1.1 Chemistry & Measurements Scientists throughout the world use the International System of Units, abbreviated SI, for their

measurements. There are seven base units in the SI system.

Measurements can involve very large or very small units that do not correspond with the base SI

units. SI units are conveniently modified by a series of prefixes that represent multiples of the

base units. Thus 1/1000th of a meter (or 0.001m) becomes 1 millimeter or simply 1 mm.

Because the numbers chemists use are often very small or very large, it is convenient to express

these numbers in scientific notation. Notice also that all measurements contain both a number

and the unit of measure.

1.2 Measuring Mass Mass is defined as the amount of matter in an object. The standard SI unit of mass is the

kilogram (1 kg weighs 2.205 lb). Smaller mass units are frequently used in chemistry.

1 gram = 0.001 kg = 1.0 x 103

kg

1 milligram (1 mg) = 1.0 x 103

g

1 microgram (1g) = 1.0 x 106

g

Note that the terms mass and weight are used interchangeably, but they do have different

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meanings. Mass is the amount of matter in an object, while weight is a measure of the pull of

gravity on an object.

1.3 Measuring Length The standard unit of length in the SI system is the meter, abbreviated m (1 m = 39.37 in). A

meter, like a kilogram, is too large a unit of measure for most chemistry work. Chemists

frequently use the following smaller units of measure:

1 centimeter (1 cm) = 0.01 m = 1.0 x 102

m

1 millimeter (1mm) = 1.0 x 103

m

1 micrometer (1m) = 1.0 x 106

m

1 nanometer (1nm) = 1.0 x 109

m

1 picometer (1 pm) = 1.0 x 1012

m

1.4 Measuring Temperature Of the three common temperature scales, the Kelvin scale is generally used for scientific work.

Much of the world uses the Celsius scale, except the United States which uses the Fahrenheit

scale. The kelvin (abbreviated K), as the unit of measure is called, is the same physical increment

as a Celsius degree (abbreviated C). The difference between the two units of measure is that

the corresponding temperature scales are offset by a fixed amount. The Fahrenheit degree

(abbreviated F) is smaller than the Kelvin and Celsius degrees, and the Fahrenheit scale is offset

by a different amount.

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An understanding of the three temperature scales will help you make temperature conversions

without blindly applying a formula. Look at the number of degrees that separate the freezing

point of water and the boiling point of water on the Fahrenheit scale: There are 180 degrees.

This same temperature interval is separated by 100C and 100 K. Thus, one F is 100/180 or

5/9 the size of a kelvin or C. That is, there are fewer Celsius degrees than Fahrenheit degrees in

the same range because Celsius degrees are "fatter". The Fahrenheit scale is offset from the

Celsius scale by 32F. These facts lead to the following conversion equations:

The Kelvin scale is offset from the Celsius scale by 273.15. Thus,

Let's convert a temperature commonly used in baking350Fto C and to kelvins. Our conversion path will be to convert from F to C, then perform a second conversion from C to

K. To convert from F to C, we first subtract the offset of 32.

350F 32 = 318F Next we convert the size of the degrees.

318F (5C/9C) = 177C

To convert C to kelvins, we simply add 273.15, the Kelvin scale offset.

177 + 273.15 = 450 K

1.5 Derived Units: Measuring Volume The seven fundamental SI units are not sufficient to describe units of measurement for such things

as area, volume, density, etc. These units are called derived units because they can be

expressed using one or more of the seven base units.

Volume, the amount of space occupied by an object, is measured in SI units by the cubic

meter, abbreviated m3. Smaller, more convenient measures of volume are frequently used.

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1 dm3 = 1 liter (1 L)

1 cm3

= 1 milliliter (1 mL)

Measuring liquid volume is a common laboratory task. Some of the specialized glassware used in

chemistry labs is shown below. Which one is the buret?

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1.6 Derived Units: Measuring Density Density is an intensive physical property that relates the mass of an object to its volume.

Notice the wide range of densities of common substances listed in the table below.

The volume of many substances changes with temperature, so densities, too, are temperature

dependent.

Density provides a useful link in the laboratory between the mass of a substance and its volume.

It is sometimes simpler to use volumetric glassware to measure a particular volume of a

substance, and then to convert that volume to mass.

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ELEMENTS and COMPOUNDS

MIXTURES and their separation

CHEMICAL REACTIONS and EQUATIONS

KEYWORDS ... atom ... chemical change ... chromatography ... compound ... covalency ... distillation (simple/fractional) ... element ... equations - (word, picture, symbol, quizzes) ... formula ... impure/pure ... ionic equations ... ionic valency ... magnet ... mixture ...

molecule ... physical change ... products ... reactants ... separating mixtures ... chemical

symbols - (elements, formula, in equations) ... state symbols ... valency ... working out formulae ...

Introduction and Some keywords (pictures)

ATOM

An ATOM is the smallest particle of a substance which can have its characteristic properties. BUT remember atoms are built up of even more fundamental sub-atomic particles - the electron, proton and neutron.

A MOLECULE is a larger particle formed by the chemical combination of two or more atoms. The molecule may be an element or a

compound eg hydrogen H2 or carbon dioxide CO2 and the atoms are held

together by covalent bonds.

ELEMENT

and symbols

H I Na Al Fe C Ag U?

An ELEMENT is a pure substance made up of only one type of atom*, 92 in the Periodic Table naturally occur from hydrogen H to

uranium U. Note that each element has symbol which is a single capital

letter like H or U or a capital letter + small letter eg cobalt Co,

calcium Ca or sodium Na. Each element has its own unique set of properties but the Periodic

Table is a means of grouping similar elements together. They may

exist as atoms like the Noble Gases eg helium He or as molecules eg hydrogen H2 or sulphur S8. (more examples applied to equations

and see note about 'formula of elements') * At a higher level of thinking, all the atoms of the same

element, have the same atomic or proton number. This number determines how many electrons the atom has, and so

ultimately its chemistry. Any atom with 27 protons and electrons is cobalt!

COMPOUND and FORMULA

A COMPOUND is a pure substance formed by chemically combining at least two different elements by ionic or covalent bonding

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CH4

Compounds can be represented by a FORMULA, eg sodium chloride NaCl (ionic, 2 elements,1of sodium and 1 of chlorine), methane CH4 (covalent, shown on the left has 2 elements in it, 4 of

carbon and 1 of hydrogen*) and glucose C6H12O6 (covalent, 3 elements, 6 atoms of carbon, 12 of hydrogen and 6 of oxygen). There

must be at least two different types of atom (elements) in a

compound.(* the 1 is never written in the formula, no number means 1)

Compounds have a fixed composition and therefore a fixed ratio

of atoms represented by a fixed formula, however the compound is made or formed.

In a compound the elements are not easily separated by

physical means, and quite often not easily by chemical means either. The compound has properties quite different from the

elements it is formed from. o For example soft silvery reactive sodium + reactive green gas

chlorine ==> colourless, not very reactive crystals of sodium chloride.

The formula of a compound summarises the 'whole number'

atomic ratio of what it is made up of eg methane CH4 is composed of 1 carbon atom combined with 4 hydrogen atoms.

Glucose has 6 carbon : 12 hydrogen : 6 oxygen atoms, sodium

chloride is 1 sodium : 1 chlorine atom. When there is only one atom of the element, there is no

subscript number, the 1 is assumed eg Na in NaCl or C in CH4. When there is more than 1 atom of the same element, a subscript

number is used eg the 4 in CH4 meaning 4 hydrogen atoms. Sometimes, a compound (usually ionic), is partly made up of two or

more identical groups of atoms. To show this more accurately ( ) are

used eg o calcium hydroxide is Ca(OH)2 which makes more sense than

CaO2H2 because the OH group is called hydroxide and exists in

its own right in the compound. o Similarly, aluminium sulphate has the formula

Al2(SO4)3 rather than Al2S3O12, because it consists of two aluminium ions Al3+ and three sulphate ions

SO42-.

The word formula can also apply to elements. eg hydrogen H2,

oxygen O2, ozone O3 (2nd unstable form of oxygen), phosphorus P4, sulphur S8, have 2, 2, 3, 4 and 8 atoms in their molecules. Elements

like helium He are referred to as 'monatomic' because they exist as single uncombined atoms.

MIXTURE A MIXTURE is a material made up of at least two substances which may be elements or compounds. They are usually easily separated by physical means eg filtration, distillation, chromatography etc. Examples: air, soil, solutions.

PURE PURE means that only one substance present in the material and can be an element or compound.

A simple physical test for purity and helping identify a compound is to

measure the boiling point of a liquid. Every pure substance melts and boils at a fixed temperature.

o If a liquid is pure it may boil at a constant temperature (boiling

point). o An impure liquid could boil higher or lower than the expected

boiling point and over a range of temperature.

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o If a solid is pure, it will quite sharply at the melting point. o An impure solid melts below its expected melting point and

more slowly over a wider temperature range.

IMPURE IMPURE usually means a mixture of mainly one substance plus one or more other substances physically mixed in.

The % purity of a compound is important, particularly in drug

manufacture. Any impurities present are less cost-effective to the

consumer and they may be harmful substances.

PURIFICATION Materials are purified by various separation techniques. The idea is to separate the desired material from unwanted material. they include:

o Filtration to separate a solid from a liquid. You may want the solid or the liquid or both!

o Simple distillation to separate a pure liquid from dissolved solid impurities which have a very high boiling point.

o Fractional distillation to separate liquids with a range of different boiling points, especially if relatively close together.

o Crystallisation to get a pure solid out of a solvent solution of

it. o Chromatography can be used on a larger scale than spots'

to separate out pure samples from a mixture.

