,A1: Structure and bonding in applications of science
Electronic structure of atoms
Definitions
Proton: a positively charged subatomic particle found in the nucleus of all atoms
Neutron: an uncharged subatotmic particle found in the nucleus of all atoms except hydrogen
Electron: a negatively charged subatomic particle found orbiting the nucleus in all atoms
Bohr model: a simple atomic model with a small positively charged nucleus surrounded by defined rings of
negatively charged electrons
Shell: an energy level occupied by electrons which surrounds the nucleus
Orbital: a subshell with a certain energy level, which can hold up to two electrons
Aufbau principle: electrons fill the lowest available energy shells first, and singly occupy orbitals in the same
energy level before pairing up
The basic model of the atom has two main components – the positively charged nucleus, which contains positive protons and neutral neutrons, and the
cloud of negatively charged electrons which encircle it. In all atoms, the number of protons equals the number of electrons. Protons and neutrons are far
larger and heavier than electrons.
Mass number
The number of electrons is equal to the number of protons. The number of protons is given by the
11 electrons
in
11 protons atomic number.
12 neutrons The number of protons plus neutrons in an atom is given by the mass number.
Atomic number
In the Bohr model of the atom, electrons are arranged in shells around the nucleus, which have higher energy levels as you move further away from
the centre. These shells of electrons orbit around the nucleus at fixed distances, rather like planets orbiting the sun. Each shell has a maximum
number of electrons that can fit into it – once it is full, extra electrons must go into the next shell.
Shells and orbitals
Within a shell, not all electrons have the exact same energy – they are split into subshells which have different energies. Subshells are further
divided into orbitals. Each orbital can hold two electrons with different spins.
Orbitals are filled according to the Aufbau principle – the
lowest available energy level (the one closest to the
nucleus) is filled first, before filling ones at higher levels.
Orbitals are labelled s, p, d or f, depending on which subshell they are in – p orbitals are higher in energy than the s orbitals in the same shells, d
orbitals are higher than p orbitals, and so on.
Rules for filling orbitals
✓ Each box represents an orbital which can hold two electrons.
✓ An electron is represented as an arrow with half a head.
✓ The lowest energy levels are filled first.
✓ The 4s subshell is lower in energy than the 3d subshell, so is filled
first (important for transition metals).
✓ All orbitals in a subshell are filled singly before electrons start to
pair up (pairing up requires more energy due to repulsions between
electrons in the same orbital).
By arranging electrons according to these rules, the most stable
electron configuration is obtained.
,Ionic bonding
Definitions
Electrostatic attraction: attractive forces between opposite charges
Noble gas configuration: a full shell of outer electrons; the most energetically stable arrangement
Giant ionic lattice: an arrangement of negative ions and positive ions in a regular 3D pattern
Ionic bonds are formed from the electrostatic attraction between oppositely charged ions.
• Ions are formed by the loss or gain of electrons.
• The atom losing electrons becomes positively charged (a cation).
• The atom gaining electrons becomes negatively charged (an anion).
By exchanging electrons, both elements are able to achieve a full outer shell of electrons, which is the most stable
electronic configuration. This is also sometimes called a noble gas configuration because all noble gas atoms
already have a full outer shell (which is the reason they are so unreactive).
Sodium has one Fluorine has seven
outer electron, outer electrons,
which it needs to and needs one
lose to have a full more to gain a full
outer shell. outer shell.
Sodium gives one Fluorine accepts one electron from
electron to fluorine sodium and, therefore, becomes
and, therefore, negatively charged – both
becomes positively elements now have a full outer
charged. shell of electrons.
The ratio of positive ions to negative ions does not have to be 1:1, e.g. in MgCl2.
Mg has two outer The opposing charges must
electrons, which it cancel out, so for every Mg2+
loses to become a ion there must be two Cl–
2+ ion. ions.
Ionic bonds are most commonly observed between a metal and a non-metal, such as sodium and chlorine. Oppositely charged ions join together into a
regularly ordered giant ionic lattice.
