SEM 1 EXAM Flashcards
how does nucleus stay together
electrostatic repulsion forces protons apart, but strong nuclear force attracts nuclear particles together
if nuclear force is balanced, atoms is stable
what happens if nuclear forces are unbalanced
then the nucleus will be unstable and decay over time
isotopes
atoms of an element with the same number of protons, but different number of neutrons.
chemical properties are similar (atom behaviour in chem reaction is based on arrangement and number of electrons.
different physical properties (mass and nuclear stability)
chemical and physical properties
how elements participate in chemical reaction
feature one can measure and observe
radioactive decay of isotopes
if forces in atom are balanced, atom is stable
if not balanced, atom is unstable and nucleus undergoes radioactive decay to become stable.
in radioactive decay, radiation is emitted, which is useful for radioactive dating and medical diagnosis and treatment
half-life
time taken for half the original sample to decay
relative atomic mass
compares the mass of atom to the mass of another atom
all atoms are compared to carbon 12
isotopes have different masses and abundances on earth
weighted average of the relative masses of the isotopes of an element relative to carbon12
Democritus
disagreed with the idea of infinitely divisible matter, proposed that matter consisted of tiny particles with nothing between them but empty space called “atomos”
dalton
elements composed of atoms
an atom of same element have a same properties, atoms of different elements have different properties
atoms not created nor destroyed or changed into different types during reactions
c reaction is separation, combination , rearrangement of atoms
compounds are combinations of atoms in specific ratio
thomson
found electrons
realised that atoms could be divided further
plum pudding model: electrons embedded in positively changed mass
rutherford
proposed the atom consisted mostly of empty space
electrons orbited around nucleus. most of mass was in the nucleus, all positive charge was located in the nucleus carried by protons
bohr
electrons were in electrons shells around the nucleus, each shell has a specific amount of energy
Chadwick
discovered the existence of neutrons that accounted for around half the mass of the nucleus, had neutral charge
Electrostatic attraction
related to distance the valence electrons are from the nucleus
strength of attraction of protons and electrons
shielding effect
related to number of inner shell electrons, shield the effect of the protons from valence electrons.
reduces ionisation energy
correlates to number of inner electron shells and attraction of valence electrons to nucleus
core charge
how defective the charge of the nucleus is at holding the valence electrons in
increases going towards noble gases
atomic radii
decrease going right
increase going down
electronegativity
ability of atom to attract electron
increases going right
decrease going down as electron is further from nucleus so it is harder to attract
ionisation energy
energy required to remove loosely bound electron
depends on
nuclear charge: more nuclear charge means more protons in nucleus which increases ionisation energy because valence electrons are tightly held on.
distance form nucleus: further distance decreases electrostatic attraction between valence electrons and nucleus. valence electrons are loosely held.
increases towards right
decreases going down because valence shell is further away making it easier to lose electrons
electron affinity
ability to accept electron and from a negative ion
metallic character
set of properties associated with metals
depends on ability to lose valence electrons
decrease going right
increase going down
ground state
shells next to nucleus are taken up by electrons
when electron is at its lowest possible energy levels
how electrons jump
atoms absorb energy, electron around nucleus gain this extra energy and move up to higher energy levels.
electrons can be so excited that they leave (ionisation)
energy levels are discreet. electrons can only exist in specific energy levels
electrons cannot exist between energy levels, therefore
amount of energy absorbed and emitted = difference in one energy level and another. electrons could be excited from one energy level to another by specific amounts of energy that corresponds to the difference in energy levels.
amount of energy absorbed by any sample of one element is the same
excited state
an atom in which electrons occupy higher energy levels than the lowest possible energy levels.
as it moves back to ground state, emits light
atomic emission spectroscopy
absorption and emission of light help identify element
process of analysing light emitted by electron as it return to ground state
spectroscope: takes emitted light and separates it into its component wavelengths to produce a line emission spectrum.
sample to analysed is heated to much higher temperatures, and light emitted is passed through a prism. the prism disperses light into its component colours. the monochromatic allows single wavelengths to pass at a time, then spectra is recorded
fingerprint
AAS
uses absorption of light by electrons to measure how much of an element is present. used AAS
- element analysed is determined: there could be many element in a sample but we need to focus on one. thus the lamp is made of the same element being tested. current passes through gaseous sample of element in lamp and emits light. has unique wavelengths to other elements
- vaporised: changing substances in sample to atoms. when light passes through, only the element being tested will absorb the light because it has the same energy levels as atoms that emitted light from lamp.
- detection: light passed through sample enters monochromator-> selects one wavelength of light for analysis by detector-> measures intensity of light and turns it into a number.
compared to known samples by constructing a calibration curve. concentration against absorbance value.
absorbance value
measure of amount of light that passed through sample without being absorbed, lets you know concentration
what is mass spectrometry
based on different masses of atoms in sample
determine what elements are present in a sample or what isotope and their abundance in a sample.
determines mass relative to carbon12 and calculates relative abundance of the isotope
determine isotopic composition of an element.
method of mass spectrometry
sample vaporised then…
1. ionisation: sample is bombarded with high energy electrons/ UV LIGHT.
removes valence electrons leaving atom with positive charge
does this with ALL atoms.