Picture examples of Elements, Compounds and Mixtures

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METHODS of SEPARATING MIXTURES

Simple Distillation

Distillation involves 2 stages and both are physical

state changes. (1) The liquid or solution mixture is boiled to

vaporise the most volatile component in the mixture

(liquid ==> gas). The ant-bumping granules give a

smoother boiling action. (2) The vapour is cooled by cold water in the

condenser to condense (gas ==> liquid) it back

to a liquid (the distillate) which is collected. This can be used to purify water because the

dissolved solids have a much higher boiling point and

will not evaporate with the steam. BUT it is too simple a method to separate a mixture

of liquids especially if the boiling points are relatively close.

Fractional Distillation

Fractional distillation involves 2 main

stages and both are physical state changes. It can only work with liquids with

different boiling points. (1) The liquid or solution mixture is boiled

to vaporise the most volatile component in

the mixture (liquid ==> gas). The ant-bumping granules give a smoother boiling

action. (2) The vapour passes up through a

fractionating column, where the separation takes place (theory at the end).

This column is not used in the simple distillation described above.

(3) The vapour is cooled by cold water in

the condensor to condense (gas ==>

liquid) it back to a liquid (the distillate) which is collected.

This can be used to separate alcohol from a fermented sugar solution. It is used on a large scale to separate the components of crude oil, because the different

hydrocarbons have different boiling and condensation points FRACTIONAL DISTILLATION THEORY:

o Imagine green liquid is a mixture of a blue liquid (but. 80oC) and a yellow liquid (bpt.

100oC), As the vapour from the boiling mixture enters the fractionating column it begins to cool and condense. The highest boiling or least volatile liquid tends to

condense more ie the yellow liquid (water). The lower boiling more volatile blue liquid

gets further up the column. Gradually up the column the blue and yellow separate from each other so that yellow condenses back into the flask and pure blue distills over to be

collected. The 1st liquid, the lowest boiling point, is called the 1st fraction and each liquid distills over when the top of the column reaches its particular boiling point to give

the 2nd, 3rd fraction etc. o To increase the separation efficiency of the tall fractionating column, it is

usually packed with glass beads, short glass tubes or glass rings etc. which greatly

increase the surface area for evaporation and condensation. o In the distillation of crude oil the different fractions are condensed out at different

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points in a huge fractionating column. At the top are the very low boiling fuel gases like

butane and at the bottom are the high boiling big molecules of waxes and tar.

Paper Chromatography

This method of separation is used to see what coloured materials make up eg a food dye analysis.

The material to be separated eg a food dye (6) is

dissolved in a solvent and carefully spotted onto

chromatography paper alongside known colours on a

'start line' (1-5). The paper is carefully dipped into a solvent, which is

absorbed into the paper and rises up it. Due to different solubilities and different molecular 'adhesion' some colours move more than

others up the paper, so effecting the separation of the different coloured molecules. Any colour which horizontally matches another is likely to be the same molecule ie red (1

and 6), brown (3 and 6) and blue (4 and 6) match, showing these three are all in the food dye

(6).

It is possible to analyse colourless mixture if the components can be made coloured eg protein can be

broken down into amino acids and coloured purple by a chemical reagent called ninhydrin and many colourless organic molecules fluoresce when ultra-violet light is shone on them.

FILTRATION EVAPORATION CRYSTALLISATION

Filtration use a filter paper or fine porous ceramic to

separate a solid from a liquid. It works because the

tiny dissolved particles are too small to be filtered BUT any non-dissolved solid particles are too big to go

through! Evaporation means a liquid changing to a gas or

vapour. In separation, its removing the liquid from a

solution, usually to leave a solid. It can be done quickly with gentle heating or left out to 'dry up' slowly. The solid will almost certainly be less

volatile than the solvent and will remain as a crystalline residue. Crystallisation can mean a liquid substance changing to its solid form. However, the term

usually means what happens when the liquid from a solution has evaporated to a point beyond the solubility limit. Then solid crystals will 'grow' out of the solution.

All three of these separation methods are involved in (1) separation of sand and salt mixtures

or (2) salt preparations eg from dissolving an insoluble base in an acid.

Miscellaneous Separation Methods

MAGNET

This can be used to separate iron from a mixture with sulphur (see below). It is

used in recycling to recover iron and steel from domestic waster ie the 'rubbish' is on a conveyer belt that passes a powerful magnet which pluck's out magnetic

materials.

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PHYSICAL CHANGES

These are changes which do not lead to new substances being

formed. Only the physical state of the material changes. The substance retains exactly the same chemical composition. Examples

...

melting, solid to liquid, easily reversed by cooling eg ice and liquid

water are still the same H2O molecules. dissolving, eg solid mixes completely with a liquid to form a solution,

easily reversed by evaporating the liquid eg dissolving salt in water, on evaporation the original salt is regained.

So freezing, evaporating, boiling, condensing are all physical changes. separating a physical mixture eg chromatography, eg a coloured dye solution is easily

separated on paper using a solvent, they can all be re-dissolved and mixed to form the original

dye. So distillation, filtering are also physical changes.

CHEMICAL CHANGES - REACTIONS - reactants and products

Heating iron and sulphur is classic chemistry experiment. A mixture of silvery grey iron filings and yellow sulphur powder is made. The iron can be plucked out with a magnet ie an easily achieved physical separation because

the iron and sulphur are not chemically combined yet! They are still the same iron and sulphur. On heating the mixture, it eventually glows red on its own and a dark grey solid called iron

sulphide is formed. Both observations indicate a chemical change is happening ie a new

substance is being formed. We no longer have iron or sulphur BUT a new compound with different physical properties

(eg colour) and chemical properties (unlike iron which forms hydrogen with acids, iron sulphide

forms toxic nasty smelling hydrogen sulphide!). iron + sulphur ==> iron sulphide or in symbols: Fe + S ==> FeS AND it is no longer possible to separate the iron from the sulphur using a magnet! So signs that a chemical reaction has happened include:

o colour changes, o temperature changes, o change in mass eg

some solids when burned in air gain mass in forming the oxide eg magnesium forms magnesium oxide

some solids lose mass when heated, eg carbonates lose carbon dioxide in thermal decomposition

Therefore a chemical change is one in which a new substance is formed, by a process

which is not easily reversed and usually accompanied by an energy (temperature) change. This is summarised as reactants ==> products as expressed in chemical

equations in words or symbols.

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THE CONSTRUCTION OF CHEMICAL EQUATIONS

"How to write and understand chemical equations"

Seven equations are presented, but approached in the following way o (1a-7a) the individual symbols and formulae are explained o (1b-7b) the word equation is presented to summarise the change of reactants to products o (1c-7c) a balanced 'picture' equation which helps you understand reading formulae and atom

counting to balance the equation o (1d-7d) the fully written out symbol equation with state symbols (often optional for starter

students)

Chemical Symbols and Formula

For any reaction, what you start with are called the reactants, and what you form are called the products. o So any chemical equation shows in some way the overall chemical change of ... o REACTANTS ==> PRODUCTS, which can be written in words or symbols/formulae.

It is most important you read about formula in an earlier section of this page. empirical formula and molecular formula are dealt with on another page. In the equations outlined below several things have been deliberately simplified. This is to allow the 'starter' chemistry

student to concentrate on understanding formulae and balancing chemical equations. Some teachers may disagree with

this approach BUT my simplifications are: o the word 'molecule' is sometimes loosely used to mean a 'formula', o the real 3D shape of the 'molecule' and the 'relative size' of the different element atoms is ignored o if the compound is ionic, the ion structure and charge is ignored, its just treated as a formula

Molecular and Structural Formulas

A molecular formula gives the types and the count of atoms for each element in a compound. An example

of a molecular formula is ethane, C2H6. Here the formula indicates carbon and hydrogen are combined in

ethane. The subscripts tell us that there are 2 carbon atoms and 6 hydrogen atoms in a formula unit.

The structural formula shows the atoms in a formula unit and the bonds between atoms as lines. Single bonds are one line, Double bonds are two lines. Triple bonds are three lines. The Lewis dot structure shows

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the number of valence electrons and types of bonds in the molecule.

Lewis dot structure Ball and stick model

Electron pairs that are shared are physically between the symbols for the atoms. Electron pairs that are unshared are called lone pairs. Lone pairs are not between atom symbols.

1a A single symbol means an uncombined single atom of the element, or Fe 1 atom of iron, or S 1

atom of sulphur (2Fe would mean two atoms, 5S would mean five atoms etc.)

or the formula FeS means one atom of iron is chemically combined with 1 atom of sulphur to form the compound called iron sulphide

2a

or the formula NaOH means 1 atom of sodium is combined with 1 atom of oxygen and 1 atom of hydrogen to form the compound called sodium hydroxide

or the formula HCl means 1 atom of hydrogen is combined with 1 atom of chlorine to form 1 molecule of the compound called hydrochloric acid

or the formula NaCl means 1 atom of sodium are combined with 1 atom chlorine to form the compound called sodium chloride

or the formula H2O means 2 atoms of hydrogen are chemically combined with 1 atom of oxygen to form the compound called water.