, The strength of the ionic bond is affected by:
• ionic charge
• ionic radius
Smaller ions with higher charges result in stronger electrostatic attractions. For example:
• NaCl has stronger ionic bonding than KCl – even though sodium ions and potassium ions have the same charge, Na+ is a smaller ion
• MgCl2 has stronger bonding than NaCl – the ions are of a similar size, but magnesium ions have a charge of +2, whereas a sodium ion has a charge
of +1
Covalent bonding
Definitions
Covalent bond: a shared pair of electrons that is electrostatically attracted to both nuclei
Dative bond / coordinate bond: a covalent bond in which both electrons in the bonding pair were donated by one atom
Lone pair: pair of electrons not involved in covalent bonding
In contrast to ionic bonding, covalent bonding involves sharing of electrons rather than complete exchange from one atom to the other, and is most
common between non-metals. The electrostatic attraction is between the positively charged nuclei and the negatively charged bonding electrons
shared between them.
As in ionic bonding, the sharing of electrons is done to give each atom involved a full outer shell of electrons. In the example, both the hydrogen atom
and the oxygen atom provide one electron each to form a bonding pair. Oxygen has six outer electrons and needs eight (two more) to have a full
outer shell, so it can form covalent bonds with two hydrogen atoms.
In the covalent
bond, one
electron comes
from hydrogen
(dot) and one
electron comes
from oxygen
(cross)
Oxygen has two
pairs of electrons
left over that don’t
take part in bonding
(lone pairs).
The stronger a covalent bond is, the shorter the bond length is. Stronger bonds are a result of the two bonding atoms having a large difference in
electronegativity.
Bond strength
increases / bond
length
decreases
Multiple bonding
Sometimes, more than one pair of electrons must be shared in order to reach that noble gas configuration. One way to resolve this is to bond to more
than one atom (as shown above); another is to share more than one pair of electrons and form a multiple bond.
Double bonds and triple bonds are stronger than single bonds and, therefore, are also shorter in length.
Electronic structure of atoms
Definitions
Proton: a positively charged subatomic particle found in the nucleus of all atoms
Neutron: an uncharged subatotmic particle found in the nucleus of all atoms except hydrogen
Electron: a negatively charged subatomic particle found orbiting the nucleus in all atoms
Bohr model: a simple atomic model with a small positively charged nucleus surrounded by defined rings of
negatively charged electrons
Shell: an energy level occupied by electrons which surrounds the nucleus
Orbital: a subshell with a certain energy level, which can hold up to two electrons
Aufbau principle: electrons fill the lowest available energy shells first, and singly occupy orbitals in the same
energy level before pairing up
The basic model of the atom has two main components – the positively charged nucleus, which contains positive protons and neutral neutrons, and the
cloud of negatively charged electrons which encircle it. In all atoms, the number of protons equals the number of electrons. Protons and neutrons are far
larger and heavier than electrons.
Mass number
The number of electrons is equal to the number of protons. The number of protons is given by the
11 electrons
in
11 protons atomic number.
12 neutrons The number of protons plus neutrons in an atom is given by the mass number.
Atomic number
In the Bohr model of the atom, electrons are arranged in shells around the nucleus, which have higher energy levels as you move further away from
the centre. These shells of electrons orbit around the nucleus at fixed distances, rather like planets orbiting the sun. Each shell has a maximum
number of electrons that can fit into it – once it is full, extra electrons must go into the next shell.
Shells and orbitals
Within a shell, not all electrons have the exact same energy – they are split into subshells which have different energies. Subshells are further
divided into orbitals. Each orbital can hold two electrons with different spins.
Orbitals are filled according to the Aufbau principle – the
lowest available energy level (the one closest to the
nucleus) is filled first, before filling ones at higher levels.
Orbitals are labelled s, p, d or f, depending on which subshell they are in – p orbitals are higher in energy than the s orbitals in the same shells, d
orbitals are higher than p orbitals, and so on.