- acceleration: cations accelerated through electrical field at very high speeds according to mass-charge ratio
- deflection: pass through a magnetic field where they undergo deflection according to their mass. ions separated according to mass and charge
lighter= more deflection - detection: detectors measure amount of ions that strike them
graphed in a mass spectrum.
why does lamp pulsate in atomic absorption spectroscopy
light from the lamp is pulsed so that the detector can distinguish between the light from the cathode lamp that is left over after absorbance and the continuous light naturally emitted gaseous atoms in the flame/burner returning back to their ground state.
finds extent of lamp light being absorbed and to what extent
material
substances that make other objects. often mixtures of many substances. can be pure (elements and compounds)
pure substance
Pure substances are defined as substances that are made of only one type of atom or molecule.
properties cannot be altered
distinct and measurable properties for any given arrangement of atoms or molecules
pure substance vs homogeneous mixture
homogeneous mixture: components can be separated
pure substances: can’t be separated
elements
substances made of one type of atom
compounds
pure substance made up of more than one type of atoms
chemically combined, uniform composition
pure substances
properties cannot be altered. properties are distinct and measurable for any given arrangement of atoms or molecules
mixtures
properties can be changed depending on how much of each component is added mixture.this useful as properties can be controlled.
individual components keep their properties
variable composition
alloy
mixture of metal with other metals or small amount of non metals
e.g. iron is soft an sprone to corrosion. but adding some carbon creates steel which is stronger and corrosion resistant
polymer
material with a molecular structure that is composed of many repeating smaller units bonded together.
compared to metals, they are less dense, electrical insulator, corrosion resistant
natural: wool, silk, paper
synthetic: polystyrene
ceramics
inorganic, non metallic solid. composed of metals, nonmetals, metalloids held together by ionic and covalent bonds.
can be highly ordered (crystalline)or highly irregular (amorphous)
natural: kaolinite (makes porcelain)
synthetic: silicon carbide (used as an abrasive)
properties: hard, high compressive strength, able to withstand high temperatures. most are good insulators, but some have semi/super conducting properties
composite materials
combination of two or more distinct materials with significantly different chemical and physical properties.
creates a range of properties that would be unobtainable with using one of the individual components.
e.g. reinforce concrete = concrete matrix +embedded steel bars
low tensile strength of concrete counteracted with high tensile of steel
maintains high compressive strength of concrete
heterogenous mixture
non uniform mixtures that contain physically separate materials
not pure or uniform
mixtures can have different proportions of the same components .
eg. granite has 3 different minerals, each piece has a different amount but it is still granite
homogenous material
materials that have uniform composition throughout
breaking it smaller = would be identical
homogenous solution
consists of solute and solvent, solute is distributed throughout the solvent as very small particles, so it appears uniform
nanoscale
structure between 1 and 100 nm
nanometer = 10 ^-9
nanomaterial
substances (natural and synthetic) that are composed of single units that exist in ten nanoscale
very large surface area compared to the volume they occupy
adsorption of nano-materials
where molecules stick to surface of of a solid or liquid
can be used to remove unwanted chemicals and gases
large surface area means that a small volume of nanoparticles can adsorption a large number of molecules
transportation of nano materials
can transport molecules that are adsorbed
small size and large surface area mean that they can transport chemicals through air, skin and even cells
used in chemotherapy treatments for cancer as it can transport drugs to specific cells
nano materials as catalysts
can speed up rate of reaction without being altered or destroyed in the process. provide surface for reaction to occur
reactant molecules absorb onto the surface of the nanoparticles which allows reactant molecules to combine to form the product
large sa means many reactions can occur at once. ->increases rate of reaction
gold as nanoparticles
properties can change as nanoparticles
gold changes its colour
melt at lower temperatures and appear red in solution
zinc oxide in sunscreen
good absorbers of uv radiation and appear transparent because the particle size is much smaller than the wavelength of visible light, so light can go through it with very little being affected
other metal oxides tend to be milky white
having nanoparticles means that it’s distributed evenly so uv radiation is absorbed/scattered by zinc oxide. if they were bigger it means that it wouldn’t be evenly distributed and up rays could pass through
Nanoparticles in medicine
colloidal gold: labeling antigens, optimises distribution of drugs to cells in difficult areas (brain, retina, tumours), targets tumors and provides detection
silver: antibacterial and antifungal properties, due realease of silver ion from surface of metal
sa enhances antibacterial effectiveness
safety risks of nanoparticles
can travel through air, bloodstream, skin, cells
can interact with inside body (biomolecules) to cause unwanted reactions
composite nanomaterial
small size and unique properties are useful for them to be composite nanoparticles
stain resistant cotton: cotton fibres covered with water resistant nanoparticles
tyres: added carbon improved resistance and abrasion. increase electrical conductivity, prevents build up of static electricity
fullerenes
3d structures formed by networks or carbon atoms
nanotube: cylindrical tube, reinforcement in composite nanomaterial
graphene: flat 2d layer of carbon atoms arranged in hexagons
buckyball: soccer ball shaped 3d structure
top down method to make nanoparticles
size of material is progressively reduced , by grinding until size is achieved.
+: large quantities produced, cheap, uniform
-: limited to simple structures
sieving
based on particle size
separates mixtures of different particles, smaller particles pass through holes, large ones are trapped
filtration
particles not dissolved in solvent
separates suspended solid from a liquid using a filter funnel and paper.
