class notes - educalabserbal.pntic.mec.es/jnavar13/e3fq/classnotes/e3fq-classnotes.pdf · 1....

PHYSICS & CHEMISTRY DEPT.

CLASS NOTES

I E S

MA

RIANO BAQUERO

José A. Navarro Ramón

Contents

1 Measurement and units 51.1 Physical quantity . . . . . . . . . . . . . . . . . . . . . . . . . . . 51.2 Measurement . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 51.3 Unit . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 51.4 Base and derived quantities . . . . . . . . . . . . . . . . . . . . . 61.5 Systems of units. The International System (SI) . . . . . . . . . 61.6 Base and derived quantities of the SI . . . . . . . . . . . . . . . . 61.7 SI Unit prefixes . . . . . . . . . . . . . . . . . . . . . . . . . . . . 61.8 Imperial units . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 71.9 Measuring instruments . . . . . . . . . . . . . . . . . . . . . . . . 71.10 Scientific notation . . . . . . . . . . . . . . . . . . . . . . . . . . 81.11 Significant figures . . . . . . . . . . . . . . . . . . . . . . . . . . . 81.12 Unit conversion . . . . . . . . . . . . . . . . . . . . . . . . . . . . 91.13 Errors . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 111.14 Vocabulary . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 11

2 The states of matter 132.1 Properties of matter. Density . . . . . . . . . . . . . . . . . . . . 13

2.1.1 Volume . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132.1.2 Mass . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132.1.3 Matter . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 132.1.4 General and characteristic properties . . . . . . . . . . . . 132.1.5 Density . . . . . . . . . . . . . . . . . . . . . . . . . . . . 14

2.2 The aggregation states of matter . . . . . . . . . . . . . . . . . . 142.2.1 Solid, liquid and gas . . . . . . . . . . . . . . . . . . . . . 142.2.2 Kinetic theory of matter . . . . . . . . . . . . . . . . . . . 142.2.3 Kinetic theory and aggregation states . . . . . . . . . . . 142.2.4 Kinetic theory and temperature . . . . . . . . . . . . . . . 15

3 Material systems 193.1 Classifying matter . . . . . . . . . . . . . . . . . . . . . . . . . . 19

3.1.1 Pure substances . . . . . . . . . . . . . . . . . . . . . . . 193.1.2 Elements . . . . . . . . . . . . . . . . . . . . . . . . . . . 203.1.3 Compounds . . . . . . . . . . . . . . . . . . . . . . . . . . 203.1.4 Mixtures . . . . . . . . . . . . . . . . . . . . . . . . . . . 203.1.5 Heterogeneous mixtures . . . . . . . . . . . . . . . . . . . 213.1.6 Homogeneous mixtures . . . . . . . . . . . . . . . . . . . 213.1.7 Colloids and suspensions . . . . . . . . . . . . . . . . . . . 22

3

3.2 Solutions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 223.2.1 Solvent and solute (Textbook p.64) . . . . . . . . . . . . . 223.2.2 Classification of solutions attending to physical state . . . 223.2.3 Classification of solutions attending to composition . . . . 22

3.3 Concentration of a solution . . . . . . . . . . . . . . . . . . . . . 233.3.1 Percent by mass . . . . . . . . . . . . . . . . . . . . . . . 233.3.2 Percent by volume . . . . . . . . . . . . . . . . . . . . . . 233.3.3 Grams per litre . . . . . . . . . . . . . . . . . . . . . . . . 24

3.4 Solubility . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 243.4.1 Solubility curve . . . . . . . . . . . . . . . . . . . . . . . . 243.4.2 Solubility of gases in liquids . . . . . . . . . . . . . . . . . 24

3.5 Osmosis . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 243.6 Separation of mixtures . . . . . . . . . . . . . . . . . . . . . . . . 24

3.6.1 Separation of heterogeneous mixtures . . . . . . . . . . . 243.6.2 Separation of homogeneous mixtures . . . . . . . . . . . . 24

4 Atoms 254.1 Dalton’s atomic theory . . . . . . . . . . . . . . . . . . . . . . . . 25

4.1.1 The model . . . . . . . . . . . . . . . . . . . . . . . . . . 254.2 Thomsom’s atomic theory . . . . . . . . . . . . . . . . . . . . . . 25

4.2.1 Cathode rays . . . . . . . . . . . . . . . . . . . . . . . . . 254.2.2 The plum cake model . . . . . . . . . . . . . . . . . . . . 25

