bonding, structure and properties of matter Flashcards
ions
atom with a positive or negative charge
atom that has lost/gained an electron
- can be single atoms or groups of atoms
when atoms lose/gain electrons to form ions, they are trying to get a full outer shell
- atoms with a full outer shell are very stable
the number of electrons lost/gained is the same as the charge on the ion
metals form ions
lose electrons from their outer shell to form positive ions (cations)
non-metals form ions
gain electrons in their outer shell to form negative ions (anions)
groups lost likely to form ions
1 & 2, 6 & 7
elements in the same group have the same number of
outer electrons
elements in the same period have the same number of
electron shells
ionic bonding
transfer of metals
metal and non-metal
electrostatic attraction between oppositely charged ions
when metal + non-metal react, the metal atom loses electrons to form a positively charged ion and the non-metal gains these electrons to form a negatively charged ion
- these oppositely charged ions are strongly attracted to one another by electrostatic forces (an ionic bond)
dot and cross diagrams
ionic bonding
draw one
show how ionic compounds are formed
show the arrangement of electrons in an atom or ion
- each electron is represented by a dot or cross
- diagrams can show which atom the electrons in an ion originally came from
pros
- show how ionic compounds are formed
cons
- don’t show the structure of the compound
- don’t show the size of the ions or how they’re arranged
ionic compounds
a giant structure of ions - giant ionic lattice
- held together by electrostatic forces of attraction between oppositely charged ions
- ions form a closely packed regular lattice arrangement
- strong ionic bonds require lots of energy to break
- there are very strong electrostatic forces of attraction between oppositely charged ions in all directions in the lattice
- the structure is regular and repeating (a ‘lattice’)
- crystal structure
e.g. sodium chloride (table salt)
ionic compound properties
- high melting + boiling points - many strong bonds between ions lots of energy required
- solids at room temp
- regular crystal structures
- lots of energy required to break the strong forces of electrostatic attraction that make the ionic bonds between oppositely charged ions
- don’t conduct electricity when solid - the ions are held in place
- -conduct electricity when liquid - the ions are free to move + will carry electric charge
- some ionic compounds also dissolve in water - the ions separate and are all free to move in the solution, so will carry electric charge
ball and stick model
shows the regular pattern of a giant ionic lattice (ionic crystal)
pros
shows how all the ions are made
suggests that the crystal extends beyond what’s shown in the diagram
cons
model isn’t to scale (relative sizes of the ions aren’t shown)
in reality there aren’t any gaps between ions
covalent bonds
the sharing of electrons between 2 non-metal atoms
- the positively charged nuclei of the bonded atoms are attracted to the shared pair of electrons by electrostatic forces, making covalent bonds very strong
- atoms only share electrons in their outer shells (highest energy levels)
- each single covalent bond provides one extra shared electron for each atom
- each atom involved generally makes enough covalent bonds to fill its outer shell - the electronic structure of a noble gas (very stable)
- covalent bonding happens in compounds of non-metals and in non-metal elements
dot and cross diagram
covalent bonds
draw one
electrons drawn in the overlap between the outer orbitals of two atoms are shared between those atoms
pros
show which atoms the electrons in a covalent bond come from
cons
don’t show the relative sizes of the atoms
don’t show how the atoms are arranged in space
displayed formula diagram
covalent bonds
draw one
shows the covalent bonds as single lines between atoms
pros
show how atoms are connected in large atoms
cons
don’t show the 3D structure of the molecule, or which atoms the electrons in the covalent bond have come from
3D model diagram
covalent bond
draw one
pros
show the atoms, the covalent bonds and their arrangement in space next to each other
cons
3D models can get confusing for large molecules where there are lots of atoms to include
- don’t show where the electrons in the bonds have come from
ammonia
NH₃
simple molecular covalent structures
made up of molecules containing a few atoms joined together by covalent bonds
examples:
H₂ - hydrogen
Cl₂ - chlorine
O₂ - oxygen
N₂ - nitrogen
H₂O - water
simple molecular substances properties
- atoms within the molecules are held together by very strong covalent bonds
- by contrast, the forces of attraction between these molecules are very weak
- low melting and boiling points - small intermolecular forces to break and you don’t have to break the covalent bonds - molecules are easily parted from one another
- most are liquids/gases at room temp
- as molecules get bigger, the strength of the intermolecular forces increase, so more energy is needed to break them - melting + boiling points increase
