Atomic Number of protons (and consequently electrons) in the nucleus Atomic Structure
number of an atom
Subatomic Relative Relative
Mass number Total number of protons and neutrons in the nucleus of an
particle mass charge
atom
Isotope An atom that has the same number of protons (so they are
atoms of the same element) and electrons (therefore are still Proton (p+) 1 +1/+ve
neutral) but have a different number of neutrons (so they have
a different mass number)
Neutron (n0) 1 0
Cation Positive ion
Anion Negative ion
Atom The smallest particle of an element that can exist on its own in Electron (e-) 1/1840 -1/-ve
a stable environment – always neutral
History of the atom
2000 years ago:
Greek philosophers Democritus and Leucippus first suggested matter was
made up of individual particles called “atoms”, however Aristotle believed matter
was made of the “four elements”: air, earth, fire, water. Aristotle was more
renowned and considered correct.
1808:
John Dalton published his Atomic Theory, where he stated that “all elements
are made up of small indivisible particles called atoms”, when the idea was
taken seriously.
1897:
Electron discovered by J.J Thomson, new model of atomic structure proposed
called the Plum Pudding Model – rings of negative electrons embedded in a
sphere of positive charge, just like currants in a Christmas pudding. He also put
forward that the atom was neutral as it contained positive and negative
charges.
1913:
Ernest Rutherford and his researchers used alpha particles to probe inside
atoms which led to discovery of the nucleus. His research shows that the atom
consisted of electrons revolving around a positively charged nucleus.
Positives particles are called protons, 2000x heavier than electrons, which cancel
each other and make the atom neutral. Electrons are contained in much larger
region around nucleus, and most weight is in the nucleus.
1932:
James Chadwick discovers where the extra mass came from – as mass was
larger than expected – he showed that the nucleus of an atom contained two
no. of protons = different types of particles, protons and neutrons, which have the same
atomic number
weight as protons but no charge. All atoms except hydrogen have neutrons
no. of electrons =
atomic number Present day:
no. of neutrons =
Atom is made up of a small, dense, positive nucleus made of protons and
mass no. - atomic neutrons with negatively charged electrons orbiting nucleus. The radius of
no. an atom is very small and is about 10-8cm
Electronic configuration The electronic configuration of an atom can be
predicted from its atomic number. For example, the
atomic number of sodium is 11. Sodium atoms have
11 protons and so 11 electrons:
Electron shell Max electrons
●
2 electrons occupy the first shell
1st 2 ●
8 electrons occupy the second shell
●
1 electron occupies the third shell
2nd 8
3rd 8 This electronic configuration can be written as 2.8.1
(each dot separates one shell from the next). This
electronic configuration can also be shown as a
diagram. In these diagrams:
●
each shell is modelled as a circle
●
each electron is modelled as a dot or a
cross
, Isotopes
Size of an atom See definition on other side. The number of neutrons will be written on the
Atoms have a radius of about 0.1nm (1x10-10m) end, e.g Cl-35 and Cl-37. They are still arranged in the same way, which
and the nucleus is less than 1/10000 of that of means their chemical reactivity is identical. The only difference is that
an atom (1x10-14m). To find out how much that is less neutrons = less weight. Consequently, there are different physical
in m, just make sure there are the same number properties: e.g. Cl-37 has greater density and different melting and boiling
of zeroes (including the one before the decimal) point than Cl-35.
as the number in subscript: e.g 10-10m would be Calculate a “RAM” or “Relative Atomic Mass” number with this equation, in
0.0000000001m this case used for Cl-35 and Cl-37
1 2 3 4 5 6 7 8 9 10
Compounds
There are two definitions for a compound that
are used in CCEA:
●
A compound is a substance made up of 2
or more different types of atoms The term "relative" in the context of atomic mass refers to the fact that
atomic masses are measured relative to a standard reference point, which is
●
A substance made up of 2 or more different the mass of a specific isotope of carbon known as carbon-12. The relative
element chemically combined together atomic mass of an element is defined as the average mass of the naturally
occurring isotopes of that element, weighted by their relative abundances,
Bonding compared to mass of C-12 which is defined as 12 (no units). This is because
of the tiny size of atoms which are far too small to measure.
