topic 2 atoms, elements and compounds

� INTRODUCTION

In this new topic, you will be introduced to atoms, elements and compounds. Let us recap what we have learnt so far; we learnt that the things that we see around us are all matter. They occupy space and have mass. They also have different forms and appearances. Some substances are gases, some are liquids and some are solids; some are hard and shiny but others are soft and dull. Different substances behave differently. For example, iron rusts but gold does not, and copper conducts electricity while sulphur does not. How can these

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� Atoms, Elements and Compounds

LEARNING OUTCOMES

By the end of this topic, you should be able to:

1. Identify how matter is classified based on its atoms;

2. Differentiate between metals and non-metals;

3. Write the chemical symbols for elements;

4 Write the formula of molecules for elements and compounds;

5. Identify similarities and differences between elements, compoundsand mixtures;

6. Describe alloys; and

7. Differentiate between solution, solute and solvent.

TOPIC 2 ATOMS, ELEMENTS AND COMPOUNDS � 23

observations be explained? What are these substances made up of that give them different forms and appearances? Well, you will find the answers later as we begin this topic by identifying how matter is classified based on its atoms, and learning how to differentiate between metals and non-metals. Then, we will learn how to write the chemical symbols for elements and the formula of molecules for elements and compounds. Later, we will learn how to identify similarities and differences between elements, compounds and mixtures followed by an explanation of alloys. Lastly, we will look at solution, solute and solvent, and learn how to prepare them. Are you ready? Let us start the lesson!

NATURE OF ATOMS, ELEMENTS, COMPOUNDS AND MIXTURES

In Topic 1, we learned that matter is anything that has mass and occupies space. Matter, whether it is living or non-living, is made up of aatoms � the almost tiny and ssmall building blocks oof matter. The properties of matter relate not only to the kinds of atoms it contains (ccomposition) but also to the arrangement of these atoms (sstructure). This helps us to classify and describe the many different kinds of matter that can be found around us. All the many kinds of matter can be classified and described in two ways: (i) according to its pphysical state as a gas, liquid or solid (which has been discussed in Topic 1); and (ii) according to its composition � either pure substances or mixtures, as shown in Figure 2.1.

Figure 2.1: Classification of matter

Let us look at these classifications further, starting with pure substances.

2.1

� TOPIC 2 ATOMS, ELEMENTS AND COMPOUNDS 24

2.1.1 Pure Substances

What does a pure substance stand for?

For example, water and ordinary table salt (sodium chloride), which make up the primary components of seawater, are pure substances. Pure substances, in turn, can either be eelements or ccompounds. These two types of pure substances will be discussed further in the following sections. (a) EElement

What is an element?

An element is the simplest substance with the following three features:

(i) It consists of oonly one type of atom;

(ii) It ccannot be broken down into simpler substances either by physical or chemical means; and

(iii) It can exist either as individual atoms or mmolecules as shown in Figure 2.2.

A ppure substance (usually referred to simply as a ssubstance) is matter that has a ffixed composition and ddistinct properties.

An eelement is a ssubstance which ccannot be broken down into simpler substances by chemical or physical methods.


Figure 2.2: Atoms and molecules of an element

Can you give some examples of individual atoms? Examples of elements that consist of individual atoms include aluminium, zinc, iron, calcium and gold. What is your definition of a molecule?

Can you give some examples of molecules? Examples of elements that consist of molecules include oxygen, hydrogen and nitrogen. For example, an oxygen molecule consists of two oxygen atoms whereas a hydrogen molecule consists of two hydrogen atoms which are held together in specific shapes.

A mmolecule consists of ttwo or more atoms of the ssame element, or ddifferent elements, which are cchemically bound together.


Do you know that only about 90 of the 115 presently known elements occur naturally? The remaining ones have been produced artificially by nuclear chemists using high energy particle accelerators. These elements can be further grouped into metals, non-metals and semi-metals as explained in Table 2.1.

Table 2.1: Three Group of Elements

Group Description Example

Metal There are 90 types of metals. Potassium, mercury, lead, magnesium, silver and sodium.

Non-metal There are 18 types of non-metals. Hydrogen, chlorine, bromine, phosphorus, carbon and oxygen.

Semi-metals

There are seven semi-metals or metalloids whose properties are intermediate between metals and non-metals.

Boron, silicon, germanium, arsenic, antimony, tellurium and astatine.

We will learn more about metals and non-metals in subtopic 2.3.

