Grade 12 Chemistry (Last Concept map ;)
Organic Chemistry
Amines
nitrogen
naming
add -bi, -tri to -amine to indicate multiple of same type of alkyl
add suffix -amine
each alkyl group in alphabetical order
ammonia molecules
amines
hydrogen is replaced by an alkyl chain
3 possible covalent bonds
1 alkyl chain
2 alkyl chain
3 alkyl chain
primary amine
secondary amine
tertiary amine
esters
H2O as a product
formed by
reaction between
alchool
carboxylic acid
are carboxylic acids
OH is replaced with O and alkyl
OCO
naming
first
second
alchool
carboxylic acid
anoic acid --> oate
ol --> yl
carboxylic acid
named using suffix -anoic acid
COOH group
contain carboxyl group
C=O
contains hydroxyl group
OH
the end of chain
most polar molecule
aldehydes and ketones
lower boiling points than alchools
ketone
-one
carbonyl group in the middle of the chain
not at the end
aldehyde
carbonyl group at the end of carbon chain
suffix -al
alcohols and ethers
ether
naming ethers
use -oxy to the end of name
shorter chain first
shorter chains name
organic molecules that contain oxygen bonded to carbon
can't participate in hydrogen bonding
hydrogen replaced by alkyl group
alchool
OH group
hydroxyl
naming alcohols
-ol
number of carbons closest to functional group
OH
Alkyl chains
-yl
not main chain
alkane chain branched from alkane chain
same prefixes as alkanes
number longest carbon chain
lowest number at the bonds
then the alkyl chain
functional groups
grouped together when a group of atoms contribute in an interesting way
carbon bonded to a more electronegative atom
double/triple bond
carbon chains are unsaturated
kinks in chains
C=O
C-OH
C-Cl
impact properties of molecules
naming alkenes and alkynes
more than 3 carbon atoms
count the carbons
count closest to the bond
same as alkanes
first part
how many carbons in the longest chain
second part
double bond
triple bond
-ene
-yne
alkanes
alkanes have double or triple bonds
name it
count number of carbon atoms in longest chain
add -ane to the end
this will be prefix
first part
how many carbons in longest chain
second part
if there is any bonds in molecule
drawing organic molecules
lewis structures
referred as structural formula
condensed structural formula
CH3CH2CH3
shorthand
each end of a line is a carbon atom
no symbols
each line represents a bond
each carbon has hydrogen atoms bonded to it
carbon bonds
hydrocarbons
only hydrogen and carbon atoms
organic compound
compounds based on carbon chains
carbon
triple
from covalent bonds
single
double
structural isomers
molecular formula can be the same as another compund
structure is always different
non-polar
longer the chain...higher boiling point
halides
hydrogen atom switch with halogen
unsaturated
stereoisomers
different atom arrangement
same backbone
isomers
cis
same side of double bond
trans
opposite side of double bond
markovnikov's rule
how atoms add to the double bond
"hydrogen halide or water molecule reacts with an alkene, the hydrogen atom will generally bond to the carbon atom in the multiple bond that has the most hydrogen atoms already bonded to it"
aromatic
unsaturated
ring structure
naming
alkyl group
-benzene
benzene ring that lost a hydrocarbon
phenyl group
"phenyl" as substituent
properties
non-polar
insoluble in water
polymers
built from monomers
monomer is repeating units to form polymer
polyethene
homopolymer
formed by reaction of simple monomer
simplest one
copolymer
2+ monomers combined
addition or condensation reaction
protiens
silk
natural
living things
starch
cellulose
RNA + DNA
synthetic
look at natural polymers for properties
sourced from plants
polyester
rubber (tires)
disposing synthetic properties
harsh
don't break down quickly
engineering biodegradable synthetic materials
addition
very long
organic molecule