Cation A positive ion
Ions
Anion A negative ion
An ion is a particle formed when an atom
Compound A substance formed when two or more elements/atoms are loses or gains one or more electrons
chemically combined
Mixture Two or more substances mixed together, which are usually easy
to separate; not chemically combined
Diatomic Two atoms covalently bonded in a molecule
Molecule Two or more atoms covalently bonded together – (specific)
smallest particle of a substance that retains the chemical and
physical properties of the substance
Element A substance made up of only one type of atom and cannot be
broken down into anything simpler by chemical means
●
To have a full outer shell of electrons
Ionic bonding ●
This makes them stable
●
Achieve a noble gas configuration
Ionic bonding occurs when a metal transfers electrons to a non-metal,
producing positive and negative ions.
●
The ionic bonds that form are the electrostatic attractions between When an atom loses electrons it becomes
positive and negative ions positively charged and forms a cation
●
These bonds are strong and require a lot of energy to break them
When it gains electrons, it becomes
These can be represented in dot and cross negatively charged and forms an anion
diagrams, for example:
Electrostatic Always write the charge e.g. Mg2+
Sodium Chloride attraction
Na 2:8:1 between
Cl 2:8:7 oppositely
charged ions
, Covalent bonding Water Lone pair of
electrons – not
involved in bonding
Covalent bonding occurs when non-
metals share an electron pair to obtain a
stable electron structure
These can be presented like H-H or H=H,
with each line representing a shared pair of
electrons
Hydrogen Ammonia
Chloride
Methane Oxygen – double
covalent bond
Carbon dioxide –
two double covalent
bonds
Nitrogen – triple
covalent bond
Metallic bonding
Metallic bonding occurs when metals use
their outermost electrons to form a sea of
electrons, which move in all directions
throughout the metallic structure
These free moving/delocalized electrons
act like glue, holding the metal ions
together in a giant metallic lattice – with
regular rows of tightly packed cations. The e-
electrostatic attraction between the sea of
delocalized electrons and positive metal
ions/cations is known as the metallic bond
, Structures Ionic structures
Malleable A substance that can be hammered into shape
Ductile A substance that can be drawn out into wires
Alloy A mixture of two or more elements, at least one of which a metal and the
resulting mixture has metallic properties
Allotrope Different forms of the same element in the same physical state Each positive (metal) and
negative (non-metal) ion is
surrounded by 6 oppositely
charged ions – forms a lattice
●
High melting and boiling point ●
Electrical conductivity (always giant) – strong ionic
Strong ionic bonds hold ions firmly Ionic compounds ONLY conduct when bonds that give ionic compounds
together in compounds – solids with molten or dissolved in water. They their typical physical properties:
high melting and boiling points – large do not conduct in solid state. When
amounts of energy required to break an ionic compound is melted or ●
Solubility in water
down strong electrostatic ionic dissolved in water the strong Ionic compounds usually dissolve
attractions between oppositely electrostatic attractions are in water. When an ionic compound
charged ions. Bond breaking is overcome and ions are free to move dissolves in water, the ionic lattice
endothermic and energy must be and carry the charge – they cannot breaks down and the ions become
provided to break them conduct when solid as ions are fixed free to move about in the
and not free to move solution. Water molecules surround
both positive and negative ions.
The ions are now said to be
hydrated.
Simple covalent
molecules
Although the covalent
bonds within the
Bonding between non-metals. Covalent molecules are strong
molecules like water, methane, ammonia there are only weak
etc. can be classified as simple covalent forces of attraction
substances. In these molecules, the atoms between the
are held together by strong covalent separate molecules
bonds called van der Waals’
forces
Also known as molecular covalent
●
Low melting and boiling point ●
Do not conduct
Simple covalent substances can be electricity
elements or compounds and they exist as As simple molecular
solids, liquids or gases at room substances are composed of
temperature. Although there are strong small molecules with no
covalent bonds within molecules there charges or free electrons
are only weak attractive forces between they cannot conduct
the molecules and this results in simple electricity
molecular substances having low melting
points and boiling points. Little energy
●
Low solubility in
is needed to separate the molecules as water
it does not involve breaking the strong As simple molecular
covalent bonds which require a great substances have no
amount of energy to break. charges present, they
are only very slightly
soluble in water, if at all.