(b) CCompound Let us move on to learn about compound. Firstly, let us define it. Can you

give its definition?

A ccompound is a substance which consists of ttwo or more elements chemically combined together.


We can also say that a compound is a pure substance that is formed when two or more different elements combine chemically; thus, it contains two or more kinds of elements bonded together as shown in Figure 2.3.

Figure 2.3: Molecules of a compound

We can determine a compound by inspecting these three features:

(a) It ccan be broken down into a simpler type of matter (elements) by chemical means (but not by physical means);

(b) It has properties that are ddifferent from its component elements; and

(c) It has a cconstant composition throughout and always contains the same ratio of its component atoms.

Can you provide the examples of compounds? Some examples of compounds include carbon dioxide and sodium chloride. When one atom of carbon combines with two atoms of oxygen, carbon dioxide is formed. Such transformation is said to be a chemical reaction. Similarly, when an atom of sodium combines with an atom of chlorine, sodium chloride is formed. Sodium chloride has different properties from sodium and chlorine.


2.1.2 Mixtures

Before we discuss further, let us first look at the definition of mixture. Can you define it?

Or to give a more detailed definition, we can say that mixtures are combinations of two or more substances that are mixed physically in which each substance retains its own chemical identity and hence its own properties, just like the one shown in Figure 2.4.

Figure 2.4: Mixture of elements and compounds

While pure substances have fixed composition, the ccomposition of mixtures can vary. For example, a cup of sweetened tea can contain either a little sugar or a lot. The substances making up a mixture such as sugar and water are called components of the mixture. Mixture can be further classified as either hhomogenous or hheterogeneous (refer Figure 2.1). What are the differences between them? Well, let us look at homogenous mixture first.

Mixtures are ttwo or more substances that are mmixedtogether but nnot chemically joined.


(a) HHomogeneous Mixture

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For instance, the air is a homogenous mixture of the gaseous substances such as nitrogen, oxygen and smaller amounts of other substances. The nitrogen in the air has all the properties of pure nitrogen because both the pure substance and the mixture contain the same nitrogen molecules. Salt, water and many other substances dissolve in water to form homogenous mixtures. Do you know that homogenous mixtures are also called solutions? This will be further discussed in subtopic 2.7.

(b) HHeterogenous Mixture How about a heterogeneous mixture? What can you say about it?

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Can you think of some examples of heterogeneous mixtures? Some examples are when you mix sand with sugar, water with petrol, dust with air, and sulphur with iron filings.

2.1.3 Separating Components of a Mixture

Are you aware that a mixture can be separated into its components by physical means? This is because each component of a mixture retains its own properties. So, how do we do that? We can do that by using some of these methods:

� Filtration;

� Evaporation;

� Distillation;

� Fractional distillation;

� Crystallisation; and

� Chromatography.

A hheterogeneous mixture is a mixture in which the mixing is nnot uniform and therefore has rregions of different compositions. The components are ddistinguishable.

A hhomogenous mixture is a mixture in which the mixing is uuniform and therefore has a cconstant composition throughout the mixture. The components are iindistinguishable. �


(a) FFiltration Firstly, this method is suitable for an iinsoluble solid and lliquid mixture as

shown in Figure 2.5.

Figure 2.5: Filtration process

Source: http://www.saskschools.ca

As you can see in Figure 2.5, the mixture is passed through a filter. The residue is the substance that remains on the filter paper. The filtrate is the substance that flows through the filter paper. A mixture of sand and water is a good example of this method. Filtration yields sand as the residue and water as the filtrate.

(b) Evaporation The second method of separation is evaporation. Let us define this method

first.

Evaporation is a method of sseparating a ssolid that has been dissolved in a solvent.


When do we use this method? This method is suitable for a ssoluble solid and lliquid mixture. Figure 2.6 shows you how to create evaporation. If a mixture is heated or left over a few days, the solvent or liquid evaporates, leaving the solid as rresidue.

Figure 2.6: Evaporation process

Source: http://www.allrefer.com

For example, a mixture of salt and water can be separated by evaporating the water and leaving the solid salt behind.

(c) Distillation The next method of separation is distillation. Let us define its meaning first.

Distillation is a process to sseparate a ssubstance (in the form of a solution) from its ssolvent.


When do we use this method? This method is suitable for a homogenous mixture or ssolution. The liquid to be separated is evaporated by boiling, and its vapour is then collected through condensation as illustrated in Figure 2.7. The condensed vapour, which is in the form of purified liquid, is called the ddistillate.

