This document is filled with all of my notes as an IB Chemistry HL student where many topics are summarised in 1-2 pages with the most important parts except fro chapter 10 and option D which required a bit more due to their elevated importance in paper 2 and 3. This has helped me study for my fina...
Chapter 11 & 21, Measurements & Data Processing (IB Chemistry)
Chapter 10&20 , Organic chemistry (IB Chemistry)
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Homolegeous mixture Moles & Particles
A
• uniform composition & properties throughout Ideal Gas equation Titrations
• ex; air, bronze particles • to determine the concentration of unknown solution
Heterogenous mixture
•
•
non-uniform composition, properties are not the same throughout
ex; concrete, orange juice
-Te constant
· •
•
equipment: graduated pipettes, volumetric pipettes,
brunette (± 0.05 cm^3)
EQUIVALENCE POINT: when the 2 solutions have
I
Gas law relationships reacted completely
Separating mixtures
• air —-> fractional dissipation Moles & Mass
Smaa 2
• salt & sand —-> solution & filtration ·
P Pat
• pigments in food colours —-> paper chromatography
• sulfur & iron —-> use a magnet ~
moles
1
Ionic equations 1 gmol- V T
• shows only the ions or other particles taking part in a reaction, without spectator
ions
Excess & Limiting reactants Temp
State changes
• excess of one or more reactant (excess)
·
vaporising
melting
#
GAS
D
·
Fiening
concensin
• LIMITING REACTANT: not excess reactant
e
• Calculating limiting reactant: number of moles of
reactants, ration of reactants in the equation p
Deposition
The Mole Percentage Yield Real gases
• Avogadro’s constant (Na or L): number of particles equivalent to the percentage -
actual yield
-100
-
relative atomic mass or molecular mass (in grams) yield yield A
theoretical 200k
~
- = 6.02 x 10^23 g/mol Avogadro's law
• Molar mass: the mass of a substance with this number of particles • STP = 273 K, 100 kPa
• 1 mole = atoms in 12.00g of nC • 1 mole (STP) of gas = 22.7 dm^3 I 500k
-
- 6.02 x 10^23 of 12-C = mass of 12.00g • units: dm^3/mol K
Relative mass Molar Gas volume -100
e
----------- ideal
RELATIVE ATOMIC MASS:
-
gas
• weighted average mass of one atom compared to 1/12 of 12-C
• Determined by weighted average mass of isotopes
moles Te Pressure
Heightedaverage massofono ↳
er is
Ar =
&d high pressure
,
at low temp
RELATIVE ISOTOPIC MASS: Concentrations of Solutions
• mass of a particular atom of an atom compared to 1/12 mass of 12-C real gases deviate significantly
·
• ISOTOPES: atoms of the same element with a different neutron number
Relative
from ideal gas
Elisotope abundances relative isotopic mas o
atomic =
100
RELATIVE MOLECULAR MASS, Mr
• weighted average mass of a molecule compared to 1/12 mass of 12-C tration X solution
(dm3)
Mr
weighted average mass of one molecule (g/dm3))
I
=
Ye mass of one atom of "C
, Isotopes Electromagnetic spectrum *
• are atoms of the same element that contain the same number of • a range of frequencies that covers all electromagnetic radiation • heat or electricity can be used to EXCITE an electron to a higher energy level
protons & electrons but a different number of neutrons & their respective wavelengths and energy • when electrons ‘fall’ back down they must lose the energy different between 2 energy
• FREQUENCY: how many waves pass per second levels
Relative atomic mass • WAVELENGTH: distance between 2 consecutive peaks • this loss of energy releases ELECTROMAGNETIC ENERGY
• the average mass of all the isotopes
high energy • gives evidence for Bohr’s model of discrete energy levels so an exact amount of energy
new enery length short wavelength
(t high frequency • limitations to this model: assumes electrons are fixed, spherical, only works for hydrogen
total mass (% abundancexmassa) +
BX ......
B) low fre quency
(with 1 electron)
......
~wwGAMT
=
of 100 atoma 100
Electron energy levels
RADIO
WAVES MICRO-INFRARED
VISIBLE UV X-
RAYS
Ionisation Energy
LIGHT
WAVES • IE: the energy required for en electron to escape the atom, or reach the upper limit of
• ELECTRONIC CONFIGURATION: the arrangement of electrons in an atom
3x00 m/s convergence
• PRINCIPAL QUANTUM NUMBER (n): number of energy levels / quantum shells 2
va C = v .
