Shad Ahmad Chemistry
Topic 1 – Key Concepts in Chemistry
1) Describe how the Dalton model of an atom has changed over time because of the discovery of
subatomic particles
1. Dalton – Atoms were solid spheres.
2. Thompson – Plum-pudding model; ball of positive charge with electrons dotted inside it.
3. Rutherford – Fired alpha particles (positive helium ions) at gold foil. Neutral plum pudding
should get alpha pass through it. However, some alpha particles were deflected. Atom was
concluded to be mostly empty space, most mass in centre, nucleus is positive, and electrons
orbit nucleus.
4. Bohr – Electrons orbit at specific distances; the further an electron is from the nucleus, the
more energy it has. When electron loses energy, EM radiation is emitted and it jumps down
a shell (and vice-versa).
5. Chadwick – Neutrons.
2) Describe the structure of an atom…
…as a nucleus containing protons and neutrons, surrounded by electrons in shells.
3) Recall the relative charge and relative mass of a proton, neutron, and electron.
4) Explain why atoms contain equal numbers of protons and electrons
Because they have no charge – are neutral, so the charge of electrons and protons cancel each other
out (opposite charges).
5) Describe the nucleus of an atom in terms of size compared to the whole atom.
Nucleus is very small compared to the overall size of the atom.
6) Recall that most of the mass of an atom…
…is concentrated in the nucleus.
,Shad Ahmad Chemistry
7) Recall the meaning of the term mass number of an atom.
Mass number = Total number of protons and neutrons in the atom.
8) Describe atoms of a given element as having…
…the same number of protons in the nucleus and that this number is unique to that element.
9) Define isotopes
Isotopes are different atoms of the same element containing the same number of protons but
different numbers of neutrons in their nuclei.
10) Calculate the numbers of protons, neutrons and electrons in atoms given the atomic number
and mass number.
Number of protons is also the number of electrons (the atomic number).
Number of neutrons if (the mass number – the atomic number).
11) Explain how the existence of isotopes results in relative atomic masses of some elements not
being whole numbers.
Because the relative atomic mass of an element is the average mass of one atom of the element,
compared to 1/12 of the mass of one atom of carbon-12.
Relative atomic mass takes two isotopes into account (including how much of it there is). EG:
chlorine-35 is much more abundant than chlorine-37, so chlorine’s RAM is 35.5.
12) Calculate the relative atomic mass of an element from the relative masses and abundances of
its isotopes.
Multiply each relative isotopic mass by its isotopic abundance, and add up the results.
Then divide by the sum of the abundances.
13) Describe how Mendeleev arranged the elements, known at that time, in a periodic table by
using properties of these elements and their compounds.
He sorted elements into groups, based on their properties of these elements and their compounds.
As he did this, he realised if he put the elements in the order of atomic mass, a pattern appeared –
he could put elements with similar chemical properties in columns.
,Shad Ahmad Chemistry
14) Describe how Mendeleev used his table to predict the existence and properties of some
elements not then discovered.
He used the properties of the other elements in the columns with gaps to predict the properties of
undiscovered elements.
15) Explain that Mendeleev thought he had arranged elements in order of increasing relative
atomic mass but this was not always true because of…
…the relative abundance of isotopes of some pairs of elements in the periodic table.
16) Explain the meaning of atomic number of an element in terms of position in the periodic table
and number of protons in the nucleus.
The periodic table shows the elements in order of ascending atomic number.
The group to which the element belongs to corresponds to the number of electrons it has in its outer
shell (group 7 have 7).
The period (row) to which the element belongs to corresponds to the number of shells of electrons it
has.
17) Describe that in the periodic table elements with similar properties are placed…
…in the same vertical columns called groups.
18) Identify elements as metals or non-metals according to their position in the periodic table,
explaining this division in terms of the atomic structures of the elements.
Non-metals on the far right (where the stair-line is that separates).
20) Explain how the electronic configuration of an element is related to its position in the periodic
table.
The number of shells which contain electrons is the same as the period of the element.
The group number tells you how many electrons occupy the outer shell of the element.
21) Explain how ionic bonds are formed by the transfer of electrons between atoms to produce
cations and anions, including the use of dot and cross diagrams.
Negative ions (anions) form when atoms gain electrons – have more electrons than protons.
Positive ions (cations) form when atoms lose electrons – have more protons than electrons.
22) Define an ion.
Ions are charged particles – can be single atoms or groups of atoms.
, Shad Ahmad Chemistry
23) Calculate the numbers of protons, neutrons and electrons in simple ions given the atomic
number and mass number. Use FE (2+) as an example.
24) Explain the formation of ions in ionic compounds from their atoms, limited to compounds of
elements in groups 1, 2, 6 and 7.
Group 1 and 2 elements are metals. They lose electrons to form cations.
Group 6 and 7 elements are non-metals. They gain electrons to form anions.
25) Explain the use of the endings –ide and –ate in the names of compounds.
Ions ending with –ide are negative ions containing only one element.
Ions ending with –ate are negative ions contains oxygen and at least one other element.
28) Explain how a covalent bond is formed.
A covalent bond is a strong bond that forms when a pair of electrons is shared between two non
metals.
29) Recall that covalent bonding results in…
…the formation of molecules.
30) Recall the typical size (order of magnitude) of atoms and small molecules.
Simple molecules generally have sizes of 10 -10 m.
The bonds that form between them are generally about 10 -10 m too.