Figure 2.7: Distillation process

Source: http://cbskilkenny.ie/quickrevise/Experiments/Chemistryvids/ Distillation/

For example, seawater can be separated by distillation. Water has a much lower boiling point than table salt as water is more volatile. If we boil a solution of salt and water, water will evaporate and leave behind salt. The water vapour is converted back to liquid form on the walls of the condenser.


(d) Fractional Distillation Now, let us move on to the fourth method which is fractional distillation.

What do you know about it? Can you define this method?

When do we use this method? This method is suitable for lliquid-lliquid mixtures with ddifferent boiling points. When heated, the component of the mixture with the lower boiling point will evaporate first and be distilled, followed by the component with the next higher boiling point and so on, as shown in Figure 2.8.

Figure 2.8: Fractional distillation

Source: http://www.chemistrydaily.com

Fractional distillation is a method to separate a mixture of compounds by their bboiling points. This is done by heating them to certain temperatures.


For example, the liquid-liquid mixtures in crude oil can be separated by fractional distillation into its components: petrol at 70�C, followed by naphtha at 140�C, kerosene at 180�C, diesel at 260�C, and so forth.

(e) Crystallisation

Now, let us learn about crystallisation. What does it mean?

For example, a copper (II) sulphate solution can be separated into its components, that is, copper (II) sulphate and water, by heating the solution until it is concentrated. Then, we do filtering, and cool down the hot filtrate to obtain solid copper (II) sulphate in the form of crystals. To obtain the best crystals, the crystallisation should be conducted slowly. A cold filtration separates the crystals from the solvent, which is water.

(f) Chromatography

Lastly, let us look at the final method, which is chromatography. The differing abilities of substances to adhere to the surfaces of various solids such as paper and starch make it possible to separate mixtures. This is the basis of chromatography. But what is the formal definition of chromatography?

Crystallisation is a process of fforming crystals from a uniform solution.

Chromatography refers to a sset of methods used to separate different compounds which normally involve sseparating chemicals and iidentifying them by ccolour.


Ink is a good example for this method. The components in ink, which is a dye mixture, can be separated by paper chromatography as shown in Figure 2.9.

Figure 2.9: Chromatography

Source: http://images.google.com

Identifying Elements, Compounds and MixturesGiven the following substances, (i) circle the substances which are elements, (ii) box up the substances which are compounds, and (iii) tick off the substances which are mixtures. Good luck!

Bar of soap

Sulphur powder

Iron filings and

gold filings

Magnesium ribbon

Sand Sugar Petrol

Water Plastic Pewter Bronze Copper (II) oxide and

carbon Sodium

Calcium chloride

Ethanol Air Iodine Sulphuric

acid Aluminium

foil Charcoal

Marble chips

Vinegar Wood Gold ring Silver ring Rice and

salt

Dilute hydrochloric

acid Crude oil

ACTIVITY 2.1


DIFFERENCES BETWEEN METALS AND NON-METALS

In the previous subtopic, you have learnt that elements can be grouped into three groups: metal, non-metal and semi-metal. However, for this subtopic, we wil only cover the two main groups, which are metals and non-metals. What can we say about metals? Metals are the llargest category of elements and are easy to characterise by their appearance. AAll except mercury are ssolid at rroom temperature and most have the silvery shine. In addition, metals are generally mmalleable, rather than brittle (can be pounded into tthin sheets) and dductile, which means it can be ttwisted and drawn into wwires without breaking. Metals are also ggood conductors of hheat and electricity. Metals react with non-metals to form iionic compounds. For example, the reaction of aluminium with bromine produces aluminium bromide, an ionic compound.

2Al (s) + 3Br2 (l) 2AlBr3 (s)

Most metal oxides are bbasic oxides, which when ddissolved in water react to form metal hydroxides, as in the following example:

CaO (s) + H2O (l) Ca(OH)2 (aq)

Metal oxides also demonstrate their basic ability by reacting with acid to form salt and water as illustrated in the reaction of magnesium oxide with hydrochloric acid. This forms magnesium chloride and water.

MgO (s) + 2HCl (aq) MgCl2 (aq) + H2O (l)

Can you figure out some examples of metal? You can refer to Figure 2.10, which shows some examples of metal. Are you familiar with these metals?

