Chemistry AOS#1 Flashcards
atomic theory
- -> matter is made of atoms
- -> atoms cannot be broken down into smaller substances
- -> elements contain only one type of atom in fixed ratios
nano meter
one billionth (10^-9) of a metre
nanotechnology
the use of technology to manipulate and investigate properties of substances on a nano scale
conversion chart
multiply by 1000 (10^3)
m–> mm–> μm –> nm
nanomaterials
have different properties than the same material in bulk form
they have a greater surface area to volume ratio (increased exposure to the environment)
nanotubes
allow for water to leave but keeps the salt inside of them used in salt water –> filtration plants
atomic number
number of protons in an atom
assigned the letter Z
mass number
mass number is represented by the letter A
the number of protons and neutrons
isotopes
when elements have same protons but different neutrons
eg. carbon 12, carbon 13, carbon 14
isotopes have same atomic number but different mass number
they have similar chemical properties –> electron configuration is the same
they have different physical properties –> different mass
atomic emission spectrum
when atoms are heated they give off electromagnetic radiation or light
gives clues about electonic structure of an element
- -> atoms of the same element produce identical line spectra
- -> each element has a unique line spectrum because of their unique electronic structure
bohrs model
electrons revolve around nucleus in fixed circular orbits
- ->electrons orbits correspond to specific energy levels
- -> electrons cannot exist between two energy levels
- -> electrons in outershell have more energy than that of closer ones
excited state
electrons absorb energy and jumps to a higher energy level
- -> the closest the electron is to the nucleus, the more energy is needed to move them to a higher state
- ->when going back to the ground state, electrons release photons which is the light we see
subshells
within each shell, there are certain subshells with different energy levels
subshell fill with the lowest energy first
1s<2s<2p<3s<4s etc.
1s= lowest energy
4s = highest energy
the subshell ‘s’
one orbital
max number of electrons = 2
the subshell ‘p’
3 orbitals
max number of electrons = 6
subshell ‘d’
5 orbitals
max number of electrons = 10
subshell ‘f’
7 orbitals
max electrons = 14
electron configuration of chromium (24)
…. 3p^6 4s^1 3d^5
at the higher energy level, some electrons don’t want to be paired up, - negative charges repel - so they jump to the empty space in the next subshell
electron configuration of copper (29)
3p^6 4s^1 3d^10
one electron from 4s^2 moves to 3d subshell to avoid being parted up because both are negatively charged.
copper and chromium will only EVER have 4s^1
ions
have an imbalance between protons and electrons
groups
have the same physical and chemical properties. groups go down the periodic table
periods
represents the number of electron shells present
–> period 2 = 2 electron shells
core charge aka. effective nuclear charge
the measure of attractive force felt by valence electrons towards the nucleus
core charge = number of protons - number of innershell electrons (not valence)
group 1 has a low core charge –> want to give away their electrons
group 18 has a high core charge –> doesn’t want to give away their electrons
atomic radius
measurement of the size of the atom
- -> the distance from the nucleus to outermost electrons
- -> small core charge = larger atoms as electrons are not held as tightly towards the nucleus
electronegativity
the ability of an element to attract electrons towards itself
core charge trends
increases across a period, and no change down a group
electronegativity trends
decreases down a group, increases across a period
atomic radius trends
increases down a group (as there is more electrons), and decreases across a period
first ionisation energy
amount of energy required to remove one electron from an element in the gas phase
first ionisation energy trends
decreases down a group, increases across a period
mirrors the core charge going left –> right
first ionisation energy trends
decreases down a group, increases across a period
mirrors the core charge going left –> right
metallic characteristic trends
moving left to right on the periodic table, the metallic characteristics decrease
–> the number of electrons found in valence shell increases, therefore there is a less possibility of forming a positive ion
zinc
zinc is unlike other transition metals as it has a full 3d subshell. transition metals also form more than one cation, although zinc will only form a 2+ ion when two electrons drop from the 4s shell.
reducing strength of elements
the ability for an element to lose electrons
–> increases with the core charge along a period
trends of reactivity
reactivity along a period is like a parabola
–> chlorine and sodium is more reactive than iron
when it is easiest for atoms to lose electrons?
