sem 2 exam (unit 1 content) Flashcards
avogadro’s number
6.022 x 10^23 particles
mole
a precisely defined quantity of matter, equal to Avogadro’s number of particles/atoms/molecules/formula units
atoms
the smallest, indivisible building block of matter that can exist stably and independently, uniquely defining a chemical element
subatomic particles
Particle Location Relative Mass Charge
Proton (p+) Nucleus 1 +1
Neutron (n0) Nucleus 1 0
Electron (e-) Electron Cloud 1/1836 -1
isotopes
variations of same element with different numbers of neutrons
- same chemical properties e.g. reactivity, bonding, electron structure
- different physical properties e.g. density, mass, half-life, size, MP/BP
relative atomic mass
the weighted average of all elemental isotopes (calculated using relative isotopic abundance), as relative to 1/12 of the mass of a carbon-12 atom
mass number
the mass of one specific isotope of an element, generally in whole numbers (unlike relative atomic mass which is in decimal)
mass spectrometry
an accurate instrumental technique used to measure relative isotopic mass and relative abundance within a sample, and thus calculate relative atomic mass
steps of mass spectrometry
- vaporisation: sample enters the spectrometer in a gaseous form, after being vaporised in a vacuum chamber
- ionisation: the vapour is passed through a high-energy electron beam, where collisions with the beam result in loss of one (or sometimes two) electron/s, thus forming cations
- acceleration: resulting cations are accelerated by an electric field to form a high speed beam of positive ions
- detection: the high speed beam of positive ions are directed through a strong magnetic field, perpendicular to the ion’s path, as generated by the electromagnet, where ions are deflected into circular paths of different radii based on mass; lower mass / lighter ⇒ more deflection ⇒ smaller radius
- deflection: ions are collected with the current measured, then being graphed as relative abundance (y) over m/z mass to charge ratio (x) ⇒ most cations formed have +1 charge so the m/z ratio is usually numerically equal to mass (m) of the various ions; height of peak in graph is actually the relative intensity (proportional to relative abundance)
atomic absorption spectra
element-specific frequencies of electromagnetic radiation (light) at which energy is absorbed when transitioning up from a ground to an excited state
electron promoted to a higher level
continuous spectrum of light has specific frequencies (black lines) of light missing; gaps between the energy levels (absorbed by atom)
atomic emission spectra
element-specific frequencies of electromagnetic radiation (light) at which energy is emitted when transitioning down from an excited to a ground state
only specific frequencies of light emitted as coloured lines on a black background; the frequencies shown = energy emitted
flame test (basic emission spectroscopy)
heat a sample of chosen substance using a flame, thus exciting the e-
since the excited state is unstable, the electrons eventually drop back to ground state
energy is emitted as photons at characteristic frequencies
combinations of photon frequencies produce coloured light
the light is viewed through a spectroscope monochromator
emission lines can be matched to identify the element (QUALITATIVE)
uses: identification of unknown metals (metallic cations), firework displays, flares
atomic absorption spectroscopy
element must be previously determined through other qualitative methods (since this method determines concentration using a hollow cathode lamp of the same element)
sample is vaporised through an atomiser, resulting in a flame with hydrocarbon fuel, oxidants, and a gaseous form of the tested element
light from the lamp passes through the atomised sample; only the element tested would be able to absorb light; frequencies corresponding to energy levels as atoms
unabsorbed light is focused through a slit and enters a monochromator, separating wavelengths of interest
the selected wavelength goes through a detector, which numerically depicts the intensity of the light (i.e. measures the unabsorbed light), producing an absorbance value
absorbance value is directly proportional to elemental concentration
using many samples with known elemental concentration, a standard calibration curve is created using line of best fit, and the unknown sample’s absorbance value is compared, using interpolation to determine its concentration
strengths of flame tests
- quick and easy test for metal atoms
- convenient (easy to access materials)
weakness of flame tests
- qualitative data only (subjective)
- only a small range of metals are detectable with a flame test
(emissions may not be on the visible light spectrum) - metals in low concentrations may be difficult to observe
- mixtures of metals will produce confusing results
- used a standard flame
=> luminous with orangish hue that may obscure emitted colours
=> perhaps not hot enough to achieve proper excitation of metal
advantages of AAS
- quantitative data (comparable, easy to analyse)
- can handle/process mixtures of many metals
- highly selective for one metal to be tested