result of addition reactions
unsaturated carbon-carbon bonds
plastic bags
radical
break double bond
continues to get next monomer
chain continues
continues for the number depicted by 'n'
plastics
can be moulded
under heat/pressure
thermoplastics
from petrochemicals
properties
unreactive
stable single bonded
weak intermolecular forces
flexible when heated (usually)
heat increases molecular motion
unwanted plastics
condensation polymers, can be biodegradable
cross linking
form networks
allow hydrogen bonding between chains
more cross linking = tighter the chains are linked together
condensation
very long
organic molecule
result of condensation reactions
2 monomers with 2 functional groups
ester or amide links
polyesters
monomers with 2 functional groups
-OH group
-COOH group
must attach to end
ester linkage
growing chain. monomer keep attaching
polyester + water
2 hydroxyl groups and 2 carboxylic acid groups
dacron
PET
condensation copolymer
polyamides
condensation reaction
alcohol and carboxylic acid
nylon lab
natural
cellulose
polysaccharides
monosaccharide
ketone or aldehyde
with hydroxyl substituents
large polymer of monosaccharides
bond of glycogen
peptides
2+ amino acids linked
same as amide linkage
protiens
amino acids
-COOH
-NH2
Hydrogen
R-aditional group
structure
primary structure
secondary structure
hydrogen bonding between oxygen and carbonyl
shape is important
nucleic acid
DNA
polymer in cells nucleus
RNA
polymer in cell cytoplasm
makes up DNA and RNA
nucleotides
5-carbon sugar
nitrogen-containing organic base
phosphoric acid molecule
DNA or RNA combines with phosphoric acid
quantum matter
electron
J.J. Thomson
negative pole repelled the ray
all atoms must contain electrons
electrodes
plum pudding model
randomly 
radioactivity
spontaneous decay or disintegration of nucleus
uranium produces image on plate
elements that emit radiation
alpha, betta, gamma
atom
rutherford
alpha particle at gold foil
mostly empty space
nucleus
dense centre of atom with positive charge
proton
positive subatomic particle
neutron
neutral subatomic particle
isotopes
same number of protons different neutrons
nearly all the volume of an atom
radioisotopes
emits radioactive gamma rays
alpha or beta particles
nature
light
electromagnetic radiation
maxwell's theory
continuous wavelengths that form spectrum
photoelectric
electrons emitted by matter
then absorbs energy
Hertz
frequency of light was important in determining energy of emitted electrons
Planck's
matter can gain or lose energy
E=nhf
6.63x10^-34
measured and is a constant
quantum
burst/packet of energy
photons
unit of light energy
fundamental particle
electron can escape a surface
electron absorbs photon
atomic spectra
spectroscopy
analysis of spectra to determine properties of source
measure intensity of light
no light=spectrum
spectrum of hydrogen atom
advancing atomic theory
excited hydrogen atoms
release energy by emitting light in wavelengths
emission spectrum
electromagnetic radiation emitted by an atom
when atom is returned to a lower energy state from a higher one
continuous spectrum
contains all wavelengths in specific region
line spectrum
contains only wavelengths characteristic of the element being studied
unique to each element
when excited electrons move lower energy level
they emit a photon of light
transition
electron moves up or down energy level
Bohr model
first 18 elements
2, 8, 18 electrons
does not fully describes structure
electrons can't actually orbit nucleus
flame test lab
mechanical model
schrödinger's standing wave
electron has wave lengths
electron bound to a nucleus
resembles a standing wave 
nodes
starting and ending of each cycle
antinode
amplitude of wave
must be whole number
applied standing waves to hydrogen atom
n=1,2,3,4...