3a

or the symbol Mg means 1 atom of the element called magnesium

or 2HCl means two separate molecules of the compound called hydrochloric acid (see example 2)

or the formula MgCl2 means 1 formula of the compound called magnesium chloride, made of one atom of magnesium and two atoms of chlorine.

or the formula H2 means 1 molecule of the element called hydrogen made up of two joined hydrogen

atoms

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4a

or the formula CuCO3 means one formula of the compound called copper carbonate, made up

of one atom of copper is combined with one atom of carbon and three atoms of oxygen to form the compound copper carbonate

or the formula H2SO4 means one formula of the compound called sulphuric acid, which is made up of two atoms of hydrogen, one atom of sulphur and four atoms of oxygen

or the formula CuSO4 means one formula of the compound called copper sulphate which is made up of one atom of copper, one atom of sulphur and four atoms of oxygen

H2O (example 2)

or the formula CO2 means one molecule of the compound called carbon dioxide which is a chemical combination of one atom of carbon and two atoms of oxygen.

5a

or the formula CH4 means one molecule of the compound called methane which is made of one

atom of carbon combined with four atoms of hydrogen

or 2O2 means two separate molecules of the element called oxygen, and each oxygen

molecule consists of two atoms of oxygen CO2 (see also example 4)

or 2H2O means two separate molecules of the compound called water (see

also example 2)

6a

or the formula Mg(OH)2 is the compound magnesium hydroxide made up of one magnesium, two

oxygen and two hydrogen atoms BUT the OH is a particular combination called hydroxide within a compound, so it

is best to think of this compound as a combination of an Mg and two OH's, hence the use of the ( ).

or 2HNO3 means two separate molecules of the compound nitric acid, each

molecule is made up of one hydrogen atom, one nitrogen atom and three oxygen atoms.

or the formula Mg(NO3)2 is the compound magnesium nitrate, it consists of a

magnesium (ion) and two 'nitrates' (ions), each nitrate consists of one nitrogen and three oxygen atoms, again the

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nitrate is a particular combination of atoms within a compound and hence the use of ( ) again.

or 2H2O meaning two molecules of the compound water (see also examples 2 and 5)

7a or the formula Al2O3 means one formula of the compound called aluminium oxide, made up

of two atoms of aluminium Al and three atoms of oxygen O

or 3H2SO4 meaning three molecules of the compound called sulphuric acid (see also example 4)

or the formula Al2(SO4)3 means one formula of the compound called

aluminium sulphate, it consists of two aluminium, three sulphur and twelve oxygen atoms BUT the SO4 is a particular grouping called sulphate, so it is best to think of the compound as a combination of two Al's and three

SO4's

or 3H2O means three separate molecules of the compound

called water (see also examples 2 and 5)

Chemical word equations

==> means the direction of change from reactants =to=> products no symbols or numbers are used in word equations always try to fit all the words neatly lined up from left to right, especially if its a long word equation eg for

clarity in example 4, some names are split in two parts using two lines, one under the other, this 'style' helps

understanding when it comes to revision!

1b iron + sulphur ==> iron sulphide

2b sodium hydroxide + hydrochloric acid ==> sodium chloride + water

3b magnesium + hydrochloric acid ==> magnesium chloride + hydrogen

4b copper + sulphuric ==> copper + water + carbon o carbonate acid sulphate dioxide

5b methane + oxygen ==> carbon dioxide + water

6b magnesium hydroxide + nitric acid ==> magnesium nitrate + water

7b aluminium oxide + sulphuric acid ==> aluminium sulphate + water

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Chemical picture equations

There are three main points to writing and balancing equations

Writing the correct symbol or formula for each equation component. Using numbers if necessary to balance the equation. if all is correct, then the sum of atoms for each element should be the same on both side of the equation

arrow ..... o in other words: atoms of products = atoms of reactants

This is a chemical conservation law of atoms and later it may be described as the 'law of conservation of mass.

o the 7 equations are first presented in 'picture' style and then written out fully with state symbols o The individual formulas involved and the word equations have already been presented above.

PRACTICE QUESTIONS - words and symbols o Multiple choice quiz on balancing numbers o Word-fill exercises o Reactions of acids with metals, oxides, hydroxides and carbonates.

1c

on average one atom of iron chemically combines with one atom of iron forming one molecule of iron

sulphide atom balancing, sum left = sum right: 1 Fe + 1 S = (1 Fe + 1S) two elements chemically combining to form a new compound

2c

the reactants are one molecule of sodium hydroxide and one molecule of hydrochloric acid the products are one molecule of sodium chloride and one molecule of water all chemicals involved are compounds atom balancing, sum left = right: ( 1 Na + 1 O + 1 H) + (1 H +1 Cl) = (1 Na + 1 Cl) + (2 H's + 1 O)

3c

one atom of magnesium reacts with two molecules of hydrochloric acid the products are one molecule of magnesium chloride and one molecule of hydrogen Mg and H-H are elements, H-Cl and Cl-Mg-Cl are compounds atom balancing, sum left = right: (1 Mg) + (1 H + 1 Cl) + (1 H + 1 Cl) = (1 Mg + 2 Cl's) + (2H's)

4c

the reactants are one formula of copper carbonate and one molecule of sulphuric acid the products are one formula of copper sulphate, one molecule of water and one molecule of carbon

dioxide all molecules are compounds in this reaction

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atom balancing, sum left = sum right: (1 Cu + 1 C + 3 O's) + (2 H's + 1 S + 4 O's) = (1 Cu + 1 S + 4

O's) + (2 H's + 1 O) + (1 C + 2 O's)

5c

one molecule of methane is completely burned by two molecules of oxygen to form one molecule of carbon dioxide and two molecules of water atom balancing, sum left = sum right: (1 C + 4 H's) + (2 O's) + (2 O's) = (1 C + 2 O's) + (2 H's + 1 O)

+ (2 H's + 1 O)

6c

one formula of magnesium hydroxide reacts with two molecules of nitric acid to form one formula of

magnesium nitrate and two molecules of water (all compounds) atom balancing, sum left = sum right: (1 Mg + 2O's + 2 H's) + (1 H + 1 N + 3O's) + (1 H + 1 N +

3O's) = (1 Mg + 2 N's + 6 O's) + (2 H's + 1 O) + (2 H's + 1 O)

7c

one formula of aluminium oxide reacts with three molecules of sulphuric acid to form one formula of aluminium sulphate and three molecules of water note the first use of numbers (3) for the sulphuric acid and water! so picture three of them in your head, otherwise the picture gets a bit big! atom balancing, sum left = sum right: (2 Al's + 3 O's) + 3 x (2 H's + 1 S + 4 O's) = (2 Al's + 3 S's + 12

O's) + 3 x (2 H's + 1 O)

Chemical symbol equations (rules already stated above)

1d Fe(s) + S(s) ==> FeS(s) atom balancing, sum left = sum right: 1 Fe + 1 S = (1 Fe + 1S) all the reactants (what you start with) and all the products (what is formed) are all solids in this case. When first learning symbol equations you probably won't use state symbols at first (see end note).

2d NaOH(aq) + HCl(aq) ==> NaCl(aq) + H2O(l) atom balancing, sum left = right: (1 Na + 1 O + 1 H) + (1 H +1 Cl) = (1 Na + 1 Cl) + (2 H's + 1 O)

3d Mg(s) + 2HCl(aq) ==> MgCl2(aq) + H2(g) atom balancing, sum left = right: (1 Mg) + 2 x (1 H + 1 Cl) = (1 Mg + 2 Cl's) + (2H's)

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4d CuCO3(s) + H2SO4(aq) ==> CuSO4(aq) + H2O(l) + CO2(g) balancing sum left = sum right: (1 Cu + 1 C + 3 O's) + (2 H's + 1 S + 4 O's) = (1 Cu + 1 S + 4 O's) +

(2 H's + 1 O) + (1 C + 2 O's)

5d CH4(g) + 2O2(g) ==> CO2(g) + 2H2O(l) atom balancing, sum left = sum right: (1 C + 4 H's) + 2 x (2 O's) = (1 C + 2 O's) + 2 x (2 H's + 1 O)

6d Mg(OH)2(aq) + 2HNO3(aq) ==> Mg(NO3)2(aq) + 2H2O(l) atom balancing, sum left = sum right: (1 Mg + 2O's + 2 H's) + 2 x (1 H + 1 N + 3 O's) = (1 Mg + 2 N's

+ 6 O's) + 2 x (2 H's + 1 O)

7d Al2O3(s) + 3H2SO4(aq) ==> Al2(SO4)3(aq) + 3H2O(l) atom balancing, sum left = sum right: (2 Al's + 3 O's) + 3 x (2 H's + 1 S + 4 O's) = (2 Al's + 3 S's + 12

O's) + 3 x (2 H's + 1 O)

NOTE 1: means a reversible reaction, it can be made to go the 'other way' if the conditions are changed. Example:

o nitrogen + hydrogen ammonia

o N2(g) + 3H2(g) 2NH3(g) o balancing: 2 nitrogen's and 6 hydrogen's on both sides of equation

Note 2 on the state symbols X(?) of reactants or products in equations

(g) means gas, (l) means liquid, (s) means solid and (aq) means aqueous solution or dissolved in water

eg carbon dioxide gas CO2(g), liquid water H2O(l), solid sodium chloride 'salt' NaCl(s) and copper sulphate solution CuSO4(aq)

VALENCY - COMBINING POWER - FORMULA DEDUCTION

(2nd draft) The valency of an atom or group of atoms is its numerical combining power with other atoms or groups

of atoms. The theory behind this, is all about stable electron structures!

o The combining power or valency is related to the number of outer electrons. o You need to consult the page on "Bonding" to get the electronic background.