4.3 Rutherford’s atomic theory . . . . . . . . . . . . . . . . . . . . . 264.3.1 The gold foil experiment . . . . . . . . . . . . . . . . . . . 264.3.2 Rutherford’s model . . . . . . . . . . . . . . . . . . . . . . 264.3.3 Subatomic particles . . . . . . . . . . . . . . . . . . . . . 274.3.4 Atomic number . . . . . . . . . . . . . . . . . . . . . . . . 274.3.5 Mass number . . . . . . . . . . . . . . . . . . . . . . . . . 274.3.6 Chemical element . . . . . . . . . . . . . . . . . . . . . . . 274.3.7 Isotopes number . . . . . . . . . . . . . . . . . . . . . . . 274.3.8 Ions . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 28

1. Measurement and units

1.1 Physical quantity (Textbook p.23)

A physical quantity is any property that can be measured.

Examples: velocity, length, pressure, volume, density, etc.Non-physical properties: A table, the sky, intelligence, beauty, etc.

1.2 Measurement (Textbook p.23)

Measurement is the process by which we can determine the value of phys-ical quantities.

We generally measure a physical quantity by comparing it with a unit of thesame kind in order to find the number of units that they contain.

1.3 Unit (Textbook p.23)

A physical unit is a value that is adopted as standard for a physical quant-ity.

Unit examples are: metre, litre, Newton, km/h, mph, inch, foot, etc.

A good unit must be:

• Constant: It must not depend on the person that measures. It must notchange with time or with temperature.

• Universal: It must be the same everywhere.

• Reproducible: It can be made whenever needeed.

As a result of every measurement we must write a number followed by the unitthat has been used. Examples: 1 m; 10.4 km/h.

5

6 CHAPTER 1. MEASUREMENT AND UNITS

1.4 Base and derived quantities (Textbook p.23)

Base quantities are the simplest ones and most of the time they aredefined by themselves.

More generally, they are quantities that the scientific community arbitrarilydecides them to be so.Examples: Length, time, etc. We do not provide many examples now becausedifferent system of units may choose different sets of base quantities.

Their units are called base units. Base unit examples may be metres (m),seconds (s), etc.

Derived quantities are defined in terms of the base quantities.

Example: Velocity.It may be defined as length per unit time.So, it is a derived property because it depends on length and time.

As it was the case with base properties, we do not provide here many ex-amples of derived properties.

Derived unit examples may be: Newton (N), kilometres per hour (km/h),square centimetres (cm2), etc.

1.5 Systems of units. The SI (Textbook p.24)

A system of units is a set of arbitrary base and derived properties alongwith their units.

There are a number of different systems of units, as the metric system, CGS,MKS, etc.The most widely used is the International System of Units, abbreviated SI (fromthe French Système International d’Unités.)

In the following table we find the seven base properties and their units:

1.6 Base and derived quantities of the SI

The SI has seven base properties, as you can see in table 1.1.The rest of properties: force, electric charge, pressure, density, volume, ve-

locity, acceleration, energy, etc., are derived.

1.7 SI Unit prefixes (Textbook p.25)

In table 1.2 you will find the most important SI unit prefixes.

1.8. IMPERIAL UNITS 7

# Property Unit Unit’s symbol1 Length metre or meter m2 Mass kilogramme or kilogram kg3 Time second s4 Temperature kelvin K5 Electric current ampere A6 Luminous intensity candela cd7 Susbstance amount mole mol

Table 1.1: Base quantities and base units of the SI

Prefix Symbol Factor

Greater than one

Exa E 1018

Peta P 1015

Tera T 1012

Giga G 109

Mega M 106

Kilo k 103

Hecto h 102

Deca da 101

Equal to one 100 = 1

Less than one

Deci d 10−1

Centi c 10−2

Mili m 10−3

Micro µ 10−6

Nano n 10−9

Pico p 10−12

Femto f 10−15

Atto a 10−18

Table 1.2: Most important SI unit prefixes

1.8 Imperial units

In most English speaking countries people use the imperial units. They are veryimportant in those countries. You may find some of them in table 1.3.

1.9 Measuring instruments

In this section we are going to review some measuring instruments:

• To measure length we use a ruler or a tape measure.

• To measure mass we use a balance or a set of scales (scales in the UK,scale in the USA) or a balance scale or a digital scale.