- don’t conduct electricity as they aren’t charged so no free electrons or ions
polymers
large long chain molecules made up of lots of small monomers joined together by covalent bonds
long chains of repeating units
atoms joined by strong covalent bonds
- intermolecular forces are larger than between simple covalent molecules, however weaker than ionic or covalent bonds (lower boiling points than ionic or giant molecular compounds)
- most are solid at room temp
giant covalent structures
+ examples
macromolecules
- all the atoms are bonded to each other by strong covalent bonds
- very high melting + boiling points as lots of energy is required to break the covalent bonds between the atoms
- don’t conduct electricity - don’t contain charged particles - not even when molten (excluding graphite)
examples:
Diamond - made from carbon atoms only - each carbon atom forms four covalent bonds in a very rigid giant covalent structure
Graphite - made from carbon atoms only - each carbon atom forms three covalent bonds to create layers of hexagons. each carbon atom also has one delocalised electron
Silicon dioxide - sometimes called silica, what sand is made of. each grain of sand is one giant structure of silicon and oxygen
allotropes of carbon
diamond
graphite
graphene
fullerenes
diamond
has a giant covalent structure made up of carbon atoms that each form four covalent bonds - makes diamond very hard
- strong covalent bonds take a lot of energy to break and give diamond a very high melting point
- doesn’t conduct electricity because it has no delocalised electrons
graphite
contains sheets of electrons
- each carbon atom only forms three covalent bonds, creating sheets of carbon atoms arranged in hexagons
- aren’t any covalent bonds between the layers - they’re only held together weakly, so they’re free to move over each other - makes graphite soft and slippery, so it’s an ideal lubricating material
- high melting point - the covalent bonds in the layers need loads of energy to break
- only three out of each carbon’s four outer electrons are used in bonds, so each carbon atom has one electron that’s delocalised (free) and can move
- therefore graphite** conducts** electricity and thermal energy
graphene
one layer of graphite
- a sheet of carbon atoms joined together in hexagons
- the sheet is one atom thick, making it a two-dimensional substance - thin
- the network of covalent bonds makes it very strong
- incredibly light, so can be added to composite materials to improve their strength without adding much weight
- contains delocalised electrons so can conduct electricity through the whole structure - has the potential to be used in electronics
fullerenes
from spheres and tubes
- molecules of carbon, shaped like closed tubes or hollow balls
- mainly made up of carbon atoms arranged in hexagons, but can also contain pentagons or heptagons
uses:
- medicine - drug delivery into the body - ‘cage’ other molecules
- fullerenes have a huge surface area, so could help make great industrial catalysts - individual catalyst molecules could be attached to the fullerenes
- make great lubricants
e.g. buckminsterfullerene - C₆₀ - forms a hollow sphere
carbon nanotubes
tiny carbon cylinders
-high tensile strength (don’t break when stretched)
-conduct both electricity and thermal energy (delocalised electrons)
uses:
- electronics
- strengthen carbon fibre tennis rackets (strengthen materials without adding much weight)
- lubricants
metallic bonding
metal + metal
- very strong
- the electrons in the outer shell of the metal atoms are delocalised
- strong forces of electrostatic attraction between the positive metals ions and the shared negative electrons
- the forces of attraction hold the atoms together in a regular structure
- delocalised electrons are free to move throughout the structure + are shared throughout the structure so the metallic bonds are strong
substances held together by metallic bonding include metallic elements and alloys
- the delocalised electrons produce all the properties of metals
properties of metals
most are solid at room temp - the electrostatic forces between the metal atoms and the delocalised electrons are very story so need lots of energy to be broken - most compounds with metallic bonds have high melting + boiling points
good conductors - contain many delocalised electrons which carry charge and thermal energy through the whole structure, so metals are good conductors
most are malleable - the layers of atoms in metals can slide over each other - means they can be bent/hammered/rolled into flat sheets
alloys + pure metals
a mixture of two or more metals or a metal and another element
- harder so more useful than pure metals
- different elements have different sized atoms so when another element is mixed with a pure metal, the new metal atoms will distort the layers of metal atoms, making it more difficult for them to slide over each other - makes alloys harder than pure metals
pure metals (consist of a single element) are often too soft when they’re pure so mixed with other metals to make them harder
states of matter
solid, liquid, gas