- can test larger range of elements
- very sensitive to low concentrations
ionisation energy
energy required for the process by which atoms lose electrons and ionise IN A GASEOUS STATE
electronegativity
the ability of an atom in a molecule to attract a pair of electrons in a covalent bond towards itself, depending on atomic radius and number of unshielded protons
first ionisation energy
energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of ions (M → M+ + e-); measure of strength of attraction between valence electrons and the nucleus
atomic radius
distance between nucleus and valence electrons
valency
a measure of an atom’s bonding capacity (combining power)
bonding
the forming of chemical bonds (either ionic, covalent network, covalent molecular, or metallic) to attain stability by having a full outermost valence shell
ion
atoms, or groups of atoms that are electrically charged due to the loss or gain of electrons
ionic bonding
After ions are formed (due to the gain or loss of electrons) to attain noble gas configurations, the ionic particles become electrically charged, yet are perfectly stable particles, capable of existing independently. However, whenever ions exist within a specific distance to each other (specifically a positively charged cation and a negatively charged anion), the electrostatic attraction between positive and negative charges is what holds the ion together to form an ionic lattice. An ionic lattice is a giant, theoretically endless crystalline structure, consisting of a consistently repeated pattern of cations, surrounded by anions, surrounded by anions, and so on.
properties of ionic compounds
- high MP/BP
- hard, brittle
- thermal and electric conductivity (when liquid, not solid)
water of crystallisation
water chemically bonded into a crystal structure of ionic salts (completely embedded as a full H2O molecule between gaps)
hydrous salts
contain water
anhydrous salts
don’t contain water
metallic bonding
metals atoms held together by electrostatic attraction in a rigid 3D lattice of positively charged, metallic cations, surrounded by a nondirectional, delocalised, mobile electron sea
properties of metals
- malleable, ductile
- thermal and electric conductivity
- high MP/BP
- metallic lustre
covalent bonding
occurring between two or more non-metals, involves the directional sharing of electron pairs, resulting in the electrostatic attraction between the positive nuclei of the atoms and the shared electron pair
properties of covalent molecular compounds
- low MP/BP
- non-conductivity
- soft and brittle
covalent networks
covalently bonded lattice structures, formed with Group 14 elements, that exist in theoretically endless, repeating patterns.
common examples are diamond (C), graphite (C), and silicon dioxide/sand (SiO2).
properties of covalent network compounds
- high MP/BP
- non-conductivity (exception: graphite)
- hard, brittle (exception: graphite)
vsepr theory
each atom in a molecule will achieve a geometry that minimises the repulsion between electrons in the valence shell of the atom, specifically by maximising the 3D angle of repulsion
allotropes
variations of elements with different physical forms, and thus significantly different physical and structural properties
nanomaterials
materials that utilise nanoparticles (any particle with at least one dimension in the 1 - 100nm size range, where 1nm = 1 × 10^(-9) metres)
examples of uses of nanoparticles
UV Blocking (ZnO, TiO2): photostable, used in sunscreen lotion
=> visible white layer, but invisible/colourless when nanoparticles
Nanosilver: antibacterial, antifungal properties
=> easier to infiltrate cellular processes and destroy bacteria
=> bandages, masks, filtration, personal health products, cosmetics
Quantum Dots (ZnS, ZnSe, CdSe): biological tracers
=> emit size-dependent, wavelengths of life
=> only fluorescent and coloured when nanoparticles
Carbon Nanotubes (network structure made of graphene):
=> thin sheet, excellent conductivity => effective at carrying currents
=> smaller, cooler, more efficient computers
safety concerns of nanoparticles
unknown extents of risks with nanoparticles
possibly contaminating waterways, soil, even accumulating in cells
intermolecular forces (vs intramolecular forces)
intermolecular forces (e.g. dipole-dipole forces, dispersion forces, hydrogen bonding) occurs BETWEEN molecules, whereas intramolecular bonding (e.g. ionic, covalent, metallic bonding) occurs WITHIN the molecules
intermolecular forces are significantly weaker, and influence properties such as melting points, boiling points, solubility
ionic: incredibly strong electrostatic attractive forces between cations and anions (INTRAmolecular forces); made of formula units (lattice)
metallic: incredibly strong electrostatic attractive forces between cations and anions (INTRAmolecular forces); metallic lattice
covalent network: sheer number of covalent bonds between particles are INTRAmolecular forces, the entire substance basically being one whole molecule
covalent molecular: the only substances in which both intramolecular (very strong) and intermolecular forces (very weak) are at play
dispersion forces