Rules for filling orbitals
✓ Each box represents an orbital which can hold two electrons.
✓ An electron is represented as an arrow with half a head.
✓ The lowest energy levels are filled first.
✓ The 4s subshell is lower in energy than the 3d subshell, so is filled
first (important for transition metals).
✓ All orbitals in a subshell are filled singly before electrons start to
pair up (pairing up requires more energy due to repulsions between
electrons in the same orbital).
By arranging electrons according to these rules, the most stable
electron configuration is obtained.
,Ionic bonding
Definitions
Electrostatic attraction: attractive forces between opposite charges
Noble gas configuration: a full shell of outer electrons; the most energetically stable arrangement
Giant ionic lattice: an arrangement of negative ions and positive ions in a regular 3D pattern
Ionic bonds are formed from the electrostatic attraction between oppositely charged ions.
• Ions are formed by the loss or gain of electrons.
• The atom losing electrons becomes positively charged (a cation).
• The atom gaining electrons becomes negatively charged (an anion).
By exchanging electrons, both elements are able to achieve a full outer shell of electrons, which is the most stable
electronic configuration. This is also sometimes called a noble gas configuration because all noble gas atoms
already have a full outer shell (which is the reason they are so unreactive).
Sodium has one Fluorine has seven
outer electron, outer electrons,
which it needs to and needs one
lose to have a full more to gain a full
outer shell. outer shell.
Sodium gives one Fluorine accepts one electron from
electron to fluorine sodium and, therefore, becomes
and, therefore, negatively charged – both
becomes positively elements now have a full outer
charged. shell of electrons.
The ratio of positive ions to negative ions does not have to be 1:1, e.g. in MgCl2.
Mg has two outer The opposing charges must
electrons, which it cancel out, so for every Mg2+
loses to become a ion there must be two Cl–
2+ ion. ions.
Ionic bonds are most commonly observed between a metal and a non-metal, such as sodium and chlorine. Oppositely charged ions join together into a
regularly ordered giant ionic lattice.
, The strength of the ionic bond is affected by:
• ionic charge
• ionic radius
Smaller ions with higher charges result in stronger electrostatic attractions. For example:
• NaCl has stronger ionic bonding than KCl – even though sodium ions and potassium ions have the same charge, Na+ is a smaller ion
• MgCl2 has stronger bonding than NaCl – the ions are of a similar size, but magnesium ions have a charge of +2, whereas a sodium ion has a charge
of +1
Covalent bonding
Definitions
Covalent bond: a shared pair of electrons that is electrostatically attracted to both nuclei
Dative bond / coordinate bond: a covalent bond in which both electrons in the bonding pair were donated by one atom
Lone pair: pair of electrons not involved in covalent bonding
In contrast to ionic bonding, covalent bonding involves sharing of electrons rather than complete exchange from one atom to the other, and is most
common between non-metals. The electrostatic attraction is between the positively charged nuclei and the negatively charged bonding electrons
shared between them.
As in ionic bonding, the sharing of electrons is done to give each atom involved a full outer shell of electrons. In the example, both the hydrogen atom
and the oxygen atom provide one electron each to form a bonding pair. Oxygen has six outer electrons and needs eight (two more) to have a full
outer shell, so it can form covalent bonds with two hydrogen atoms.
In the covalent
bond, one
electron comes
from hydrogen
(dot) and one
electron comes
from oxygen
(cross)
Oxygen has two
pairs of electrons
left over that don’t
take part in bonding
(lone pairs).
The stronger a covalent bond is, the shorter the bond length is. Stronger bonds are a result of the two bonding atoms having a large difference in
electronegativity.
Bond strength
increases / bond
length
decreases
Multiple bonding
Sometimes, more than one pair of electrons must be shared in order to reach that noble gas configuration. One way to resolve this is to bond to more
than one atom (as shown above); another is to share more than one pair of electrons and form a multiple bond.
Double bonds and triple bonds are stronger than single bonds and, therefore, are also shorter in length.