number of an atom
Subatomic Relative Relative
Mass number Total number of protons and neutrons in the nucleus of an
particle mass charge
atom
Isotope An atom that has the same number of protons (so they are
atoms of the same element) and electrons (therefore are still Proton (p+) 1 +1/+ve
neutral) but have a different number of neutrons (so they have
a different mass number)
Neutron (n0) 1 0
Cation Positive ion
Anion Negative ion
Atom The smallest particle of an element that can exist on its own in Electron (e-) 1/1840 -1/-ve
a stable environment – always neutral
History of the atom
2000 years ago:
Greek philosophers Democritus and Leucippus first suggested matter was
made up of individual particles called “atoms”, however Aristotle believed matter
was made of the “four elements”: air, earth, fire, water. Aristotle was more
renowned and considered correct.
1808:
John Dalton published his Atomic Theory, where he stated that “all elements
are made up of small indivisible particles called atoms”, when the idea was
taken seriously.
1897:
Electron discovered by J.J Thomson, new model of atomic structure proposed
called the Plum Pudding Model – rings of negative electrons embedded in a
sphere of positive charge, just like currants in a Christmas pudding. He also put
forward that the atom was neutral as it contained positive and negative
charges.
1913:
Ernest Rutherford and his researchers used alpha particles to probe inside
atoms which led to discovery of the nucleus. His research shows that the atom
consisted of electrons revolving around a positively charged nucleus.
Positives particles are called protons, 2000x heavier than electrons, which cancel
each other and make the atom neutral. Electrons are contained in much larger
region around nucleus, and most weight is in the nucleus.
1932:
James Chadwick discovers where the extra mass came from – as mass was
larger than expected – he showed that the nucleus of an atom contained two
no. of protons = different types of particles, protons and neutrons, which have the same
atomic number
weight as protons but no charge. All atoms except hydrogen have neutrons
no. of electrons =
atomic number Present day:
no. of neutrons =
Atom is made up of a small, dense, positive nucleus made of protons and
mass no. - atomic neutrons with negatively charged electrons orbiting nucleus. The radius of
no. an atom is very small and is about 10-8cm
Electronic configuration The electronic configuration of an atom can be
predicted from its atomic number. For example, the
atomic number of sodium is 11. Sodium atoms have
11 protons and so 11 electrons:
Electron shell Max electrons
●
2 electrons occupy the first shell
1st 2 ●
8 electrons occupy the second shell
●
1 electron occupies the third shell
2nd 8
3rd 8 This electronic configuration can be written as 2.8.1
(each dot separates one shell from the next). This
electronic configuration can also be shown as a
diagram. In these diagrams:
●
each shell is modelled as a circle
●
each electron is modelled as a dot or a
cross
, Isotopes
Size of an atom See definition on other side. The number of neutrons will be written on the
Atoms have a radius of about 0.1nm (1x10-10m) end, e.g Cl-35 and Cl-37. They are still arranged in the same way, which
and the nucleus is less than 1/10000 of that of means their chemical reactivity is identical. The only difference is that
an atom (1x10-14m). To find out how much that is less neutrons = less weight. Consequently, there are different physical
in m, just make sure there are the same number properties: e.g. Cl-37 has greater density and different melting and boiling
of zeroes (including the one before the decimal) point than Cl-35.
as the number in subscript: e.g 10-10m would be Calculate a “RAM” or “Relative Atomic Mass” number with this equation, in
0.0000000001m this case used for Cl-35 and Cl-37
1 2 3 4 5 6 7 8 9 10
Compounds
There are two definitions for a compound that
are used in CCEA:
●
A compound is a substance made up of 2
or more different types of atoms The term "relative" in the context of atomic mass refers to the fact that
atomic masses are measured relative to a standard reference point, which is
●
A substance made up of 2 or more different the mass of a specific isotope of carbon known as carbon-12. The relative
element chemically combined together atomic mass of an element is defined as the average mass of the naturally
occurring isotopes of that element, weighted by their relative abundances,
Bonding compared to mass of C-12 which is defined as 12 (no units). This is because
of the tiny size of atoms which are far too small to measure.