A group of atoms, which is part of a formula, with a definite composition, is sometimes referred to as a radical. In the case of ions, the charge on the ion is its valency or combining power (list below). To work out a formula by combining 'A' with 'B' the rule is:

o number of 'A' x valency of 'A' = number of 'B' x valency of 'B', However it is easier perhaps? to grasp with ionic compound formulae.

o In the electrically balanced stable formula, the total positive ionic charge must equal the total negative ionic

charge. Example: o Aluminium oxide consists of aluminium ions Al3+ and oxide ions O2- o number of Al3+ x charge on Al3+ = number of O2- x charge on O2- o the simplest numbers are 2 of Al3+ x 3 = 3 of O2- x 2 (total 6+ balances total 6-) o so the simplest whole number formula for aluminum oxide is Al2O3

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Examples of ionic combining power of ions (left, valency = numerical

charge value)

Examples of covalent combining power of atoms (valencies below)

Hydrogen H (1) Chlorine Cl and other halogens (1)

Oxygen O and sulphur S (2) Boron B and aluminium Al (3)

Nitrogen (3, 4, 5) Carbon C and silicon Si (4)

Phosphorus (P 3,5)

Examples of working out covalent formulae

'A' (valency) 'B' (valency) deduced formula

1 of carbon C (4) balances 4 of hydrogen H (1) 1 x 4 = 4 x 1 = CH4

1 of nitrogen (3) balances 3 of chlorine Cl (1) 1 x 3 = 3 x 1 = NCl3

1 of carbon C (4) balances 2 of oxygen O (2) 1 x 4 = 2 x 2 = CO2

The diagram on the left illustrates the

three covalent examples above for

methane CH4

nitrogen trichloride NCl3

carbon dioxide CO2

Examples of working out ionic formulae

'A' (charge=valency) 'B' (charge=valency) deduced formula

2 of Na+ (1) balances 1 of O2- (2) 2 x 1 = 1 x 2 = Na2O

1 of Mg2+ (2) balances 2 of Cl- (1) 1 x 2 = 2 x 1 = MgCl2

1 of Fe3+ (3) balances 3 of F- (1) 1 x 3 = 3 x 1 = FeF3

2 of Fe3+ (3) balances 3 of SO42- (2) 2 x 3 = 3 x 2 = Fe2(SO4)3

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The Structure of Atoms

Atoms are the smallest particles of matter whose properties we study in Chemistry. However from experiments

done in the late 19th and early 20th century it was deduced that atoms were made up of three fundamental sub-atomic particles (listed below)

A Portrait of an Atom

The diagram below gives some idea on the structure of an atom, it also includes some important

definitions and notation used to describe atomic structure. The atomic number (Z) is also known as the proton number of the nucleus of a particular element. It is the proton number that determines the specific identity of a particular element and its electron

structure. The mass number (A) is also known as the nucleon number (N), that is the number of particles in

the nucleus of a particular isotope. Protons and neutrons are the nucleons present in the positive nucleus and the negative

electrons are held by the positive nucleus in 'orbits' called energy levels or shells. In a neutral atom the number of protons equals the number of electrons.

ISOTOPES

Isotopes are atoms of the same element with different numbers of neutrons. This gives each

isotope of the element a different mass or nucleon number but being the same element they have

the same atomic or proton number. There are small physical differences between the isotopes eg the heavier isotope has a greater

density or boiling point. However, because they have the same number of protons they have the same electronic structure and

identical chemically. Examples are illustrated below. Do NOT assume the word isotope means it is radioactive, this depends on the stability of the nucleus

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i.e. unstable atoms (radioactive) might be referred to as radioisotopes. Many isotopes are stable and NOT radioactive i.e. most of the atoms that make up you and the world around you!

, and are the three isotopes of hydrogen with mass numbers of 1, 2 and 3, with 0, 1 and 2 neutrons respectively, but all have 1 proton. Hydrogen-1 is the most common, there is a trace of hydrogen-2 naturally but hydrogen-3 is very unstable and is used in atomic fusion weapons.

and are the two isotopes of helium with mass numbers of 3 and 4, with 1 and 2

neutrons respectively but both have 2 protons. Helium-3 is formed in the Sun by the initial nuclear

fusion process. Helium-4 is also formed in the Sun and as a product of radioactive alpha decay of an unstable nucleus. An alpha particle is a helium nucleus, it picks up two electrons and becomes the

atoms of the gas helium.

and are the two isotopes of sodium with mass numbers of 23 and 24, with 12

and 13 neutrons respectively but both have 11 protons. Sodium-23 is quite stable e.g. in common salt (NaCl, sodium chloride) but sodium-24 is a radio-isotope and is a gamma emitter used in medicine as a

radioactive tracer e.g. to examine organs and the blood system. The relative atomic mass of an element is the average mass of all the isotopes present compared to

1/12th of the mass of carbon-12 atom (12C = 12.00000 ie the standard).

The Electronic Structure of Atoms

(electron configuration, arrangement in shells or energy levels)

The electrons are arranged in energy levels or shells around the nucleus and with increasing distance from the nucleus.

Each electron in an atom is in a particular energy level (or shell) and the electrons must occupy the lowest available energy level (or shell) available nearest the nucleus.

When the level is full, the next electron(s) go into the next highest level (shell) available. There are rules about the maximum number of electrons allowed in each shell and you

have to be able to work out the arrangements for the first 20 elements (see the Periodic Table diagrams further down).

o The 1st shell has a maximum of 2 electrons o The 2nd shell has a maximum of 8 electrons o The 3rd shell has a maximum of 8 electrons o The 19th and 20th electrons go into the 4th shell (limit of GCSE knowledge)

If you know the atomic (proton) number, you know it equals the number of electrons in a neutral atom, you then apply the rules to work out the electron arrangement (configuration).

Examples: diagram, symbol or name of element (Atomic Number = number of electrons in a neutral atom), shorthand electron arrangement

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On Period 1 (2 elements only)

On Period 2 .... .... (3 of the 8 elements)

On Period 3 .... (3 of the 8 elements)

On Period 4 ===> Kr [2.8.18.8],

The Periodic Table and Electronic Structure

Below are the electron arrangements for elements 1 to 20 set out in Periodic Table format

(Hydrogen and The Transition metals etc. have been omitted). When you move down to the next period you start to fill in the next shell according to the maximum electrons in a shell rule (see

previous section)

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o The first element in a period has one outer electron (eg sodium Na 2.8.1), and the last

element has a full outer shell (eg argon Ar 2.8.8) o Apart from hydrogen (H, 1) and helium (He, 2) the last electron number is the group

number. o and the number of shells used is equal to the Period Number.

Atomic structure diagrams

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Structure, Bonding and properties

1. Why do atoms bond together?

Some atoms are very reluctant to combine with other atoms and exist in the air around us as single atoms. These are the Noble Gases and have very stable electron arrangements eg 2, 2.8 and

2.8.8 because their outer shells are full. The first three are shown in the diagrams below and explains why Noble

Gases are so reluctant to form compounds with other elements.

(atomic number) and electron arrangement.

All other atoms therefore, bond together to become electronically more stable, that is to become like Noble Gases in electron arrangement. Bonding produces new substances and usually involves only the 'outer shell' or 'valency' electrons

The phrase CHEMICAL BOND refers to the strong electrical force of attraction between the atoms or ions in the structure. The combining power of an atom is sometimes referred to as its valency and its value is linked to the number

of outer electrons of the original uncombined atom

COVALENT BONDING - sharing electrons to form molecules with covalent bonds, the bond is usually formed between two non-metallic elements in a molecule.

IONIC BONDING - By one atom transferring electrons to another atom.

COORDINATE COVALENT BONDING A coordinate covalent bond is special because it involves a shared pair of electrons that came from a single atom.

METALLIC BONDING The crystal lattice of metals consists of ions NOT atoms surrounded by a 'sea of electrons' forming giant lattice. These free or 'delocalised' electrons are the 'electronic glue' holding the particles

together. There is a strong electrical force of attraction between these mobile electrons (-) and the 'immobile' positive metal ions (+) and this is the metallic bond.

An ion is an atom or group of atoms carrying an overall positive or negative charge

o eg Na+, Cl-, [Cu(H2O)]2+, SO4

2- etc. If a particle, as in a neutral atom, has equal numbers of protons (+) and electrons (-) the particle charge is zero

ie no overall electric charge. The proton number in an atom does not change BUT the number of associated electrons can! If negative electrons are removed the excess charge from the protons produces an overall positive ion. If negative electrons are gained, there is an excess of negative charge, so a negative ion is formed. The charge on the ion is numerically related to the number of electrons transferred. For any atom or group of atoms, for every electron gained you get a one unit increase in negative charge, for

every electron lost you get a one unit increase in the positive

The atom losing electrons forms a positive ion (cation) and is usually a metal.

The atom gaining electrons forms a negative ion (anion) and is usually a non-metallic element.

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NOBLE GASES are very reluctant to share, gain or lose electrons to form a chemical bond. For most other elements the types of bonding and the resulting properties of the elements or compounds are described in detail below. In all the

electronic diagrams ONLY the outer electrons are shown.

