Chemistry - Topic 2 - Bonding, Structure and Properties of Matter
Nanotechnology
Effects of Nano particles on health aren't fully understood.
The way they affect the body isn't fully understood so everything needs to be tested to minimise risk.
Could have long term effects that we do not know about yet for the particles that are already out .
People think that products containing nanoscale particles should be clearly labelled.
What they are used for.
Some Nano particles conduct electricity so they can be used in tiny electronic circuits.
Silver Nano particles have antibacterial properties, they can be added to polymer fibres that are then used to make surgical masks and wound dressings and can also be added to deodorants.
Nano medicine, particles are absorbed more easily by the body meaning they could deliver drugs right to the cells where they are needed.
Nano particles are also being used on cosmetics. For example they're used to improve moisturisers without making them really oily.
Help make new catalysts.
Sun creams as they have been shown to protect skin better, not yet clear whether they can get into your body and kill your cells also possible that when they are washed away they might damage the environment.
Have a large surface area to volume ratio, surface area to volume ratio = surface area/volume. As particles decrease in size the size of their surface area increases in relation to their volume. Can cause properties of a material to be different depending on whether it's a nanoparticle or whether its in bulk.e.g need less of a material that's made up of nanoparticles to work as an effective catalyst.
Giant covalent structures and polymers
Macromolecules (Giant Covalent Structures)
They have very high melting and boiling points as lots of energy is needed to break the covalent bonds between the atoms.
They don't contain charged particles, so they don't conduct electricity - not even when molten (except for a few weird exceptions such as graphite).
In giant covalent structures, all the atoms are bonded to each other by strong covalent bonds.
The main examples that you need to know about are diamond and graphite, which are both made from carbon atoms only, and silicon dioxide (silica).
Diamond
Each carbon atom forms four covalent bonds in a very rigid giant covalent structure.
Graphite
Each carbon atom forms three covalent bonds to create layers of hexagons. Each carbon atom also has one delocalised (free) electron.
Silicon Dioxide
Sometimes called silica, this is what sand is made of. Each grain of sand is one giant structure of silicon and oxygen.
Polymers
To find the molecular formula of a polymer write down the molecular formula of the repeating unit in brackets, and put and 'n' outside.
So for Poly(ethene), the molecular formula of the polymer is (C2H4)n.
The intermolecular forces between polymer molecules are larger then between simple covalent molecules, so more energy is needed to break them. This means most polymers are solid at room temperature.
All the atoms in a polymer are joined by strong covalent bonds.
The intermolecular forces are still weaker than ionic or covalent bonds, so they generally have lower boiling points than ionic or giant molecular compounds.
In a polymer, lots of small units are linked together to form a long molecule that has repeating sections.
Instead of drawing out a whole long polymer molecule (which can contain thousands or even millions of atoms), you can draw the shortest repeating section, called the repeating unit.
Covalent bonding
Sharing Electrons
Atoms only share electrons in their outer shells (highest energy levels) .
Each single covalent bond provides one extra shared electron for each atom.
The positively charged nuclei of the bonded atoms are attracted to the shared pair of electrons by electrostatic forces making covalent bonds very strong.
Each atom involved generally makes enough covalent bonds to fill up its outer shell. Having a full outer shell gives them the electronic structure of a noble gas, which is very stable.
When non-metal atoms bond together, they share pairs of electrons to make covalent bonds.
Covalent bonding happens in compounds of non-metals (e.g. H20) and in non-metal elements (e.g. Cl2)
Different ways of drawing covalent bonds
Displayed Formula
The displayed formula shows the covalent bonds as single lines between atoms.
This is a great way of showing how atoms are connected in large molecules. However, they don't show the 3-D structure of the molecule, or which atoms the the electrons in the covalent bond have come from.
3-D model
The 3-D model shows the atoms, the covalent bonds and their arrangement in space next to each other. But 3-D models can quickly get confusing for large molecules where there are lots of atom to include. They don't show where the electrons in the bonds come from either.
Dot and cross diagrams
You can use dot and cross diagrams to show the bonding in covalent compounds.
Electrons drawn in the overlap between the outer orbitals of two atoms are shared between those atoms.
Dot and cross diagrams are useful for showing which atoms the electrons in a covalent bond come from, but they don't show the relative sizes of the atoms, or how the atoms are arranged in space.
You can find the molecular formula of a simple molecular compound from any of these diagrams by counting up how many atoms of each element there are.
Formation of Ions
Groups 1&2 and 6&7.
Ions are made when electrons are transferred
When metals form ions, they lose electrons from their outer shell to form positive ions.
When non-metals form ions, they gain electrons into their outer shell to form negative ions.
When atoms lose or gain electrons to form ions, all they're trying to do is get a full outer shell like noble gases. atoms with a full outer shell are very stable.
The number of electrons lost or gained is the same as the charge on the ion. E.g. if 2 electrons are lost the charge is 2+. If 3 electrons are gained the charge is 3-.