temporary attractive forces that arise in all substances, since electrons are constantly moving, and thus not always symmetrically distributed, resulting in a very temporary/momentary, weak “instantaneous dipole”
factors that influence dispersion forces
larger surface area ⇒ maximised (potential) proximity between molecules ⇒ stronger dispersion forces
heavier (more protons ⇒ more electrons), more polarisable:
greater number of electrons ⇒ more chance / likelihood of an instantaneous dipole, AND the instantaneous dipole would be stronger due to a greater negative charge
heavier (in a container) ⇒ denser, more likely to sink or settle together ⇒ increased proximity ⇒ stronger dispersion forces
number of atoms ⇒ increased proximity
bond polarity
electronegativity = electron-attracting power of an atom when involved in a covalent bond
- covalent molecular bonding occurring between different atoms of different properties
- a difference in electronegativity between atoms
- uneven distribution of electrons (elements with higher electronegativity tend to “steal”)
- causes an unequal distribution of charge
- results in a bond being polar (areas of partial, temporary charge)
factors that influence molecular polarity
two identical atoms ⇒ same electronegativity ⇒ equal distribution of electrons ⇒ equal distribution of charge ⇒ “pure covalent” bonds ⇒ nonpolar bonds
molecular shape ⇒ symmetrical polyatomic molecule with polar bonds ⇒ polarity “cancels” itself out ⇒ bonds are polar but overall molecule is nonpolar
dipole-dipole forces
interactions between polar molecules (with a partial charge): electrostatic attraction (and repulsion) between physically aligned covalent molecules, experiencing a temporary, relatively weak Dipole Moment
factors that influence dipole-dipole forces
stronger molecular dipole ⇒ stronger dipole-dipole force
larger molecule ⇒ more electrons ⇒ more chance of dipole moment ⇒ stronger forces
greater surface area ⇒ maximised physical alignment
hydrogen bonding
hydrogen bonding is a specific type of dipole-dipole bonding, containing NOF-H (nitrogen/oxygen/fluorine)-hydrogen bonds, with incredibly high polarity since these bonds are specifically very polar (NOF extremely electronegative)
thus, these are much more stronger, due to high electronegativities AND how the hydrogen is attracted to highly negative lone electron pairs of other molecules
factors that influence hydrogen bonding
smaller atomic radius ⇒ electrons confined to smaller spaces ⇒ incredibly dense ⇒ stronger partial charge ⇒ stronger polarity ⇒ stronger hydrogen bonding
relative strength of intermolecular forces
hydrogen bonding > dipole-dipole forces > dispersion forces
melting and boiling point (with intermolecular forces)
stronger dispersion forces (sometimes dipole-dipole forces, hydrogen bonding) ⇒ more energy required to overcome this electrostatic attraction and intermolecular forces ⇒ increased kinetic energy ⇒ higher melting and boiling point temperatures
vapour pressure (with intermolecular forces)
weaker intermolecular forces ⇒ lower boiling points ⇒ increased vapour pressure and volatility (more in gaseous form)
solubility (with intermolecular forces)
“like dissolves like” (similar types of intermolecular forces in solvent and solute)
the sum of intermolecular bonds in the mixture is at least as strong as the sum of the intermolecular bonds of solute ⇒ polar dissolves well in water (polar)
pure substance
substances made of only one type of molecule with a fixed / constant / uniform composition
well-defined, constant physical and chemical properties
e.g. fixed boiling and melting points, predictable reactivity
elements
pure substances made of 1 type of unique atom, defined by number of protons in nucleus
molecules
any particle comprised of two or more elements that are chemically bonded (doesn’t have to be different elements e.g. diatomic molecules)
compounds
any particle comprised of two or more different elements that are chemically bonded
mixtures
substances containing 2 or more different substances in varying proportions, physically combined
properties vary with composition, depending on identity and quantity of components → “average” or properties
homogenous mixtures
uniform composition (same phase, looking even)
e.g. solutions (solute doesn’t settle in solvent, yet can be physically separated)
heterogenous mixtures
nonuniform composition (physically separate, visible sections)
e.g. granite (mineral grains), milk (fat globules suspended in water), toothpaste (solid particles in liquid), cereal in milk, oil+water
solutions
homogenous mixtures with uniform dispersion of solutes (that dissolves) and solvents (that it dissolves in; larger quantity)
transparent/see-through, though not always colourless
explain how differences in physical properties of substances in a mixture can be used to separate them
can utilise the difference in physical properties (e.g. solubility, melting/boiling points for state changes, particle sizes, magnetism) of mixture components to separate components
e.g. filtration, decanting, use of separating funnel, recrystallization, distillation, fractional distillation.
