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Year 11 Chemistry
Atomic Structure and The Periodic Table
Pracs, Notes and Worksheets
Atomic Structure
1. Explain the meanings of the numbers 35 and 80 with respect to 80
Br
35
_____________________________________________________________________________________
_____________________________________________________________________________________
2. The letters D, E, G, and so on have been used in place of the usual symbols for the elements.
1
20
7
1D
10E
3G
16
18
8J
14
9L
7M
7
12
3Q
12
6R
5T
23
9
11V
4X
14
6Y
22
12W
(a)
How many different elements are listed?
_____________
(b)
Write the names and symbols for these elements _________________________________________
_______________________________________________________________________________________
(c)
List two letters representing different isotopes of the same element
_______________
(d)
List all sets of two atoms which have the same number of neutrons __________________________
_______________________________________________________________________________________
107
3.
For the palladium isotope
46 Pd
state:
a)
the atomic number
_______
c)
the number of protons
e)
the number of electrons in an atom ____ f)
g)
which of the following atoms are isotopes of palladium:
107
47 Y
b) the mass number
________
108
46 X
107
61 Z
_______
d) the number of neutrons
_______
the number of electrons in a Pd2+ ion ___
105
46 W
__________________________
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4. Using s, p, d notation, write electron configurations for:
a)
an atom of the element with atomic number 16
_______________________________________
b)
an atom of the element with atomic number 25
_______________________________________
c)
the copper(II) ion
_______________________________________
d)
an atom of the element in period three, Group II
_______________________________________
5. Identify the following electron configurations as atoms either in the ground state (neutral) or in an excited
state (cation or anion). Then identify the atom/ion
a)
1s22s22p3
_______________________________________________
b)
1s22s22p63s13p3
_______________________________________________
c)
1s22s22p63s13p63d1
_______________________________________________
d)
1s22s22p63s23p63d64s2
_______________________________________________
6. Write the electron configuration and the appropriate chemical symbol for each of the following species:
a)
the alkali metal in period 2 __________________________________________________________
b)
the third noble gas _________________________________________________________________
c)
the transition metal with 8 electrons in the 3d-subshell ____________________________________
d)
the element with 5 electrons in its fourth shell as its outer shell _____________________________
e)
the ions present in common salt, NaCl _________________________________________________
f)
the ion with a 3+ charge which has the same number of electrons as neon _____________________
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g)
the ion with a 3– charge which has the same number of electrons as argon
_______________________________________________________________________________________
h)
the halogen in period 3 _____________________________________________________________
i)
a magnetic element ________________________________________________________________
j)
a colourless gas in period 2 __________________________________________________________
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The Periodic Table
Dimitri Mendeleev (1839-1907) developed the modern form of the Periodic Table. He arranged the
elements known in 1869 in order of ‘atomic weight’ and began a new row so those elements with similar
chemical properties were grouped together. This work was ground-breaking in that Mendeleev recognised
the importance of chemical properties of the elements and left gaps for elements yet to be discovered.
The Periodic Table as a framework for the study of chemistry
In 1913, Charles Moseley, a research student working under Ernest Rutherford at the University of
Cambridge, determined the atomic number (number of protons) of all of the known elements and it was
realised that this corresponded exactly to the order of the elements on the Periodic Table.
Further work clarified the links between atomic structure, chemical properties and the structure of the
Periodic Table.
Atomic Structure
Remember that the number of protons __________ (equals, is greater than, is less than) the number of
electrons in a (neutral) atom. The electrons fill into shells and subshells in order of increasing energy.
The Period Number of an element (state definition) _____________________________________________
The Periodic Table breaks into blocks, which correspond to the highest energy subshell being filled.
Groups
Periods
s
block
d block
p block
f block
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The Group Number of an element (state definition) _____________________________________________
_______________________________________________________________________________________
The Roman numerals I to VIII (or 0 for the eighth Group) are traditionally only applied to the s and p block
elements. In VCE Chemistry the groups are numbered I to XVIII.
Electron Configuration
This represents the order of filling electrons into shells and subshells of increasing energy.
Some Groups show the relationship between electronic configurations and physical properties very clearly.
Write the electronic configurations for the first 3 members of Group I.