Ions are charged particles - they can be single atoms (e.g. Cl-) or groups of atoms (e.g. N03)
Group 6 and 7 elements are non-metals. They gain electrons to form negative ions (anions).
Group 1 and 2 elements are metals and they lose electrons to from positive ions (cations).
You don't have to remember what ions most elements form - nope, you just look at the periodic table.
The elements that most readily form ions are those in Groups, 1,2,6 and 7.
Simple Molecular Substances
Allotropes of Carbon
Ionic Compounds
Metallic Bonding
Ionic Bonding
States of Matter
Ionic Bonding — Transfer of Electrons
Dot and Cross Diagrams Show How Ionic Compounds are Formed
Ionic Compounds All Have Similar Properties
Look at Charges to Find the Formula of an Ionic Compound
Ionic Compounds Have A Regular Lattice Structure
Ionic compounds are formed into a structure called a giant atomic lattice.
The ions form a tightly packed, regular lattice.
When the compounds are solid, the ions are relatively stationary apart from some vibration (thermal energy). When ionic compounds melt, the ions are free to move so an electric current can be carried.
Ionic compounds dissolve easily in water. The ions separate and are free to move, so they can carry electric current.
A lot of energy is needed to overcome the strong electrostatic attraction, so ionic compounds have high melting and boiling points.
If it's a dot and cross diagram, count how many atoms there are of each element. This should give you the empirical formula, e.g. 1 magnesium and 2 chlorine = MgCl2.
If it's a 3D diagram, find what elements are in the compound. Then balance the charges so that the overall charge is 0.
You can find the empirical formula of an ionic compound from a diagram of the compound.
Examples of Simple Molecular Substances
Properties of Simple Molecular Substances
Nitrogen (N2): Nitrogen forms a triple bond with other nitrogen atoms.
Methane (CH4): Carbon has four outer electrons, which is half a full shell. It can form four covalent bonds with hydrogen atoms to fill up its outer shell.
Oxygen (O2): Oxygen atoms need two more electrons to complete its outer shell, so two oxygen atoms form a double bond by sharing two pairs of electrons.
Water (H2O): In water molecules, the oxygen shares a pair of electrons with two hydrogen atoms to form two covalent bonds.
Chlorine (Cl2): Like hydrogen atoms, chlorine atoms form a diatomic molecule with a single bond.
Hydrogen Chloride (HCl): This is very similar to H2 and Cl2. Again, both atoms only need one more electron to complete their outer shells.
Hydrogen (H2): Hydrogen atoms only have one electron, so they only need one more. This means they often form single covalent bonds with other hydrogens or other elements.
The atoms within the molecules are held together by very strong covalent bonds. By contrast, the forces of attraction between these molecules are very weak.
To melt or boil a simple molecular compound, you only need to break these feeble intermolecular forces and not the covalent bonds. So the melting and boiling points are very low, because the molecules are easily parted from each other.
Substances containing covalent bonds usually have simple molecular structures, like the examples above.
Most molecular substances are gases or liquids at room temperature.
As molecules get bigger, the strength of the intermolecular forces increases, so more energy is needed to break them, and the melting and boiling points increase.
Molecular compounds don’t conduct electricity, simply because they aren’t charged, so there are no free electrons or ions.
Graphite Contains Sheets of Hexagons
Graphene is One Layer of Graphit
Diamond is Very Hard
Fullerenes Form Spheres and Tubes
Each carbon atom forms four covalent bonds in a very rigid giant covalent structure.
Each carbon atom forms three covalent bonds to create layers of hexagons. Each carbon atom also has one delocalised (free) electron.
Fullerenes are hollow and big enough to "cage" other molecules. The fullerene structure forms around another atom or molecule, which is then trapped inside.
To to their large surface area, fullerenes can be used to make great industrial catalysts— individual catalyst molecules could be attached to the fullerenes. Fullerenes also make great lubricants.
They are mostly made up of hexagons of carbon atoms, but can consist of pentagons and heptagons also.
Carbon Nanotubes
Fullerenes are molecules of carbon, shaped like closed tubes or hollow balls.
The tubes are very long in comparison to their diameter.
They are good electrical and thermal conductors along the tube due to their free electrons.
One kind of fullerene are nanotubes: Tiny carbon cylinders.
They have a high tensile strength. This means they don't break easily when stretched.
Metals are Good Conductors of Electricity and Heat
Most Metals are Malleable
Most Metals are Solid at Room Temperature
Alloys are Harder Than Pure Metals
Metallic Bonding Involves Delocalised Electrons
These forces of attraction hold the atoms together in a regular structure and are known as metallic bonding. Metallic bonding is very strong
Compounds that are held together by metallic bonding include metallic elements and alloys.
The electrons in the outer shell of the metal atoms are delocalised There are strong forces of electrostatic attraction between the positive metal ions and the shared negative electrons.
It’s the delocalised electrons in the metallic bonds which produce all the properties of metals.
Metals also consist of a giant structure.
