Metal + Water Metal Hydroxide + Hydrogen

Year 11 Chemistry Atomic Structure and The Periodic Table

Pracs, Notes and Worksheets

Atomic Structure 80 1. Explain the meanings of the numbers 35 and 80 with respect to 35 Br

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2. The letters D, E, G, and so on have been used in place of the usual symbols for the elements.

1 20 7 16 18 14 7 12 12 23 9 14 22 1D 10E 3G 8J 9L 7M 3Q 6R 5T 11V 4X 6Y 12W

(a) How many different elements are listed? ______

(b) Write the names and symbols for these elements ______

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(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 ______

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107 3. For the palladium isotope 46 Pd state: a) the atomic number ______b) the mass number ______c) the number of protons ______d) the number of neutrons ______e) the number of electrons in an atom ____ f) the number of electrons in a Pd2+ ion ___ g) which of the following atoms are isotopes of palladium: 107 108 107 105 47 Y 46 X 61 Z 46 W ______

4. Using s, p, d notation, write electron configurations for: a) an atom of the element with atomic number 16 ______

2/40 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 ______g) the ion with a 3– charge which has the same number of electrons as argon

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3/40 h) the halogen in period 3 ______i) a magnetic element ______j) a colourless gas in period 2 ______

The Periodic Table

4/40 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 The Group Number of an element (state definition) ______

5/40 ______

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 ______

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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. Properties of the elements

6/40 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’ ______

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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 The arrow points in the Core Charge Core Charge 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.

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

7/40 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:

8/40 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?

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Show the trends for Atomic Radii across a Period and within a Group.

Atomic Radii 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.

11 Measurements are made in the gas phase so that we are only considering an unbonded atom and hence the stability of the electron configuration.

9/40 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 Na Mg Al Si P S Cl Ar First Ionisation 502 744 584 793 1017 1006 1257 > 1526 Energy (kJmol-1)

1800 First Ionisation 1600 Energy 1400 (kJ mol-1) 1200

Atomic Number 11 12 13 14 15 16 17 18

How does this pattern relate to the electronic configurations of the elements?

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10/40 Show the trends for Ionisation Energy across a Period and within a Group.

Ionisation Energy The arrow points in the direction corresponding to an increase in the property.

Electronegativity

Define the term ‘Electronegativity’ ______

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11/40 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 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

12/40 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.

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.

13/40 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:

Complete the following sentence:

In terms of chemical properties, metals tend to ______(lose/gain) electrons readily to form ______(anions/cations).

 Insert the appropriate word

14/40 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.

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

15/40 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.

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 ______

16/40  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?

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 Aluminium Silicon Phosphorus Sulfur Chlorine

Formula Na2O MgO Al2O3 SiO2 P4O10 SO3 Cl2O7 Appearance White White solid White solid Colorless White solid Colorless Colorless

17/40 (at 20C) solid solid liquid liquid Melting Temp 920 3802 2027 1710 422 17 -92 (C) Bonding Type Ionic Ionic Ionic Covalent Covalent Covalent Covalent Network Behaviour in Basic Basic Insoluble Insoluble Acidic Acidic Acidic Water

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:

+ 3+ Al2O3(s) + 6H (aq)  2Al (aq) + 3H2O(l)

- - Al2O3(s) + 2OH (aq) + 3H2O(l)  2Al(OH)4 (aq)

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)

18/40 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.

19/40 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

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60 Series1

0 0 2 4 6 8 10

20/40 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

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b) sulphur trioxide and water

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c) aluminium oxide and an acid

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d) aluminium oxide and a base

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21/40 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?

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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?

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b) Which Group of elements was not discovered until the 1890s by William Ramsay?

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c) Suggest a reason for this.

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22/40 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.

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b) the atomic radius of an atom decreases from left to right across a period.

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c) the metallic character of an element decreases from left to right across a period.

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d) the oxidising strength of an element decreases down a group.

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e) first ionisation energy increases from left to right across a period.

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23/40 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. 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

a) Explain why the second ionisation energy is greater than the first ionisation energy.

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b) Explain why the sixth ionisation energy is so much greater than the fifth ionisation energy.

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24/40 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. ______

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Q2. Describe the appearance of the surface after a few minutes. ______

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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? ______

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Q4. Why are the metals sodium and potassium not stored under water? ______

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You will see the reaction of sodium and potassium with water in Part B

25/40 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.

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Q7. Is the formation of an oxide on aluminium usually to our advantage or disadvantage? Explain your answer.

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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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26/40 Q9. The white powder is magnesium oxide. MgO. Write an equation for the reaction which produces MgO

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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.

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Q11. Describe what happens when iron filings are sprinkled into a flame. Give a possible explanation for your observation.

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Silver, platinum and gold do not react with oxygen. Metals such as these, which are not very reactive, are called NOBLE METALS.

27/40 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?

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Q14. Which of the metal(s) used in this experiment is/are less dense than water? How can you tell?

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28/40 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

29/40 Copper

Q15. Write a list of the metals that you have observed in this series of experiments:

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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. ______

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b) The trend in reactivity of metals within a group. ______

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c) The trend in reactivity of metals within a period. ______

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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): Zn, Al, Ag, Ca, Cu, Fe, K, Pb, Mg and Na.

Reactivity on heating Reactivity on addition of Reactivity on addition Metal with oxygen water of acid (eg 2M HCl) 1. Burn forming oxides and Explosive mixture I peroxides. React to form ______forming many products. N 2. C ______R 3. E eg. __Na(s) + ____ A Burning forming oxides. React to form “salt” and S 4. → ______I eg __ Mg(s) + ______N 5. Oxide coating must be G → ______removed. eg __ Zn(s) + ______6. R Does not react with water. → ______E 7. Reacts with water in the presence A eg __ Fe(s) + ______C of oxygen. T Do not burn but oxidize → 8. I on surface. ______V Do not react with water. I 9. Do not react. T Does not react. Y

31/40 10.

Write balanced chemical equations for the following reactions:

(i) Aluminium with hydrochloric acid.

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(ii) Aluminium with sulfuric acid (assuming the protective oxide coating ahs been removed).

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(iii) Zinc with sulfuric acid.

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(iv) Lithium with water.

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32/40 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 Hydroxide + Hydrogen

Metal + Acid 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

33/40 Period – reactivity decreases as you go from left to right across a period. Why?

Group – reactivity increases as you go down a group. Why? 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.

34/40 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.

35/40 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.

36/40 37/40 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 acidic

38/40 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 Oxide Cation oxidation number Na Na2O Mg MgO Al Al2O3 Si SiO2 P P4O10

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.

39/40 40/40