NOTES ON WEEK 2 REVISION – ATOMIC STRUCTURE AND BONDING
Atom = the smallest part of an element that could possibly exist.
Element = all atoms in the element are the SAME.
Subatomic particle Mass Charge
Proton 1 +1
Electron 0 (1/1836) -1
Neutron 1 0
• All of the atoms mass is in the nucleus.
• Overall charge of an atom is 0 as the protons and neutrons cancel each other out.
• Atomic number à shows the number of protons and electrons (Z)
• Mass number à the number of protons and neutrons added together (A)
Isotopes
Isotope = atoms of the same element with the same number of protons and neutrons but a different number of neutrons.
à SAME chemical properties. as electrons are the subatomic particle that cause chemical reactions!
à DIFFERENT physical properties. As atomic mass has increased OR decreased.
Cation = an atom with a higher number of protons and a lower number of electrons (POSITIVE)
Anion = an atom with a higher number of electrons and a lower number of protons (NEGATIVE)
Relative Mass
• We use the carbon-12 atom isotope to base ALL atomic masses.
• The mass of carbon-12 isotope is defined as 12 atomic mass units (12u)
• Standard mass for atomic mass is 1u (1/12th of a carbon-12 atom)
• 1u is approximately the mass of a proton or a neutron.
Relative isotopic mass = the mass of an isotope relative to 1/12th of the mass of an atom of carbon-12.
Relative atomic mass = the weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of
carbon-12.
TAKES ACCOUNT FOR:
• Percentage abundance of each isotope Mass to charge ratio =
• Relative isotopic mass of each isotope.
M/Z (relative mass of ion/relative
Mass spectrometry charge on ion)
à works out the percentage abundance of isotopes in a sample.
1) Sample placed in the mass spectrometer.
2) The sample is vaporised and then ionised to form positive ions.
3) Ions are accelerated. (Heavier ions move slowly and more difficult to deflect than lighter ions, so the ions of each
isotope are separated).
4) Ions detected on a mass spectrum as mass to charge ratio. Each ion reaching the detector adds to the signal, so
the greater the abundance the larger the signal).
Electrons and bonding
Shells = energy levels, energy level increases as the shell number increases.
à Principal quantum number = the shell number / energy level number.
Orbital = a region of space around the nucleus in which an electron is likely to be found. (two electrons in an orbital which
spin in opposite directions)
, ORBITAL Characteristics
S • Sphere shape (cloud within the shape of a sphere)
• From n=1, there is one S orbital in each shell.
P • Dumb bell shape (cloud within the shape of a dumbbell)
• From n=2 there are 3 p orbitals
D/F • Each shell from n=3 has 5 d-orbitals
• Each shell from n=4 has 7 f-orbitals
Sub shell = a group of the same type of orbitals in a shell. E.g. shell 3 would have 5-d orbitals (acts as one subshell)
Filling of orbitals
• Goes in the order S, P, D….
• BUT… the 3 d-subshell is at a higher energy level than the 4 s-subshell so the 4s fills before the 3d.
Electron pairs
• Must have arrows in opposite directions (opposite spins) because they are negative, hence they repel each other.
• You must occupy ALL orbitals, even if there is only one arrow in the box (one electron)
Electron configuration
electron
Subshell
Blocks and the periodic table
• S-BLOCK – highest energy electrons in the s-subshell = left block of two groups (1+2)
• P-BLOCK – highest energy electrons in the p-subshell = right block of 6 groups.
• D-BLOCK – highest energy electrons in the d-subshell = centre block of 10 groups
Bonding
Ionic bonding = the electrostatic attraction between positive and negative ions. (between cations and anions)
• Giant ionic lattice structure with alternating positive and negative ions.
Melting and boiling points Solubility Electrical conductivity
Solids at room temperature Dissolve in Polar substances. Solid = no conductivity
High MP. And BP. As lots of energy is Polar molecules break down the Liquid (molten) / aqueous = conducts
required to break the strong ionic lattice and surround each ion in as the ions are free to move so are
bonds. solution. able to carry electrical charge
As ionic charge increases, there is
stronger attractions between the
ions… so higher MP and BP.
IONIC BONDING OCCURS BETWEEN A METAL AND A NON-METAL
Processes of solubility:
n Ionic lattice must be broken down
n Water molecules must attract and surround ions
Atom = the smallest part of an element that could possibly exist.
