- Elements are made of tiny particles of matter called atoms.
- Each atom is made of subatomic particles called protons, neutrons and electrons.
- Their size is so tiny that we can’t really compare their masses in conventional units such as kilograms or grams, so a unit called the relative atomic mass is used.
- One relative atomic mass unit is equal to 1/12 the mass of a carbon-12 atom.
- All other elements are measured relative to the mass of a carbon-12 atom and since these are ratios, the relative atomic mass has no units.
- Hydrogen for example has a relative atomic mass of 1, meaning that 12 atoms of hydrogen would have exactly the same mass as 1 atom of carbon.
- The relative mass and charge of the subatomic particles are shown below:
Defining Proton Number
sub-atomic particle relative mass relative charge
proton 1 +1
neutron 1 0
electron 1/1836 -1
- Define proton number (atomic number) as the number of protons in the nucleus of an atom
- The atomic number (or proton number) is the number of protons in the nucleus of an atom. The symbol for this number is Z.
- It is also the number of electrons present in an atom and determines the position of the element on the Periodic Table.
Defining Nucleon Number
- Define nucleon number (mass number) as the total number of protons and neutrons in the nucleus of an atom
- (or mass number) is the total number of protons and neutrons in the nucleus of an atom. The symbol for this number is A.
- The nucleon number minus the proton number gives you the number of neutrons of an atom.
- Note that protons and neutrons can collectively be called nucleons.
- The atomic number and mass number for every element is on the Periodic Table
THE PERIODIC TABLE
- Elements are arranged on the Periodic table in order of increasing atomic number where each element has one proton more than the element preceding it.
- Hydrogen has 1 proton, helium has 2 protons, lithium has 3 etc.
- The table is arranged in vertical columns called Groups numbered I – VIII and in rows called Periods.
- Elements in the same group have the same amount of electrons in their outer shell, which gives them similar chemical properties.
● Atoms of the same element which have the same proton number but a different
nucleon number (different number of neutrons)
(Extended only) Understand that isotopes have the same properties because
they have the same number of electrons in their outer shell
● Isotopes have the same chemical properties because they have the same
number of electrons in their outer shell and the number of electrons in the outer
shell is responsible for chemical reactions
● they also have the same number of protons so are still classed as the same element
State the two types of isotopes as being
● Radioactive and non-radioactive
State one medical and one industrial use of radioactive isotopes
● Medical uses:
o Sterilising equipment
● Industrial uses:
o Smoke alarms
- Radioactive isotopes (radioisotopes) are unstable due to the imbalance of neutrons and protons, which causes the nucleus to decay over time through nuclear fission and emit radiation. Examples of radioisotopes include tritium and carbon-14.
- Decay occurs at a different rate for each isotope, but the time taken for the radioactivity of an isotope to decrease by 50% is constant for that particular isotope and is known as the half-life.
- Radioactive isotopes have numerous medical and industrial uses.
- Non-radioactive isotopes are stable atoms which really only differ in their mass.
- Radiation is extremely harmful and kills cells so isotopes are used to treat cancer. The isotope cobalt-60 is frequently used for this purpose.
- Medical tracers as certain parts of the body absorb isotopes and others do not. In this way an isotope can be injected into the blood and its path through the body traced with a radiation-detecting camera, revealing the flow of blood through bodily systems.
- Medical instruments and materials are routinely sterilized by exposure to radiation, which kills any bacteria present.
- Radioactive dating uses the carbon-14 isotope to date carbon-containing materials such as organic matter, rocks and other artefacts. The half-life of C-14 is 5730 years and so this technique is often used to date very old historical objects.
- Similar to the medical use, radioactive tracers are deployed for detecting leaks in gas or oil pipes.
- The radioactive isotope uranium-235 is used as fuel for controlled fission reactions in nuclear power plants.
Describe the build-up of electrons in ‘shells’ and understand the significance
of the noble gas electronic structures and of the outer shell electrons (The
ideas of the distribution of electrons in s and p orbitals and in d block
elements are not required.) Note: a copy of the Periodic Table, will be
available in Papers 1, 2, 3 and 4.
● Electrons are arranged around the nucleus in shells. Starting with the first shell
(closest to the nucleus), each shell is filled with electrons before any further
shells gain any electrons
o First shell can have up to 2 electrons
o Second shell can have up to 8 electrons
o Third shell can have up to 8 electrons
● When reacting, all atoms will try to acquire this perfect arrangement of electrons
– i.e. having the maximum number of electrons as possible in their outer shell –
therefore, all atoms try to have 8 electrons in their outer shell (unless they only
have one shell then they will try to have only 2) because this is the most stable
● Nobles gases have 8 electrons in their outer shells
already (except helium, which has 2), making them
very stable so unreactive
- The arrangement of electrons in shells can also be explained using numbers.
- There is a clear relationship between the outer shell electrons and how the Periodic Table is designed.
- The number of notations in the electronic configuration will show the number of shells of electrons the atom has, showing the Period in which that element is in.
- The last notation shows the number of outer electrons the atom has, showing the Group that element is in.
- Elements in the same Group have the same number of outer shell electrons.