Li (Z = 3)
______________________________________________
Na (Z = 11)
______________________________________________
K (Z = 19)
______________________________________________
State the similarity in these electronic configurations ___________________________________________
_______________________________________________________________________________________
Group I is also called the alkali metals.
You should be able to write the electron configuration of the first thirty elements in terms of subshells,
given their atomic number. For example helium has the electronic configuration of 1s2.
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Properties of the elements
Throughout this section, remember that both the physical and chemical behaviour of an element involves the
use of its outershell electrons. As such, the behaviour is related to the electronic configuration.
Patterns can be seen in the properties of members of:
 the same Group
 the same Period
Two important concepts that influence the properties of the elements are:

core charge

electron-electron repulsion
These properties, in particular core charge, are important in explaining trends in the atomic radius of the
elements. Trends in the atomic radius of the elements can be used to explain trends in ionization energies
and electronegativity. These concepts can in turn be used to describe the behaviour of the outershell
electrons and, hence, the properties of the elements.
Define ‘Core Charge’ _____________________________________________________________________
_______________________________________________________________________________________
Trends for properties are often shown on the short form of the Periodic Table. The rectangle indicating the
short form of the Periodic Table typically refers to the Group I to Group XVIII elements (the ‘s’ and ‘p’
block elements)
For example the trend for
Core Charge
(*)
Core Charge
(*)
The arrow points in the
direction corresponding to
an increase in the property.
Better termed nuclear charge or the effective electric field strength on the outer electrons when
looking at the trend within a Group. However, it is adequate in VCE Chemistry to just use the term
Core Charge.
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Electron-electron repulsion is a measure of the force of the outer-shell electrons pushing against each other.
This tends to spread the electrons further apart. However the increase in electron-electron repulsion is
outweighed by the increase in the core charge as you move from Group I to Group XVII. The result is that
the atomic radii of the elements decrease as you move from left to right across the Periodic Table.
Atomic Radius
The atomic radius of an atom is defined as the distance of closest approach to another atom and is the
distance at which the mutual repulsion of the electron clouds and the mutual attraction of the nuclear charge
of each for the electrons of the other are in equilibrium. The size of an atom in a molecule is the covalent
radius. The size in a metallic crystal is the metallic radius. The values quoted in most sets of data are the
covalent radii for non-metals and metallic radii for metals.
In general it is adequate to think of the atomic radius as the distance from the centre of the atom to the
furthermost electron of the atom.
In your own words briefly explain the pattern in the atomic radii down each Group.
_______________________________________________________________________________________
_______________________________________________________________________________________
_______________________________________________________________________________________
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Looking at the patterns that emerge across a Period, using Period 3 as an example.
Complete the following table:
Element
Electronic configuration
Atomic radius (10-9m)
Na
0.191
Mg
0.160
Al
0.130
Si
0.118
P
0.110
S
0.102
Cl
0.099
Ar
0.095
How does the atomic radius relate to the electronic configuration as you move across a Period?
_____________________________________________________________________
Show the trends for Atomic Radii across a Period and within a Group.
Atomic Radii
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The arrow points in the
direction corresponding to
an increase in the property.
Ionisation energy (IE)
The energy required to remove one electron from a neutral atom in the gas phase1 is referred to as the first
ionisation energy. An atom of an element has as many ionisation energies as there are electrons.
e.g. the energy required for the process: Na(g)  Na+(g) + e- is the first ionisation energy,
and then the second ionisation energy would be: Na+(g)  Na2+(g) + e- and so on.
Use the date below to plot the first ionisation energies of the Period 3 elements. ‘Join the dots’ to plot the
pattern.
Element
First
Ionisation
Energy
(kJmol-1)
Na
Mg
Al
Si
P
S
Cl
Ar
502
744
584
793
1017
1006
1257
> 1526
2400
2200
2000
1800
First
Ionisation 1600
Energy
1400
(kJ mol-1)
1200
1000
800
600
400
200
11
12
13
14
15
16
1
17
18
Atomic Number
Measurements are made in the gas phase so that we are only considering an unbonded atom and hence the stability of the
electron configuration.
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How does this pattern relate to the electronic configurations of the elements?
_______________________________________________________________________________________
_______________________________________________________________________________________
Show the trends for Ionisation Energy across a Period and within a Group.
Ionisation Energy
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The arrow points in the
direction corresponding to an
increase in the property.
