1. Atomic structure and the Periodic Table Modul03a(iii)
Materials
Atomic structure and the Periodic Table
Modul03a(iii)

2. Compounds and mixtures - Mixtures
Look at this site as an introduction

3. At the centre of an atom is a nucleus containing protons and neutrons.
Describe the structure of an atom in terms of electrons and a nucleus containing protons and neutrons
At the centre of an atom is a nucleus containing protons and neutrons.
Electrons are arranged around the nucleus in energy levels or shells.
Make sure you can label a simple diagram of an atom

4. Atoms are made up of 3 types of particles electrons , protons  and neutrons . 
Electrons are tiny, very light particles that have a negative electrical charge (-).
Protons are much larger and heavier than electrons and have the opposite charge, protons have a positive charge (+). 
Neutrons are large and heavy like protons, however neutrons have no electrical charge. 

5. Define proton number and nucleon number.
PROTON NUMBER (also called ATOMIC NUMBER)
The proton number is equal to the number of protons in the nucleus of an atom.
The proton number is also equal to the number of electrons in orbit around a neutral atom.
NUCLEON NUMBER (also called MASS NUMBER)
The nucleon number is equal to the total number of nucleons (protons and neutrons) in the nucleus of an atom.
This gives an idea of the total mass of the atom

6. Representing atoms Atoms/elements can be represented like this: AXZ
Where X is the chemical symbol for the element,
Z is the number of protons ie the atomic number,
A is the number of neutrons and protons combined ie the mass number.

7. State the relative charges and approximate relative masses of protons, neutrons and electrons
PARTICLE
RELATIVE CHARGE RELATIVE MASS
proton
neutron
electron
ENERGY LEVEL OR SHELL
MAXIMUM NUMBER OF ELECTRONS
first 2
second 8
third

8. Describe the build-up of electrons in ‘shells’ and understand the significance of the noble gas electronic structures and of valency electrons

9. Electron Shells Electrons are added in order
Begin adding to the first shell
When this is full (2 electrons) start on the next shell
When the second shell is full (8 electrons) start on the third shell
A full outer shell means the atom /particle is very stable and won’t react with other atoms/particles

10. in

11. A Carbon atom has 6 protons

13. A nitrogen atom has 7 protons

15. 23

16. ELEMENTS Elements are made of atoms, which are extremely small.
Sometimes in elements 2 or more atoms join together to form molecules
Because all the atoms are the same though it is still an element
All atoms of a given element have the same chemical properties and contain the same number of PROTONS
The number of PROTONS tells us which element the atom/particle has come from

17. Use proton number and the simple structure of atoms to explain the basis of the Periodic Table (see C9), with special reference to the elements of proton number 1 to 20.
You should be able to draw out the electronic structures for the first 20 elements if you know the proton and nucleon number

18. Define isotopes
Atoms of the same element can have different numbers of neutrons
These are called isotopes
Adding neutrons makes the atom heavier
Some isotopes are radioactive because they are very heavy and spontaneously break down
Not all isotopes are radioactive

19. Ions and ionic bonds

20. CHEMICAL BONDS H – hydrogen will form one bond
O - oxygen will form two bonds
C – carbon will for four bonds
N – nitrogen will form three bonds
Atoms will form bonds by sharing electrons (covalent bonds)
Atoms will form bonds by losing or gaining electrons (ionic bonds)
The number of bonds an atom will form is called its valency

21. Describe the formation of ions by electron loss or gain.
Some atoms need to gain 1 or 2 electrons to fill up their outer shell
These form new positively charged particles called ions
Other atoms need to lose electrons to leave the next full shell exposed as the outer shell
These form new negatively charged particles called ions

22. IONIC BONDING
A giving and taking of electrons
Positively or negatively charged particles called ions are formed

23. They form positively charged ions
Describe the formation of ionic bonds between metals and non-metals as exemplified by elements from Groups I and VII.
Metals usually have 1 or 2 (sometimes 3 eg Al)electrons in the outer shell
If they can donate these to another element then they are left with a full outer shell (8 electrons or 2 if it is the first shell)
They form positively charged ions
Non-metals usually have 3 or more electrons
They need to fill up this shell by accepting or sharing electrons from other elements
If they have 6 or more then they can accept electrons and for negatively charged ions

24. THE IONIC BOND
formed between elements whose atoms need to “lose” electrons
and those which need to gain electrons
electrons are transferred from one atom to the other.