orbitals
region around nucleus where electron= high probability
heisenberg's uncertainty principle
impossible to know
exact position
exact speed
describe probability of finding electron in location
wave function
mathematical probability of finding an electron in a certain region of space
electron isn't moving around the nucleus in a circle
electron probability density
finding electron at given point
derived from wave equations
determine shape of orbitals
also a distribution
quantum mechanical model
electrons can be in different orbitals
based on quantum theory
probability for the location of electrons
quantum numbers
describe quantum mechanical properties of orbitals
principal (n)
describes size
describes energy
atomic orbital
whole-number values
spaces between atomic shells aren't equal
n increases, energy required for an electron to occupy orbital increases as well
subshells
s, p, d, f
different shapes
secondary (l)
shape of orbital
from 0 to n-1
n=1 only s orbitals exist
n=3 there can be s, p and d orbitals
no stable element higher than subshell f
magnetic (ml)
orientation
relative to other orbitals
whole number between +1 and -1 including 0
ml=number of orbitals possible
shapes/orientations of orbits
2s and 3s contain separate areas of 0 probability
nodes
p orbitals
not spherical
2 lobes separated by node at nucleus
xyz system
d orbitals
l=2
occur when n=3+
spin (ms)
+1/2 or -1/2
electron can spin in 1 of 2 opposite directions
pauli exclusion
"in a given atom, no 2 electrons can have same set of 4 quantum numbers (n,l,ml,ms))"
Atomic structure and periodic table
elements arranged according to way electrons arrange around the nuclei of atom
multi-electronic
determines chemical behaviour
kinetic energy of electrons as moving about nucleus
potential energy attraction between nuclei and electrons
potential energy repulsion between 2 electrons
can make aproximations
more effective electrons go through shielding electrons, lower energy of electrons in said orbital
electron configuration
location of electron energy levels
number of electrons in energy level
aufbau principle
theory that an atom is built by addition of electrons
fill orbitals starting at lowest available orbital
examine electron configuration of elements
describes relative energies of the electrons in an atom
hund's rule
orbitals of same energy
lowest energy configuration
pauli exclusion
maximum number of unpaired electrons
unpaired electrons are represented as having parallel spins
put electron in each p,d,f before paring them
all unpaired electrons should have the same spin
electron configuration
use periodic table to determine number of electrons in the atom or ion
assign electrons by main energy level
then by sublevel
using an energy-level diagram or an aufbau diagram
distribute electrons by main energy according to hund's rule
fill each sublevel before starting next sublevel
for anions
for cations
remove electrons
add electrons
periodic table
valence electrons 4s has lower energy than 4d orbitals
outermost principal quantum level of atom
transition metals
highest level are d orbitals
chromium is an exception
2 half s and d
4s^13d^5
lower energy
more stable
copper is an exception
n+1 fill before nd orbitals
chemical bonding
ionic bond
electrostatic attraction between oppositely charged ions
strong attraction
higher electronegative in one
crystal lattice
examine the valance electron configuration
isoelectric
having the same number of electrons per atom, ion, or molecule
molecular elements
hydrogen atoms are certain distance apart
covalent bonding
proton-electron force
proton-proton balance
electron-electron forces
chemical bond in which atoms share bonding electrons
bonding electron pair
pair involved in bonding
found in space between atoms
lewis theory of bonding
atoms and ions are stable
full valence shell
stable electron
when paired
chemical bonds
create full valence shell
sharing electrons
exchange of electrons
non-metals and metals
covalent bond
coordinate covalent bonding
lewis structure
arrangement of covalent electrons
polyatomic ion
duet rule
hydrogen and period 2 metals
octet rule
atoms tend to form most stable substances
surrounded by 8 electrons in valence shell
lone electron pair
valance electrons that are localized to given atom
exceptions
fewer or more than 8 electrons around central atom
electrons involved are from 1 atom
H+ with ammonia forms covalent bond with remaining pair of electrons
electrons come from nitrogen
3-D structure
arrangement of ions or atoms making pure substance
determines how pure substance will behave
VSEPR theory
determines geometry of molecule
electron pairs far apart as possible
electron-pair repulsion
occurs between electron pairs
positioned as far apart as possible
minimize repulsive force between electron pairs
2 bonding pairs
far apart as possible
3 bonding pairs
farthest at 120 degrees
lone pairs dont play factor in determining shape
multiple bonds
double and triple
shorter
stronger
pure substances are the same single or triple
1 more bonding pair of electrons
work
ionic and molecules
phoshine fails
add more rules
bond polarity
non-polar covalent bond
electrons are shared equally
2 identical molecular atoms
hydrogen
nitrogen