Electronegativity
Define the term ‘Electronegativity’ _________________________________________________________
_______________________________________________________________________________________
Electronegativity cannot be directly measured and must be calculated from other atomic or molecular
properties. The most commonly used method of calculation is that originally proposed by Pauling. This
gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from
0.7 to 4.0 (hydrogen = 2.2).
Refer to the electronegativity values in the table above to show the trends for electronegativity across a
Period and within a Group.
Electronegativity
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The arrow points in the
direction corresponding to an
increase in the property.
In your own words briefly explain how the electronegativity pattern for each Group is related to atomic
structure
How can the electronegativity pattern across Period 3 be explained in terms of atomic structure?
The concept of electronegativity is a very useful one to explain general patterns in strong bonding forces
across the Periodic Table.
The percent ionic character of a bond can be determined by looking at the difference in
electronegativities of the atoms bonded together.
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Metallic/Non-metallic Character
The main trends are related to whether the element can be classified as a metal or non-metal. Metallic
bonding is covered in VCE Chemistry Unit 1 and in Unit 2 you study the reactions of metals with the
atmosphere, with acids and in redox reactions. The reactivity of metals and their position in the Periodic
Table is also explored in Unit 1/2.
On the outline of the Periodic Table below, use three different colours to shade the elements that are
classified as:

Metals

Metalloids or semi-metals

Non-metals
Over 80% of the elements are classified as metals. List the general physical properties of a typical metal:
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Complete the following sentence:
In terms of chemical properties, metals tend to ___________ (lose/gain) electrons readily to form
____________ (anions/cations).
Oxidising/Reducing Strength
When a metal ___________ (loses/gains) electrons, it is _____________ (oxidised/reduced).
Hint: OIL RIG (Oxidation Is Loss [of electrons] and Reduction Is Gain [of electrons])
If an atom loses electrons, it causes another atom to gain electrons (i.e. it causes reduction). The atom that
loses electrons is therefore acting as a reducing agent or reductant.
If an atom gains electrons, it causes another atom to lose electrons (ie it causes oxidation). The atom that
gains electrons is therefore acting as an oxidising agent or oxidant.

Insert the appropriate word
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Patterns down a Group
Think about the patterns of the atomic radius (AR) and electronegativity (EN) that occur in the Group I
elements:
Li
Na
K
Rb
Cs
Fr
On the basis of this, would you expect the elements of Group I to act as oxidants or reductants? Explain
your response.
Explain any trend that might be expected in this property as you go down the Group from Li to Fr.
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Patterns across a Period
Consider how the atomic radius and electronegativity might affect the redox properties of the elements in
the same Period. For example, Period 3:
Na Mg Al Si P S Cl Ar
List the elements that you would expect to act as:

Reductants
_______________________

Oxidants
_______________________

Unclear or neither
_______________________
The metals (Na, Mg, Al) all tend to lose their outer valence electrons (be oxidised) and hence act as
reductants. The non-metals (P, S, Cl) are increasingly strong oxidants as EN increases and AR
decreases, causing them to attract electrons more strongly. Si is a metalloid and hence any redox
properties are not easy to predict. Ar is a Noble Gas and as such is unreactive with no redox properties
at all.
What general trend might be seen in the redox properties of elements belonging to the same Period?
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Trends across the oxides of Period 3
Each of the elements will react with oxygen to form an oxide but the bonding and chemical properties of
the oxides change across the Period. Highest (most oxidised form of the element) oxide only is given in
the following table.