2. Covalent Bonding Covalent bonds are formed by atoms sharing electrons to form molecules. This type of bond usually formed between two non-metallic elements. The molecules might be that of an element ie one type of atom only OR from

different elements chemically combined to form a compound.

The covalent bonding is caused by the mutual electrical attraction between the two positive nuclei of the two atoms of the

bond, and the electrons between them.

One single covalent bond is a sharing of 1 pair of electrons, two pairs of shared electrons between the same two atoms gives a double bond and it is possible for two atoms to share 3 pairs of electrons and give a triple bond.

The bonding in Small Covalent Molecules

The simplest molecules are formed from two atoms and examples of their formation are shown below. The electrons are shown as dots and crosses to indicate which atom the electrons come from, though all electrons are the same. The diagrams may only show the outer electron arrangements for atoms that use two or more electron shells. Examples of

simple covalent molecules are

Example 1: 2 hydrogen atoms (1) form the molecule of the element hydrogen H2

and combine to form where both atoms have a pseudo helium structure of 2 outer

electrons around each atom's nucleus. Any covalent bond is formed from the mutual attraction of two positive

nuclei and negative electrons between them.

Example 2: 2 chlorine atoms (2.8.7) form the molecule of the element chlorine Cl2

and combine to form where both atoms have a pseudo argon structure of 8

outer electrons around each atom. All the other halogens would be similar eg F2, Br2 and I2 etc. Valency of halogens is 1 here.

Example 3: 1 atom of hydrogen (1) combines with 1 atom of chlorine (2.8.7) to form the molecule of the compound hydrogen chloride HCl

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and combine to form where hydrogen is electronically like helium and chlorine like argon. All the other hydrogen halides will be similar eg HF, HBr and HI etc.

Example 4: 2 atoms of hydrogen (1) combine with 1 atom of oxygen (2.6) to form the molecule of the compound we call water H2O

and and combine to form so that the hydrogen atoms are electronically like

helium and the oxygen atom becomes like neon. The molecule can be shown as with two hydrogen - oxygen single covalent bonds. Hydrogen sulphide will be similar, since sulphur (2.8.6) is in the same Group 6 as oxygen. Valency

of oxygen and sulphur is 2 here.

Example 5: 3 atoms of hydrogen (1) combine with 1 atom of nitrogen (2.5) to form the molecule of the compound we call ammonia NH3

three of and one combine to form so that the hydrogen atoms are electronically

like helium and the nitrogen atom becomes like neon. The molecule can be shown as with three nitrogen -

hydrogen single covalent bonds (

Example 6: 4 atoms of hydrogen (1) combine with 1 atom of carbon (2.4) to form the molecule of the compound we call methane CH4

four of and one of combine to form so that the hydrogen atoms are electronically

like helium and the nitrogen atom becomes like neon. The molecule can be shown as with four carbon -

hydrogen single covalent bonds. SiH4 will be similar because silicon (2.8.4) is in the same group as carbon.

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All the bonds in the above examples are single covalent bonds. Below are three examples 7-9, where there is a double bond in the molecule, in order that the atoms have stable Noble Gas outer electron arrangements around each

atom. Carbon has a valency of 4.

Example 7: Two atoms of oxygen (2.6) combine to form the molecules of the element oxygen O2.

The molecule has one O=O double covalent bond . Oxygen valency 2.

Example 8: One atom of carbon (2.4) combines with two atoms of oxygen (2.6) to

form carbon dioxide CO2. The molecule can be shown as with two carbon = oxygen double covalent bonds Valencies of C and O are 4 and 2 respectively.

Example 9: Two atoms of carbon (2.4) combine with four atoms of hydrogen (1) to

form ethene C2H4. The molecule can be shown as with one carbon = carbon double bond and

four carbon - hydrogen single covalent bonds The valency of carbon is still 4.

Examples 10-13: The scribbles below illustrate some more complex examples. Ex. 10 nitrogen; Ex. 11 ethane; Ex. 12 chloromethane and Ex. 13 methanol. The valencies or combining power in theses examples are N 3, H 1, C

4, Cl 1 and O 2.

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MULTIPLE BONDS

Why multiple bonds form: Nonmetal atoms bond to reach a stable or low energy condition. This happens when a main-group atom shares enough electrons to achieve a rare gas valence shell. Sometimes the number of electrons needed cannot be provided by sharing electrons simply in single bonds(pairs).

Nitrogen molecules are an example of this situation. Nitrogen atoms have five valence electrons. We

know that molecules of N2 exist based on a great deal of measurements. If two nitrogen atoms

simply formed a single bond the dot structure would look like the illustraton below. Each nitrogen atom would have only six electrons not an octet. The single bond doesn't provide the two nitrogen atoms with an octet. If the octet rule is to be followed, the nitrogen atoms must share more than two electrons.

The trial and error method is used to decide how many shared electrons are needed to create a structure where the octet rule is met. Since one bond didn't work the next thing to try is a double bond where the atoms share four electrons. Unfortunately the double bond structure only provides 7 valence shell electrons not eight.

Because the single and double bonds didn't do the trick the next thing to try is a triple bond where the nitrogen

atoms share six electrons. The count for both atoms when a triple bond is used in the structure shows that the

octet rule is met.

Multiple bonds are very common in carbon compounds. The carbon atom can form four bonds. These

four bonds can be all single,CH3CH3, two single and one double, CH2CH2, two double, one single and

one triple, CHCH.

Ethane has a single bond carbon-

carbon bond

Ethene has a double carbon-carbon

bond

Acetylene has a triple carbon-carbon

bond

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The carbon-carbon single bond has a

bond length of 154 picometers.

The carbon-carbon double bond has a

bond length of 133 picometers.

The carbon-carbon triple bond has a

bond length of 120 picometers.

The carbon-carbon bond lengths decrease as the number of shared electrons and bonds increase. This is reasonable

because there are more electrons attracted to each nucleus. The repulsions between positive nuclei decrease as more

electrons are shared. The relative lengths for the three types of bonds are summarized here. REMEMBER these

compare bonds between pairs of atoms like ;

C:::C, C::C, C:C or C:::N, C::N, C:N or N:::N, N::N N:N or N::O , N:O

The Properties of simple covalent molecular substances -

small molecules!

The electrical forces of attraction, that is the chemical bond*, between atoms in any molecule are

strong and most molecules do not change chemically on moderate heating.(* sometimes referred to as the

intramolecular bond) However, the electrical forces** between molecules are weak and easily weakened further on

heating. These weak attractions are known as **intermolecular forces and consequently the bulk material is not

usually very strong. Consequently small covalent molecules tend to be volatile liquids, easily vapourised, or low melting

point solids. On heating the inter-molecular forces are easily overcome with the increased kinetic energy gain of the particles

and so have low melting and boiling points. They are also poor conductors of electricity because there are no free electrons or ions in any state to

carry electric charge. Most small molecules will dissolve in a solvent to form a solution.

Large Covalent Molecules and their Properties

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(macromolecules - giant covalent networks and polymers)

It is possible for many atoms to link up to form a giant covalent

structure or lattice. The atoms are usually non-metals. This produces a very strong 3-dimensional covalent bond network or

lattice. This gives them significantly different properties from the small

simple covalent molecules mentioned above. This is illustrated by carbon in the form of diamond. Carbon can form

four single bonds so each carbon bonds to four others etc. This type of structure is thermally very stable and they have high melting

and boiling points. They are usually poor conductors of electricity because the electrons are

not usually free to move as they can in metallic structures. Also because of the strength of the bonding in all directions in the structure,

they are often very hard, strong and will not dissolve in solvents like

water.

Silicon dioxide (silica, SiO2) has a similar 3D structure and properties,

shown below diamond. The hardness of diamond enables it to be used as the 'leading edge'

on cutting tools. Diamond is an allotrope of carbon. Two other allotropes of diamond are

described below. Allotropes are different forms of the same element in the same physical state (3 solid forms of carbon are described

here).

Carbon also occurs in the form of graphite. The carbon atoms form

joined hexagonal rings forming layers 1 atom thick. There are three strong covalent bonds per carbon, BUT, the fourth

bond carbon can form from its four outer electrons, is shared between the

three bonds shown (this requires advanced level concepts to fully explain, and this bonding situation also occurs in fullerenes described below).

The layers are only held together by weak intermolecular forces

shown by the dotted lines NOT by strong covalent bonds. Like diamond and silica (above) the large molecules of the layer ensure

graphite has typically very high melting point because of the strong 2D bonding network (note: NOT 3D network)..

Graphite will not dissolve in solvents because of the strong bonding BUT there are two crucial differences compared to diamond ...

o Electrons, from the 'shared bond', can move freely through each layer, so graphite is a conductor like a metal (diamond is

an electrical insulator and a poor heat conductor). Graphite is used in electrical contacts eg electrodes in electrolysis.

o The weak forces enable the layers to slip over each other so

where as diamond is hard material graphite is a 'soft' crystal, it feels slippery. Graphite is used as a lubricant.

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Bonding in polymers and 1-3 'dimension' concepts in macromolecules

The bonding in polymers or plastics is no different in principle to the examples described above, but there is quite

a range of properties and the difference between simple covalent and giant covalent molecules can get a bit

'blurred'. o Bonds between atoms in molecules, eg C-C, are called intra-molecular bonds. o The much weaker electrical attractions between individual molecules are called inter-molecular forces.