• To measure time we use a clock or a stopwatch.

• To measure temperature we use a thermometre.


Unit Unit Imperial SIQuantity name symbol equivalence equivalence

Length

inch in – 1 in = 2.54 cmfoot ft 1 ft = 12 in 1 ft = 30.48 cmyard yd 1 yd = 3 ft 1 yd = 0.9144 mmile mi 1 mi = 1 760 yd 1 mi = 1.6093 km

Massounce oz – 1 oz = 28.35 gpound lb 1 lb = 16 oz 1 lb = 0.4536 kgstone stone 1 stone = 15 lb 1 stone = 6.3503 kg

Volumefluid ounce fl oz – 1 oz = 28.414 mLpint pt 1 pt = 20 floz 1 pt = 0.5683 Lgallon gal 1 gal = 8 pt 1 gal = 4.5461 L

Table 1.3: Imperial units

• To measure liquid volume we use a measuring cylinder.

• To measure velocity we use a speedometer.

• To measure force or weight we use a dynamometer.

• To measure atmospheric pressure we use a barometer.

1.10 Scientific notation (Textbook p.27)

Sometimes we need to write a very large (or very small) number, i.e. the number0.000000045 is very small and it is best written as: 4.5 · 10−8 using scientificnotation.

Scientific notation must be written as the product of two factors. The firstbeing a number greater than 0 and less than 10, times a power of ten.

The following numbers are written using scientific notation:

3 · 101, 1.9 · 100 and 8.10 · 10−3

These numbers are not in scientific notation:

30 · 101, 12.9 and 0.57 · 10−3

1.11 Significant figures (Textbook p.27)

There are three rules. The only tricky numbers are zeros:

I All non-zero digits are significant digits:

⇒ 5 has one significant digit.

⇒ 2.9 has two significant digits.

1.12. UNIT CONVERSION 9

⇒ 129.3468 has seven significant digits.

I Zeros that occur between significant digits are significant digits:

⇒ 1009 has four significant digit.

⇒ 80.9 has three significant digits.

⇒ 100.0008 has seven significant digits.

I Zeros to the right of the decimal point and to the right of anon-zero digit are significant:

⇒ 0.100 has three significant digit.

⇒ 0.00010 has two significant digits.

⇒ 830 has two significant digits.

⇒ 830.0 has four significant digits.

A trick: All the digits of the number part in scientific notation always hassignificant digits.

Examples:

Number Significant figures9107 four401 three

0.0060 two1 800 two

0.3000 four1.00 · 10−6 three

1.12 Unit conversion (Textbook p.25)

The method we are going to learn is called unit conversion factors. It is im-portant to note that this is not the easiest method, but it works in most dificultsituations:

We can make a unit conversion factor by multiplying the value times aunity fractiona(a fraction in which the numerator and the denominatorhave the same value).

aDon’t confuse a unity fraction with a unit fraction which is any fraction hav-ing 1 as numerator.

A unity fraction always equals 1.

Example 1: This is a very simple example, so simple that you may think thatthere is no need to use it. Anyway, let’s pretend that we want to convert 35hours to minutes. We are going to solve it in one step:

hours 1=⇒ minutes

The method:


1. First we multiply the value times an undefined unity fraction.

35 h ·

2. We want to convert hours into minutes, so we need to get rid of thehours. It can be done by writing hours in the denominator because hoursis multiplying:

35 h ·h

3. Minutes must go in the numerator:

35 h · min

h

4. Next we need the numerator being equal to the denominator (the fractionequals 1). This can be easily done because we know that 1 hour equals60 minutes:

35 h · 60 min

1 h

5. We are almost done. Then the hours cancel out:

35 �h ·60 min

1 �h

6. Then, we multiply 35 times 60 and write min:

35 �h ·60 min

1 �h= 2 100 h

That’s it: 35 hours are the same as 2 100 min.