Evaporations
particles dissolved in liquid
boiling/evaporating solvent away from solute
distillation
differences in boiling point
separates two liquids with different boiling points or a soluble solid and the solvent
distillation flask, condenser, receiving flask
fractional distillation
vapour passes through a column packed with glass beads. less volatile component component condenses on the glass beads and drips back to distillation flask
can be used to separate miscible liquids
separating funnel
difference in density
immiscible liquids separate into layers due to density
tap is opened to let the more dense component out of the mixture
metallic bonding
chemical bonding that results from attraction between metal atoms and the surrounding sea of electrons
metals have low ionisation energy that is why they lose electrons easily
high MP and BP of metals
forces of electrostatic attraction between metal cation and delocalised electrons are very strong so it requires more energy to break them
why do metals conduct electricity
delocalised electrons are free to flow through the metal and so carry a current.
why are metals malleable
bonding in metals is not rigid. as metals are metal with a force, the atoms slide through electron sea to new positions while continuing to maintain their connections to each other. this also makes the ductile.
attractive forces stronger than repulsive.
non directional bonding
why are metal dense
particles are closely packed together
why are metals shiny and opaque
freely moving and delocalised electrons are present so metals can reflect light and appear shiny.
close packing of cations prevent light from slipping through making it opaque
what does a greater core charge in a metal mean in terms of bonding
atoms packed tightly with stronger bonds
metallic bond structure
- cations are closely packed in 3d network. cations occupy fixed positions in lattice
- delocalised electrons moving freely. they belong to the lattice as a whole, not an individual atom
- these electrons come from valence electrons not inner shell ones
- cations held in position due to ESF of attraction between cations and electrons (metallic bonding)
limitations to metallic bonding
structure can't explain: variations in properties magnetic nature differences in electrical conductivities *more complex model needed
ways to modify metals
alloy production
heat treatment
formation of nanoparticles
metal alloys
mixing metals with other substances (metal/carbon). substances are melted, mixed and cooled.
harder and lower MP: since atoms of different sizes are now included, the properties may differ. lattice doesn’t move in the same way (lower MP)
iron and carbon
steel
harder and less corrosive, but less malleable because atoms are slightly different in size and lattice can’t move past each other as easily.
ionic bonds
when a metal and non metal come into contact and a electron is lost from the metal and given to non metal. cations and anions formed. held together by electrostatic forces of attraction between ions (ionic bonding)
ions have inert gas electron configuration
each cation is surrounded by anion and vice versa (attractive forces outweigh repulsive forces)
high MP and BP of ionic bonds
large amount of thermal energy to overcome attraction between oppositely charged ions and allow them to move freely
bricks in furnace made of MgO
why are ionic compounds brittle and hard
strong forces of electrostatic attraction hold ions together, so a strong force is required to break them
brick in houses CaPO4
when a force is applied ions move in the direction of the force, like ions are forced together causing them to repel (repulsion causes it to shatter)
ionic compounds and electrical conductivity
solid: not free to move, can’t conduct current
ceramic insulators
solution/liquid: lattice dissociates in water into charged ions that move freely in solution
cations -> cathode
anions -> anode
ammonium chloride in dry cell batteries as electrolytes
ionic compounds and solubility
soluble: ions break away and mix with water molecules
insoluble: ions remain bonded and don’t form a solution
depends on:
1. attraction between cation and anion
2. ion and water molecule
why are metals good conductors of heat
cations vibrate vigorously and are able to transfer thermal energy to each other
delocalised electrons readily transfer thermal energy as they move through the lattice
why do metals form cations
low ionisation energy
interstitial alloy
when atoms of element that is supposed to be mixed with metal is much smaller and fill spaces between lattice
Substitutional alloy
larger, similar property metal added to lattice
covalent molecular bonds
sharing of electrons to attain full valence shell. molecules are discrete (individual)
if atoms have similar electronegativities than they are like to form covalent bonds because they have the same affinity for electrons and don’t want to donate.
positively charged nuclei attracted to shared electrons
low MP of covalent molecular bonds
Intermolecular forces are weak so not much energy is required to break them
non conductive of covalent molecular
no mobile particles
electrons locked in covalent bonds
softness of covalent molecular
intermolecular forces are weak
molecules can be moved out of position easily
hence soft
covalent networks and properties
only intramolecular forces made form network of repeating lattices of covalently bonded atoms (intramolecular) eg silicon dioxide, carbon, silica properties: hard high MP + BP solid @room temp non conductive (except graphite)
allotrope
different forms of an element
atoms bonded together in different specific ways
have slightly different properties
diamond
carbon allotrope
hard -> not discrete in covalent network lattice
each carbon surrounded by 4 other carbon
high thermal conductivity because atoms are held together strongly
graphite
carbon allotrope
slippery, soft, greasy
conductive (free moving electrons)
covalent layer lattice ( each carbon has 3 other carbons and 1 delocalised electron per carbon atom)
each sheet is slightly positive and pushes down electrons which means there is weak attraction.
% composition
how much M, m, n of each element make up the total chemical
Properties of gas
can be compressed diffuse low density no fixed shape or volume (diffuse) measured by volume in L
things that affect behaviour of gases
temp
high temp = high energy
increasing temp increases volume and pressure if volume is kept the same
pressure (kPa)
increase in pressure decreases volume and increase temp as particles now have higher energy
volume (space taken up by gas)
increase in volume decreases pressure and temperature because of lower energy and less collision as particles are further apart
relationship between volume and pressure at a constant temp
Boyles law
inversely proportional
as one increases, the other decreases
bike pump example
STP
standard temp and pressure
t: 0ºC or 273.15ºK
p: 100 kPa
1 mol of any gas at STP will occupy 22.71 L
flame test procedure
- Place a small amount of salt in a clean watch glass.