2.2


Figure 2.10: Metals Source: http://images.google.com

Like metals, non-metals are easy to characterise by their appearance. NNon-metals are ggases, lliquids or ssolids at rroom temperature. They are not silvery in appearance and several are bbrightly coloured. The ssolid non-metals are bbrittle rather than malleable, and they are ppoor conductors of hheat and eelectricity. The melting points of non-metals are generally llower than those of metals. Like metals, non-metals react with metals to form iionic compounds. Compounds composed entirely of non-metals are mmolecular compounds that tend to be ggases, lliquids, or llow melting point solids. Among the examples are hydrogen chloride and carbon dioxide. Most nnon-metal oxides are aacidic oxides, which when ddissolved in water react to form aacids, as in the following example (carbon dioxide):

CO2 (g) + H2O (l) H2CO3 (aq) Carbon dioxide dissolves in water to form carbonic acid. Non-metal oxides also dissolve in basic solutions to form salt as shown in the following example:

CO2(g) + 2NaOH(aq) Na2CO3(aq) + H2O(l)


Carbon dioxide reacts with a base, sodium hydroxide, to form a salt, sodium carbonate, and water. The different physical properties of metals and non-metals are summarised in Table 2.2.

Table 2.2: Properties of Metals and Non-Metals

Physical Properties Metals Non-Metals

Appearance of the surface Shiny Dull

Conductivity of electricity Good Poor

Conductivity of heat Good Poor

Melting point High Low

Density High Low

Malleability Malleable Non-malleable

Ductility Ductile Not ductile

CHEMICAL SYMBOLS OF ELEMENTS

How do we represent elements? A set of symbols written in the form of one or two letters are used to represent the atoms of a particular element, just as shown in the examples in the earlier subtopic. The first letter of an elementÊs symbol is always capitalised, and the second letter, if any, is lowercase. For example, the chemical symbol for the element calcium is Ca. Many of the symbols comprise only one or two letters of the elementÊs English name such as H for hydrogen, C for carbon, S for sulphur and so forth. Other symbols are derived from Latin or other languages such as Na for sodium (Latin, natrium), Pb for lead (Latin, plumbum), W for tungsten (German,wolfram).

2.3


Table 2.3 shows you examples of elements that are represented by the first letter of its name.

Table 2.3: One-Letter Symbols of Elements

Element SSymbol

Hydrogen H

Nitrogen N

Phosphorus P

Fluorine F

Iodine I

Sulphur S

Oxygen O

Table 2.4, on the other hand, shows you examples of elements that are represented by two-letter symbols.

Table 2.4: Two-Letter Symbols of Elements

Element SSymbol

Bromine Br

Magnesium Mg

Manganese Mn

Calcium Ca

Chlorine Cl

Neon Ne

Nickel Ni


Lastly, Table 2.5 shows examples of elements whose symbols are derived from Latin names.

Table 2.5: Latin-Based Symbols of Elements

Element LLatin Name SSymbol

Silver Argentum Ag

Copper Cuprum Cu

Mercury Hydragyrum Hg

Potassium Kalium K

Tin Stannum Sn

Iron Ferrum Fe

Lead Plumbum Pb

FORMULAE OF MOLECULES FOR ELEMENTS AND COMPOUNDS

As discussed earlier, the atom is the smallest representative sample of an element. Atoms can be combined to form molecules. Many elements found in nature are in molecular form, that, is two or more of the same type of atoms bounded together. For example, the oxygen normally found in air consists of molecules that contain two oxygen atoms. Any molecule that is made up of ttwo atoms is called a ddiatomic molecule. This molecular form of oxygen can be represented by a chemical formula. The chemical formula for a substance shows the chemical composition of the elements present and the ratio in which the atoms of the elements occur. For a substance composed of molecules, the chemical formula that indicates the aactual

2.4

Identify the symbols for the following elements:

Xenon Osmium Nickel Fluorine Astatine

Barium Radium Platinum Plutonium Silicon

SELF-CHECK 2.1


number and types of atoms in the molecule is called the mmolecular formula. The molecular formula for oxygen is O2. The subscript in the formula tells us that two oxygen atoms are present in each molecule. Other examples of elements that normally occur as diatomic molecules are shown in Table 2.6.

Table 2.6: Molecular Formula of Elements

Element MMolecular formula

Bromine Br2

Nitrogen N2

Fluorine F2

Iodine I2

Chlorine Cl2

Hydrogen H2

Compounds that are composed of molecules are called mmolecular compounds and they contain mmore than one type of atom. For example, a molecule of water consists of two hydrogen atoms and one oxygen atom. Its molecular formula is H2O. An absence of subscript on the O indicates one atom of oxygen per water molecule. Other examples of compounds that exist as molecules are shown in Table 2.7.