Sodium
Magnesium
Formula
Na2O
MgO
Appearance
(at 20C)
Melting Temp
(C)
Bonding Type
White
solid
920
Behaviour in
Water
Aluminium
Silicon
Phosphorus
Sulfur
Chlorine
Al2O3
SiO2
P4O10
SO3
Cl2O7
White solid
White solid
White solid
3802
2027
Colorless
solid
1710
422
Colorless
liquid
17
Colorless
liquid
-92
Ionic
Ionic
Ionic
Covalent
Covalent
Covalent
Basic
Basic
Insoluble
Covalent
Network
Insoluble
Acidic
Acidic
Acidic
The reaction of Group I and II oxides with water produces a basic solution due to the reaction of water
with O2- ions:
Na2O(s) + H2O(l)  2Na+(aq) + 2OH-(aq)
Write the equation for the reaction of magnesium oxide with water:
Aluminium oxide is commonly known as alumina. Whilst alumina is insoluble in water, it can be classed as
an amphoteric oxide as it reacts slowly with dilute acids and bases:
Al2O3(s) + 6H+(aq)  2Al3+(aq) + 3H2O(l)
Al2O3(s) + 2OH-(aq) + 3H2O(l)  2Al(OH)4-(aq)
18/40
Sulfur forms sulfur trioxide, which is composed of chains or rings of SO3 molecules in the solid state,
and SO2 molecules in the gas phase. It reacts with water to form sulfuric acid:
SO3(g) + H2O(l)  H2SO4(aq)
The lower oxide of sulfur, SO2, also reacts with water, forming the weaker sulfurous acid:
SO2(g) + H2O(l)  H2SO3(aq)
19/40
Summary of Trends within the Periodic Table
In this summary of the major trends of the elements the transition elements are not included. Also the noble
gases are not included where indicated by a shaded area.
Arrows within boxes indicate direction of increase in the trend while arrows between boxes indicate
correlations between different measures.
(*)
Better termed nuclear charge or the effective electric field strength on the outer electrons.
20/40
Questions – Atomic Structure and The Periodic Table
1. Identify the following elements:
a)
the element with the smallest atomic radius in period 3. ___________________
b)
the element with the greatest electronegativity in period 3. _____________________
c)
the most reactive metal in period 3. _____________________
d)
the element with the lowest first ionisation energy in period 3. _____________________
e)
the element with the highest first ionisation energy in period 3. _____________________
f)
the element with the smallest (stable)ionic radius in period 3 _____________________
2. Below is a graph of successive ionisation energies for one of the first 30 elements.
a.
Identify the element. ___________________________
b.
State two distinguishing features that lead you to this conclusion
_______________________________________________________________________________
_______________________________________________________________________________
120
100
80
60
Series1
40
20
0
0
2
4
6
21/40
8
10
c.
In the space below use a pencil and ruler to neatly sketch a graph of successive ionisation
energies for magnesium. (The general trend is important. It does not have to be accurate)
3. Write balanced chemical equations for the following reactions:
a)
sodium oxide and water
__________________________________________________________________________
b)
sulphur trioxide and water
_______________________________________________________________________________
c)
aluminium oxide and an acid
_______________________________________________________________________________
d)
aluminium oxide and a base
_______________________________________________________________________________
22/40
4. List the common oxidation states for each of the elements in period 3:
Na
Mg
Al
Si
P
S
Cl
5. Mendeleev organised the known elements into the first Periodic Table in 1869. How does the
arrangement of the elements in the modern Periodic Table differ from the version used by Mendeleev?
_______________________________________________________________________________________
_______________________________________________________________________________________
6. Mendeleev left gaps in his Periodic Table and predicted the discovery of several unknown elements.
Germanium was subsequently discovered and its properties were very similar to those predicted by
Mendeleev.
a)
What information did Mendeleev use to predict the properties of germanium?
_______________________________________________________________________________________
_______________________________________________________________________________________
b)
Which Group of elements was not discovered until the 1890s by William Ramsay?
_______________________________________________________________________________________
c)
Suggest a reason for this.
_______________________________________________________________________________________
_______________________________________________________________________________________
23/40
7.
Using the concept of core charge explain the following trends which occur within the Periodic Table:
a)
electronegativity increases from left to right across a period.
_______________________________________________________________________________________
_______________________________________________________________________________________
b)
the atomic radius of an atom decreases from left to right across a period.
_______________________________________________________________________________________
_______________________________________________________________________________________
c)
the metallic character of an element decreases from left to right across a period.
_______________________________________________________________________________________
_______________________________________________________________________________________
d)
the oxidising strength of an element decreases down a group.
_______________________________________________________________________________
_______________________________________________________________________________
e)
first ionisation energy increases from left to right across a period.
_______________________________________________________________________________
_______________________________________________________________________________
24/40
8.
Circle the correct response and give a reason as to why this trend is seen,
Atomic radius increases/decreases down a group.
Electronegativity increases/decreases down a group.
Ionisation energy increases/decreases down a group.