- Add to 1-2 mL of ethanol to the sample and stir with a toothpick
- Place the watch glass on a heat proof mat
- Carefully light the ethanol
- When the flame burns low observe the colour produced
- Repeat for each salt using a clean watch glass each time
examples of things that aren’t materials
chemicals like: HCl, chlorophyll, carbon dioxide
sedimentation
centrifugation
gravitational separation, pouring of liquid from undissolved settled solid
speeds up sedimentation by spinning very quickly, even very small dissolved solids are settle
solutions and compounds are both?
homogenous
solutions can be physically separated because its not chemically bonded, just dissolved. however it does have a uniform composition
Solutions consists of solute dissolved in solvent
why do nanomaterial properties change
as they become smaller, surface area to volume ratio increases which changed properties because majority of properties occurs at the surface
non octet covalent molecules
Be and B are small atoms with only 2 and 3 electron pairs, because it is small enough, arrangements are stable enough for molecules to still form without following octet rule.
boron has 3 valence, but still makes BF3, now has 6 valence, molecule forms even though it doesn’t follow octet rule.
P and S are larger atoms that can accomodate 5 or 6 bonding pairs
kinetic energy and gas
KE can be transferred from one particle to another but the total KE remains the same
collisions between gas particles are elastic collisions, KE is conserved
pressure (gases)
force exerted on surface by gas particles as they collide with the surface
force/area
uses of AES
analyse metallic elements in solid samples
analyse trace elements in soil, water ,etc
uses of AAS
identify toxic metals in urine or blood
pollutants in soil or water
quantity an prescience of precious metals such as gold and silver
uses of MS
drug testing and discovery
food contamination detection
sample of hydrogen atoms absorbing light
if exposed to continuous spectrum of visible light, hydrogen will only absorb certain colours, these colours have energies that perfectly correspond to energy needed for electron to move up energy levels.
dark lines are colours absorbed
why are there several absorption/emission lines, if hydrogen only has 1 electron
spectrums represent the collective emissions of many individual hydrogen atoms
chemical reactions
when particles collide and are rearranged to form new particles. chemical reactions involve energy changes.
as the reactant particles are rearranged, the chemical energy of the reactants is changed also.
energy can be absorbed or released
chemical energy
stored in chemical bonds between atoms and molecules. energy results from:
attractions between electrons and protons
repulsion between nuclei
repulsion between electrons
movement of electrons
vibration and rotation around bonds
law of conservation of energy
states energy cannot be created or destroyed.
chemical energy stored in a substance reduce and the energy must go elsewhere
a substance can’t gain in chemical energy without absorbing that energy from another source
system
the chemical reaction
refers to energy changes that occur as bonds and formed between the atoms of the elements involved in the reaction
surroundings
everything but the chemical reaction
energy goes from the surrounding to the reaction or energy is released into the surroundings
energy changes
the reactants in a reaction have a certain amount of energy in bonds, since the products are a rearrangement of the particles so they have different bonds and thus different amount of energy
for particles to separate,
energy is required, the separated articles have more energy than when they were together. particles coming together the products will have less energy than the separate particles, this lost energy is transferred to the surroundings
specific heat capacity
energy needed to change temperature of 1kg of a substance by 1 degree
exothermic
energy of products is less than energy of reactants and the lost energy is transferred to its surroundings
endothermic
when energy of the products is greater than the energy of the reactants
energy is absorbed from the surroundings
melting or boiling
endothermic
freezing
exothermic
activation energy
energy required to break the bonds of reactants so that a reaction can proceed
energy barrier that must be overcome before a reaction can start
unless this minimum energy amount is met, reactants rebound and move away from each other without reacting
combustion
the release of chemical energy when the fuel is burnt in the presence of oxygen
exothermic
complete combustion
when oxygen is plentiful and products are carbon dioxide and water
incomplete combustion
when oxygen supply is limited
not all carbon can be converted into carbon dioxide so carbon monoxide and or carbon is produced instead
energy content of fuels
heat of combustion of a fuel is defined as enthalpy change when a specified amount of fuel burns completely in oxygen
energy content: chemical energy available from a substance
fuels have high energy contents
factors that affect reaction rate
surface area of solid reactants concentration of reactants in a solution gas pressure Temperature presence of catalysts
rate of reaction
change in concentration of a reactant or product per unit time
need to measure amount of product being formed or amount of reactants being used up in a given ti8me period
depends on activation energy
magnitude of activation energy determine how easy it is for a reaction to occur and what proportion of collisions result in a successful reaction
collision theory
for collision theory to occur reactant particles must
collide with each other
collide with each other with sufficient energy to break bonds between the reactants
collide with the correct orientation to break the bonds between reactants and so allow formation of new products, if not correct orientation, particles bounce off each other
transition state
a new arrangement of atoms when activation energy is absorbed
occurs at the stage of maximum potential energy in the reaction (activation energy)
bond breaking and forming occur at this stage and arrangement of atoms is unstable
increase reaction rate
increasing a number of collisions that can occur in a given time
increasing proportion of collisions with an energy equal or greater than the activation energy
increasing frequency of collisions
increasing concentration of reactants (increasing pressure increase concentration)
increasing surface area of a reactant (more particles, more exposed surfaces of particles )
