definitions of reduction
- the loss of oxygen
- the gain of hydrogen
- the gain of electrons
write an equation to show the first ionisation of sulfuric acid
H2SO4 = HSO4- + H+
acid + ammonia =
sulfuric acid + NH3 =
forms only a salt, no water
- (NH4)2SO4
give a use for barium chloride solution
used to test for sulfate ions
what is the effect when
- funnel is left in the burette
- jet space not filled correctly
- difficulty when seeing the base of the meniscus
- decreases final reading
- increases final reading
- decreases the final reading
state how many orbitals there are in a p-sub shell and how the electrons are arranged if this sub-shell is full
- three atomic orbitals
- each orbital holds two electrons
-with opposite spins
why are 2s orbitals filled before 2p orbitals
- 2p orbitals have a higher energy compared to 2s orbitals
the group one and group 2 metals are also known as s-block elements
- explain this statement
- the outer electrons are s-sub-shell electrons for all group one and two elements
explain the strength of the ionic bond in sodium chloride
- strong electrostatic attraction
- between Na+ and Cl- ions
explain why it is easier to use aqueous MgF2 than molten MgF2 in a lab setting
- MgF2 has a high melting point
- as there are strong electrostatic attractions between oppositely charged ions in all directions
predict whether the enthalpy of the OH bond in Both h20 and CH3OH will be the same
no
as bond enthalpy is a mean taken across a range of compounds
state how the bond enthalpy of the C=O will compare to the bond enthalpy of the C-O
- the bond enthalpy of the C=O bond will be stronger/higher
explain how ionic bonding holds the particles together in an ionic compound
- strong electrostatic forces of attraction between the positively charged ions and negatively charged ions
explain the solubility of lithium and the lithium halides
- lithium in insoluble, but reacts with water to form lithium hydroxide solution
- all the lithium halide salts are water soluble
- lithium iodide is the most water soluble
- as you move down the group, the attraction between the lithium ion and the halide ion decreases
describe the bonding process that forms lithium halides, with respect to electrons
- lithium donates one electron
- the halogen accepts one electron
- the lithium forms a 1+ ion and the chloride forms as 1- ion
describe how the shape of the periodic table is linked to electronic structure, with reference to sub-shell structure (6 marks)
- elements in the same period have the same number of partially filled shells
- elements in the same group have the same number of valence electrons
- the s block has two columns as it only takes 2 electrons to completely fill a sub shell
- the p block has 6 columns as it takes 6 electrons to completely fill the p-sub shells
- the d block has 10 columns as it takes 10 electrons to completely fill the d sub shell
- the d block beings on period 4 as d-orbitals are filled after 4s orbitals
suggest why there are three possible p-sub-shells but only one possible s-sub-shell in an atom
- s orbitals are spherical so multiple sub shells are not possible as they would occupy the same space
- p orbitals are dumbell shaped and can be positioned along the x.y or z axis in three different directions
- three p-orbitals would not overlap significantly
justify why hydrogen is positioned in the middle of the periodic table and not as the first element in group one
- hydrogen has very different physical properties to group 1 metals
- has very different chemical properties to group 1 metals - does not form cations as readily
- despite having an outer shell configuration ending in s1
suggest a way in which the bond angle in ammonia could become 109.5 and explain
- the nitrogen can form a dative bond
- there will then no longer be a lone pair
- no extra repulsion
predict and explain if the halogen molecules are able to conduct electricity in any state
no as they have no delocalised electrons
explain why CCl4 is non polar whilst CH3Cl is polar
- C-Cl bond has a difference in electronegativity / polar
- CCl4 is symmetrical so polarity of the Cl-Cl bonds cancels out