Oxidation strength increases/decreases down a group
Reducing strength increases/decreases down a group
9.
a)
The nitrogen atom has the following six successive ionisation energies (I.E.) measured in kJmol–1:
I.E.1 = 1400
I.E.2 = 2850
I.E.3 = 4560
I.E.4 = 7450
I.E.5 = 9460
I.E.6 = 53100
Explain why the second ionisation energy is greater than the first ionisation energy.
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
b)
Explain why the sixth ionisation energy is so much greater than the fifth ionisation energy.
_______________________________________________________________________________
_______________________________________________________________________________
_______________________________________________________________________________
25/40
Prac - The Reactivity Series of Metals
PART A: Reactions of Metals with Oxygen
Aim: To observe reactions between metals and oxygen to form metal oxides.
This series of demonstrations will show you the reactions of a variety of different metals with the oxygen of
the air (20% oxygen) and also with pure oxygen.
*
*
Your teacher will cut a piece of sodium.
Observe and describe the freshly cut surface.
Q1.
Describe the appearance of the freshly cut surface of sodium. _______________________________
_______________________________________________________________________________________
Q2.
Describe the appearance of the surface after a few minutes. ________________________________
_______________________________________________________________________________________
The action of air on iron is slow compared with the action of air on sodium and potassium. A freshly cut
surface of sodium is shiny. Exposed to the air, the surface tarnishes rapidly as oxides of sodium are formed.
An equation for the formation of the main product is:
4Na(s) + O2(g) → 2Na2O(s)
Q3.
Why are sodium and potassium stored under kerosene or paraffin oil? ________________________
_______________________________________________________________________________________
Q4.
Why are the metals sodium and potassium not stored under water? ___________________________
_______________________________________________________________________________________
You will see the reaction of sodium and potassium with water in Part B
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Rust is formed on iron in the presence of moist air. Rust (hydrated iron (III) oxide) is porous to oxygen and
moisture and so rusting can proceed until there is no metal left. Some metals react with oxygen to form a
surface layer of oxide that is so strongly bound to the metal that no further reaction occurs. This is called a
PROTECTIVE OXIDE LAYER. Zinc, magnesium, aluminium and chromium, for example, have
protective oxide coatings.
Q5.
Write equations for the formation of:
a) zinc oxide _____________________________________________________________________
b) aluminium oxide ________________________________________________________________
Gold remains shiny because it does not form an oxide layer on its surface.
Q6.
Is the formation of iron oxide in air usually to our advantage or disadvantage? Explain your answer.
_______________________________________________________________________________________
_______________________________________________________________________________________
Q7. Is the formation of an oxide on aluminium usually to our advantage or disadvantage? Explain your
answer.
_______________________________________________________________________________________
_______________________________________________________________________________________
The reaction of metals with oxygen in air is more rapid if the metal is heated.
DO NOT LOOK DIRECTLY AT BURNING MAGNESIUM.
Q8.
Describe what happens when magnesium is heated in a flame until it begins to burn.
_______________________________________________________________________________________
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Q9.
The white powder is magnesium oxide. MgO. Write an equation for the reaction which produces MgO
_______________________________________________________________________________________
Some metals will burn in air only if they are in a finely divided form.
Q10. Describe the effect of heating an iron sheet rod in a flame.
_______________________________________________________________________________________
Q11. Describe what happens when iron filings are sprinkled into a flame. Give a possible explanation for
your observation.
_______________________________________________________________________________________
Silver, platinum and gold do not react with oxygen. Metals such as these, which are not very reactive, are
called NOBLE METALS.
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PART B: Reactions of Metals with Cold Water
APPARATUS SET UP
Metal
Observations
Sodium
Calcium
Magnesium
Potassium
Q12. Name the gas collected in these experiments. _____________________________________
Q13. Note the colour of phenolphthalein with the other product formed when each of the metals used
reacted with water. What type of substance is the other product?
_________________________________________________________________________________
Q14. Which of the metal(s) used in this experiment is/are less dense than water? How can you tell?
_______________________________________________________________________________________
_______________________________________________________________________________________
29/40
PART C: Reactions of Metals with Dilute Acids
Make observations about the reaction of the metals with dilute hydrochloric acid solution.