In thermosoftening plastics like poly(ethene) the bonding is like ethane except there are lots of carbon

atoms linked together to form long chains. They are moderately strong materials but tend to soften on heating and are not usually very soluble in solvents. The structure is basically a linear 1 dimensional strong

bonding networks. Graphite structure is a layered 2 dimensional strong bond network made of layers of joined

hexagonal rings of carbon atoms with weak inter-molecular forces between the layers. Thermosetting plastic structures like melamine have a 3 dimensional cross-linked giant covalent

structure network similar to diamond or silica in principle, but rather more complex and chaotic! Because of the

strong 3D covalent bond network they do not dissolve in any solvents and do not soften on heating and are much stronger than thermoplastics.

More on polymers in Oil Notes and Extra Organic Chemistry Notes.

3. Ionic Bonding Ionic bonds are formed by one atom transferring electrons to another atom to form ions. Ions are atoms, or groups of atoms, which have lost or gained electrons.

The atom losing electrons forms a positive ion (a cation) and is usually a metal. The overall charge on the ion is positive due to excess positive nuclear charge (protons do NOT change in chemical reactions).

The atom gaining electrons forms a negative ion (an anion) and is usually a non-metallic element. The overall charge on the ion is negative because of the gain, and therefore excess, of negative electrons.

The examples below combining a metal from Groups 1 (Alkali Metals), 2 or 3, with a non-metal from Group 6 or Group 7

(The Halogens)

Example 1: A Group 1 metal + a Group 7 non-metal eg sodium + chlorine sodium chloride NaCl or ionic formula Na+Cl- In terms of electron arrangement, the sodium donates its outer electron to a chlorine atom forming a single positive sodium ion and a single negative chloride ion. The atoms have become stable ions, because electronically, sodium becomes like neon and chlorine like argon.

Na (2.8.1) + Cl (2.8.7) Na+ (2.8) Cl- (2.8.8)

can be summarised electronically as [2,8,1] + [2,8,7] [2,8]+ [2,8,8]-

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ONE combines with ONE to form

The valencies of Na and Cl are both 1, that

is, the numerical charge on the ions. NaF, KBr, LiI etc. will all be electronically similar.

Example 2: A Group 2 metal + a Group 7 non-metal eg magnesium + chlorine magnesium chloride MgCl2 or ionic

formula Mg2+(Cl-)2 In terms of electron arrangement, the magnesium donates its two outer electrons to two chlorine atoms forming a double positive magnesium

ion and two single negative chloride ions. The atoms have become stable ions, because electronically, magnesium

becomes like neon and chlorine like argon.

Mg (2.8.2) + 2Cl (2.8.7) Mg2+ (2.8) 2Cl- (2.8.8)

can be summarised electronically as [2,8,2] + 2[2,8,7] [2,8]2+ [2,8,8]-2

ONE combines with TWO to form see *

* NOTE you can draw two separate chloride ions, but in these examples a number subscript has been used, as in

ordinary chemical formula.

Example 3: A Group 3 metal + a Group 7 non-metal eg aluminium + fluorine aluminium fluoride AlF3 or ionic formula Al3+(F-)3 In terms of electron arrangement, the aluminium donates its three outer electrons to three fluorine atoms forming a triple positive aluminium ion and three single negative fluoride ions. The atoms have become stable ions, because electronically, aluminium and fluorine becomes electronically like neon. Valency of Al is, F is 1.

Al (2.8.3) + 3F (2.7) Al3+ (2.8) 3F- (2.8)

can be summarised electronically as [2,8,3] + 3[2,7] [2,8]3+ [2,8]-3

ONE combines with THREE to form

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Example 4: A Group 1 metal + a Group 6 non-metal eg potassium + oxygen potassium oxide K2O or ionic formula (K+)2O

2- In terms of electron arrangement, the two potassium atoms donates their outer electrons to one oxygen atom. This results in two single positive potassium ions to one double negative oxide ion. All the ions have the stable electronic structures 2.8.8 (argon like) or 2.8 (neon like). Valencies, K 1, oxygen 2. Na2O, Na2S, K2S etc. will be

similar.

2K (2.8.8.1) + O (2.6) 2K+ (2.8.8) O2- (2.8)

can be summarised electronically as 2[2,8,8,1] + [2,6] [2,8,8]+2 [2,8]2-

TWO combine with ONE to form

Example 5: A Group 2 metal + a Group 6 non-metal eg calcium + oxygen calcium oxide CaO or ionic formula Ca2+O2- In terms of electron arrangement, one calcium atom donates its two outer electrons to one oxygen atom. This results in a double positive calcium ion to one double negative oxide ion. All the ions have the stable electronic structures

2.8.8 (argon like) or 2.8 (neon like). the valency of both calcium and oxygen is 2. MgO, MgS, or CaS will be similar electronically (S and O both in Group 6)

Ca (2.8.8.2) + O (2.6) Ca2+ (2.8.8) O2- (2.8)

can be summarised electronically as [2,8,8,2] + [2,6] [2,8,8]2+ [2,8]2-

ONE combines with ONE to form

Example 6: A Group 3 metal + a Group 6 non-metal eg aluminium + oxygen aluminium oxide Al2O3 or ionic formula (Al3+)2(O

2-)3 In terms of electron arrangement, two aluminium atoms donate their three outer electrons to three oxygen atoms. This results in two triple positive aluminium ions to three double negative oxide ions. All the ions

have the stable electronic structure of neon 2.8. Valencies, Al 3 and O 2.

2Al (2.8.3) + 3O (2.6) 2Al3+ (2.8) 3O2- (2.8)

can be summarised electronically as 2[2,8,3] + 3[2,6] [2,8]3+2 [2,8]2-

3

TWO combines with THREE to form

The Properties of Ionic Compounds

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The alternate positive and negative ions in an ionic solid are arranged in an orderly way in a giant ionic lattice

structure shown on the left. The ionic bond is the strong electrical attraction

between the positive and negative ions next to each other in the lattice.

The bonding extends throughout the crystal in all

directions. Salts and metal oxides are typical ionic compounds. This strong bonding force makes the structure hard (if brittle)

and have high melting and boiling points, so they are not

very volatile! The bigger the charges on the ions the stronger the bonding

attraction eg magnesium oxide Mg2+O2- has a higher melting point than sodium chloride Na+Cl-.

Unlike covalent molecules, ALL ionic compounds are

crystalline solids at room temperature. They are hard but brittle, when stressed the bonds are

broken along planes of ions which shear away. They are NOT

malleable like metals. Many ionic compounds are soluble in water, but not all,

so don't make assumptions. The solid crystals DO NOT conduct electricity because

the ions are not free to move to carry an electric current.

However, if the ionic compound is melted or dissolved in water, the liquid will now conduct electricity, as the ion

particles are now free.

4. BONDING IN METALS

Electron-Sea Model of Metals

In the electron-sea model, a metal crystal is considered to be a

three-dimensional array of metal cations immersed in a sea of

valence electrons. The delocalized valence electrons are free to

move throughout the crystal and are not associated with any

one particular metal cation. The mobility of the electrons

accounts for the high electrical conductivity of metals.

Thermal conductivity can also be ascribed to the mobile

electrons that conduct heat by carrying kinetic energy from

one part of the crystal to another. Three-dimensional

delocalized bonding allows the metal to be both malleable and

ductile.

The crystal lattice of metals consists of ions NOT atoms surrounded by a 'sea of electrons' forming another

type of giant lattice.

The outer electrons (-) from the original metal atoms

are free to move around between the positive metal ions formed (+).

These free or 'delocalised' electrons are the 'electronic

glue' holding the particles together. There is a strong electrical force of attraction

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between these mobile electrons (-) and the 'immobile' positive metal ions (+) and this is the metallic bond.

This strong bonding generally results in dense, strong materials

with high melting and boiling points. Metals are good conductors of electricity because these 'free'

electrons carry the charge of an electric current when a potential difference (voltage!) is applied across a piece of metal.

Metals are also good conductors of heat. This is also due to the free

moving electrons. Non-metallic solids conduct heat energy by hotter more strongly vibrating atoms, knocking against cooler less strongly vibrating

atoms to pass the particle kinetic energy on. In metals, as well as this effect, the 'hot' high kinetic energy electrons move around freely to

transfer the particle kinetic energy more efficiently to 'cooler' atoms. Typical metals also have a silvery surface but remember this may be

easily tarnished by corrosive oxidation in air and water. Unlike ionic solids, metals are very malleable, they can be readily bent, pressed or hammered into shape. The

layers of atoms can slide over each other without fracturing the structure. The reason for this is the mobility of

the electrons. When planes of metal atoms are 'bent' or slide the electrons can run in between the atoms and maintain a strong bonding situation. This can't happen in ionic solids.

Note on Alloy Structure

1. Shows the regular arrangement of the atoms in a metal crystal and the white spaces show where the free electrons are (yellow circles actually positive metal ions).

2. Shows what happens when the metal is stressed by a strong force. The layers of atoms can slide over each other and the bonding is maintained as the mobile electrons keep in contact with atoms, so the metal remains intact BUT a different shape.

3. Shows an alloy mixture. It is NOT a compound but a physical mixing of a metal plus at least one other material (shown by red circle, it can be another metal eg Ni, a non-metal eg C or a compound of carbon or manganese,

and it can be bigger or smaller than iron atoms). Many alloys are produced to give a stronger metal. The

presence of the other atoms (smaller or bigger) disrupts the symmetry of the layers and reduces the 'slip ability' of one layer next to another. The result is a stronger harder less malleable metal.