25. Sodium + Chlorine Sodium Chloride
Sodium needs to lose one electron and Chlorine needs to gain one electron
This leaves both Na and Cl with a full outer shell

26. FORMATION OF MAGNESIUM CHLORIDECl Mg Cl e¯
Mg needs to lose two electrons
Each Cl needs to gain 1 electron
This produces one Mg2+ ion and two Cl- ions

27. IONIC BONDING
Animations

28. SODIUM ATOM 2,8,1 CHLORINE ATOM 2,8,7
SODIUM CHLORIDE Na Cl
SODIUM ATOM
2,8,1
CHLORINE ATOM
2,8,7

29. SODIUM ION 2,8 CHLORIDE ION 2,8,8
SODIUM CHLORIDE + Na Cl
SODIUM ION
2,8
CHLORIDE ION
2,8,8
both species now have ‘full’ outer shells; ie they have the electronic configuration of a noble gas

30. SODIUM ION 2,8 CHLORIDE ION 2,8,8
SODIUM CHLORIDE + Na Cl
SODIUM ION
2,8
CHLORIDE ION
2,8,8
Na Na e¯
2,8, ,8
ELECTRON TRANSFERRED
Cl e¯ Cl¯
2,8, ,8,8

31. CHLORINE ATOMS 2,8,7 MAGNESIUM ATOM 2,8,2
MAGNESIUM CHLORIDE Cl Mg
CHLORINE ATOMS
2,8,7
MAGNESIUM ATOM
2,8,2 Cl

32. CHLORIDE IONS 2,8,8 MAGNESIUM ION 2,8
MAGNESIUM CHLORIDE Cl 2+ Mg
CHLORIDE IONS
2,8,8
MAGNESIUM ION
2,8 Cl

33. IONS Charged particle formed during ionic bonding
Positively charged ions have lost electrons
Negatively charged ions have gained electrons

34. COMPOUNDS
Compounds are formed by the chemical combination of two or more different kinds of atoms. (Whole numbers only)
Covalent compounds are made up of molecules
Ionic compounds are made up of ions

35. Explain the formation of ionic bonds between metallic and non-metallic elements.
Metallic elements generally need to lose small numbers of electrons to gain a full outer shell
Non-metallic elements usually need to gain a small number of electrons to gain afull outer shell
In reacting together electrons are donated by the metal and accepted by the non-metal

36. FORMATION OF MAGNESIUM CHLORIDECl Mg Cl e¯
Mg needs to lose two electrons
Each Cl needs to gain 1 electron
This produces one Mg2+ ion and two Cl- ions

37. Describe the lattice structure of ionic compounds as a regular arrangement of alternating positive and negative ions, exemplified by the sodium chloride structure.
GIANT IONIC CRYSTAL LATTICE – NaCl sodium chloride
Oppositely charged ions held in a regular
3-dimensional lattice by electrostatic attraction
The Na+ ion is small enough relative to a Cl¯ ion to fit in the spaces so that both ions occur in every plane.
Cl-
Chloride ion
Na+
Sodium ion

38. Each Na+ is surrounded by 6 Cl¯
and each Cl¯ is surrounded by 6 Na+

39. Physical properties of ionic compounds
Melting point very high
A large amount of energy must be put in to overcome the strong electrostatic attractions and separate the ions.
Strength Very brittle
Any dislocation leads to the layers moving and similarly charged ions being next to each other. The repulsion splits the crystal.
Electrical
Do not conduct when solid - ions are held strongly in the lattice.
Conduct when molten or in aqueous solution - the ions become mobile and conduction takes place.
Solubility
Insoluble in non-polar solvents but soluble in water.
Water as it is a polar solvent and stabilises the separated ions.

40. IONIC COMPOUNDS - ELECTRICAL PROPERTIES
Solid ionic compounds do not conduct electricity
Na+
Cl-
Ions are held strongly together
Positive ons can’t move to the cathode
Negative ions can’t move to the anode
Molten ionic compounds do conduct electricity
Solutions of ionic compounds in water do conduct electricity
Ions have more freedom in a liquid so can move to the electrodes
Na+
Cl-
Cl-
Na+
ANODE
CATHODE
Cl-
Na+
Cl-
Cl-
Na+
Cl-
Na+
Na+

41. Molecules and covalent bonds

42. Covalent bonding is a sharing of electrons to give full outer shells
State that non-metallic elements form non-ionic compounds using a different type of bonding called covalent bonding.
Covalent bonding is a sharing of electrons to give full outer shells
This occurs if larger numbers of electrons are needed to fill the outer shell
Or the element is not reactive enough to donate or accept the electrons
Its all to do with the amount of energy required and how tightly the electrons are held to the atom