increase pressure
increasing energy of reactions
increasing temperature
catalyst lower activation energy required
importance of shape of molecule
affects vapour pressure, melting point, boiling point and solubility. shape determines how it interacts with other molecules
VSEPR
uses knowledge of valence shell to predict shape
based on principle that negatively charged electron pairs in valence shell repel each other. meaning that these electrons are arranged as far away from each other as possible
electron pairs in the single covalent bonds repel each other
lone pairs are treated the same as electron pairs in covalent bonds because they still take up space int he atom
double and triple bonds are treated as a single bond
tetrahedral (methane)
electron pairs in the single covalent bonds repel each other
repulsion forces the bonds as far apart from each other as possible, so hydrogens are arranged in a tetrahedral shape at an angle of 109.5º
pyramidal (ammonia)
lone pairs are treated the same as electron pairs in covalent bonds because they still take up space int he atom
3 hydrogen atoms for a pyramidal shape with the nitrogen. the ,one pair occupies more space because it isn’t covalently bonded to a and is therefore closer to the nitrogen. this means that the 3 covalent bonds are pushed closer together making the angle less than 109.5º
trigonal planar (water)
2 lone pairs and 2 covalent bonds
4 electron pairs repel each other and 2 hydrogen atoms for a v shape with the oxygen. because of the two lone pairs the 2 covalent bonds are pushed closer together
angle around central atom less than 109.5º
120º
linear (HF)
fluorine has 3 lone pairs and 1 covalent bond
four pairs form a tetrahedral arrangement due to mutual repulsion
results in a linear model
180º
BH3
BeH2
3 electron pairs, no lone pairs
trigonal planar
2 electron pairs, no lone pairs
linear
strength of intermolecular forces
covalent molecular forces have lower melting and boiling points because forces BETWEEN molecules are weaker. these forces are disturbed when covalent substance change state
vapour pressure
pressure that gaseous molecules exert on the closed container walls when the rates of evaporation and condensation become equal
liquids with stronger intermolecular forces have lower vapour pressure molecules are held together tightly making it harder for them to escape from the surface of the liquid
increasing temp increase vapour pressure. molecules have a higher average kinetic energy that enables them to overcome intermolecular forces. as temperature increases, more particles have enough energy to escape from liquid to gas. increasing pressure
boiling point of a liquid
temperature at which the liquids vapour pressure reaches the atmospheric pressure of the surroundings
electronegativity and polarity
in intermolecular forces, the electrostatic attraction is between positive and negative charges in the molecules
charges result because of uneven electron distribution within molecules
electronegativity is the tendency of an atom in a covalent bond to attract electrons
in a covalent bond atoms are competing for electrons being shared between them
non polar diatomic molecules
when atoms have identical electronegativities, then electrons are shared equally between them. non polar because there is no charges on either end
even distribution of valence electrons between atoms
if electronegativities between atoms are very similar, then it can be considered non-polar
electron density
measure of the probability of an electron being present at a particular location within an atom
polar diatomic molecules
2 atoms with different electronegativities, electrons stay closer to the most electronegative atom because they are more strongly attracted to that atom
polar: imbalanced electron distribution. has two oppositely charged poles
all diatomic with different atoms are polar to some extent, greater difference in electronegativities the greater the polarity
delta
partial charge, adding up partial charges = 0 because charge are due to uneven sharing of electrons, so molecules are still neutral
negative: has larger share of electrons, excess of electrons
positive: lost some of its share of electrons
dipole
separation of positive and negative charges
polyatomic molecules polarity
symmetrical balanced dipoles= non-polar
methane: carbon is more electronegative but individual dipoles cancel each other out. results in a molecule with no overall dipole
in non-symmetrical molecules, individual dipoles don’t cancel each other, creating a net dipole making the molecule polar
dipole dipole forces
only in polar molecules
results from attraction between positive and negative ends of polar molecules
weak since partial charge are small
more polar the molecule, the stronger the force
stronger force the higher the melting and boiling points because force bond molecules together in solid or liquid. higher energy is required to to break stronger bonds
hydrogen bonding
only occurs in molecules in which a hydrogen atom is covalently bonded to a N,O or F
NOF are small and very electronegative so they strongly attract electrons in a covalent bond. creates a large partial positive charge on hydrogen atom, which is attracted to lone pairs of electrons in neighbouring NOF atoms
hydrogen in covalent bond with NOF is is exposed due to lack of electrons and is highly attracted to lone pairs of NOF in other molecules
strong intermolecular bond, x10 than dipole dipole, but 1/10 of ionic
dispersion forces
forces of attraction in non polar substances (existence of intermolecular forces, without them nothing will hold molecules together and they would be gases)
caused by temporary (instantaneous) dipoles in molecules that result in random movement of the electrons surrounding molecule
strength increases as size of molecules increase. larger molecules have more electrons so it is easier to create temporary dipoles. larger molecules have higher melting and boiling points
molecules that form long chains have higher dispersion forces, more contact area to interact with it neighbouring molecules to form stronger forces
how to increase polarity of a moelcule
larger difference in electronegativities
high degree of asymmetry in the shape of the molecule causing an imbalance in the bond dipole
requirement of hydrogen bonding
- hydrogen atom covalent bonded to an NOF atom
- lone pair of electron of NOF atom of neighbouring molecules
H bonds are caused by NOF’s small radius and high electronegativities and single electron of hydrogen atom
larger radii reduce concentration of negative charge around them.