5.Coordinate Covalent Bonds(Dative Bonding)

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A coordinate covalent bond is special because it involves a shared pair of electrons that came from a single atom. Ammonia had a nitrogen atom with an unshared pair of electrons. These can be shared with an electron defficient atom like H+.

ammonia ammonia proton ammonium ion

Water molecules have two unshared pairs of electrons. These form coordinate covalent bonds with cations that are dissolved in water. This is one reason why water dissolves many ionic solids. The energy released when the water molecules bond to the

cations is often enough to break up the ionic solid.

water cation water and cation

A dissolved cation will form as many as six coordinate covalent bonds.

6. INTERMOLECULAR FORCES Intermolecular forces are those forces that occur between particles (molecules, atoms, or ions).

The strength of these forces at a given temperature dictates whether a substance will have the

properties of a solid, a liquid, or a gas. The term van der Waals forces encompasses all types of

intermolecular forces. All intermolecular forces arise from electrostatic interactions governed by

the basic rule that like charges repel and unlike charges attract.

Hydrogen Bonds

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A hydrogen bond is an attractive interaction between a hydrogen that is bonded to a very electronegative atom (O, N, or F) and an unshared electron pair on another electronegative atom.

Hydrogen bonds can be quite strong. Substances that form hydrogen bonds have unusually high boiling

points due to the extra energy that must be used to separate the molecules.

DONE

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CHEMICAL TESTS

INORGANIC TESTS TEST FOR TEST

METHOD OBSERVATIONS TEST CHEMISTRY

hydrogen gas H2

lit splint or spill squeaky pop! (might see condensation on

test tube)

2H2(g) + O2(g) ==> 2H2O(l) + energy!

carbon dioxide gas CO2

bubble into

limewater (aqueous calcium

hydroxide solution)

turns cloudy fine milky white precipitate of calcium carbonate

Ca(OH)2(aq) + CO2(g) ==> CaCO3(s) + H2O(l)

oxygen gas O2 glowing splint or spill

re-ignites it - flame

C(in wood) + O2(g) ==> CO2(g)

Hydrogen chloride gas HCl, in water

hydrochloric acid

(i) blue litmus and (ii) drop of silver

nitrate on the end

of a glass rod

(i) litmus turns red, (ii) white precipitate

with silver nitrate

(i) Strongly acid gas, (ii) in water forms chloride ions - hence precipitate with silver nitrate.

Hydrogen

bromide HBr and

Hydrogen iodide HI

As above. In water

they are

hydrobromic acid and hydriodic acid.

as above but cream

HBr or yellow HI

precipitate

as above - combination of acid and halide ion

tests

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Sulphur dioxide gas SO2

freshly made potassium

dichromate(VI)

paper

paper changes from orange to green

the dichromate(VI) ion, Cr2O72-

(aq) is reduced to the green Cr3+(aq) ion

Ammonia gas NH3

strong pungent

odour*, (i) red litmus, (ii) fumes

conc. hydrochloric acid

(i) litmus turns blue,

(ii) white clouds with HCl fumes.

(i) only common alkaline gas and (ii) forms fine

ammonium chloride crystals with HCl (*volatile organic aliphatic amines give the same result,

and smell more fishy)

Chlorine gas Cl2 [test (ii) on

its own is no good, could be

HCl]

(i) blue litmus, (ii) drop silver nitrate

on the end of a glass rod

pungent green gas, (i) litmus turns red and

then is bleached white, (ii) white precipitate

(i) non-metal, is acid in aqueous solution and a

powerful oxidising agent, (ii) forms chloride ion in water

Iodine solid (i) heating, (ii) test aqueous solution or solid with starch

solution

(i) purple vapour, (ii)

blue black colour with starch solution

Nitrogen(IV) oxide

(or nitrogen dioxide)

NO2

no simple relatively unambiguous test

nasty brown gas strong oxidising agent

Water liquid H2O (i) white anhydrous

copper(II)

sulphate, (ii) dry blue cobalt

chloride paper

(i) turns from white to

blue, (ii) turns from blue to pink

(i) blue hydrated copper(II) crystals or solution

formed, (ii) hydrated cobalt ion formed [Co(H2O)6]

2+

Carbonate ion CO32-

(or hydrogencarbonate

HCO3-)

add any dilute

strong acid to the

suspected carbonate - if

colourless gas given off, test with

limewater

fizzing - colourless

gas - turns

limewater milky cloudy (see above

CO2)

carbonate/hydrogencarbonate + acid ==> salt +

water + carbon dioxide, then white precipitate

with limewater.

Sulphate ion [sulphate(VI)] SO4

2-

to a solution of the

suspected sulphate

add dilute hydrochloric acid

and a few drops of barium chloride

or nitrate solution

white precipitate of

barium sulphate Ba2+(aq) + SO4

2-(aq) ==> BaSO4(s)

any soluble barium salt + any soluble sulphate ==> barium sulphate

Sulphite ion [sulphate(IV)] SO3

2-

(i) add dilute hydrochloric acid

to the suspected sulphite, (ii) test

any gas evolved

with fresh potassium

(i) acrid choking

sulphur dioxide gas formed,

(ii) the dichromate

paper turns from orange to green

(i) sulphite salt + hydrochloric acid ==> chloride salt + sulphur dioxide, (ii) the sulphur dioxide

reduces the dichromate(VI) to chromium(III). Note: sulphites do not give ppt. with acidified

barium chloride/nitrate because sulphites dissolve

in acids.

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dichromate(VI) paper

Sulphide ion

S2- for (ii) dangerous hydrogen sulphide formed

(i) If soluble, add a

few drops lead(II) ethanoate solution.

(ii) If solid, add dil. HCl(aq) acid, test

gas with lead(II) ethanoate paper.

(i) Black ppt. of lead

sulphide.

(ii) Rotten egg smell of hydrogen sulphide

and the H2S gas turns

lead(II) ethanoate paper black.

(i) Pb2+(aq) + S2-

(aq) => PbS(s)

(ii) MS(s) + 2H+

(aq) => M2+

(aq) + H2S(g) (e.g. M = Pb, Fe, Cu, Ni etc.) Then reaction (i) above occurs.

Chloride ion

Cl-

(i) if soluble, add dilute nitric acid

and silver nitrate

solution, (ii) if insoluble salt, add

conc. sulphuric

acid, warm if necessary then

test gas as for HCl above.

(i) white precipitate

of silver chloride soluble in

dilute ammonia, (ii) get fumes of

hydrogen chloride which turn blue litmus

red and give a white precipitate with silver

nitrate solution

(i) Ag+(aq) + Cl-(aq) ==> AgCl(s) , any soluble

silver salt + any soluble chloride ==> silver

chloride precipitate, (ii) Cl-(s) + H2SO4(l) ==> HSO4

-(s) + HCl(g) , then Ag

+(aq) + Cl

-(aq) ==>

AgCl(s)

Bromide ion

Br-

(i) if soluble, add dilute nitric acid

and silver nitrate

solution, (ii) if insoluble salt, add

conc. sulphuric acid, warm if

necessary

(i) cream precipitate

of silver bromide, only soluble

in concentrated

ammonia, (ii) orange vapour, test for

sulphur dioxide.

(i) Ag+(aq) + Br-(aq) ==> AgBr(s) any soluble

silver salt + any soluble bromide ==> silver

bromide precipitate, (ii) bromide ion is oxidised to bromine and the sulphuric acid is reduced to

sulphur dioxide

Iodide ion I- (i) if soluble, add dilute nitric acid

and silver nitrate solution, (ii) if

insoluble salt can

heat with conc.

sulphuric acid, (ii) get purple fumes

of iodine and very smelly hydrogen

sulphide, (iii) if soluble, add

lead(II) nitrate

solution

(i) yellow precipitate

of silver iodide insoluble in

concentrated

ammonia, (ii) purple vapour and rotten egg smell!, (iii) a yellow

precipitate forms

(i) Ag+(aq) + I-(aq) ==> AgI(s) , any soluble silver

salt + any soluble iodide ==> silver iodide

precipitate, (ii) iodide ion is oxidised to iodine and the sulphuric acid is reduced to hydrogen

sulphide, (iii) insoluble lead(II) iodide formed,

Pb2+(aq) + 2I-(aq) ==> PbI2(s)

Nitrate ion [or nitrate(V)] NO3

-

(i) boil the

suspected nitrate with sodium

hydroxide solution

and fine aluminium powder (Devarda's

Alloy)

(i) the fumes

contain ammonia,

which turns red litmus

blue, see ammonia test details

(ii) Where the liquids

(i) the aluminium powder is a powerful reducing

agent and converts the nitrate ion, NO3-, into

ammonia gas, NH3

(ii) NO complex of iron(II) formed

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(ii) Add iron(ii) sulphate solution

and then conc. sulphuric acid (the

'brown ring' test)

meet a brown ring forms

Nitrite ion [or nitrate(III)] NO2

-

No simple test, (i) in acid solution it decomposes to give nasty brown fumes of NO2, (ii) it

decolourises (purple ==> colourless) acidified potassium manganate(VII), (iii) it liberates

iodine from acidified potassium iodide solution, (iv) forms ammonia with hot Al powder/NaOH(aq) and gives 'brown ring' test - see nitrate tests above.