43. COVALENT BONDING
A sharing of electrons

45. MOLECULES Molecules are made when two or more atoms bond covalently
A molecule of water is made of two hydrogen atoms and one oxygen atom
It’s formula is H2O

46. A covalent bond
Consists of a shared pair of electrons with one electron being supplied by each atom either side of the bond.
atoms are held together because their nuclei which have an overall positive charge are attracted to the shared electrons + +

47. Formation between atoms of the same element; N2, O2, diamond, graphite
between atoms of different element
CO2, SO2 CCl4, SiCl4 BeCl2

48. In covalent bonding
atoms share electrons to get the nearest noble gas electronic configuration
ie to gain a ‘full’ outer shell
2 electrons if it’s the first shell
Or 8 electrons if its any other shell

49. HYDROGEN H H
atoms share a pair of electrons to form a single covalent bond
A hydrogen MOLECULE is formed
Hydrogen atom needs one electron to complete its outer shell
Another hydrogen atom also needs one electron to complete its outer shell
WAYS TO REPRESENT THE MOLECULE
H H
H H

50. H Cl H Cl Cl H HYDROGEN CHLORIDE
atoms share a pair of electrons to form a single covalent bond
Chlorine atom needs one electron to complete its outer shell
Hydrogen atom also needs one electron to complete its outer shell
WAYS TO REPRESENT THE MOLECULE
H Cl
H Cl

51. H C H C H METHANE WAYS TO REPRESENT THE MOLECULE H
Each hydrogen atom needs 1 electron to complete its outer shell
H C H H H C H H
A carbon atom needs 4 electrons to complete its outer shell
Carbon shares all 4 of its electrons to form 4 single covalent bonds

52. N H N H H N H H AMMONIA WAYS TO REPRESENT THE MOLECULE H
Each hydrogen atom needs one electron to complete its outer shell
H N H H
H N H H H
Nitrogen atom needs 3 electrons to complete its outer shell
Nitrogen can only share 3 of its 5 electrons otherwise it will exceed the maximum of 8
A LONE PAIR REMAINS

53. O H O H O H WATER WAYS TO REPRESENT THE MOLECULE H
Each hydrogen atom needs one electron to complete its outer shell O
H O H
H O H
Oxygen atom needs 2 electrons to complete its outer shell
Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8
2 LONE PAIRS REMAIN

54. HYDROGEN H H H H
both atoms need one electron to complete their outer shell
atoms share a pair of electrons to form a single covalent bond
DOT AND CROSS DIAGRAM
H H
H H

55. C C H H H C H H C H H METHANE H H H H H H H H
Carbon atom needs four electrons to complete its outer shell
Each hydrogen atom needs one electron to complete its outer shell
Carbon shares all 4 of its electrons to form 4 single covalent bonds H
H C H H
DOT AND CROSS DIAGRAM
H C H H

56. N N H N H H N H H H AMMONIA H H H H H H
Nitrogen atom needs three electrons to complete its outer shell
Each hydrogen atom needs one electron to complete its outer shell
Nitrogen can only share 3 of its 5 electrons otherwise it will exceed the maximum of 8
A LONE PAIR REMAINS
H N H
H N H H H

57. WATER H O H O H H
Oxygen atom needs two electrons to complete its outer shell
Each hydrogen atom needs one electron to complete its outer shell
Oxygen can only share 2 of its 6 electrons otherwise it will exceed the maximum of 8
TWO LONE PAIRS REMAIN
H O H
H O H

58. OXYGEN O O O O
Each oxygen atom needs two electrons to complete its outer shell
each oxygen shares 2 of its electrons to form a
DOUBLE COVALENT BOND
O O

59. Draw dot-and-cross diagrams to represent the sharing of electron pairs to form single covalent bonds in simple molecules, exemplified by (but not restricted to) H2, Cl2, H2O, CH4 and HCl
Try Cl2 for yourself

60. SIMPLE COVALENT MOLECULES
Bonding Atoms are joined together within the molecule by covalent bonds
Electrical Don’t conduct electricity as they have no mobile ions or electrons
Solubility Tend to be more soluble in organic solvents than in water

61. Boiling point
Low
because intermolecular forces (van der Waals’ forces) are weak;
as the intermolecular forces are weak, little energy is required to separate molecules from each other so boiling points are low
these forces increase as molecules get larger
e.g. CH °C; C2H °C; C3H °C;
some boiling points are higher than expected for a given mass
because you can get additional forces of attraction eg in water