hydrogen bonding in water
ice floats because it is less dense
in ice water molecules are held in a crystal lattice. in which molecules are held further part then in liquid water
when ice melts crystal lattice collapses and molecules pack together more tightly
crystal has greater volume, 4 hydrogen bonds to 4 neighbouring molecules
electrostatic attraction
strong in water because of large electronegativity difference meaning larger partial charge
water properties
can dissolve many substances easily high specific heat capacity high latent heat of vaporisation expands on freezing react with CO2 to make glucose and oxygen
water shape
lone pairs of water have greater repulsion than electron pairs in covalent bonds so they push the hydrogen atoms closer together
going down hydrides (except water)
melting and boiling points increase
Intermolecular forces are stronger (dispersion)
dispersion forces increase with strength as mass increases
water vs hydrides
hydrogen bonds in water give it a much higher boiling and melting point
high surface tension of water
water molecules at the surface are not completely surrounded by other water molecules, so they form hydrogen bonds everywhere BUT above.
molecules in body of water have no net force but surface level molecules have a downwards force. water molecules at surface are so attracted to each other by strong hydrogen bonds
surface tension meaning
measure of the resistance of a liquid to increasing its surface area
heat capacity of water
measure of a substances capacity to absorb and store heat energy, how much energy a substance absorbs as temp increases
SHC: measure of how much energy is required to increase temp by 1º of 1 gram of substance
SHC of water is relatively high because of hydrogen binds. these bonds can absorbs and store large amount of heat energy before they break
latent heat
when a something reaches its boiling/melting point the temp remains constant until all of substance changes state even though further energy is being absorbed
energy still being absorbed while state is changing is latent heat
energy absorbed by a fixed amount of substance as it changes state at its melting (fusion)/boiling (vaporisation)point
depends on strength of intermolecular forces
LHF of water is 6.0KJ/mol. energy required to disrupt ice lattice by breaking some hydrogen bonds
LHV of water is 44KJ/mol.
dissolution
particles of solute separate
particles of solvent separate
particles of solute and solvent are attracted to each other
forces:
holding solute, holding solvent (H Bonds), form between solute and solvent
like dissolves like
H bonds between water i such stronger that dispersion forces in non polar. attraction between water molecules cannot be overcome and water doesn’t separate to form a solution
water dissolving covalent
molecular compounds that form H bonds:
strength of intermolecular forces are similar, so they readily interact with each other
the more polar, the more likely it will dissolve
molecular compounds that ionise:
highly polar molecular compounds form ions when dissolved in water
polar covalent bonds break producing ions, covalent bond forms between ion and water molecule (makes hydronium in HCl), ion dipole attraction between ions and water
water dissolving ions
ion dipole: attraction between ion and polar molecule
if ion dipole attraction are strong enough, the water molecules can pull out ions from crystal lattice and into the solution
ionic compound–>hydrated ions (dissociation)
energy required to separate ions from lattice must be less than energy released when ions are hydrated
solubility
maximum amount of solute that will dissolve in a given amount of solvent
increasing temp increases solubility, at higher temps solute and solvent have more energy to overcome forces of attraction that hold particles together in the solid
saturated: no more solute can be dissolved at a certain temp
unsaturated: contains less solute needed to make solution saturated
supersaturated: contains more dissolved solute than a saturated solution
crystallisation
when an unsaturated solution becomes saturated and crystals form.
1. cooling: may reduce the solubility of a dissolved solute to the point where not all of the substance available is soluble. slow cooling = larger crystals
2. evaporating the solvent from the solution
falser rate of evaporation = smaller crystals
gases in water
Solubility decreases with rising temp
deoxygenates water so water must be cooled before released into environment
chromatography
technique used to separate and analyse the substances present in a mixture. it can be used to analyse numerous organic and inorganic substances.
eg: contaminants in water, toxic gases in air, additives and impurities in food, drugs in blood
chromatography only works for small amounts of different substances because the solvent can only carry small amount of the mixture with it.
how chromatography works
has stationary phase (solid/liquid), and a mobile phase (liquid/gas). mobile phase moves up stationary phase due to capillary action.
as components of the mixture are swept upwards over the stationary phase by the solvent, they undergo a continual process of adsorption onto the solid stationary phase, followed by desorption and dissolving into the mobile phase. the ability of component molecules to adsorb depend on the polarity of stationary phase and molecules. attraction between component and solvent molecules depend on the polarity as well.
components in chromatography
chemicals in a mixture
chromatogram
output of a chromatography procedure
in TLC and paper: patterns of bands or spots formed on a plate or on the paper
in HPLC: graph produced
rate of movement in chromatography depends on
how strongly the component adsorbs onto the stationary phase
how readily the component dissolves in the mobile phase
paper and TLC
qualitative analysis
paper: the stationary phase is high quality absorbent paper
TLC: the stationary phase is a thin layer of a fine powder (aluminium oxide) spread on a glass or plastic plate
small spot of solution of sample is placed on one end of chromatography paper/plate, spot is called the origin. the paper/plate is placed in a container with the solvent. origin is a bit above the solvent level, so that components can be transported up the paper/plate and not be dissolved in the solvent.