chlorine
polar covalent bond
electrons aren't shared equally
1 atom atracts more than other
electronegativity
ability to attract shared electron to itself
different elements have different electronegativities
bonds are all mixtures of covalent and ionic character
spectrum values
diople
separation of positive and negative charges in a region
arrow points to negative end
molecular polarity
polar
non-polar
net dipole
only non-polar bonds
bond dipole sum of zero
geometric shape
determines polarity
quantum mechanics
hybridization
electron density decreases uniformly
further from the center
electrons are concentrated in centre of benzene
valence bond theory
atomic orbitals overlap to form new orbital
with pair of opposite-sign electrons
lowest energy state is obtained
2s orbitals
covalent
1s orbitals now have opposite signs
any unfilled orbital can overlap with another unfilled
hybrid orbitals
process of forming hybrid orbitals
combination of 2 different orbitals
carbon atom
1s and 3 p
end with 4sp^3
all identical to each other
tetrahedral arrangement
can form bonds
cant form bonds
triple and double covalent bonds
between unpaired electrons from
s orbitals
p orbitals
sp
sp^2
sp^3
directly overlap
sigma bond
s orbitals
p orbitals
multiple bond
pi bond
a bond that is formed when sides of the lobes of 2 orbitals overlap
sp and p orbitals
90 degrees
triple bond
1 sigma
2 pi bonds
intermolecular forces
causes 1 molecule to interact with another
bond
van der waals forces
chemical bond within molecule
between 2 molecules
diople diople
london dispersion
hydrogen bonding
polar
position positive and negative ends near each other
attractive and repulsive
1% as strong as ionic and covalent
weaken rapidly
strong diople-diople
hydrogen bonded to
oxygen
nitrogen
fluorine
partially negative atom on nearby molecule
very high boiling points
remain liquids even at high temperatures
non-polar molecules
increase as molecular mass increases
all undergo this
electrons move around nucleus
non-symmetrical electron distribution
weak
dosent last long
polarization
form a dipolar charge distribution
liquids
not compressible like gas
higher densities
surface tension
resistance of liquid to increase surface area
capillary action
spontaneous rising of a liquid in a narrow tube
cohesive
adhesive
solids
composite material
composed of 2 or more distinct materials that remain separate from each other when solid
ionic crystals
crystal lattice
packed positive and negative ions
hard
brittle solids
electricity
high melting points
metallic crystals
closely packed atoms
electrostatic
free-moving electrons
not all have same properties
electron sea theory
electrons move freely
around positive nuclei
metallic bonding
hold nuclei and electrons together
metals
molecular crystals
composed of molecules
intermolecular forces keeping together
complex
neutral molecules
not conduct current
covalent network crystals
interwoven network
high melting point
super hard
buckyball
60 carbon atoms
hollow cage
carbon nanotubes
carbon atoms
graphite rolled in cylinder
energy
energy
thermochemistry
energy changes
physical or chemical change
ability to do work
amount of energy transformed over a distance (j)
potential
body or system
position
composition
kinetic energy
object doe to motion
relesed
combustion of gasoline
products have less potential energy than reactants
thermal energy
kinetic and potential
heat
transfer of thermal to a cooler object
temperature
average kinetic energy
law of conservation of energy
"cannot be created or destroyed"
no new energy can be made
chemical system
reactants and products being studied
open
closed
energy can leave
calorimeter
matter and energy can leave
matter cannot leave
isolated
matter and energy cant leave
endothermic
exothermic
releasing energy
absorbing energy
nuclear energy
fussion
2 nuclei combine
fission
neutron to split atomic mass
heat capacity
(c)
amount of thermal energy required to raise temperature of 1 g by 1 degree celsius
calorimetry
measuring thermal energy change
chemical or physical change
calorimeter
measure thermal energy changes
well insolated chamber
temperature change of water
change of thermal energy
bomb calorimeter
calculations
q=mc(delta T)
delta T=T final - T inital
enthalpy
total amount of thermal energy in substance
delta H
change
energy released/absorbed
h>0
h<0
endothermic
exothermic
molar
change associated with physical, chemical, nuclear change
bond energies
dissociation energy
energy required to break given chemical bond
positive value
type of atom dependent
as number of bonds increase
bond length shortens
hess's law
change in enthalpy in chemical process is independent
change in enthalpy is same regardless of conversion
"change of conversion of reactants to products is the same whether the conversion occurs in several steps"
rules
reversed chemical reaction
reverse delta H sign
formation
change in enthalpy that accompanies formation of 1 mol of a compound in standard states
(1) remains unchanged
(2) is reversed
usually solids
sometimes gasses
exothermic
endothermic
sources
fossil fuels
crude oil
coal
natural gas
efficiency
ratio of energy output to the energy input of any system
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reactions