62. Draw dot-and-cross diagrams to represent the multiple bonding in N2, C2H4 and CO2
Try these for yourselves

63. Answers http://www.chemprofessor.com/bonding.htm

64. Try to set up a table comparing ionic and covalent compounds
Describe the differences in volatility, solubility and electrical conductivity between ionic and covalent compounds.
Try to set up a table comparing ionic and covalent compounds

65. Ionic and molecular (covalent) compounds

66. Giant structures

67. COVALENT NETWORKS
GIANT MOLECULES
MACROMOLECULES
They all mean the same!

68. Describe the giant covalent structures of graphite and diamond.
Many atoms joined together in a regular array by a large number of covalent bonds

69. DIAMOND, GRAPHITE MELTING POINT Very high
structures are made up of a large number of covalent bonds,
all of which need to be broken if the atoms are to be separated.
ELECTRICAL Don’t conduct electricity –
have no mobile ions or electrons
but... Graphite conducts electricity

70. Graphite consists of layered planes of carbon atoms
The hexagonal carbon rings provide the delocalised electrons, allowing easy conduction within the planes.

71. DIAMOND, GRAPHITE STRENGTH Diamond is Hard
exists in a rigid tetrahedral structure
Graphite is soft
In the form of hexagonal layers which move over each other

72. DIAMOND MELTING POINT VERY HIGH
many covalent bonds must be broken to separate the atoms
STRENGTH STRONG
each carbon is joined to four others in a rigid structure
ELECTRICAL NON-CONDUCTOR
No free electrons - all four carbon electrons are used for bonding

73. GRAPHITE MELTING POINT VERY HIGH
many covalent bonds must be broken to separate the atoms
STRENGTH SOFT
each carbon is joined to three others in a layered structure
layers are held by weak van der Waals’ forces
can slide over each other
ELECTRICAL CONDUCTOR
Only three carbon electrons are used for bonding which leaves the fourth to move freely along layers
layers can slide over each other
used as a lubricant and in pencils

74. DIAMOND
GRAPHITE

75. Relate their structures to the use of graphite as a lubricant and of diamond in cutting.
Graphite – lubricant
Layers of atoms slide over each other
Diamond – cutting
Atoms strongly bonded gives a very strong, hard substance

76. Describe the structure of silicon(IV) oxide (silicon dioxide).
Crystalline silicon has the same structure as diamond.
To turn it into silicon dioxide, all you need to do is to modify the silicon structure by including some oxygen atoms.

77. SILICON DIOXIDE MELTING POINT VERY HIGH
many covalent bonds must be broken to separate the atoms
STRENGTH STRONG
each silicon atom is joined to four oxygen atoms -
each oxygen atom are joined to two silicon atoms -
ELECTRICAL NON-CONDUCTOR - no mobile electrons

78. METALLIC
BONDING

79. METALLIC BONDING

80. Involves a lattice of positive ions surrounded by delocalised electrons
Metal atoms achieve stability by “off-loading” electrons to attain the electronic structure of the nearest noble gas.
These electrons join up to form a mobile cloud which prevents the newly-formed positive ions from flying apart due to repulsion between similar charges.

81. Atoms arrange in regular close packed 3-dimensional crystal lattices.
The outer shell electrons of each atom leave to join a mobile “cloud” or “sea” of electrons which can roam throughout the metal. The electron cloud binds the newly-formed positive ions together.

82. Stoichiometry

83. Use the symbols of the elements to write the formulae of simple compounds
This is a relatively simple way to write formulae
This site helps you learn the chemical symbols

84. Complete the table Element or polyatomic group Compound
Chemical formula
Chemical name
NaCl
Sodium chloride
MgCl2
Sodium sulphate
Al2(SO4)3
Aluminium sulphate

85. Deduce the formula of a simple compound from the relative numbers of atoms present
This may be useful to you

86. Deducing chemical formulae
It is possible to work out the chemical formulae of a compound when given the elements present.
For example, if we are told that a certain compound contains both sodium and chlorine, we can deduce its formulae to be NaCl.
Remember when working out the chemical formulae to take into consideration the valency of the elements present.