Elements, compounds and mixtures
- All substances can be classified into one of these three types.
- A substance made of atoms that all contain the same number of protons (one type of atom) and cannot be split into anything simpler.
- There is a limited number of elements and all elements are found on the Periodic Table.
- E.g. hydrogen, carbon, nitrogen.
- A pure substance made up of two or more elements chemically combined together.
- There is an unlimited number of compounds.
- Compounds cannot be separated into their elements by physical means.
- E.g. copper (II) sulphate (CuSO4), calcium carbonate (CaCO3), carbon dioxide (CO2).
- A combination of two or more substances (elements and/or compounds) that are not chemically combined.
- Mixtures can be separated by physical methods such as filtration or evaporation.
- E.g. sand and water, oil and water, sulphur powder and iron filings.
Metals and nonmetals
- The Periodic Table contains over 100 different elements.
- They can be divided into two broad types: metals and nonmetals.
- Most of the elements are metals and a small number of elements display properties of both types. These elements are called metalloids or semimetals.
Properties of metals
- Conduct heat and electricity.
- Are malleable and ductile (can be hammered and pulled into different shapes).
- Tend to be lustrous (shiny).
- Have high density and usually have high melting points.
- Form positive ions through electron loss.
- Form basic oxides.
Properties of nonmetals
- Do not conduct heat and electricity.
- Are brittle and delicate when solid and easily break up.
- Tend to be dull and nonreflective.
- Have low density and low melting points (many are gases at room temperature).
- Form negative ions through electron gain (except for hydrogen).
- Form acidic oxides.
Describe an alloy, such as brass, as a mixture of a metal with other elements
● Most metals in everyday uses are alloys. Pure copper, gold, iron and aluminium
are all too soft for everyday uses and so are mixed with small amounts of other
elements (in these cases with similar metals) to make them harder for everyday
o Gold in jewellery is usually an alloy with silver, copper and zinc
o brass is an alloy of copper and zinc
● alloys are harder than pure metals because:
o in a pure metal, all the + ions are the same size
and in a regular arrangement so can easily slide
over each other
o in an alloy, there are + ions from different
metals, meaning they are different sizes, which
disrupts the regular arrangement and prevents
the layers from sliding as easily
Ions and Ion formation
- An ion is an electrically charged atom or a group of atoms formed by the loss or gain of electrons.
- This loss or gain of electrons takes place to gain a full outer shell of electrons.
- The electronic structure of an ion will be the same as that of a noble gas – such as helium, neon and argon.
Ionisation of metals and non-metals
- Metals: all metals lose electrons to another atom and become positively charged ions.
- Non-metals: all non-metals gain electrons from another atom to become negatively charged ions.
- Describe the formation of ions by electron loss or gain
- ● an ion is an atom or group of atoms with a positive or negative charge
- ● ions are formed by an atom losing or gaining electrons (which have a -1 charge)
- ● if an atom gains electrons, it becomes a negative ion
- ● if an atom loses electrons, it becomes a positive ion
- o Cation = positive ion (+ -> ca+ion)
- o Anion = negative ion (Negative -> aNion)
- Describe the formation of ionic bonds between elements from Groups I and
- ● an ionic bond is formed when an electron is transferred from one atom to
- ● when ionic bonds are formed between group 1 and 7:
- o group 1 atom loses one electron and forms a +1 ion
- o group 7 atom gains the electron the group 1 atom lost and becomes a -1
- ● Electron transfer during the formation of an ionic compound can be represented
- by a dot and cross diagram
- (Extended only) Describe the formation of ionic bonds between metallic and
- non-metallic elements
- ● Metal + Non-metal: electrons in the outer shell of the metal atom are
- o Metal atoms lose electrons to become positively charged ions
- o Non-metal atoms gain electrons to become negatively charged ions
- (Extended only) Describe the lattice structure of ionic compounds as a
- regular arrangement of alternative positive and negative ions
- ● Held together by strong electrostatic forces of attraction between oppositely
- charged ions, which are regularly arranged
- ● The forces act in all directions in the lattice, and this is called ionic bonding.
- An example is sodium chloride (salt):
- Na+ (small blue particles) and Cl- (larger green ones)
- ● Strong electrostatic forces of attraction between alternating positive and
- negative ions
- ● Requires a lot of energy to overcome these forces of attraction
- Covalent compounds are formed when electrons are shared between atoms.
- Only non-metal elements participate in covalent bonding.
- As in ionic bonding, each atom gains a full outer shell of electrons.
- When two or more atoms are chemically bonded together, we describe them as ‘molecules’.
Ionic and Covalent Compounds
- Describe the differences in volatility, solubility and electrical conductivity between ionic and covalent compounds
- Ionic compounds:
- Have high melting and boiling points so ionic compounds are usually solid at room temperature.
- Not volatile so they don’t evaporate easily.
- Usually water soluble as both ionic compounds and water are polar (see polarity in Glossary).
- Conduct electricity in molten state or in solution as they have ions that can move and carry charge.
- Covalent compounds:
- Have low melting and boiling points so covalent compounds are usually liquids or gases at room temperature.
- Usually volatile which is why many covalent organic compounds have distinct aromas.