Table 2.7: Molecular Formula of Compounds

Compound MMolecular Formula

Hydrogen peroxide H2O2

Hydrogen chloride HCl

Carbon dioxide CO2

Carbon monoxide CO

Methane CH4

Ethene C2H4

Notice how the composition of each compound is given by its chemical formula. Also, notice that these substances are composed only of non-metallic elements.


SIMILARITIES AND DIFFERENCES OF ELEMENTS, COMPOUNDS AND MIXTURES

You have learnt the differences of metals and non-metals. How about elements, compounds and mixtures? Let us study Table 2.8, which shows the similarities and differences between elements, compounds and mixtures.

Table 2.8: Differences between Elements, Compounds and Mixtures

Aspect Element Compound Mixture

Formation Can be obtained by breaking down a compound.

Formed when elements are combined in a chemical reaction.

Formed by mixing different components together physically.

Energy involved in the formation

� Heat or light energy is usually released or absorbed

No energy in the form of light or heat is absorbed or released.

Constituents A pure substance that is made up of one type of atom.

A substance made up of two or more different elements which are chemically combined.

Consists of elements, compounds or both which are combined physically.

Ratio of constituents

The proportion of elements is fixed.

The components or elements of a compound combine in a fixed proportion by mass.

The components or substances in a mixture can be mixed in any proportion by mass.

Ability to be broken down into simpler substances

Cannot be broken down into simpler substances.

Can be broken down into simpler substances by chemical means using heat or electricity. The products are either simpler compounds or the elements which make up the compound.

Can be separated into its components by physical means. The components are only physically separated. They are not converted or broken down into other substances.

Properties of the constituents

Each element has its own distinct properties.

The properties of a compound are different from their constituent elements.

Has the properties of its constituent elements.

2.5


You are given the following: aluminium foil, iron filings, sulphur, magnet, evaporating dish, bunsen burner, and test tube holder.

(i) Put a small amount of iron filings into a beaker. Note the colour of the iron filings. Wrap one end of a magnet in a paper towel and dip it into the iron filings. Record your observations. (Silvery-white metal attracted to the magnet.)

(ii) Put a small amount of sulphur in a beaker. Note the colour of the sulphur. Wrap one end of a magnet in a paper towel and dip it into the sulphur. Record your observations. (Yellow crystals not attracted to the magnet.)

(iii) Mix the contents of the two beakers. Note the colour of the mixture. Wrap one end of a magnet in a paper towel and dip it into the mixture. Record your observations. (A combination of silvery-white metal and yellow crystals. Iron filings still attracted to the magnet but not the sulphur.)

(iv) Line the inside of an evaporating dish with aluminium foil and set it aside for later use.

(v) Put a small amount of the mixture into a test tube. Heat the mixture. Record your observation. (Blackish grey compound observed).

(vi) Pour the contents of the test tube into the evaporating dish lined with aluminium foil. Observe the product and fill in the table below.

Substance Observation Response to Magnet

Iron “

Sulphur X

Mixture “

Compound

Compare the properties of the individual elements, mixture and compound. Give examples to support your answer.

ACTIVITY 2.2


ALLOYS

Now, let us move on to learn about alloys. What is the definition of alloys?

These mixtures are hhomogenous and are prepared by hheating and bbonding the metals together. The resultant alloy has completely ddifferent properties from the starting metals. Figure 2.11 shows you the structure of atoms in a metal which are packed together very closely.

Figure 2.11: A metal structure

As a result, most metals have a hhigh density. Besides that, the layers of atoms in a metal can sslide over each other easily and cause the properties of metals such as being mmalleable and dductile as shown in Figure 2.12.

Figure 2.12: The metal structure before and after a force is applied to it

2.6

An aalloy is a mixture of ttwo or more elements; both can be metals, or one can be a metal and another a non-metal.


In an alloy, the atoms of different metals have different sizes. This makes the layers of atoms slide over each other even harder due to the disruption of the initial orderly layers of metal atoms. Usually a small quantity of other metals need to be added to a pure metal to make it even harder, stronger and tougher. Figure 2.13 shows the structure of an alloy in which the big foreign metal atoms disrupt the orderly atom distribution of the initial metal and stop them from sliding.