Example 2: Sometimes we can write the unit fraction using prefixes (only ifthese prefixes operate with the same quantity): We want to find how many µm(micrometers) are there in 5 · 108 dam (decameters). This can be achieved inone simple step:

µmstep 1=⇒ dam

Procedure:

1. First we multiply the value times an undefined fraction:

5 · 108 µm ·

2. We want to convert µm into dam, so we need to get rid of the µm. This canbe done by writing µm in the denominator because they are multiplyingin the value:

5 · 108 µm ·µm

1.13. ERRORS 11

3. dam must go in the numerator:

5 · 108 µm · dam

µm

4. Next we need the numerator being equal to the denominator (fractionequals 1). Now there is a trick: decametres and micrometres are prefixesof the same unit: metre. So we can cross their values. We know that theµ prefix equals 10−6, so we write this power of ten in the other side:

5 · 108 µm · 10−6 dam

µm

5. Now we know that the da prefix equals 101, so we write this power of tenin the other side:

5 · 108 µm · 10−6 dam

101 µm

6. We are almost done. We cancel out the µm:

5 · 108 ��µm ·10−6 dam

101 ��µm

7. Finally we work out the result:

5 · 108 ��µm ·10−6 dam

101 ��µm= 5 · 108−6−1 dam = 5 · 101 = 50 dam

And that’s it: 5 · 108 µm are the same as 50 dam.

1.13 Errors (Textbook, p.26)

• Las medidas no son exactas.

• Precisión de un aparato de medida.

1.14 Vocabulary

amount/@"maUnt/ - (n.) cantidad.barometer/ba"r6mIt@/ - (n.) barómetro.base quantity/"beis "kw6nt@ti/ - magnitud fundamental.centi- /"senti/ - (prefix) una centésima.conversion /"k@n"v3:S@n/ - (n.) conversión, transformación.country /"k2ntri/ - (n.) país.deca- /"dek@/ - (prefix) diez veces.derived quantity /dI"raivd "kw6nt@ti/ - magnitud derivada.dynamometer /daIn@"m6mIt@/ - (n.) dinamómetro.electric current/I"lectrIk k2r@nt/ - corriente eléctrica.feet /fi:t/ - (n.) pies.foot /fUt/ - (n.) pie.


gallon /"gæll@n/ - (n.) galón.giga- /"gIga/ - (prefix) mil millones (ES); one billion (EN).hecto- /"hekt@U/ - (prefix) cien vecess.imperial /Im"pI@ri@l/ - (adj.) imperial.inch /IntS/ - (n.) pulgada.kilo- /ki:l@U/ - (prefix) mil veces.length /leNT/ - (n.) longitud.liquid /"lIkwId/ - (n.) líquido.luminous /lu:mIn@s/ - (adj.) luminoso/a.micro- /"maIkr@U/ - (prefix) una millonésima.mass /"mæs/ - (n.) masa.measure /"meZ@/ - (n.) medida; (v.) medir.measurement /"meZ@m@nt/ - (n.) medida.mega- /"meg@/ - (prefix) masa.meter /"mi:t@/ - (n.) metro.mile /"maIl/ - (n.) milla.notation /n@U"teIS@n/ - (n.) notación.numerator /"nju:m@,reIt@/ - (n.) numerador.ounce /aUns/ - (n.) onza.physical quantity /"fIzIkl "kw6nt@ti/ - magnitud física.physics /"fIzIks/ - (n.) física.pico- /paInt/ - (prefix) una billonésima (ES); one million millionth (EN).pint /paInt/ - (n.) física.pound /paUnd/ - (n.) libra.prefix /"pri:fIks/ - (n.) prefijo.pressure /"preS@/ - (n.) presión.process /"pr@Uses/ - (n.) proceso.reproducible /ri:pr@"dju:sIb@l/ - (n.) reproducible.ruler /"ru:l@/ - (n.) regla, metro.scientific /sai@n"tIfIk/ - (adj.) científico/a.set /set/ - (n.) conjunto.speedometer /spI"d6mIt@/ - (n.) velocímetro.temperature /"tempr@tS@/ - (n.) temperatura.tape measure /teip "meZ@/ - (n.) cinta métrica.unit /"ju:nIt/ - (n.) unidad.universal /"ju:nI"v3:s@l/ - (adj.) universal.value /"vælju:/ - (n.) valor.velocity /v@"l6s@ti/ - (n.) velocidad.volume /"v6lju:m/ - (n.) volumen.yard /ja:d/ - (n.) yarda.

2. The states of matter

2.1 Properties of matter. Density

2.1.1 Volume

Volume is the space that an object occupies.

Volume is a derived quantity in the SI. Its SI unit of is m3.The following identities regarding m3 and litres are very useful:

1 m3 = 103 L1 cm3 = 1 mL1 dm3 = 1 L

2.1.2 Mass

Mass is quantitative measure of an object’s resistance to change in velocity.