Ammonium ion NH4+

no smell at first, add COLD sodium

hydroxide solution to the suspected

ammonium salt

and test any gas with red litmus

smelly ammonia

evolved! and red

litmus

turns blue

ammonia gas is evolved: NH4+

(aq) + OH-(aq) ==>

NH3(g) + H2O(l)

Hydrogen ion ie acids! H+ or H3O

+ (note: to completely identify acids you need to test for the

anion eg chloride for HCl

etc.)

(i) litmus or universal indicator

or pH meter, (ii) add a little sodium

hydrogencarbonate powder

(i) litmus turns red, variety of colours with

univ. ind. strong - red,

weak - yellow/orange, (ii) fizzing with any carbonate - test for

CO2 as above

(i) pH meter gives a value of less than 7, the lower the pH number the stronger the acid, the

higher the H+ concentration, (ii) HCO3-(aq) + H

+(aq)

==> H2O(l) + CO2(g)

Hydroxide ion ie an alkali OH- (note: to

completely identify

alkalis you need to test

for the cation eg sodium for NaOH etc.)

(i) litmus or

universal indicator or pH meter, (ii)

add ammonium

salt

(i) turns litmus blue,

variety of colours univ. ind. dark green -

violet for weak -

strong, (ii) if strongly alkaline ammonia

should be released, see ammonia test for rest of details

(i) pH meter gives a value of more than 7, the

higher the pH number the stronger the alkali, the higher the OH- concentration, (ii) ammonia gas is

evolved: NH4+

(aq) + OH-(aq) ==> NH3(g) + H2O(l)

Positive metal cations via flame tests (see below for

NaOH and NH3 for metal ion

tests too)

The metal salt or

other compound is mixed with

concentrated hydrochloric acid

and a sample of the mixture is

heated strongly in

a bunsen flame on the end of a

cleaned nichrome wire (platinum if

you can afford it!)

lithium Li+ crimson All colours are due to electronic excitation to a higher level. You see the light emitted as the electron returns to its lower more stable level.

This is the basis of atomic emission and absorption spectroscopy. Aluminium, magnesium,

iron and zinc do not produce a useful identifying flame colour.

sodium Na+ yellow

potassium K+ lilac

calcium Ca2+ brick red

barium Ba2+ apple green

copper(II) Cu2+

blue/green

Positive metal cations via sodium hydroxide (NaOH)

Dilute sodium hydroxide

solution is added to a solution

aluminium ion: Al3+(aq) + 3OH-(aq) ==> Al(OH)3(s) white precipitate

* The ppt. is not soluble in excess of the weak alkali ammonia, but

dissolves in the strong alkali sodium hydroxide: Al(OH)3(s) + 3OH-(aq)

==> [Al(OH)6]3-

(aq) (amphoteric behaviour)

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or ammonia (NH3) solutions (both alkalis, giving hydroxide ions,

OH-, in their solutions)

containing the suspected ion.

Both the

precipitate formed and the

effect of excess alkali are

important observations.

All precipitates white, unless

otherwise stated and all tend to be

gelatinous in

nature.

The test can be repeated with

aqueous

ammonia solution

(sometimes wrongly called

'ammonium

hydroxide'). The observations are

usually, but not always, similar.

ppt. = precipitate.

calcium ion: Ca2+(aq) + 2OH-(aq) ==> Ca(OH)2(s) white ppt. * The

ppt. is not soluble in excess of NH3 or NaOH.

magnesium ion: Mg2+(aq) + 2OH-(aq) ==> Mg(OH)2(s) white ppt. *

The ppt. is not soluble in excess of NH3 or NaOH. You could distinguish Mg from Ca with a flame test.

copper(II) ion: Cu2+(aq) + 2OH-(aq) ==> Cu(OH)2(s)

***blue/turquoise ppt. - this does dissolve in excess ammonia to

give a deep blue solution.

iron(II) ion: Fe2+(aq) + 2OH-(aq) ==> Fe(OH)2(s) dark green ppt.*

The ppt. is not soluble in excess of NH3 or NaOH.

iron(III) ion: Fe3+(aq) + 3OH-(aq) ==> Fe(OH)3(s) brown ppt.* The

ppt. is not soluble in excess of NH3 or NaOH.

zinc ion: Zn2+(aq) + 2OH-(aq) ==> Zn(OH)2(s) white ppt. The ppt.

dissolves in both excess sodium hydroxide or ammonia to give a clear

colourless solution.

MISCELLANEOUS CATION TESTS:

(i) Lead(II) ion

(i) add potassium

iodide solution => yellow precipitate

(i) Pb2+(aq) +2I-(aq) ==>PbI2(s) lead(II) iodide ppt.

Metal Carbonates

Sometimes heating a metal carbonate

strongly to

decompose it provides some

clues to its identity. Adding

acid => CO2 and

the colour of the resulting solution

(eg blue Cu2+(aq),

may also provide

clues. The metal ion solution might

also give a flame

colour or a

copper(II) carbonate==> copper(II) oxide + carbon dioxide: CuCO3(s) ==> CuO(s) + CO2(g)

[green] ==> [black] + [colourless gas, test with limewater, white

precipitate]

zinc carbonate==> zinc oxide + carbon dioxide

ZnCO3(s) ==> ZnO(s) + CO2(g)

[white] ==> [yellow hot, white cold] +[colourless gas, test with

limewater, white precipitate]

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hydroxide precipitate with

sodium hydroxide

eg copper.

ORGANIC TESTS TEST FOR TEST

METHOD OBSERVATIONS TEST CHEMISTRY

ALKENE or alkyne any other non-aromatic

unsaturated hydrocarbons

bubble gas through, or add

liquid to, a solution

of bromine in hexane or water

the orange/brown bromine, decolourises,

as a saturated

colourless organic bromo-compound is

formed (saturated alkanes give no fast

reaction with bromine)

R2C=CR2 + Br2 ==> BrR2C-CR2Br

RC CR + 2Br2 ==> Br2RC-CRBr2

R = H, alkyl or aryl

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The Three States of Matter

KEYWORDS: Boiling Boiling point Brownian motion Condensing Cooling curve Diffusion Evaporation Freezing Freezing point Gas particle picture Heating curve

Liquid particle picture Melting Melting point Properties of gases Properties of liquids Properties of solids sublimation Solid particle

picture

The particle model of a Gas

Almost no forces of attraction between the particles so they are completely free of each other. Particles widely spaced and scattered at random throughout the container so there is no order in the

system. Particles move rapidly in all directions, frequently colliding with each other and the side of the

container. With increase in temperature, the particles move faster as they gain kinetic energy.

Using the particle model to explain the properties of a Gas

Gases have a very low density (light) because the particles are so spaced out in the container. o Density order: solid > liquid >>> gases

Gases flow freely because there are no effective forces of attraction between the particles. o Ease of flow order: gases > liquids >>> solids (no real flow in solid unless you powder it!)

Gases have no surface, and no fixed shape or volume, and because of lack of particle attraction,

they always spread out and fill any container (so gas volume = container volume). Gases are readily compressed because of the empty space between the particles.

o Ease of compression order: gases >>> liquids > solids (almost impossible to compress a solid)

If the container volume can change, gases readily expand* on heating because of the lack of particle attraction, and readily contract on cooling.

o On heating, gas particles gain kinetic energy, move faster and hit the sides of the

container more frequently, and significantly, they hit with a greater force. o Depending on the container situation, either or both of the pressure or volume will increase

(reverse on cooling). o Note: * It is the gas volume that expands NOT the molecules, they stay the same size!

The natural rapid and random movement of the particles means that gases readily spread or diffuse. Diffusion is fastest in gases where there is more space for them to move.

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o The rate of diffusion increases with increase in temperature as the particles gain kinetic energy and move faster.

o Other evidence for random particle movement: When smoke particles are viewed under a microscope they appear to 'dance around'

when illuminated with a light beam at 90o to the viewing direction. This is because the

smoke particles show up by reflected light and 'dance' due to the millions of random hits from the fast moving air molecules. This is called 'Brownian motion'. At any given

instant of time, the hits will not be even, so the smoke particle get a greater bashing in a random direction.

If a long glass tube is filled at one with a plug of cotton wool soaked in conc. hydrochloric acid, and a similar plug of conc. ammonia solution at the other end. If left undisturbed and horizontal, despite the lack of tube movement (eg shake to mix), a

white cloud forms about 1/3rd along from the conc. acid tube end. What happens is the colourless gases ammonia and hydrogen chloride diffuse

down the tube and react to form fine white crystals of the salt ammonium

chloride. NH3(g) + HCl(g) ==> NH4Cl(s) Note the rule: The smaller the molecular mass, the faster the molecules

move. Therefore the smaller the molecular mass, the faster the gas diffuses. eg Mr(NH3) = 14 + 1x3 = 17, moves faster than Mr(HCl) = 1 + 35.5 =

36.5 AND that's why they meet nearer the HCl end of the tube! So the experiment is not only evidence for molecule movement, its

also evidence that different molecular masses move on at different speeds.

A coloured gas, that is heavier than air, is

put into a gas jar and a second gas jar is placed over it separated with a cover.

If the cover is removed than coloured gas diffuses into the colourless air above. It

can't be due to convection because the more dense gas starts at the bottom!

No 'shaking' or other means of mixing is required. The random movement of both

lots of particles is enough to ensure that both gases are completely mixed

eventually.

This is clear evidence for the process of

diffusion due to particle movement.

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The particle model of a Liquid

Much greater forces of attraction between the particles in a liquid compared to gases, but not

quite as much as in solids. Particles quite close together but still arranged at random throughout the container, there is a little

close range order as y