-
C =
ligh ↑
.
of • Layman series
• quantum number decreases the closer the shell to the nucleus speed
• across a PERIOD: IE increases due to —> increase in nuclear charge, decrease in atomic
• quantum number increases the greater the energy of electrons
Continuous vs Line spectrum radius, shielding remains constant, increase in energy needed to remove an electron
Aufbau principle • CONTINUOUS: in the visible region contained all the colours • down a GROUP: IE decreases due to —-> the outermost electron is farther from the
• ground state: most stable electrons configuration, lowest energy amount of the spectrum nucleus, held less tighter and so less energy is required to remove the electron
• AUFBAU PRINCIPLE: filling the subshells of energy with the lowest energy first • LINE: only shows certain fixed frequencies of electrons
• Exceptions; Cr - [Ar] 3d^5, Cu - [Ar] 3d^10 3s^1 (this is energetically favourable Dips in the IE trend
Line emission spectra
• between Beryllium & Boron there is a light decrease in IE
• each line is specific energy value, electrons posses a limited choice
Hund’s rule • Be -
It : 900k5/mol ,
Is'Is
of allocated energies 25
• B
K5/mo is
• SPIN PAIR REPULSION: spin repel each other when they spin in the
800
-It,
=
↳isfurther awayfromede aan
,
• CONVERGENCE: lines get closer towards blue end of spectrum,
same direction
lines converge to higher energy, electrons reach max energy
• electrons have small spinning charges which rotate on their own axis in • between Nitrogen & Ozygen slight decrease in IE, due to spin pair repulsion
• max electrons reach = IONISATION ENERGY
CLOCKWISE or ANTI-CLOCKWISE • N - It, =
1400k5/mol ,
Is 2s"2px'2py' 2px
• O- It 1310h5/mol 132 Is 2 pc2py' 2 pc
Hydrogen spectrum
=
,
Pauli Exclusion principle
-
↳ 2 é = repulsion between them ,
for one e
to be
• an orbital can only hold 2 electrons
makes it easier
A
L
removed
8
7 Successive ionisation energies
Sublevels & energy Periodic table
• IE of an element increases
S-block p-block
i -V
PLN L
IR C • this is because removing an electron from a position ion , attractive forces increase due to
s
↑ *
decrease in shielding & an increase in the proton to electron ratio
BREFeTT
VvV
Joop
~
(IR) • electrons are harder to remove when there is: less spin-pair repulsion, the closer the shell is
to the nucleus
4S (IR)
ENERGY
>
-
n3
4d n
e
-VVVVV
Deducing the IE group
BALMER
i 1
5d (visible)
• were there is the biggest difference in IE
+ n2
no
f block
Y
-
IS
5f
10000000
I
LYMAN
(ru) Shighogy
n
+ 11
, Periodic table trends
Transition Metals
electron affinity *
ionisation
energy VARIABLE OXIDATION STATES
D
A A • when orbitals are occupies, the expulsion between electrons pushes 4s into a higher energy
METALLIC CHARACTERISTICS
state, so it becomes lightly higher in energy then 3d, losing its electron first
• METALS: 1-3 outer shell
• transition metal ions with +3 tend to be POLARISING, have a HIGH CHARGE Density and
Alkaline
metals vigorously
be * electrons, low IE, low
pull on electrons
-react Het d
electronegativities (due to
-
-
S
with halogens
h A
delocalised electrons COMPLEX IONS
• NON-METALS: 4-7 electrons in • central metal ion + ligand
h T character 8 outer shell, high electronegativity, • LIGAND: a molecule or ion that form complexes which consist of a central metal ion and ligands
E
+ E & high EA (tendency to share —> LEWIS BASE, NUCLEOPHILE
non-metallic
S -'
W
: character - i P
5
↑
electrons and form covalent bonds • CO-ORDINATION NUMBER: number of co-ordinate bonds to the central metal atom or ion
r - • NAMING COMPLEXES: prefix for number of ligands/ligands name/element/oxidation number
metallic
8
OXIDES:
-
d e • Na and Mg oxides are basic
V
=
• Al oxides are amphoteric CATALYTIC PROPERTIES
#
• Si to Cl oxides are acidic
• can catalyse certain REDOX REACTIONS —> can be readily oxidised and reduced
*
* • HETEROGENOUS catalyst has a different physical state (phase) from the reactants
atomic radius
# • HOMOGENOUS catalyst is the same physical state (phase) as the reactants
ELECTRON AFFINITY • BIOLOGICAL catalysts
ATOMIC RADIUS
• distance between nucleus of an atom & outermost electron shell • amount of energy released when 1 mole of
electrons is gained by 1 mole of atoms of gaseous MAGNETIC PROPERTIES
• ACROSS A GROUP: nuclear charge increases = greater pull
state to form gaseous ion (- charge) • DIAMAGNETISM: the property of all materials and produces a very weak opposition to an