it is easiest when going down a group as the atomic radius increases, meaning the valence electrons are less attracted to the nucleus (further away) and electronegativity decreases
nuclear charge
aka. the actual charge on a nucleus
total charge found within a nucleus
–> increases across a period
–> number of protons
ionisation energy vs electronegativity
ionisation is the energy required to remove an electron from a gaseous atom, whilst electronegativity is the ability for the nucleus to attract an electron. it is essentially the opposite
why do metals have lower electro negativities?
as you move from left to right along a period, the number of protons increases by one, meaning the electronegativity will need to increase with the positive charge of the nucleus
characteristics of metals
- -> high melting points
- -> conducts electricity
- -> malleable (punchable and it holds its shape)
- -> ductile (can be stretched - think wire)
- -> high density (more matter)
- -> lustrous (shiny)
- -> low electronegativities (gives electrons away/low core charge)
- -> low ionisation energy (energy required to lose an electron)
properties of transition metals
d - block elements
crystals
for metal to bend, atoms must slide over each other.
Metals with neatly packed atoms will bend easier than those that are not.
A metal with large grains (perfect close packing as shown on ball-bearing model) will have fewer dislocations and bend easier than small grains
properties of s block metals
highly reactive
- -> s1 metals are highly reactive
- -> s2 metals are less reactive than s1 but still very reactive
properties of p block metals
have varying properties
properties of d block metals
- -> formed coloured compounds
- -> display magnetism
- -> act as catalysts
- -> high electric conductivity
- -> high melting points
reactivity of metals
can be determined by how readily they react with oxygen, water, steam, and dilute acids.
reactivity trends
less reactive as they move along a period
more reactive as they move down a group
metal + oxygen
metal + oxygen = metal oxide
–> 2Mg(s) + O2 (g) –> 2MgO(s)
metal + water
metal + water –> metal hydroxide + hydrogen gas
–> 2Li (s) +2H2O (l) –> 2LiOH (aq) + H2 (g)
metal + acid
metal + acid –> ionic salt + hydrogen gas
–> 2Li (s) + 2HCl (aq) –> 2LiCl (aq) +H2 (g)`
lustre of metals
the mobile electrons within metals allow for light to be reflected back into your eyes, creating a shiny, lustrous appearance
malleabiliy and ductile
- -> positive cations repel from each other, but the negative electrons ensure attraction.
- -> when bent hammered or pressed, the metal will still want to stay together because of the sea of electrons
- -> layers of atoms can move past one another without disrupting the force between cations and electrons when force is applied
- -> the non-directional nature of metallic bonds allows for this
electrical conductivity
one end of the metal becomes positive and the other negative. all electrons are attracted to the positive end.
–> the electron flow is an electric current
heat conductivity
when electrons gain kinetic/heat energy they are able to quickly transfer the heat to colder areas due to freedom of electrons
density
metallic lattices are closely packed together creating a higher density
melting point and hardness
metallic bonds are very tightly bound together so it is harder to break and boil.
–> melting points and hardness increase with the number of outer shell electrons
metals in the earths crust
are found in the form of mineral compounds
–> oxides, sulfides, silicates
only metals found in elemental states
only metals found in earths crust in their elemental states are platinum and gold
–> found in their raw form
iron ore
occurs naturally in the earths crust
–> in australia, iron oxides in iron ore are usually in the form of haematite (Fe2O3)
extracting iron from its ore
iron ore is mined from the groud and taken to the trucks for processing
–> causes damage to habitats, environment and culture
steps in a blast furnace
- coke reacts with oxygen to make carbon dioxide
- limestone decomposes to calcium oxide
- carbon dioxide reacts with more coke to form carbon monoxide
- carbon monoxide reacts with iron oxide in ore to make liquid iron
- calcium oxide formed in step two reacts with sand in the ore, leaving slag.
environmental impacts of iron production
loss of landscape
air pollution
noise pollution
disposal of slag
economic impacts of iron production
significant financial benefit
significant wage gaps
social impacts of iron mining
negotiations with indigenous people
land conflicts
workers have to fly to work in mines –> impact on family life
methods used to minimise environmental impact of iron ore
reducing water and energy consumption
limiting waste production
recycling waste materials
ways metals can be modified
through alloy production
heat treatment
coating
why do metals need to be modified?