Cation A positive ion
Ions
Anion A negative ion
An ion is a particle formed when an atom
Compound A substance formed when two or more elements/atoms are loses or gains one or more electrons
chemically combined
Mixture Two or more substances mixed together, which are usually easy
to separate; not chemically combined
Diatomic Two atoms covalently bonded in a molecule
Molecule Two or more atoms covalently bonded together – (specific)
smallest particle of a substance that retains the chemical and
physical properties of the substance
Element A substance made up of only one type of atom and cannot be
broken down into anything simpler by chemical means
●
To have a full outer shell of electrons
Ionic bonding ●
This makes them stable
●
Achieve a noble gas configuration
Ionic bonding occurs when a metal transfers electrons to a non-metal,
producing positive and negative ions.
●
The ionic bonds that form are the electrostatic attractions between When an atom loses electrons it becomes
positive and negative ions positively charged and forms a cation
●
These bonds are strong and require a lot of energy to break them
When it gains electrons, it becomes
These can be represented in dot and cross negatively charged and forms an anion
diagrams, for example:
Electrostatic Always write the charge e.g. Mg2+
Sodium Chloride attraction
Na 2:8:1 between
Cl 2:8:7 oppositely
charged ions
, Covalent bonding Water Lone pair of
electrons – not
involved in bonding
Covalent bonding occurs when non-
metals share an electron pair to obtain a
stable electron structure
These can be presented like H-H or H=H,
with each line representing a shared pair of
electrons
Hydrogen Ammonia
Chloride
Methane Oxygen – double
covalent bond
Carbon dioxide –
two double covalent
bonds
Nitrogen – triple
covalent bond
Metallic bonding
Metallic bonding occurs when metals use
their outermost electrons to form a sea of
electrons, which move in all directions
throughout the metallic structure
These free moving/delocalized electrons
act like glue, holding the metal ions
together in a giant metallic lattice – with
regular rows of tightly packed cations. The e-
electrostatic attraction between the sea of
delocalized electrons and positive metal
ions/cations is known as the metallic bond
, Structures Ionic structures
Malleable A substance that can be hammered into shape
Ductile A substance that can be drawn out into wires
Alloy A mixture of two or more elements, at least one of which a metal and the
resulting mixture has metallic properties
Allotrope Different forms of the same element in the same physical state Each positive (metal) and
negative (non-metal) ion is
surrounded by 6 oppositely
charged ions – forms a lattice
●
High melting and boiling point ●
Electrical conductivity (always giant) – strong ionic
Strong ionic bonds hold ions firmly Ionic compounds ONLY conduct when bonds that give ionic compounds
together in compounds – solids with molten or dissolved in water. They their typical physical properties:
high melting and boiling points – large do not conduct in solid state. When
amounts of energy required to break an ionic compound is melted or ●
Solubility in water
down strong electrostatic ionic dissolved in water the strong Ionic compounds usually dissolve
attractions between oppositely electrostatic attractions are in water. When an ionic compound
charged ions. Bond breaking is overcome and ions are free to move dissolves in water, the ionic lattice
endothermic and energy must be and carry the charge – they cannot breaks down and the ions become
provided to break them conduct when solid as ions are fixed free to move about in the
and not free to move solution. Water molecules surround
both positive and negative ions.
The ions are now said to be
hydrated.
Simple covalent
molecules
Although the covalent
bonds within the
Bonding between non-metals. Covalent molecules are strong
molecules like water, methane, ammonia there are only weak
etc. can be classified as simple covalent forces of attraction
substances. In these molecules, the atoms between the
are held together by strong covalent separate molecules
bonds called van der Waals’
forces
Also known as molecular covalent
●
Low melting and boiling point ●
Do not conduct
Simple covalent substances can be electricity
elements or compounds and they exist as As simple molecular
solids, liquids or gases at room substances are composed of
temperature. Although there are strong small molecules with no
covalent bonds within molecules there charges or free electrons
are only weak attractive forces between they cannot conduct
the molecules and this results in simple electricity
molecular substances having low melting
points and boiling points. Little energy
●
Low solubility in
is needed to separate the molecules as water
it does not involve breaking the strong As simple molecular
covalent bonds which require a great substances have no
amount of energy to break. charges present, they
are only very slightly
soluble in water, if at all.