87. Determine the formula of an ionic compound from the charges on the ions present
You can also use the charges on the ions to write the formulae

88. Charges/valency on ions
Group 1 2 3 4 5 6 7 8
Valency 1+ 2+ 3+ 3- 2- 1-
 Periodic table H Li Be
Transition block B C N O F Ne
 First 20 elements Na Mg
variable Al Si P S Cl Ar K Ca
valencies

89. Chemical formula (ignore ions) Sodium chloride NaCl Magnesium chloride
Chemical compound
Make the simplest neutral compound from the ions
Chemical formula (ignore ions)
Sodium chloride
Na+ Cl-
Check: = 0
NaCl
Magnesium chloride
Mg2+ Cl-           Cl-
Check: = 0
MgCl2
Aluminum chloride
Al3+  Cl-          Cl-          Cl-
Check: = 0
AlCl3
Calcium oxide
Ca2+ O2-
CaO
Sodium oxide
Na+ O2- Na+
Na2O
Lithium phosphide
Li+ P3- Li+ Li+
Li3P
Aluminum oxide
Al3+ O2- Al3+ O2-        O2-
Check: = 0
Al2O3

90. Deduce the formula of a simple compound from a model or a diagrammatic representation
You should be able to write down the formula from structural diagrams eg
C8H18
Try these

91. Define relative atomic mass, Ar
The relative atomic mass is a measure of the mass of one atom of the element

92. To calculate the relative formula mass:
Define relative molecular mass, Mr, as the sum of the relative atomic masses (relative formula mass or Mr will be used for ionic compounds).
To calculate the relative formula mass:
Write the formula of the compound.
Write the numbers of each atom in the formula.
Insert the relative atomic mass for each type of atom.
Calculate the total mass for each element.
Add up the total mass for the compound.

93. Calculate the relative formula mass of the compound with the formula: H2SO4
Answer
(H = 1, S = 32, O = 16)
  H2SO4
  (2 x H)  (1 x S)  (4 x O)
  (2 x 1)  (1 x 32)  (4 x 16)
  2    32    64
   = 98

94. This is known as the molar mass, M, and has the units g mol-1
Define the mole in terms of a specific number of particles called Avogadro’s constant. (Questions requiring recall of Avogadro’s constant will not be set.).
1 mole of a pure substance has a mass equal to its molecular mass expressed in grams.
This is known as the molar mass, M, and has the units g mol-1
One mole of any substance conatins the same number of molecules or atoms
This is called
Avogadro's number = × 1023
One mole of gas occupies 24 litres at room temperature and pressure

95. So 2 moles of a substance would have a mass = 2 x molar mass
How many molecules would it contain?
2x( × 1023)
3 moles of a substance would have a mass = 3 x molar mass etc
3 x( × 1023)
If you search this site there are quizzes etc you might like to do     

96. This leads to the formula: mass = moles x molar mass   If we let:       m = mass of substance in grams,       n = moles of pure substance,       M = molar mass of the pure substance in g mol-1  we can write the equation: n = m x M
This equation can be rearranged to give the following:       n = m ÷ M   (moles = mass ÷ molar mass)       M = m ÷ n   (molar mass = mass ÷ moles)

97. Calculate the mass of 0.25 moles of water
Write the equation: m (in grams) = n x M
Extract the data from the question: n = 0.25 mol
Calculate the molar mass of the substance using the periodic table: M(H2O) = (2 x 1.008) = g mol-1
Substitute the values into the equation and solve: mass = 0.25 x = g

98. Use the molar gas volume, taken as 24 dm3 at room temperature and pressure.

99. What is the volume of 3.5g of hydrogen?
[Ar(H) = 1]
hydrogen exists as H2 molecules, so Mr(H2) = 2,
1 mole or formula mass in g = 2g
Number of moles of hydrogen  = 3.5/2 = 1.75 mol H2
volume H2 = mol H2 x molar volume = 1.75 x 24 = 42 dm3 (or cm3)

100. You may find your textbook useful here
Calculate stoichiometric reacting masses and reacting volumes of solutions; solution concentrations will be expressed in mol/dm3. (Calculations involving the idea of limiting reactants may be set.)
Look at this website and the following pages it shows you how to do some of the calculations
You may find your textbook useful here

101. Given the equation MgCO3(s) + H2SO4(aq) ==> MgSO4(aq) + H2O(l) +CO2(g) What mass of magnesium carbonate is needed to make 6 dm3 of carbon dioxide? [Ar's: Mg = 24, C = 12, O = 16, H =1 and S = 32]
since 1 mole = 24 dm3, 6 dm3 is equal to 6/24 = 0.25 mol of gas
From the equation, 1 mole of MgCO3 produces 1 mole of CO2, which occupies a volume of 24 dm3.
0.25 moles of MgCO3 is need to make 0.25 mol of CO2
formula mass of MgCO3 = x16 = 84
required mass of MgCO3 = mol x formula mass = 0.25 x 84 = 21g

102. You might find some other useful stuff in here