Mass is a base quantity in the SI. Its SI unit of is kg.The following identities regarding m3 and litres are very useful:

2.1.3 Matter

Matter is anything that has mass and generally occupies a volume.

Examples: Air, salt, water, a book, an atom, etc.We shall see later that matter is formed of small particles called atoms which,in turn, are made up of subatomic particles as protons, neutrons and electrons1.

2.1.4 General and characteristic properties

General properties are those common to all kinds of matter.

Examples: volume, mass, length, temperature, etc. General properties cannotbe used to help identify a substance.

1An electron has mass, but it is considered to be a point particle (no volume).

13

14 CHAPTER 2. THE STATES OF MATTER

Characteristic properties are not the same to all kinds of matter.

Examples: density, electric conductivity, thermal conductivity, viscosity, colour,melting point, boiling point, solubility, magnetism, etc.

They can be used to help identify and classify substances (i.e. many physicaland chemical analysis are based on these properties).

2.1.5 Density

Density is mass per unit volume.

We can calculate density by dividing mass by volume:

density =mass

volumeor d =

m

V

It is a derived quantity. Its SI unit is kg/m3.Other non-SI units may be g/mL, g/L, mg/mm3, etc.

2.2 The aggregation states of matter

2.2.1 Solid, liquid and gas

Matter can be found in different physical states (or aggregation states). Weare going to study the three most common ones: solid, liquid and gas.

You can find some properties in figure (2.1).Liquids and gases are called fluids because they can flow.We can find other states of matter, such as plasma, which is similar to gas butmany electrons are outside their atoms. The Sun and the stars are made out ofplasma. Another state is the Bose-Einstein condensate.

2.2.2 Kinetic theory of matter

Matter is made up of tiny particles animated with a random and continuousmotion.

2.2.3 Kinetic theory and aggregation states

Kinetic theory explains how particles arrange themselves to form this threestates of matter.

2.2. THE AGGREGATION STATES OF MATTER 15

• Volume: constant. + Solids are quite dense.• Shape: constant. + Incompressible. They cannot be squashed.

+ They keep their shape, but some solids may change itif a force is applied to them.

SOLIDS

• Volume: constant. + They keep to the shape of the container.• Shape: variable. + Incompressible. They cannot be squashed.

+ They can flow.+ They can be poured.+ Liquids are quite dense (usually less dense than solids

although water is an exception, because ice is lessdense than liquid water).

LIQUIDS

• Volume: variable. + They can flow.• Shape: variable. + Compressible. Pressure or temperature changes may

modify their volume.+ Gases have a very low density

GASES

Figure 2.1: The three most common states of matter and some of their properties

SOLIDS In figure (2.2) you can find how kinetic theory explains solids.

There are two types of solids:

• Crystals: Particles are arranged in order.Examples: salt, quartz, metals, etc.

• Vitreous solids or glasses: Particles have random fixed positions.Examples: glass, plastic, silex, quartz, etc.

LIQUIDS In figure (2.3) you can find how kinetic theory explains liquids.

GASES In figure (2.4) you can find how kinetic theory explains gases.

2.2.4 Kinetic theory and temperature

Temperature is a quantity that measures the average kinetic energy of theparticles that form an object.

Kinetic energy is the energy that a particle or object has because it is movingabout.

Particles of matter can move slower, but they never can be stopped com-pletely.Temperature is a base quantity. Its SI unit is K.Other units are degrees Celsius (◦C) and degrees Fahrenheit (◦F ).


There is a lower limit to temperature: As we lower the temperature, particlesof matter move slower. If particles could be stopped, temperature could havereached its lower limit: O K.

On the other hand, there is no upper limit to temperature: Particles canalways move faster (for example, temperature in the inner part of the Sun ismillions of degrees).

Thermometers

They are devices used to measure temperature. Some of them are basedon the expansion of a liquid (as mercury or ethanol) when temperature in-creases. Other, more modern ones, are digital thermometres, they have atemperature sensor which measures temperature and writes it on a digitaldisplay.

Scales of temperature

We are going to study three scales of temperature:

• Centigrade or Celsius scale: Its unit is the degree Celsius, (◦C).It may be built by choosing two fixed points:

– The freezing point of water (or melting point of ice) equals 0 ◦C.– The boiling point of water equals 100 ◦C.