- Usually not water soluble as covalent compounds tend to be nonpolar but can dissolve in organic solvents.
- Cannot conduct electricity as all electrons are involved in bonding so there are no free electrons or ions to carry the charge.
- Describe the giant covalent structure of graphite and diamond
- Diamond and graphite are allotropes of carbon which have giant covalent structures.
- These classes of substance contain a lot of non-metal atoms, each joined to adjacent atoms by covalent bonds forming a giant lattice structure.
- Giant covalent structures have high melting and boiling points as they have many strong covalent bonds that need to be broken down.
- Large amounts of heat energy are needed to overcome these forces and break down bonds.
Diamond, graphite and fullerene are examples of Giant Covalent Structures
Uses of Giant Covalent Structures
- Relate their structures to their uses, e.g. graphite as a lubricant and a conductor, and diamond in cutting tools
- Each carbon atom bonds with four other carbons, forming a tetrahedron.
- All the covalent bonds are identical and strong with no weak intermolecular forces.
- Diamond thus:
- Does not conduct electricity.
- Has a very high melting point.
- Is extremely hard and dense (3.51 g/cm3).
- Diamond is used in jewellery and as cutting tools.
- The cutting edges of discs used to cut bricks and concrete are tipped with diamonds.
- Heavy-duty drill bits and tooling equipment are also diamond tipped.
- Each carbon atom is bonded to three others forming layers of hexagonal shaped forms, leaving one free electron per carbon atom.
- These free electrons exist in between the layers and are free to move and carry charge, hence graphite can conduct electricity.
- The covalent bonds within the layers are very strong but the layers are connected to each other by weak intermolecular forces only, hence the layers can slide over each other making graphite slippery and smooth.
- Graphite thus:
- Conducts electricity.
- Has a very high melting point.
- Is soft and slippery, less dense than diamond (2.25 g/cm3).
- Graphite is used in pencils and as an industrial lubricant, in engines and in locks.
- It is also used to make non-reactive electrodes for electrolysis.
Extended Subject Content
The Structure of Silicon(IV) Oxide (Silicon Dioxide)
- Describe the macromolecular structure of silicon(IV) oxide (silicon dioxide)
- SiO2 is a macromolecular compound which occurs naturally as sand and quartz.
- Each oxygen atom forms covalent bonds with 2 silicon atoms and each silicon atom in turn forms covalent bonds with 4 oxygen atoms.
- A tetrahedron is formed with one silicon atom and four oxygen atoms, similar as in diamond.
Diagram showing the structure of SiO2 with the silicon atoms in dark grey and the oxygen atoms in red
Diamond and Silicon(IV) Properties
- Describe the similarity in properties between diamond and silicon(IV), related to their structures
- SiO2 has lots of very strong covalent bonds and no intermolecular forces so it has similar properties as diamond.
- It is very hard, has a very high boiling point, is insoluble in water and does not conduct electricity.
- SiO2 is cheap since it is available naturally and is used to make sandpaper and to line the inside of furnaces.
(Extended only) Describe the similarity in properties between diamond and
silicon(IV) oxide, related to their structures
● Similar properties:
o Very hard
o Very high melting and boiling points
o Insoluble in water
o Does not conduct electricity
● These are due to the strong covalent bonds that hold the atoms in a rigid structure
Extended Subject Content
Electrical Conductivity and Malleability of Metals
- Describe metallic bonding as a lattice of positive ions in a ‘sea of electrons’ and use this to describe the electrical conductivity and malleability of metals
- Metal atoms are held together strongly by metallic bonding.
- Within the metal lattice, the atoms lose their valence electrons and become positively charged.
- The valence electrons no longer belong to any metal atom and are said to be delocalised.
- They move freely between the positive metal ions like a sea of electrons.
- Metallic bonds are strong and are a result of the attraction between the positive metal ions and the negatively charged delocalised electrons.
(Extended only) Describe metallic bonding as a lattice of positive ions in a
‘sea of electrons’ and use this to describe the electrical conductivity and
malleability of metals
● metallic bonding: a regular lattice of + metal ions surrounded by a sea of delocalised electrons
● Metals consist of giant structures of atoms arranged in a regular pattern.
● The electrons in the outer shell of metal atoms are delocalised and so are free to
move through the whole structure.
● The sharing of delocalised electrons gives rise to strong metallic bonds.
● electrical conductivity: metals can conduct el
The link between metallic bonding and the properties of metals:
- Metals have high melting and boiling points:
- There are many strong metallic bonds in giant metallic structures.
- A lot of heat energy is needed to overcome forces and break these bonds.
- Metals conduct electricity:
- There are free electrons available to move and carry charge.
- Electrons entering one end of the metal cause a delocalised electron to displace itself from the other end.
- Hence electrons can flow so electricity is conducted.
- Metals are malleable and ductile:
- Layers of positive ions can slide over one another and take up different positions.
- Metallic bonding is not disrupted as the valence electrons do not belong to any particular metal atom so the delocalised electrons will move with them.
- Metallic bonds are thus not broken and as a result metals are strong but flexible.
- They can be hammered and bent into different shapes without breaking.