Figure 2.13: Structure of an alloy

Why do we need an alloy? The main reason of having an alloy is to eenhance the physical properties of metals. For example, although iron is a strong and malleable metal, it suffers from a disadvantage of being prone to rust. Rust is an oxide of iron. Rust destroys the upper layers of iron and makes the metalÊs surface crumbly and weak. To avoid rust, iron can be mixed with nickel and chromium to make steel. Steel or stainless steel is a highly malleable and strong substance, which is rust proof. Thus, mixing two or more metals can bring advantages and allow for more applications. Some other common alloys are steel, brass, and duraluminium. There are three types of common alloys: aaluminium, iiron and ccopper as explained in Table 2.9.


Table 2.9: Three Types of Alloys

Type of Alloy Description Example

Aluminium Aluminium is a bright metal and conducts heat and electricity well. But it is not a strong metal. An alloy of aluminium that retains its good properties and makes it strong would be ideal for many applications. Alloys of aluminium that are light but strong are duraluminium and magnalium.

� Duraluminium: Also called duralium or duralumin.

The proportions of the metals in duralium are Aluminium � 95%, Copper � 4%, Magnesium � 0.5%, and Manganese � 0.5%. Duralium comprises mostly aluminium, but it is found to be very strong and corrosion resistant. Duralium is used in aircraft engines, car engines, pressure cooker, and industrial cauldrons, under sea vessels and ships.

� Magnalium: Contains Aluminium � 95%, and

Magnesium � 5%. Magnalium is a very hard alloy and also light in weight. Magnalium can be machined easily. It is used in many instruments and structures.

Iron To make use of ironÊs good properties and eliminate the possibility of rusting, iron has to be alloyed with other metals. The most important alloy of iron is steel.

� Steel: This alloy contains 99.5% iron, and

0.5% carbon. The different carbon contents give steel a grade; sometimes it can be as high as 1.5%. Steel is a much harder substance than iron. Steel is used for making nails, screws, railway lines, bridges, buildings, and so on. The applications of steel are limitless.

� Stainless steel: Nickel and chromium are added to

steel to make steel shine and attractive in appearance. Varying proportions of nickel and chromium can give different grades of steel. Besides being shiny, stainless steel is strong and corrosion resistant. It is widely used in making utensils, equipment and tools, and extensively used in many industries as containers.


Copper Copper is a relatively soft metal and is prone to be oxidised in the air. An oxidised copper surface is dull and unattractive. To overcome these drawbacks of copper, it can be alloyed with stannum, zinc and nickel to give brass, bronze and German silver.

� Brass: Brass is an alloy of copper and zinc

which generally has 80% copper and 20% zinc. By varying these proportions, we can obtain various grades of brass. Brass is stronger and more malleable than copper. It is golden in colour. Brass is used to make nuts and bolts, tubes, decorative items such as vases, jewellery, lamps, and so on.

� Bronze: Bronze contains 90% copper and 10%

stannum (tin). Bronze is strong and is used to make coins, medals, statues, decorative items, and so on.

� German silver: German silver has 60% copper, 20%

zinc and 20% nickel. It has a silvery shiny look, hence the name German silver. It is used for electroplating so that items look decorative. This also prevents atmospheric corrosion.

SOLUTIONS

Now, we come to the final subtopic of this topic, which is about solutions. What can you expect to learn in the following? Well, there are three segments in this subtopic, and they focus on solution, solute and solvent, saturated solution, and factors that affect solubility. Let us begin the lesson with solution, solute and solvent.

2.7


2.7.1 Solution, Solute and Solvent

Do you know what solution stands for? Can you define it?

How does it form? A solution is formed when tiny individual particles (<1 mm in diameter) of one substance are uniformly dispersed among the individual particles of the other substance. An example of a solution is sugar water. Individual molecules of sugar are uniformly distributed among the molecules of water. Sugar dissolves or breaks down in water. The dissolved substance, which is sugar, is called the solute while the liquid which dissolves the sugar, which is water, is called the solvent. Water is actually a universal solvent which can dissolve many substances. Based on the given example, we can say that a ssolution is a mmixture obtained by dissolving substances, called the ssolute, iin another substance called the ssolvent. Solutions may be mixtures of two or more solids, liquids or gases, or any one of these in another. Figure 2.14 describes the formation of a solution.

A ssolution is a hhomogeneous mixture of ttwo or mmore substances.