So we subdivide these two points into 100 equal parts. Each part is adegree Celsius. See figure (2.5).

• Kelvin or Absolute scale: Its unit is the Kelvin. We can convert betweendegrees Celsius and Kelvin using this equation:

T (K) = t(◦C) + 273

• Centigrade or Celsius scale: Its unit is the degree Celsius, (◦C).It may be built by choosing two fixed points:

– The freezing point of water (or melting point of ice) equals 32 ◦F .– The boiling point of water equals 212 ◦F .

So we subdivide these two points into 180 equal parts. Each part is adegree Fahrenheit. See figure (2.5).Next we write a useful equation for converting to and from degrees Celsiusand Fahrenheit:

t(◦C)− 0

100− 0=t(◦F )− 32

212− 32

t(◦C)

100=t(◦F )− 32

180

2.2. THE AGGREGATION STATES OF MATTER 17

Particles in solids are closely packedtogether (very close to each other) dueto intense attracting forces that the ran-dom motion cannot overcome).

Volume is constant be-cause particles cannot getmuch closer nor they canmove further away (incom-pressible)

These particles occupy a fixed positionalthough they vibrate a little around thisposition.

Shape is constant be-cause particles have a fixedposition.

Figure 2.2: Solids and kinetic theory.

Cohesion forces in liquids are not sostrong as in solids, so liquids formgroups of particles separated by smalldistances.

Volume is constant be-cause particles cannot getmuch closer nor they canmove further away.

These particles do not have a fixed pos-ition. These groups of particles are freeto move about within the vessel.

Shape is variable.Liquids can flow.They can be poured.

Figure 2.3: Liquids and kinetic theory.

Attraction forces are very small, so gasparticles are far apart from each other.

Volume is variable be-cause particles are faraway from the others sothe space between themmay change a lot (com-pressible).

These particles do not have a fixed po-sition and they have an independentand random motion, so they collide witheach other and with the walls of thecontainer.

Gas particles originate apressure.Gases can flow.

Figure 2.4: Gases and kinetic theory.


t(◦C) t(◦F)

10

20

30

40

50

60

70

80

90

100subd

ivisions

180subd

ivisions

Freezing point of water

Boiling point of water

0 ◦C

100 ◦C

32 ◦F

212 ◦F

Figure 2.5: Celsius and Fahrenheit temperature scales.

3. Material systems

3.1 Classifying matter (Textbook p.62)

Material systems can be classified in different ways. According to the state ofmatter: solids, liquids and gases. We are now interested in the classification ofmaterial systems according to composition, as shown in figure (3.1):

Material systems

Pure substances Mixtures

Elements Compounds Heterogeneous Homogeneous

Figure 3.1: Classification of matter according to its composition.

3.1.1 Pure substances (Textbook p.62)

A pure substance is a piece of matter that consists of just one component.

Examples: oxygen: O2, water (or distilled water), H2O, iron: Fe, carbon dioxide:CO2, . . .

Properties

• They have constant physical and chemical properties (boiling point,melting point, density, colour, hardnessa at given conditions.

• They also have a fixed composition.• Their characteristic properties (density, conductivity, etc.) may be

used to identify the substance.aIt is the scratch resistance of a solid. The hardest known mineral is diamond.

19

20 CHAPTER 3. MATERIAL SYSTEMS

3.1.2 Elements (Textbook p.63)

An element is a pure substance that is made of just one kind of atoms.

Examples: oxygen: O2, hydrogen: H2, ozone: O3, iron: Fe, lead: Pb, graphite:C, diamond: C, tin: Sn, silver: Ag, chlorine: Cl2, . . .

Properties

• We cannot get simpler substances from them.• They have constant physical and chemical properties at given

conditions.• They also have a fixed composition.• Their characteristic properties may be used to identify the substance.

3.1.3 Compounds (Textbook p.63)

A compound is a pure substance consisting of two or more different typesof atoms.

Examples: water (distilled water): H2O, sodium chloride (table salt): NaCl,carbon dioxide: CO2, . . .

Properties

• We can get simpler substances (elements) from them by means ofchemical reactionsa.

• They have constant physical and chemical properties at givenconditions.

• They also have a fixed composition.• Their characteristic properties may be used to identify the substance.

aWe can get hydrogen, H2, and oxygen, O2 from water by passing an electringcurrent through water.

