Module 2 Flashcards
Relative isotopic mass
Mass of an isotope relative to 1/12 mass of an atom of carbon 12
Relative atomic mass
Weighted mean mass of an atom relative to 1/12 of an atom of carbon-12
Relative molecular mass
simple molecules
Relative formula mass
Giant ionic compounds
Isotopes
Atom of same element with different number of neutrons and different mass.
Isoelectronic
Different physical properties
Avogadros constant
6.02 x 10 ^23 mol-1
Isoelectronic
Same electron structure
Molar mass
Mass in g of one mole of a substance
mass=
Mr moles
Number of particles =
moles x avagadros constant
Empirical formula
Simplest whole number ration of atoms of each element present in a compound
Water of crystallisation
Water that is chemically bonded into a crystalline structure
Calculating the formula of a hydrated salt
1) moles of anhydrous and H2O
2) calculate ratio of amounts and formulas
Check water of crystallisation has been removed by …..
Heating to a constant mass
Concentration
The amount (in mol) of a dissolved substance in 1 dm3 of solution
Moles =
Concentration x volume
Molar gas volume increases as
Temperature increases
Molar gas volume decreases as
Pressure increases
Ideal gas equation
pV= nRT
T in Kelvin (+273)
V (m3)
R = 8.314 Jmol-1K-1
Calculating moles of a gas (RTP)
n = vol/24 (dm3)
Ideal gas equation Conversions
kPa-> Pa X1000
cm3-> m3 X10^-6
Stoichiometry
Ratio of moles in a chemical reaction
Why is a high atom economy good?
-efficient
-produce little waste
-less raw materials used
-sustainable
Atom economy=
(Sum of Mr of desired products/ sum of molar mass of all products) x100
Percentage yield=
(Actual/ theoretical) x100
acids
Release H+ ions when dissolved in water
Proton donor
Strong acids
Fully dissociate when dissolved in water
Weak acids
Partially dissociate when dissolved in water
Bases
Accept H+ ions
Proton acceptors
Examples of bases
-metal oxides
-metal hydroxides
-metal carbonates
-alkalis
alkalis
Dissolve in water releasing OH- ions
Neutralisation
Reaction between acid and base to produce a salt
Neutralisation by carbonate
Acid + carbonate -> salt + water + carbon dioxide
Neutralisation by metal oxide
Acid + metal oxide -> salt + water
Neutralisation by alkali
Acid + alkali -> salt + water
Acid + ammonia
Ammonium salt
Standard solution
Solution of known concentration
Preparing standard solution of NaOH
-mass (e.g. 1.00 g) weighed out and added to beaker
-dissolve NaOH with distilled water use stirring rod
-pour into 250 cm3 volumetric flask
-rinse beaker with distilled water, wash rinsings into flask
-add distilled water until bottom of meniscus on graduation line
-stopper, invert, thoroughly mix
How to find mass needed to prepare standard solution
Find moles (n = cv)
Find mass (m = Mn)
Oxidation number ions
Monatomic- oxidation no= charge
Carrying out acid-base titration
(Finding conc H2SO4 reacting with 25cm3 NaOH)
-Pipette, add 25.0 cm3 NaOH into conical flask
-white tile, few drops of indicator (phenolphthalein or methyl orange)
-burette, sulfuric acid, initial reading, nearest 0.05 cm3
- add sulfuric acid, swirl
-colour change, final reading
- final- Initial = trial titre
-repeat, dropwise near end point until concordant results (within 0.1)
-mean using concordant results
-calculations
Ethanoic acid
CH3COOH
Oxidation number elements
0
Titration calculation finding concentration
Balance equation
Mol known
Mol unknown
Conc unknown
Oxidation no Compounds
Halides -1
H +1
O -2
Oxidation no Exceptions
Metal hydride H-1
F2O O +2
Peroxide (H2O2) O -1
Oxidation no compounds + polyatomic ions
Sum of oxidation numbers in compound = 0
Sum of oxidation no in polyatomic ion = overall charge
How are oxidation numbers shown in elements with variable oxidation numbers?
Roman numerals
E.g. sodium chlorate (I) cl +1
Oil rig
Oxidation is loss of electrons
Reduction is gain of electrons
Redox in terms of hydrogen
Oxidation is loss of hydrogen
Redox in terms of oxygen
Oxidation is gain of oxygen
Redox in terms of oxidation no
Oxidation is increase in oxidation number
What are the 4 types of orbitals?
S,p,d,f
Shape of s orbital
Spherical
Shape of p orbital
Dumbell
(3D axis)
What is an orbital?