Element = all atoms in the element are the SAME.
Subatomic particle Mass Charge
Proton 1 +1
Electron 0 (1/1836) -1
Neutron 1 0
• All of the atoms mass is in the nucleus.
• Overall charge of an atom is 0 as the protons and neutrons cancel each other out.
• Atomic number à shows the number of protons and electrons (Z)
• Mass number à the number of protons and neutrons added together (A)
Isotopes
Isotope = atoms of the same element with the same number of protons and neutrons but a different number of neutrons.
à SAME chemical properties. as electrons are the subatomic particle that cause chemical reactions!
à DIFFERENT physical properties. As atomic mass has increased OR decreased.
Cation = an atom with a higher number of protons and a lower number of electrons (POSITIVE)
Anion = an atom with a higher number of electrons and a lower number of protons (NEGATIVE)
Relative Mass
• We use the carbon-12 atom isotope to base ALL atomic masses.
• The mass of carbon-12 isotope is defined as 12 atomic mass units (12u)
• Standard mass for atomic mass is 1u (1/12th of a carbon-12 atom)
• 1u is approximately the mass of a proton or a neutron.
Relative isotopic mass = the mass of an isotope relative to 1/12th of the mass of an atom of carbon-12.
Relative atomic mass = the weighted mean mass of an atom of an element relative to 1/12th of the mass of an atom of
carbon-12.
TAKES ACCOUNT FOR:
• Percentage abundance of each isotope Mass to charge ratio =
• Relative isotopic mass of each isotope.
M/Z (relative mass of ion/relative
Mass spectrometry charge on ion)
à works out the percentage abundance of isotopes in a sample.
1) Sample placed in the mass spectrometer.
2) The sample is vaporised and then ionised to form positive ions.
3) Ions are accelerated. (Heavier ions move slowly and more difficult to deflect than lighter ions, so the ions of each
isotope are separated).
4) Ions detected on a mass spectrum as mass to charge ratio. Each ion reaching the detector adds to the signal, so
the greater the abundance the larger the signal).
Electrons and bonding
Shells = energy levels, energy level increases as the shell number increases.
à Principal quantum number = the shell number / energy level number.
Orbital = a region of space around the nucleus in which an electron is likely to be found. (two electrons in an orbital which
spin in opposite directions)
, ORBITAL Characteristics
S • Sphere shape (cloud within the shape of a sphere)
• From n=1, there is one S orbital in each shell.
P • Dumb bell shape (cloud within the shape of a dumbbell)
• From n=2 there are 3 p orbitals
D/F • Each shell from n=3 has 5 d-orbitals
• Each shell from n=4 has 7 f-orbitals
Sub shell = a group of the same type of orbitals in a shell. E.g. shell 3 would have 5-d orbitals (acts as one subshell)
Filling of orbitals
• Goes in the order S, P, D….
• BUT… the 3 d-subshell is at a higher energy level than the 4 s-subshell so the 4s fills before the 3d.
Electron pairs
• Must have arrows in opposite directions (opposite spins) because they are negative, hence they repel each other.
• You must occupy ALL orbitals, even if there is only one arrow in the box (one electron)
Electron configuration
electron
Subshell
Blocks and the periodic table
• S-BLOCK – highest energy electrons in the s-subshell = left block of two groups (1+2)
• P-BLOCK – highest energy electrons in the p-subshell = right block of 6 groups.
• D-BLOCK – highest energy electrons in the d-subshell = centre block of 10 groups
Bonding
Ionic bonding = the electrostatic attraction between positive and negative ions. (between cations and anions)
• Giant ionic lattice structure with alternating positive and negative ions.
Melting and boiling points Solubility Electrical conductivity
Solids at room temperature Dissolve in Polar substances. Solid = no conductivity
High MP. And BP. As lots of energy is Polar molecules break down the Liquid (molten) / aqueous = conducts
required to break the strong ionic lattice and surround each ion in as the ions are free to move so are
bonds. solution. able to carry electrical charge
As ionic charge increases, there is
stronger attractions between the
ions… so higher MP and BP.
IONIC BONDING OCCURS BETWEEN A METAL AND A NON-METAL
Processes of solubility:
n Ionic lattice must be broken down
n Water molecules must attract and surround ions