The electrostatic forces between the metal atoms and the delocalised sea of electrons are very strong, so need lots of energy to be broke
This means that most compounds with metallic bonds
have very high melting and boiling points, so they’re generally solid at room temperature.
The delocalised electrons carry electrical current and thermal (heat) energy through the whole structure, so metals are good conductors of electricity and heat.
The layers of atoms in a metal can slide over each other, making metals malleable — this means that they can be bent or hammered or rolled into flat sheets.
Pure metals often aren’t quite right for certain jobs often too soft
when they’re pure so are mixed with other metals to make them harder. Alloys are harder and so more useful than pure metals.
Different elements have different sized atoms. So when mixed it distorts they layers making it difficult for them to slide over each other. This makes alloys harder than pure metals
The Three States of Matter — Solid, Liquid and Gas
Solids
Liquids
How strong the forces are depends on THREE THINGS:
Gases
State Symbols Tell You the State of a Substance in an Equation
(s) — solid (l) — liquid (g) — gas (aq) — aqueous
Changing State
Substances Can Change from One State to Another
6) At the boiling point, the particles have enough energy to break their bonds. When a liquid turns into a gas at its boiling point, it is called boiling.
7) As a gas cools, its particles don't have enough energy to overcome the forces of attraction between them.
5) This energy makes the particles move faster, which further weakens and even breaks the bonds holding the liquid particles together.
8) The particles form bonds.
4) Like a solid, when a liquid is heated its particles gain more energy.
9) At the boiling point, so many particles have formed bonds that the gas becomes a liquid. This is called condensing.
3) At the melting point, the particles in the substance have enough energy to break free from their positions. When a solid turns into a liquid is called melting.
10) When a liquid cools, the particles lose energy and therefore move around less.
2) This causes particles to vibrate more, weakening the forces that hold them together.
11) Again, there is not enough energy to overcome the force of attraction so more bonds are formed.
1) When a solid is heated, it's particles gain more energy.
12) At the melting point so many bonds have formed between the particles that they’re held in place. The liquid becomes a solid. This is FREEZING.
You Have to be Able to Predict the State of a Substance
If the temperature of a substance is below the melting point, it is a solid.
When a metal and non-metal bond, it is known as ionic bonding. This is because ions are formed.
The easiest way for the metal to get to a full outer shell is to lose an electron(s), and the easiest way for the non-metal to get a full outer shell is to gain an electron(s).
The atoms gain or lose electrons to each other, becoming ions in the process.
The metal becomes positively charged and the non-metal becomes negatively charged. This results in electrostatic attraction which pulls the two ions together, bonding them.
Sodium Chloride (NaCl)
The sodium atom gives up its outer electron, becoming an Na+ ion. The chlorine atom picks up the electron, becoming a Cl– (chloride) ion.
Magnesium Oxide (MgO)
The magnesium atom gives up its two outer electrons, becoming an Mg2+ ion. The oxygen atom picks up the electrons, becoming an O2– (oxide) ion.
Magnesium Chloride (MgCl2)
The magnesium atom gives up its two outer electrons, becoming an Mg2+ ion. The two chlorine atoms pick up one electron each, becoming two Cl– (chloride) ions.
Sodium Oxide (Na2O)
Two sodium atoms each give up their single outer electron, becoming two Na+ ions. The oxygen atom picks up the two electrons, becoming an O2– ion.
There are strong electrostatic forces of attraction between oppositely charged ions, in all directions of the lattice.
The layers of carbon that form graphite are called graphene.
The network of covalent bonds makes it very strong. It's also incredibly light, so can be added to composite materials to improve their strength without adding much weight.
There are weak bonds between the layers so they can slide between each other.
Like graphite, it contains delocalised electrons so can conduct electricity through the whole structure. This means it has the potential to be used in electronics.
The use of small particles such as nanotubes is called nanotechnology. Nanotubes can be used in electronics or to strengthen materials without adding much mass and therefore weight.
the temperature
the pressure
the material
The particles don’t move from their positions, so all solids keep a definite shape and volume.
The particles vibrate about their positions - the hotter the solid becomes, the more they vibrate
Strong forces of attraction between particles, which holds
them close together in fixed positions to form a very regular lattice arrangement.
Gases don’t keep a definite shape or volume and will always fill any container.
The particles move constantly with random motion. The hotter the gas gets, the faster they move. Gases either expand when heated, or their pressure increases.
In gases, the force of attraction between the particles is very weak — they’re free to move and are far apart. The particles in gases travel in straight lines.
Liquids have a definite volume but don’t keep a definite shape,
and will flow to fill the bottom of a container.
The particles are constantly moving with random motion. The hotter the liquid gets, the faster they move. This causes liquids to expand slightly when heated.
Weak force of attraction between the particles. They’re randomly
arranged and free to move past each other, but they tend to stick closely together.
If the temperature of a substance is between the melting point and boiling point, it is a liquid.
If the temperature of a substance is above the boiling point, it is a gas.