Figure 2.14: Solute, solvent and solution

Air is a good example of gas-gas mixture consisting of other gases such as oxygen, nitrogen and carbon dioxide. Can you guess what type of solution is in a bottle of soft drink? If you open a bottle of soft drink, you will notice bubbles coming out of the liquid. All carbonated beverages have particles of carbon dioxide gas dissolved in them. A soft drink is an example of a gas-liquid solution. We have learnt about metal alloys earlier in this topic. Can you still remember? They are good examples of solid-solid solutions where the solute is a solid and the solvent is also a solid. Do you know that there are two types of solution alloys? They are ssubstitution and iinterstitial as shown in Figure 2.15.

Figure 2.15: Types of alloys


How are substitution alloys being formed? Substitution alloys are formed when atoms of the solute take the position which is normally occupied by a solvent atom. They are formed when two metallic components have similar atomic radii and chemical-bonding characteristics. For example, let us consider the alloy of silver and gold. Interstitial alloys are formed when atoms of the solute occupy the interstitial positions. For an interstitial alloy to form, the component present in the interstitial positions between the solvent atoms must have a much smaller covalent radius than the solvent atoms. For example, steel; it is the combination of an iron alloy and carbon. For your quick reference, these three types of solutions are summarised in Table 2.10.

Table 2.10: Types of Solutions

Type of Solution SSolute Solvent Examples

Gas Gas Gas Air (oxygen, nitrogen, argon and other gases)

Liquid Gas Liquid Solid

Liquid Liquid Liquid

Carbonated water (carbon dioxide in water) Petroleum (mixture of hydrocarbons) Seawater (Sodium chloride in water)

Solid Solid Liquid

Solid Solid

Metal alloys such as brass Dental amalgam (mercury in silver)

Source: Adapted from McMurray and Fay (2001), p. 432

Petroleum is a solution of a liquid in liquid. Discuss what liquids are in petroleum and how you can separate these liquids in the laboratory.

ACTIVITY 2.3


2.7.2 Saturated Solution

Do you know that a solution can be categorised into three categories? They can be a ddilute solution, cconcentrated solution or ssaturated solution. Can you differentiate them? Let us start with dilute solution. A ddilute solution has llittle solute particles dissolved in the solvent. The solvent can dissolve the solute particles more easily. When we add mmore solute into the solvent, it can sstill dissolve. At this level, we call this solution a cconcentrated solution. However, if we keep on adding the solute into the solvent until it reaches a level where the ssolute cannot dissolve any more at that particular temperature, then we call this solution a saturated solution as shown in Figure 2.16.

Figure 2.16: Dilute, concentrated and saturated solutions


If you take solid sodium chloride, NaCl or table salt, and add it to water, dissolution occurs rapidly at first but then slows down as more and more NaCl is added. Eventually, the dissolution stops because a dynamic equilibrium is reached where the number of Na+ and Cl� ions leaving the crystal or solute form to go into solution is equal to the number of ions returning from the solution form to the crystal. At this point, the solution is said to be saturated in that solute. Dissolve Solute + Solvent Solution Crystallise A saturated solution is obtained when the solution is in equilibrium with undissolved solid. Additional solute will not dissolve if added to such a solution. The amount of solute needed to form a saturated solution in a given quantity of solvent is known as the ssolubility of the solute. This is usually expressed in grams of solute in 100 g of solvent. For example, the solubility of NaCl in water is 36 g per 100 mL of water at 20�C. This is the maximum amount of NaCl that can be dissolved in water to give a stable, equilibrium solution at that temperature. If we dissolve 40 g of NaCl per 100 mL of water at 20�C, there is undissolved solute in the solution. We say the solvent has reached its saturation point. The solution is called a saturated solution as no more solute can dissolve in the solvent. It is also possible to dissolve less solute than that needed to form a saturated solution. If we dissolve 30.0 g of NaCl per 100 mL of water at 20�C, the solution is said to be unsaturated because it has the capacity to dissolve more solute as shown in Figure 2.17.


Figure 2.17: Saturated and unsaturated solutions

Source: http://images.google.com

2.7.3 Some Factors Affecting Solubility

As discussed earlier, when a solute dissolves in the solvent, a solution is formed. But when oonly a ssmall amount (or nnone at all) of a ssolute ccan be dissolved in the water, the solute is iinsoluble. Then, when the ssolute is sslightly dissolved in the water, we get a ssuspension. For example, chalkboard dust is insoluble in water and suspends on the water surface forming a suspension. So, the question is how much solute can dissolve in a solvent? Well, it all depends on the following three factors:

(a) SSize of the ssolute particles;

(b) TType of the ssolvent; and

(c) TTemperature of the ssolvent.