Metal
Evidence of reaction
Chemical equation
Magnesium
Aluminium
Zinc
Iron
Lead
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Copper
Q15. Write a list of the metals that you have observed in this series of experiments:
_________________________________________________________________________________
Based on your results list these metals into a “reactivity series” with the most reactive metal first and the
least reactive last.
Most Reactive
Least Reactive
Q16. Place the chemical symbols of the metals you have observed on the Periodic Table below.
Q17. By referring to the position of the metals in your reactivity series in the Periodic Table, comment on:
a) The general reactivity of main block metals compared with transition metals. _____________________
_______________________________________________________________________________________
b) The trend in reactivity of metals within a group. ____________________________________________
_______________________________________________________________________________________
c) The trend in reactivity of metals within a period. ____________________________________________
_______________________________________________________________________________________
Place the following metals into the reactivity series and complete the table by writing statements about the
reactivity’s of the metals and complete chemical equations.
List of metals (in random order):
Metal
1.
I
N
C
R
E
A
S
I
N
G
Zn, Al, Ag, Ca, Cu, Fe, K, Pb, Mg and Na.
Reactivity on heating
Reactivity on addition of
with oxygen
water
Burn forming oxides and
peroxides.
React to form ________
2.
____________________
3.
eg. __Na(s) + ____
4.
Burning forming oxides.
eg __ Mg(s) + ____
5.
→ _________________
6.
R
E
A
C
T
I
V
I
T
Y
7.
8.
9.
Do not burn but oxidize
on surface.
→ ____________________
Oxide coating must be
removed.
Reactivity on addition
of acid (eg 2M HCl)
Explosive mixture
forming many products.
React to form “salt” and
____________________
eg __ Zn(s) + ______
Does not react with water.
→ _________________
Reacts with water in the presence
of oxygen.
eg __ Fe(s) + ______
Do not react with water.
→
____________________
Do not react.
Does not react.
10.
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Write balanced chemical equations for the following reactions:
(i)
Aluminium with hydrochloric acid.
_______________________________________________________________________________________
(ii)
Aluminium with sulfuric acid (assuming the protective oxide coating ahs been removed).
_______________________________________________________________________________________
(iii)
Zinc with sulfuric acid.
_______________________________________________________________________________________
(iv)
Lithium with water.
_______________________________________________________________________________________
33/40
Metal Reactivity Task 1
The information on the next two pages below is jumbled up. Rearrange it to make sense according to the
correct scientific report format.
Metal + Water
Metal + Acid
Metal Hydroxide + Hydrogen
Metal Salt + Hydrogen
Metal Oxide + Water
Metal Hydroxide
Metal Oxide + Acid
Metal Salt + Water
Group one elements have one valence electron
The oxidation number of the elements in this group is +1
Group two elements have two valence electrons
The oxidation number of the elements in this group is +2
Factors to consider regarding metal reactivity:
 Nuclear/core charge
 Atomic radius
 Shielding effect
 Electron configuration
 Ionisation energy
Metal reactivity relates to an elements’ ability to be oxidised
Oxidised elements are able to form basic hydroxides and ionic compounds
Oxidation refers to the loss of electrons
In general as atomic radius increases an elements’ ability to lose electrons increases. Why?
How is an elements position in a group related to the electrostatic force of attraction between charged sub-atomic
particles and its electronegativity?
Metallic character increases down a group and decreases across a period. Draw a diagram to represent this
statement.
Metals
Period – reactivity decreases as you go from left to right across a period. Why?
Group – reactivity increases as you go down a group. Why?
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The farther to the left and down the periodic table you go, the easier it is for electrons to be removed, resulting in
higher reactivity
Non-Metals
Period – reactivity increases as you go from the left to the right across a period. Why?
Group – reactivity decreases as you go down the group
The farther up and to the right, the higher the electronegativity
Would Lithium be more reactive than Magnesium in water? Why? How could you test this?
Why is potassium more reactive than sodium?
Write the electronic configuration for all of the metals that you tested? (plus all of the elements in group one and
two)
Transition metals are much less reactive compared to group one and two metals
What is the trend going down group one for electronegativity, reactivity and atomic mass? Give reasons for your
answers.
What is the relationship between ionisation energy and metal reactivity? Use two metals as an example.
Write molecular and ionic equations for each of the reactions that you completed for this activity.