which state something is in at a certain temperature depends on how strong the forces of attraction are between the particles of the material - how strong the forces are depends on:
- the material (structure of substance + type of bonding)
- the temperature
- the pressure
particle theory
a model that explains how the particles in a material behave in each of the three states of matter by considering each particle as a small, solid, inelastic sphere
- isn’t perfect as in reality particles aren’t solid or inelastic and aren’t spheres - they’re atoms, ions or molecules
- model doesn’t show the forces between the particles, so no way of knowing how strong they are
solids
- strong forces of attraction between particles, which holds them close together in fixed positions to form a very regular lattice arrangement
- particles don’t move from their positions, so all solids keep a definite shape and volume, and don’t flow like liquids
- particles vibrate about their positions - the hotter a solid becomes, the more they vibrate (causing solids to expand slightly when heated)
liquids
- weak force of attraction between particles
- randomly arranged and free to move past each other, but tend to stick closely together
- liquids have a definite volume but don’t keep a definite shape and will flow to the bottom of a container
- particles are constantly moving with random motion. The hotter the liquid gets, the faster they move - causes liquids to expand slightly when heated
gases
- very weak forces of attraction between particles - they’re free to move and far apart
- travel in straight lines
- don’t keep a definitive shape or volume and will always fill any container
- particles move constantly with random motion
- the hotter the gas gets, the faster they move
- gases either expand when heated, or their pressure increases
state symbols
tell you the state of a substance in an equation
(s) - solid
(l) - liquid
(g) - gas
(aq) - aqueous (dissolved in water)
melting
changing state
- solid heated, particles gain more energy
- particles vibrate more, *weakens *the forces holding the solid together
- at melting point the particles have enough energy to break free from their positions (melting) and the solid turns into a liquid
boiling
changing state
- liquid heated, particles gain more energy
- particles move faster, weakens and breaks bonds holding liquid together
- boiling point - particles have enough energy to break their bonds (boiling/evaporating)
- liquid becomes a gas
condensing
changing state
- gas cools, particles no longer have enough energy to overcome the forces of attraction between them
- bonds form between the particles
- at the boiling point, so many bonds have formed between the gas particles that the gas becomes a liquid (condensing)
freezing
changing state
- when a liquid cools, the particles have less energy, so move around less
- not enough energy to overcome the attraction between particles, so bonds form between them
- at melting point, so many bonds have formed between particles that they’re held in place
- liquid becomes a solid (freezing)
amount of energy required for a substance to change state depends on…
how strong the forces between particles are
- the stronger the forces, the more energy is needed to break them, and so the higher the melting and boiling points of the substance
nanoparticles
very small particles - contain only a few hundred atoms compared to ‘normal’ billions of atoms
diameters between 1nm-100nm in size
- have a very high surface are to volume ratio - surface area is very large compared to volume
- can exhibit different properties different to those for the same material in bulk e.g. you often need less of a material that’s made up of nanoparticles to work as an effective catalyst compared to a material made up of ‘normal’ sized particles
surface area to volume ratio
surface area to volume ratio = surface area ÷ volume
- as particles decrease in size, their surface area to volume ratio increases (the size of their surface area increases in relation to their volume)
uses of nanoparticles
catalysts - huge surface area to volume ratio
nanomedicine - tiny particles (e.g. fullerenes) are absorbed much more easily by the body than most particles - could deliver drugs right into the cells where they’re needed
electric circuits - some can conduct electricity, so can be used for electric circuits in computer chips
antibacterial properties - silver nanoparticles - can be added to polymer fibres that are then used to make surgical masks, wound dressings + can be added to deodorants
cosmetics - to improve moisturisers without making them really oily
effects of nanoparticles on health
aren’t fully understood
- don’t know what the long term effects will be
- many people believe products containing nanoparticles should be clearly labelled, so consumers can choose if they want to use them
- currently used in sun creams as they have proven to be better at protecting skin from UV rays than materials in traditional sun creams
+ give better skin coverage than traditional sun creams - not clear yet if the nanoparticles g=can get into your body, and if they do, whether they might damage your cells
- possible that when they are washed away they may damage the environment