filtration and evaporation
separation based on solubility (type of sieving)
mixture of an insoluble (s1) and soluble (s2) substance:
1. add water and stir to dissolve s2, leaving s1 undissolved
2. filter residue s1 out
3. evaporate remaining filtrate to leave s2 (remove solvent/water)
decanting
process of separation of liquid from solid and other immiscible (non-mixing) liquids, by removing the liquid layer at the top from the layer of solid or liquid below; the process can be carried out by tilting the mixture after pouring out the top layer.
sieving and use of separating funnel
separation based on particle size
passing mixture through sieve of suitable mesh size (or separating funnel lined with filter paper), allowing smaller particles to pass whilst separating coarser grains
concentrates a desired component (not completely accurate)
recrystallization
separation based on varying solubilities for temperatures
heat, dissolve, then gradually decrease temperature, collect solidified pure crystals
distillation
separation based on volatility and boiling points
generally separating mixture of 2 liquids by heating (based on boiling points; volatility);
=> when heated, more volatile substance evaporates first (low BP)
=> the vapour rises, then separately condenses and is collected, whilst less volatile substance is collected in the starter flask
fractional distillation
separation based on volatility and boiling points
separating hydrocarbon components of crude oil, by heating in the base of a fractionating column (most volatile substances vaporise and rise first, collected at top)
water’s unique properties
- anomalous melting and boiling points
- relative density of solid and liquid states
- surface tension, viscosity, cohesive and adhesive forces
anomalous melting and boiling points
very strong intermolecular bonding occurring between molecules (dispersion forces, dipole-dipole forces, and specifically, hydrogen bonding with oxygen and lone electron pairs) ⇒ unusually large amount of energy required to break bonds ⇒ high MP/BP
relative density of solid and liquid states
stronger IMF => packs together more effectively
define cohesive and adhesive forces
explain water’s surface tension
cohesive forces = attractive forces that cause molecules in the same substance to “stick” together and maintain a certain shape of the liquid
[water is extremely cohesive → relatively strong hydrogen bonds, IMF)
adhesive forces = attraction between molecules of different substances
surface tension = measure of forces required to stretch or break the surface of a liquid (due to strong cohesive forces)
cohesion > adhesion at surface
water molecules on surface have net downwards cohesive force (lacking particles above)
reduced surface area, liquid forced to be as small as possible
formation of solution
solution formed by:
1. breaking bonds in individual substances
(endothermic → requires sufficient energy)
2. forming bonds between the solvent and solute
(exothermic → releases energy)
must be favourable (releasing more energy than required);
bonds formed should be equal, if not stronger, than the initial bonds
dissolution and ion-dipole forces
dissociation = ionic substance splits into separate constituents (independent mobile ions)
ionisation = molecule forms ions (creation of charges)
ion-dipole forces: attraction between fully-charged ion and partially-charged dipole => occurs in dissolution, between the formed ions and water
saturated
the theoretically possible limit of dissolved solute at a given temperature
unsaturated
contains less solute than the solution is capable of dissolving
supersaturated
contain more dissolved solute than saturated solutions (i.e. more than theoretically possible) at a given temperature.