the way a metal can be used is based on its chemical and physical properties
most metals need to be modified to be useful
alloy
metals that are melted and mixed with another substance
–> usually metal or carbon
interstitial alloy
small proportion of an element with significantly smaller atoms added to the metal
- -> added atoms sit in the gaps between cations making them not move as easy and hence stronger
- -> steel (iron [big] and carbon [small ones]) is an example of this
substitutional alloy
made of elements with similar size and chemical properties
- -> the added element takes place of metal ions in the lattice
- -> $2 coins is an example of this (copper, aluminium, and nickel)
annealed metals
metals are heated until they are red hot and then are cooled slowly
- -> results in larger crystals being formed
- -> metals are softer
- -> can be cut and shaped more easily
- -> they bend when pressure is applied
quenched metals
metals are heated until red hot and then are cooled quickly by being dunked in cold water
- -> smaller crystals are formed
- -> meaning metals become harder but more brittle
tempered metals
quenched metals are warmed again (not to red hot) and allowed to cool slowly
- -> it allows the crystals to settle
- -> reduces the brittleness of the metal but retains hardness
purpose of coating
coatings make metals more suitable for application
- -> decorative purposes
- -> functional purposes
- -> Both!!
noble coating
protection of metal by attaching a thin metallic layer to it
- -> metallic layer is less reactive than steel
- -> chromium is used as a noble coating on steel parts in cars and household items.
sacrificial coating
protection of metal from corrosion by attaching a more reactive metal to it
- -> when the metal is more reactive than steel
- -> A break in a sacrificial coating results in the formation of an electrochemical cell. The coating corrodes and the steel is protected.
- -> ex. zinc acts as a sacrificial coating on an iron roof. it reacts with CO2 in the air, forming zinc carbonate over the surface, protecting the metal and slowing the rate of corrosion
work hardened metals
beating the metal when cold
- -> causes crystal grains to become smaller
- -> bending is made more difficult now - work hardened
- -> example. bending a coathanger and then trying to bend it back to its original shape. it does not bend back in the area that has been work hardened.
- -> usually more brittle
uses of nano-silver
a useful antibacterial
gold nanoparticles
can be used as a targeted chemotherapy method
nano-iron
can remove carbon tetrachloride from rivers
–> can help control pollution
ions
atoms that gain or lose electrons to achieve more stable outer shell electron configurations
what charge does each group form
group 1 = +1 cation group 2 = +2 cation group 13 = +3 cation group 15 = -3 anion group 16 = -2 anion group 17 = -1 anion
ionic lattices
cations and anions form to make a 3D lattice
- -> held together by electrostatic forces of attraction between oppositely charged ions
- -> ions in the lattice mirror the ratio found in the formula
- -> CaF2 = 1:2 ratio of calcium to flourine
very high melting points of ionic compounds
electrostatic attraction is strong between ions and a lot of energy is required to melt them apart.
hardness and brittleness of ionic compounds
- -> cannot be easily scratched
- -> strong electrostatic attractions mean a strong force is required to disrupt the lattice
- -> once the lattice is disrupted, it will shatter easily due to repulsion of like charges
conducting electricity in an ionic compounds- solid, melted, aqueous
SOLID:
–> for a substance to conduct electricity, it must have moveable charges, but
ions in an ionic solid are fixed and have no free charged particles (ions are
charged
MELTED:
–> ions are free to move around and therefore are able to conduct electricity
because they have free charged particles
DISSOLVED:
–> ions are free to move around which means it can conduct electricity
electrolyte
a liquid substance that conducts electricity
why do ionic compounds dissolve in water?
water molecules are able to move between ions and free them, disrupting the structure of the lattice.
–> the attraction force between water and ions is stronger than the attraction
force between cations and anions
factors affecting crystal formation
sizes of crystals are caused by the diverse conditions in which they are formed
- -> the rate of evaporation affects the size of the crystal formed
- -> slow evaporation results in larger crystals
some uses of common ionic compounds
sodium hydrogen carbonate: cooking and cleaning
ammonium nitrate: fertilisers and explosives
nitrates, nitrites, sulfites: food preservation
sports drinks
as we sweat we lose electrolytes
–> electrolyte is a substance that dissociates to form ions and dissolved readily into water
hydrated ions
ionic compounds that contain water molecules bonded within the crystal
–> hydrates release water - which is a part of their structure - when they
decompose upon heating
example:
magnesium chloride hexahydrate = MgCl • 6H20