as the solvent rises up the paper/plate, the component of each sample separate depending on their attraction tot the stationary phase and their solubility in the solvent
Identifying components of a mixture
chromatography
- include standards of known chemicals on the same chromatogram as the unknown sample and comparing the resulting positions of the unknown components with those of the known samples
- calculate the retardation factor of the sample and comparing these with the values of known samples
RF= distance component travelled/ distance solvent travelled
*from the origin
component most strongly adsorbed to the stationary phase moves the shortest distance and has lower RF value
solvent front
describe the movement of the solvent during chromatography. it is visible as the wet moving edge of the solvent as the solvent travels up the stationary phase
retardation factor
each component has a characteristic RF value for the conditions under which the chromatogram was obtained
thing that change RF value:
changes in temp
type of stationary phase
amount of solvent vapour around paper/plate
type of solvent
colourless compounds
many organic compounds fluoresce and appear blue when viewed under UV light
chromatogram can be sprayed with a chemical that reacts to form coloured/fluorescent compounds
paper v TLC
paper: cheap little preparation more efficient for polar and water soluble compounds easy to handle and store
TLC:
detects smaller amounts
better separation for less polar compounds
corrosive materials can be used
wide range of stationary phases can be used
only use a give a guide to identify the chemical, since for. a particular combination of stationary phase and mobile phase, many different chemicals may have similar RF values.
column chromatography
the solid stationary phase is packed into a glass column. the sample is applied carefully to the top of the packed column, and a solvent which acts as a mobile phase is dripping slowly onto the column from a reservoir above
a tap at the bottom of the column allows the solvent (eluent) to leave the column at the same rate it enters at the other end
paper and TLC summary
components undergo a continual process of adsorption to stationary phase and desorption back into mobile phase
components undergo adsorption and desorption to different degrees depending on the strength of their attraction to the stationary and mobile phases. thus components separate as they move past the stationary phase
HPLC
based on column chromatography
allows sensitive analysis of a wide range of mixtures
separation and identification of complex mixtures of similar compounds (contaminants soluble in water, drugs in blood, hydrocarbons in oil samples, presence and concentration of dioxins; insecticide, pesticides and oil spills in water)
particles in the solid used in HPLC column are 10-20 times more smaller than column chromatography, allows more frequents adsorption and desorption of components giving better separation
small particle size in HPLC creates a considerable resistance to flow of mobile phase, so solvent is pumped through column in high pressure
range of solids can be used in HPLC, some with chemicals specially bonded to their surfaces to improve the separation of particular classes of compounds
components are detected by passing the eluent stream through Uv light beam. many organic components absorb UV light, so when an organic compound passes in front of the beam of light, a reduced signal is picked by a detector. amount of light received by detector is recorded on a chart that moves slowly at a constant speed or sent to a computer. resulting trace Is called a chromatogram. each component forms one peak in the chromatogram.
can separate compounds with relative molecular mass 1000 or even more
gas chromatography
most sensitive technique
limited to components that can be readily vaporised without decomposing (relative molecular mass <300)
analysis of trace contaminants, or tiny amounts of very potent components (drugs in urine)
mobile phase is an inert gas (nitrogen), as a carrier gas
small amount of sample is injected into the top of the column through an injection port. the port is heated to vaporise the sample which is swept into the column by the carrier gas
the column is a series of loops of glass/metal (length=2-3m, diameter=4mm), it is heated. the column is packed with a porous solid coated with an ester/liquid hydrocarbon with a high BP; or packed with and absorbent solid such as silica gel/ alumina. this is the stationary phase
components repeatedly interact with the stationary phase and are swept forward by the carrier gas. components that adsorb least strongly to the stationary phase are swept out first by the gas. as components emerge from the end of the column, they are sensed by the detector.
application of GC and HPLC
qualitative: what chemical is present in the sample
quantitative: how much of each chemical is present
qualitative: how much and purity of component
a solution of a pure compound that is thought to be one of the components is injected into column under same conditions. chromatogram is compared.
the same compound will have the same retentions time if conditions are kept the same
can also be identified by adding known compound to sample (spiking), creates a much bigger peak in the chromatogram
Quantitative: to determine concentration of a component, its peak are air compared with peak areas of samples of the same chemical at known concentrations
standard solution: has an accurately known concentration
calibration curve (peak area against concentration) determines unknown concentration
retention time
time taken for component to pass through column
characteristic of the component for the conditions of the experiment.