Region around Nucleus where an electron is most likely to be containing 2 electrons with opposite spins
Number of electrons in first four shells
2,8,18,32
Sub shells
Orbitals within energy level grouped together
Metals form ——- ions
Positive (cations)
Exceptions- electron structure
Cr and Cu
Less energy to fill 3d and have 1 4s
Or have 5 3d (all occupied singularly) and 1 4s
Electron configuration
Electrons fill orbitals in increasing energy
4s lower energy than 3d
Maximum of 2 electrons with opposite spin represents by up and down arrow
Occupied singularly before pairing
Non-metals form ——— ions
Negative (anions)
Ionic bonding definition
Electrostatic attraction between oppositely charged ions
Structure of ionic compounds
Giant ionic lattices, ions fixed in place
Properties of ionic compounds
High mp and bp (lots of energy needed to overcome strong ionic bonds)
Soluble in polar solvents (water)
Insoluble in non-polar solvents (hydrocarbons)
Strength of ionic bonds
Smaller ions form stronger bonds
Highly charged ions form stronger bonds
Conduct because
Moving charged particle (electron, ion)
Covalent bond definition
Attraction between positive nuclei and shared pair of electrons
Dative covalent bond
Both electrons contributed by one atom
Lone pairs
Pair of valence electrons nor bonded to anther atom
Able to form dative covalent bonds with atoms that have vacant orbitals
Valence
Outer shell
Average bond Enthalpy
Measure of covalent bond strength
Larger= stronger
Salt
H+ ion replaced by metal ion or NH4+ ion
Displayed formula
Bonds represented by lines
Covalent bonding happens between ….
Non-metals
Dative covalent bond displayed formula
Arrow
Electron deficient
Less than 8 electrons
Expanded octet
More than 8 electrons
valence shell electron pair repulsion theory
e- repel as far apart as possible
How much does each lone pair reduce the angle by??
≈2.5 °
LP> LP-BP> BP
repulsion
2 bonding regions
2bp, 0lp
linear 180 °
3 bonding regions
3bp, olp
trigonal planar 120 °
3 bonding regions
2bp, 1lp
non-linear 117.5 °
4 electron pairs
4bp, 0lp
tetrahedral 109.5°
4 electrons pairs
3bp, 1lp
pyramidal 107°
4 electron pairs
2bp, slp
non-linear 104.5°
e.g. H2O
5 bonding regions
5bp 0lp
trigonal bipyramidal
120°
90°
5 bonding regions
4bp, 1lp
pyramidal 119° 89°
see-saw 119° 89°
5 bonding regions
3bp, 2lp
trigonal planar 120 °
T-shaped 89°
6 electron pairs
6bp, 0lp
octahedral 90°
VSEPR
valance shell electron pair repulsion theory
6 electron pairs
4bp, 2lp
square planar 90°
6 electron pairs
5bp, 1lp
square pyramid 89°
method for determining shapes of elements
eg BrF3
central atom Br
outer e- 7
e- gained from bonds 3
e- gained from change 0
total e- 10
bp 3
lp 2
shapes of ions; SO4 2-
2 x S=O
2 x S–O
two atoms with single bond O have extra e- in outer shell resulting in 2- charge
tetrahedral shape 109.5
what can be said about the bonding affects of single bonds to double bonds?
they are similar
covalent bond H-H
attract electrons equally
electron equally shared
electronegativity
the ability of an atom to attract the bonding electrons in a covalent bond
trends in electronegativity
F most electronegative
decreases down group
increases across period
polar bond
pair of e- shared unequally
δ- slightly negative
δ+ slightly positive
dipole
separation of partial charges in a molecule
polar molecules
polar bonds (regions with different e- densities)
non-symmetrical
dipoles do not cancel
non-polar molecules
if polar bonds: symmetrical, dipoles cancel
types of intermolecular forces
induced dipole-dipole interactions (London forces)
permanent dipole-dipole interactions
hydrogen bonds
london forces
*e- move randomly, at any instant in time, distribution of e- may be uneven
*this creates weak instantaneous dipoles δδ-
*induces dipole in neighbouring molecules
permanent dipole-dipole interactions
2 polar molecules with permanent dipoles attract
hydrogen bonds
strong permanent dipole-dipole interactions
H and O/F/N
hydrogen bond diagrams
dipoles labelled
h in one molecule bonded to lone pair in another
hydrogen bond represented by dotted line
straight line
ice less dense than water
regular structure, open lattice
when ice melts, the open lattice structure collapses allowing H2O molecules to move closer together increasing density
water expands when it freezes
tetrahedral structure maximising hydrogen bonding
when ice melts, crystal structure melts as water molecules fall into empty spaces
H2O higher melting point than expected
hydrogen bonds
Aufbau principle
e- enter lowest energy level available
Hund’s rule
e- prefer to occupy orbitals on their own and can only pair up when no orbital of the same energy is available (bus analogy)
occupy singularly before pairing