rate
change in concentration of reactant
or product of a chemical reaction
chemical kinetics
area of chemistry that deals with rates of reactions
measuring
observing appearance of a product
tracking changes in gas volume
colour
mass
pH
electrical conductivity
intensityope
n vessel gas substance produced
pH meter paper or indicator
average reaction rates
change in reactant or product concentrations over given time interval
always positive number
graph
rise over run
instantaneous rate of reaction
rate of a chemical reaction at a single point in time
stoichiometric rate relationships
negative means consumption
positive means formation
nature
chemical property
relates to the behaviour of pure substance
concentration
rate of reaction increases at higher concentrations of reactants
surface area
reactants are more than one state of matter
greater surface area
faster the rate
temperature
reaction increases with temperature
double every 10 degree celsius increase and halved at every decrease of 10
presence of catalyst
substance that alters the rate of chemical reaction without permanent change
polyethene
uses metal catalyst
biological catalyst
living thing
enzymes
heterogeneous catalyst
reactants and catalyst are in different physical states
homogeneous catalyst
same physical states
collision theory
chemical reactions occur only if reactants collide
need proper orientation
enough kinetic energy
activation energy
minimum energy
activated complex
unstable arrangement of atoms
partially formed and unformed bonds
maximum potential energy point
transition state
alternative pathway for reaction
lower activation energy 
rate law
mathematical expression that allows calculation of reaction rate as function of concentration
constant
empirically
unique for single reaction at temperature
total order of reaction
sum of the exponents in rate law equation
steps
measure
test
compare
mechanics
series of elementary steps produce chemical reaction
involves entity collision
cant be simpler steps
cant be written from overall balanced equation
can be written from elementary step balance
chemical systems
equilibrium systems
until reactants run out
chemical
reactants and products reach constant concentrations
dynamic
balance between forward and reverse
occur simultaneously
as concentration increases
reactants decrease
position
relative concentrations of reactants and products
dynamic
reversible reaction
proceeds in both forward and reverse directions
closed system
final concentrations of gases at equilibrium are the same
determine concentration
ICE table
constant
K
defining law for given system
law
chemical state at equilibrium
temperature dependent
quantitative changes
chatelier principle
"chemical system at equilibrium is disturbed by a change in a property, system adjusts in a way that opposes change"
equilibrium shift
change in concentration
restore state
reversible reactions
shift when disturbed
energy as reactant or product
endothermic reactions
thermal energy removed
reactants decreased
exothermic reaction
toward product
to the right
ideal gas
no size
no attraction
obeys all gas laws
partial pressure
if gas was occupying the whole volume by mixture
inert gas
collisions that won't result in chemical reaction
quantitative changes
predict shifts
reaction quotient
product of concentration
Q
not necessary at equilibrium
instantaneous concentration
occur together at particular moment
solubility
quantity of solute
dissolves in solvent
ionic compounds
in water form dynamic equilibrium
solubility equilibrium
dynamic
between solute and solvent
acids and bases
theorys
arrhenius
acid produce hydrogen ions
base produce hydroxide ions
bronsted-lowry
acid is hydrogen ion donor
base is hydrogen acceptor
hydronium ion
water molecule that has accepted hydrogen ion
hydroxide is responsible for basic properties
pairs
conjugate acid
base accepts proton
conjugate base
acid loses hydrogen ion
acid-base pair
related by donating and accepting protons
amphiprotic
able to donate
or accept proton
acid and base
acid ionization
strong and weak
acid
strong
ionizes almost 100%
base
weak
partialy ionizes
oxyacid
acidic hydrogen atom is attatched to oxygen atom
organic
contains carbon
oxygen
hydrogen atoms
carboxylic acid
weak
partially reacts to produce hydroxide
strong
dissociates completely
water
acid or base
in the same reaction
autoionization of water
hydrogen atom transfer from one molecule to another
titration
lab preformed
addition of precise volumes
used to determine concentration of substance
buffer systems
buffer
aq solution
conjugate acid-base pair
constant pH
electrochemistry
oxidizing agents
reducing agents are oxidized when they reduce something
oxidizers are causing something to become oxidized
themselves reduced
cells and batteries
secondary
primary
fuel
battery
group of 2 or more galvanic cells connected in series
galvanic cell for which the reactants are continuous
cell that cant be recharged
needs an external source of electricity
can be recharged
activity series
more keen on loosing something less keen on gaining something
strongest oxidizing agents are weakest reducing agents
mnemoic devices
LEO says GER
OIL RIG
redox
simultaneously
reduction process
oxidation process