• DOWN A GROUP: increase in number of shells, increase in
• first EA = exothermic applied magnetic field —> from repulsion of electrons to the applied magnetic field, create a tiny
shielding, weak pull
• second EA = endothermic (overcoming repulsion magnetic dipole
between electrons and - ion) • PARAMAGENTISM: only occurs in substances which have unpaired electrons, produces
IONIC RADIUS
magnetism proportional to the applied field in the same direction
• measure of size of an ion
• FERROMAGNETISM: the alignment of the unpaired electrons in an external field can be
• DOWN A GROUP = increases ELECTRONEGATIVITY
retained so that material becomes permanently magnetised
• ACROSS A PERIOD: ionic radii increases with an increase • the ability of an atoms to attract a pair for electrons
in - charge, ionic radii decreases with an increase in + charge towards itself in a covalent bond
COLOURED COMPOUNDS:
• arises form the + nucleus ability to attract - charged
• d-block elements have unpaired electrons
electrons
IONISATION ENERGY • the d-orbitals are split into two energy levels
• PAULING SCALE - assigned electronegativity value
• DOWN A GROUP= increase in nuclear charge, increase in • electrons can transition between these energy levels
• FLUORINE most electronegative
shells, increase in shielding, increase in atomic radius, • in the meantime they can absorb energy form light at a visible wavelength and thus, one can
• DOWN A GROUP = negligible nuclear charge increase,
electrons held more loosely observe the complimentary colour
increase in shielding, increase in atomic radius, decrease
• ACROSS A PERIOD = increase in nuclear charge, shells
in attraction of nucleus and electron
remain the same, shielding remains constant, decrease in
atomic radius, electrons held more tightly
, Formals charge
Orbitals
• FC = valence electrons - 1/2 (number of bonding electrons) - (number of non bonding electrons)
• s= s orbital is closer to nucleus =
·
greater bond compared to p orbital • —> FC = V - 1/2 x B = N
• p= & • most preferred structure when formal charge is 0
• d=
• f =& Hybridization H
Cl Hy I
It - -s
294 sp3
↑ P
st
• 4 single bonds = sp^3 c = H C
J
H
C
• each atom that combine has an atomic orbital
-
p
- -
i
• 3 single bonds = sp^2
↓I I -
S
-
↳
-
P
> Oxygen has 2 lone pairs of
containing a single unpaired electron
• 2 single bonds = sp CI
C d Hip H3
electrons. This is also counted.
• when a covalent bond is formed, form a combined
CI Sp
orbital containing 2 electrons
• Consider double single bonds with hybridisation cause it is about s and p orbitals
• the greater the atomic over lap, the stronger the bonds
Shapes of Molecules 4
-bond Resonance structures Trigonal bipyramid
jop
.. 180
- CI
120
• bonding pairs = 5
• are formed from head to head (end to end) overlap to atomic orbitals F di
• lone pairs = 0
-
II
• a single covalent bond formed when 2 non-metals combine Species Resonance structure Hybrid
D&
• molecular geo = Trigonal bipyramid
&
-
• the electron density is concentrated between 2 nuclei Carbonate ions gi·
:
[o]"
Sp
Die
• s + s orbital & s + p-orbitals = sigma ortbial
di-
E > C
&
S
-
I+ &D
4) H CO is - - · • bonding pairs = 4 Fe
...... 2179
O
a-
.. :
-
S
-
Sp
• lone pairs = 1
&
-
• the electron density in a sigma bond is symmetrical about a line Benzene
•
joining in the nuclei of the atoms forming the bond
The electrostatic attraction between the electrons and nuclei bond
& • molecular geo = See saw
F
• bonding pairs = 3
the atom to each other Ozone
r I
...
↑
7
& F
• lone pairs = 2
-
D3 "
0..
.
F
-
• molecular geo = T- shape
Carbonylate ion
T -bonds
E m-
:
[
-
(RCOO )
&
c • bonding pairs = 2
R- -
m
• are formed from the sideways overlaps of adjacent orbitals -
r
o
• lone pairs = 3
r -
-
• the 2 lobes that make up the pi-bond lie above and below the plane
• molecular geo = linear
-
of a sigma bond
• this maximise overlap of p-orbitals
Octahedral
• a single pi-bond is drawn as 2 electron clouds, one arising from F
• bonding pairs = 6
F F
each lobe of p-orbitals
...........
S
·
• lone pairs = 0 #
#F
&
• p-orbital + p-orbital = pi bond if ‘lateral overlap’, from double bonds F
• molecular geo = Octahedral
O - ..... i ..... F
Ef • bonding pairs = 5
F
- c c
- plane
-
>
-
T
& • lone pairs = 1
I F>F S
S
-
-
e-cloud
• molecular geo = Square based planar
180
F
• bonding pairs = 4 ...-
F
Xe---
XX
• lone pairs = 2 F
/ 1 XX
-
• molecular geo = Square planar
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