35/40
Prac – Period 3 Group 2 Oxide Trends
The information on the next two pages below is jumbled up. Rearrange it before you do these two activities
in class to make sense according to the correct scientific report format.
Experiment One
Now add 1-2 drops of universal indicator to each test tube. Write your observations and deductions into
your table.
To show the change in properties of some oxides across period 3 by examining their appearances and
reactions
Trends across period 3: oxides
Repeat the above steps but now add dilute NaOH(aq) instead of the dilute HCl(aq). If there is a reaction
suggest an equation and hence classify the compounds as acidic if they reacted with the dilute alkalis.
To a small amount of each oxide in a test tube add about 2mL of water. Record your observations in your
results table.
Most elements can combine with oxygen to form an oxide. An oxide is classified as a binary compound –
a compound in which only two elements are present.
To each oxide with the universal indicator add 2mL of dilute HCl(aq). If there is a reaction classify the
compound as basic in the table.
36/40
Experiment Two
Group | are called the alkali metals because they react explosively with water to form alkaline solutions.
All have an outer electronic configuration of s2 and all react.
Repeat the above using calcium and cold water only.
To show the trends in the properties of Group || elements.
Place a small piece of calcium metal in an evaporating dish and apply a blue flame from the Bunsen burner.
Observe and record the results.
Although not as reactive as the Group | metals, they are still strong reducing agents forming stable positive
2+ ions. Their compounds are ionic but less soluble than their Group | counterparts.
React Mg and Ca with dilute hydrochloric acid and record your observations.
The Group || elements are called alkaline earth metals because they were first extracted from oxides found in
the earth’s crust.
Ignite a piece of magnesium in a crucible set up on a Bunsen burner. Do not look at the very bright flame
when the magnesium is burning – it can hurt your eyes. Examine the residue, add some water and
determine the pH of the resulting solution, using universal indicator paper.
Magnesium is the most used commercially of the Group || elements. It makes a very strong, light alloy, with
aluminium which is used in aircraft and automobile construction.
Group | and || make up the s-block of the periodic table. For the chemist these are the most metallic
elements.
Add a piece of magnesium to water in a test tube to which a few drops of phenolphthalein indicator have
been added. Leave for 15mins and record your observations. Then gently heat the mixture and observe and
record your results.
Examine the metals provided. Record your observations.
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Prac – Acidic and Basic Oxides
The information on the next two pages below is jumbled up. . Rearrange it before
you do this activity in class to make sense according to the correct scientific report
format.
Na2O + H2O
NaOH
Na2O + HCl
2NaCl + H2O
MgO + H2O
Mg(OH)2
MgO + HCl
MgCl2 + H2O
Al(OH)3 + 3H+
Al3+ + 3H2O
Al(OH)3 + OH-
[Al(OH)4]-
Al2O3 + HCl
AlCl3 + H2O
Al2O3 + NaOH
NaAlO2 + H2O
Generally non-metallic oxides are acidic and metallic oxides are basic
Soluble acidic oxides dissolve in water to form acids eg. CO2, SO2, SO3 and NO2
Eg. SO2 + H2O
H2SO3
sulfurous acid
SO3 + H2O
H2SO4
sulphuric acid
Soluble basic oxides dissolve in water to form alkaline solutions (containing OH- ions)
Na2O + H2O
NaOH
exothermic
Insoluble basic oxides include MgO, BaO, CuO and Al2O3
Across a period there is an increase in acidic characteristic
ie. basic
amphoteric
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acidic
Amphoteric means that a compound or molecule is able to act as either an acid or base
Sodium, magnesium, aluminium, phosphorus and sulphur all burn in oxygen to form
oxides according to the following equations:
Na + O2
Na2O
Mg + O2
MgO
Al + O2
Al2O3
P4 + O2
P4O10
S8 + O2
SO2
The acidity of an element’s oxide increase with an increase in oxidation number
The oxidation number of an element in its associated oxide compound increases
across a period
Element
Na
Mg
Al
Si
P
Oxide
Na2O
MgO
Al2O3
SiO2
P4O10
Cation oxidation number
Na2O is basic because the oxide ions have a tendency to attract protons
MgO not as basic as group one oxides because the oxide ions aren’t so free in the
compound due to the force of attraction between the doubly charged cation and anion.
As such more energy is required to break apart these two ions in the oxide compound.
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