carbon
forms more compounds than all elements combined
- 4 valence electrons, 4 needed
- strong covalent bonds with other carbons
- bonds can be single, double or triple
crude oil
when dead animals, plants and micro-organisms were buried by sand, the organisms accumulate as sediment and gradually become part of earth’s crust
increase in temp and pressure make oil and fat into hydrocarbons
crude oil is a mixture of hydrocarbons. it has a low density so it migrates up but I slacked by impervious rock
accumulation of oil and gas creates oil field
refining: not used in raw state, fractional distillation. crude oil separated to components made of range of hydrocarbons with similar BP heavier components (fractions)-> cracking breaks larger hydrocarbons molecules results in two smaller molecules one of which has a carbon-carbon double bond
alkane
hydrocarbon = only carbon and hydrogen
hydrocarbon with only single bond
it is saturated: only carbon-carbon single bonds
weak forces, non polar (increase in size increase dispersion force and MP+BP)
CnH2n+2
structural isomers same formula different arrangement
akyl group: CnH2n+1
methyl propane: methyl must be on second carbon
homologous series
series of compounds with similar properties and same general formula
where each member has one CH2 unit than the previous member
- similar structure
- pattern to physical properties (BP)
- similar chemical properties
- same general formula
alkenes
double bond (2 pairs of electrons shared) unsaturated: one more double/triple bonds
alkene: homolog series of hydrocarbons with a carbon-carbon double bond
CnH2n, more reactive, non polar (not soluble in water)
geometric isomers: same order, different arrangement. atoms in double bond are fixed unlike single bonds
benzene
C6H6
stable has a ring structure
aromatic hydrocarbon
electrons that make up double bonds aren’t in fixed positions and can move through ring
carbon-carbon bond are same length, molecule is more stable than expected of unsaturated molecule
bonds between carbon are intermediate between single and double bonds
substitution reaction
alkane and halogen (with UV light)
carbon-hydrogen bond broken replaced with carbon-halogen bond
benzene and halogen (with catalyst AlCl3 or AlBr3)
halogen can be in excess or limited
addition reaction
addition of small molecule to double bond of alkene
- 2 reactant to 1 product
- carbon-carbon double bond to a single bond
- unsaturated to saturated
- small molecule added across double bond, forms bond on each end of double bond
alkene + H2O
makes an alcohol
hydration reaction because water is a reactant
ethene + H2O -> ethanol
uses catalyst
at 300ºC, gaseous reactants, solid catalyst, gaseous products (easy to separate catalyst and product)
alkene + hydrogen gas
in presence of metal catalyst (Ni)
forms alkanes and saturated
hydrogenation
etene + H2 -> ethane
too slow to proceed at room temp without a catalyst
alkene + halogen
halogen adds across the double bond
ethene and bromine -> 1,2-dibromoethane
proceeds at room temp without a catlyst
alkene + hydrogen halide
HCL, Her, HF, HI
adds across double bond
- asymmterical alkene + asymmetrical reactant -> isomers made
unsaturation test
bromine water to determine if substance has double bonds (unsaturated)
- aqueous Br2 solution added to alkane has no reaction and no ocular change
- Br2 and alkene has an addition reaction and orange colour disappears
substitution reaction is much slower than this addition reaction at room temp without catalyst
addition reaction has an instant change
hydrocarbons and combustion
can be incomplete or complete
precipitation reaction
occurs if ions in solution combine to form a new compound that is insoluble in water
compound=precipitate
when one solution is added to another mixture formed will contain all of the ions. ions move independently and as they move, they collide with one another. if cations and anions collide they can join together to form a precipitate
hydrated ions -> attraction between ions -> ionic lattice
concentration
describes relative amount of solute and solvent present in the solution.
grams/L
ppm: mg/kg
M: mol/L
*solute/solution
dilution
adding more solvent to a solution, decreases concentration
C1V1=C2V2
soluble base
alkali
properties of acids
turns litmus indicator red corrosive taste sour react with bases ph less than 7 conduct electric current
properties of bases
turns litmus blue caustic and feel slippery bitter taste react with acids ph greater than 7 solutions conduct an electric current
acid
substance capable of making hydrogen ions in solution, or donating a hydrogen ion
base
substance capable of producing hydroxide ions in solution or accepting a proton (hydrogen ion)
Arrhenius model of acid
acid is defined as a substance that is ionised in water to produce hydrogen ions
ionisation: removal of one or more electrons from an atom or ion. reaction of a molecular substance with a solvent to form ions in solution
polyprotic acid
acids that react to form more than one hydrogen ion per molecule. molecules have more than one proton that can be ionised.
in general hydrogen atoms that are bonded to a highly electronegative atom (polar bonds are ionised in solution)
diprotic acid
triprotic acid
donates two hydrogen ions when ionised
donates three protons when ionised
hydronium ion
single proton can’t exist in water, H ion attracted to lone pair of electrons on a water molecule
(H3O+)
Arrhenius model of bases
substance that dissociates in water to form hydroxide ions. presence of hydroxide account for common properties of bases
acids and bases as electrolytes
conduct electricity because of free moving ions with charges
HCl: ionises in water to form H and CL ions
neutralisation
according to Arrhenius model: when acids and bases react the hydrogen and hydroxide ions combine to form water
adjustment of acidity
if a solution of metal hydroxide and acid is added together, hydroxide ions react with the hydrogen ions.. once all hydroxide ions have reacted with hydrogen ions (forming water) that is when acid and base have neutralised
limitations of Arrhenius model
doesn’t explain why some substances without hydrogen can form acidic solutions when mixed with water
co2 and sos make acidic solutions
nh3 and NaHCO3 mixed with water make basic solutions
this is explained by bronsted Lowry theory: acid donates hydrogen ion, and bases accepts this ion
strong and weak acids
strong (acids that ionise readily, complete ionisation, solutions contain ions with virtually no unreacted acid molecules remaining)
HCl, H2SO4, HNO3
weak (partial ionisation, only small amount of acid molecules ionise)
acetic acid
strong and weak bases
strong (completely dissociate)
Group 1+2 oxides and hydroxides
weak (only small proportion ionised at any instant)
ammonia, metal (phosphates, carbonated and hydrogen carbonate )
strength vs concentration
concentration: how much substance is dissolve in solution
strength: to what degree acid is ionised or base is dissociated
pH
range of hydrogen ion concentration on a scale to measure acidity
pH decreases as H+ ions increase
acid and metal carbonate
lime water test: confirms presence of CO2 gas, saturated solution of calcium hydroxide. when Coz is bubbled through the solution, it will turn milky/cloudy due to precipitation of CaCO3
Ca(OH)2 + CO2 -> CaCO3 + H2O