The smaller the size of the solute particles, the faster the solute dissolves in the solvent. For example, when you make coffee or tea for a drink, you can see that small crystals of sugar dissolve faster than a cube of sugar in your hot drink. The type of solvent will also affect solubility. In some cases, the solute may not dissolve in a particular solvent but may dissolve in other solvents. For example, sugar will dissolve in water but may not dissolve in other types of solvent like paraffin. Most solutes dissolve in water, hence wwater is a uuniversal solvent. The solubility of most molecular and ionic solids increases with increasing temperature of the solvent. Again, you can observe this when making hot drinks.

1. Compare elements and compounds. Give ONE example for each of them.

2. Define mixture.

3. Name and describe THREE characteristics of metals and non-metals.

4. Describe how you would prepare a saturated solution of copper (II) sulphate.

SELF-CHECK 2.2

Investigation: How much salt crystals of different sizes can be dissolved in a liquid? You are given two 150 ml beakers, glass rod, rock salt, and fine salt. Using an electronic balance, weigh about 20 g each of rock salt and fine salt. Put the rock salt in beaker A and fine salt in beaker B with 100 mL of distilled water. Record the time and stir the contents of each beaker gently using a glass rod. Each time the rock salt and fine salt have dissolved, add another 20 g each of the salts into the respective beakers. Record how many grams are used for both salts until the salts do not dissolve. Record which salt dissolves more. Record the time taken. Compare the different rates and the amount of salts dissolved.

ACTIVITY 2.4


� Atoms are the small building blocks of matter. The composition and structure

of atoms determine whether they are pure substances, or mixtures.

� Pure substances are matter that has a fixed composition and distinct properties. They can be classified into elements and compounds.

� Mixtures are two or more substances that are mixed together but not chemically joined. They can be classified as heterogeneous or homogenous.

� An element is a substance which cannot be broken down into simpler substances by chemical or physical methods. It can be classified into two main groups: metals and non-metals. Each group has its own properties.

� A compound is a substance which consists of two or more elements chemically combined together.

� A heterogeneous mixture is a mixture that does not have a uniform composition.

� A homogeneous mixture is a mixture that has a uniform composition.

� Elements, compounds and mixtures can be different in terms of formation, constituents, energy and many more.

� Metal and non-metal elements have different properties such as appearance of the surface, conductivity of electricity, conductivity of heat, melting point and so on.

� Elements can be represented by one-letter symbols, two-letter symbols, and symbols derived from their Latin names.

� The chemical molecular formula for a substance shows the chemical composition of elements present and the ratio in which the atoms of the elements occur.

� The differences between elements, compounds and mixtures can be determined in terms of formation, energy involved in the formation, constituents and so on.


� Alloys are mixtures of two or more metals.

� A solution is a homogeneous mixture in which one substance (the solute) is dissolved in another substance (the solvent).

� A saturated solution is obtained when the solution is in equilibrium with undissolved solid.

� They are three factors that affect solubility: size of the solute particles, type of the solvent, and temperature of the solvent.

Alloys

Atoms

Compounds

Elements

Metals

Mixtures

Non-metals

Saturated solution

Size

Solute

Solution

Solvent

Suspension

Temperature

Type

Brady, J. E., & Senese, F. (2004). Chemistry: Matter and its changes (4th ed.).

New York: John Wiley & Sons, Inc.

Briggs, J. G. R. (1992). Science in focus chemistry for GCE ÂOÊ Level. Singapore: Pearson Education.

Brown, T. L., Lemay, H. E., & Bursten, B. E. (2000). Chemistry: The central science (8th ed.). New Jersey: Prentice Hall.

Kots, J. C., Treichel, P. M., & Weaver, G. C. (2006). Chemistry: The chemical reactivity (2nd ed.). Victoria, Australia: Thomson Learning.


McMurray, J., & Fay, R. C. (2001). Chemistry (3rd ed.). New Jersey: Prentice Hall.

Timberlake, K. C. (2006). An introduction to general, organic, and biological chemistry (9th ed.). San Francisco, CA: Pearson-Benjamin Cummings.

Whitten, K. W., Davis, R. E., Peck, M. L., & Stanley, G. G. (2010). Chemistry (9th ed.). Belmont: Brooks/Cole.

topic 2 atoms, elements and compounds

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different substances

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classification of matter

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nature of atoms

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