Topic 1 – Key Concepts in Chemistry
1) Describe how the Dalton model of an atom has changed over time because of the discovery of
subatomic particles
1. Dalton – Atoms were solid spheres.
2. Thompson – Plum-pudding model; ball of positive charge with electrons dotted inside it.
3. Rutherford – Fired alpha particles (positive helium ions) at gold foil. Neutral plum pudding
should get alpha pass through it. However, some alpha particles were deflected. Atom was
concluded to be mostly empty space, most mass in centre, nucleus is positive, and electrons
orbit nucleus.
4. Bohr – Electrons orbit at specific distances; the further an electron is from the nucleus, the
more energy it has. When electron loses energy, EM radiation is emitted and it jumps down
a shell (and vice-versa).
5. Chadwick – Neutrons.
2) Describe the structure of an atom…
…as a nucleus containing protons and neutrons, surrounded by electrons in shells.
3) Recall the relative charge and relative mass of a proton, neutron, and electron.
4) Explain why atoms contain equal numbers of protons and electrons
Because they have no charge – are neutral, so the charge of electrons and protons cancel each other
out (opposite charges).
5) Describe the nucleus of an atom in terms of size compared to the whole atom.
Nucleus is very small compared to the overall size of the atom.
6) Recall that most of the mass of an atom…
…is concentrated in the nucleus.
,Shad Ahmad Chemistry
7) Recall the meaning of the term mass number of an atom.
Mass number = Total number of protons and neutrons in the atom.
8) Describe atoms of a given element as having…
…the same number of protons in the nucleus and that this number is unique to that element.
9) Define isotopes
Isotopes are different atoms of the same element containing the same number of protons but
different numbers of neutrons in their nuclei.
10) Calculate the numbers of protons, neutrons and electrons in atoms given the atomic number
and mass number.
Number of protons is also the number of electrons (the atomic number).
Number of neutrons if (the mass number – the atomic number).
11) Explain how the existence of isotopes results in relative atomic masses of some elements not
being whole numbers.
Because the relative atomic mass of an element is the average mass of one atom of the element,
compared to 1/12 of the mass of one atom of carbon-12.
Relative atomic mass takes two isotopes into account (including how much of it there is). EG:
chlorine-35 is much more abundant than chlorine-37, so chlorine’s RAM is 35.5.
12) Calculate the relative atomic mass of an element from the relative masses and abundances of
its isotopes.
Multiply each relative isotopic mass by its isotopic abundance, and add up the results.
Then divide by the sum of the abundances.
13) Describe how Mendeleev arranged the elements, known at that time, in a periodic table by
using properties of these elements and their compounds.
He sorted elements into groups, based on their properties of these elements and their compounds.
As he did this, he realised if he put the elements in the order of atomic mass, a pattern appeared –
he could put elements with similar chemical properties in columns.
,Shad Ahmad Chemistry
14) Describe how Mendeleev used his table to predict the existence and properties of some
elements not then discovered.
He used the properties of the other elements in the columns with gaps to predict the properties of
undiscovered elements.
15) Explain that Mendeleev thought he had arranged elements in order of increasing relative
atomic mass but this was not always true because of…
…the relative abundance of isotopes of some pairs of elements in the periodic table.
16) Explain the meaning of atomic number of an element in terms of position in the periodic table
and number of protons in the nucleus.
The periodic table shows the elements in order of ascending atomic number.
The group to which the element belongs to corresponds to the number of electrons it has in its outer
shell (group 7 have 7).
The period (row) to which the element belongs to corresponds to the number of shells of electrons it
has.
17) Describe that in the periodic table elements with similar properties are placed…
…in the same vertical columns called groups.
18) Identify elements as metals or non-metals according to their position in the periodic table,
explaining this division in terms of the atomic structures of the elements.
Non-metals on the far right (where the stair-line is that separates).
20) Explain how the electronic configuration of an element is related to its position in the periodic
table.
The number of shells which contain electrons is the same as the period of the element.
The group number tells you how many electrons occupy the outer shell of the element.
21) Explain how ionic bonds are formed by the transfer of electrons between atoms to produce
cations and anions, including the use of dot and cross diagrams.
Negative ions (anions) form when atoms gain electrons – have more electrons than protons.
Positive ions (cations) form when atoms lose electrons – have more protons than electrons.
22) Define an ion.
Ions are charged particles – can be single atoms or groups of atoms.
, Shad Ahmad Chemistry
23) Calculate the numbers of protons, neutrons and electrons in simple ions given the atomic
number and mass number. Use FE (2+) as an example.
24) Explain the formation of ions in ionic compounds from their atoms, limited to compounds of
elements in groups 1, 2, 6 and 7.
Group 1 and 2 elements are metals. They lose electrons to form cations.
Group 6 and 7 elements are non-metals. They gain electrons to form anions.
25) Explain the use of the endings –ide and –ate in the names of compounds.
Ions ending with –ide are negative ions containing only one element.
Ions ending with –ate are negative ions contains oxygen and at least one other element.
28) Explain how a covalent bond is formed.
A covalent bond is a strong bond that forms when a pair of electrons is shared between two non
metals.
29) Recall that covalent bonding results in…
…the formation of molecules.
30) Recall the typical size (order of magnitude) of atoms and small molecules.
Simple molecules generally have sizes of 10 -10 m.
The bonds that form between them are generally about 10 -10 m too.