relationship between solubility and temperature (SOLID)
solubility INCREASES with temperature (due to extra energy available)
more kinetic energy → particles able to move more → intermolecular forces weaken → can be overcome with (less) additional energy → dissociated and ionised into water more
exception = cerium (III) sulfate
relationship between solubility and temperature (GASEOUS)
solubility DECREASES with temperature (due to extra energy available)
more kinetic energy → particles able to move more → intermolecular forces weaken → more difficult to be “held” by water molecules, “leaves” the water
limewater test for presence of carbon dioxide
bubbling “mystery” gas through limewater, i.e. dilute Ca(OH)2
if carbon dioxide is present, clear colourless solution becomes cloudy, indicating the formation of a white solid precipiate [fails test; present]
Ca(OH)2 + CO2 -> CaCO3 + H2O
water pollution
any physical or chemical change in water that adversely affects the health of humans and other organisms, varying in magnitude by location
types of water pollution
sewage
disease-causing agents
sediment pollution
inorganic plant and algal nutrients
organic compounds
inorganic compounds
radioactive substances
thermal pollution
sewage
release of wastewater from drains or sewers
disease-causing agents
infectious organisms that cause diseases (from waste of infected individuals)
sediment pollution
excessive amounts of suspended soil particles
inorganic plant and algal nutrients
chemicals that stimulate growth of plants and algae (nitrogen, phosphorus)
organic compounds
chemicals containing carbon atoms e.g. oil
inorganic compounds
contaminants that contain elements other than carbon e.g. heavy metals
radioactive substances
containing unstable isotopes that spontaneously emit radiation
thermal pollution
heated water (processed during industrial processes) released into waterways
purification of drinking water:
- sewer lines bring sewage to treatment plant
- primary treatment: removing suspended/floating particles by mechanical processes
- secondary treatment: treating wastewater biologically to decompose organic material
- tertiary treatment: reducing nitrogen, phosphorous, sewage sludge (advanced)
chlorination
killing disease-causing organisms to prevent waterborne diseases
fluoridation
killing disease-causing organisms, preventing tooth decay
chromatography
group of techniques that separate individual components of a mixture by moving it through a stationary phase and a mobile phase, based on polarity (solubility)
substances of the solution either adsorb to stationary phase or desorb to mobile phase, based on intermolecular forces ⇒ polarity ⇒ solubility (like bonds with like)
stationary phase
solid (or liquid supported on solid)
mobile phase
liquid/gas flows through the stationary phase, carrying mixture components
influential factors in chromatography
- different types of polar groups [molecules preferentially bonding with phase that has the most similar bonding type]
- the amount of charged and polar groups present [more “like” bonding to SP → easier for analyte to absorb and remain in it longer → increased Rf, reduced Rt]
- molecular weight [larger → more likely to bond with stationary phase and remain in it longer→ reduced Rf, increased Rt]
- molecular geometry [greater surface area → more likely to bond with stationary phase and remain in it longer → reduced Rf, increased Rt]
paper chromatography
SP: paper (cellulose, highly polar, with hydrogen bonds)
MP: water (highly polar, with hydrogen bonds)
ink progresses up the filter paper, moved by water due to capillary action
separated by solubility and particle size → furthest up is most soluble, polar, smallest
thin layer chromatography
separation of a liquid mixture, where the stationary phase is a layer of solid particles spread on a flat plate
SP: very thin layer of highly polar absorbent material (e.g. silica gel), bonded to a glass or plastic support, with many microscopic plates (i.e. larger surface area)
MP: solvent that the mixture is in (where the analyte is dissolved)
advantages of TLC over PC
glass plate is more rigid than flexible paper, so easy to control
after separation, substances can be easily recovered: scraping the silica gel of the glass plate → adding it to solvent → dissolving → removing the silica gel by filtration to leave the analyte in solution
glass plates can be recoated
gas chromatography
separation of a gaseous mixture into individual components on passing a gas flow through a thin silica column
(small, non-volatile organic molecules - withstand high temperatures)
SP: column lined with silica gel
MP: inert gas
- sample injected into machine, where it is vaporised
- washed over the matrix by an inert gas
- some substances will be more attracted to matrix than others, having reduced speeds (adsorbing to SP, then desorbing to MP)
- substances reach detectors at varying times, which measures abundance of substance at given time, and then plots this data on a graph ⇒ QUALITATIVE analysis (Rt)
high performance liquid chromatography
separation of a liquid mixture into individual components on passing a liquid through a steel column packed with different small particles of stationary phase
(larger organic molecules, volatile and unstable to heat → no oven)
SP: column (shorter than that of gas chromatography)
MP: liquid (not a gas)
same process as GC without an oven, and a detector for QUANTITIATIVE analysis:
- running a set of standards (”knowns” of identified concentrations)
- interpolation to find estimated sample concentration corresponding to peak area
normal phase
stationary POLAR, mobile NONPOLAR
reverse phase
stationary NONPOLAR, mobile POLAR