3.1.4 Mixtures

A mixture is a piece of matter consisting of two or more different sub-stances.

Examples: iron filings and sand: Fe and SiO2, brass: Cu and Zn, bronze: Cuand Sn, salted water: NaCl and H2O, . . .

3.1. CLASSIFYING MATTER 21

Properties

• Substances that form a mixture are blended together, but they stillretain their properties (because they are not combined chemically).

• The substances that form a mixture still exist, so we can separatethem using physical transformations.

• Their composition is variable within certain limits.

3.1.5 Heterogeneous mixtures (Textbook p.62)

A heterogeneous mixture is a type of mixture that hasn’t got the sameproperties all over the system.

Examples: iron filings and sand: Fe and SiO2, sawdust and chalk, sugar (sucrose)and flour, . . .

Properties

• We can see the different substances (sometimes you may need to usea magnifying glass or a microscope).

• Their properties change abruptly from place to place.• They haven’t got a fixed composition.• Substances can be separated using some physical transformations.

3.1.6 Homogeneous mixtures (Textbook p.62)

A homogeneous mixture is a type of mixture that has the same propertiesthroughout.

When they have a liquid look, they are called solutions. If they are solid,they are called alloys.Examples: tap water: H2O and mineral salts, bronze: copper and tin, brass:copper and zinc, air: mainly nitrogen and oxygen, sucrose and water, . . . .

Properties

• You cannot see the different parts even with a microscope.• Their properties are the same throughout or they change smoothly

from place to place.• They haven’t got a fixed composition.• Substances can be separated using some physical transformations as

distillation.


3.1.7 Colloids and suspensionsColloids are homogeneous mixtures with particle sizes that consists of clumpsof molecules. The particles have dimensions between 2 and 1 000 nm. Colloidslook homogeneous to the naked eye. Fog and milk are examples of colloids.Their particles cannot be separated by filtration. Colloids have normally anopaque or translucent look because their particles are big enought to scatterlight. Colloids are stable over time.

Suspensions are mixtures whose particles are bigger than those of colloids.Their particles are visible to the naked eye. Suspensions look opaque or trans-lucent. Their particles can be separated by filtration. Suspensions separate onstanding.

3.2 Solutions (Textbook p.64)

A solution is a homogeneous mixture, which in turn, is a mixture of two ormore substances that has the same properties all over it.

3.2.1 Solvent and solute (Textbook p.64)

Solvent is a substance that dissolves another substances.

Usually it is the substance that is more abundant in a solution1.

Solute is a substance that is dissolved by a solvent to to create a solution.

Usually it is the substance that is less abundant in a solution. Note that asolution might have more than just one solute (example: tap water has waterand many mineral salts (many solutes)).

3.2.2 Classification of solutions attending to physical state(Textbook, p.65)

A solution can exist in solid, liquid or gas form depending on mixed substancesand external conditions such as temperature and pressure. See examples onpage 65 of your textbook.

3.2.3 Classification of solutions attending to composition(Textbook, p.65)

Saturated solution

A saturated solution is a solution in which the maximum amount of solutehas been dissolved.

1Sometimes we may use another criterion: If one substance has a solid state and the otheris liquid, then, if the solution has a liquid look we may say that the liquid substance is thesolvent because it has the same physical state look as the solution.

3.3. CONCENTRATION OF A SOLUTION 23

This amount depends on temperature. Any more solute added will sit as crystalson the bottom of the container.

Supersaturated solution

A supersaturated solution is a solution which has more solute than a sat-urated one.

Under certain conditions we can solve more solute. Obviously, a supersaturatedsolution is not at equilibrium, and may quickly crystallize the excess of solute.See on youtube:

https://www.youtube.com/watch?v=HnSg2cl09PI

A diluted solution is a solution in which there is a small amount of solutecompared to the total amount of possible solute that can be dissolved.

A concentrated solution is a solution which has a lot of solute dissolvedcompared to the total amount of possible solute that can be dissolved.

A saturated solution is also a concentrated one.

3.3 Concentration of a solution (Textbook, p.66)

Concentration of a solution is the abundance of solute in a certain amountof solution.

The concentration of a solution tells us about the proportion in which substanceshave been mixed up.

3.3.1 Percent by mass (Textbook, p.66)

Percent by mass of a solution is the mass of solute in 100 units of mass ofsolution:

% mass =solute masssolution mass

· 100

We can use any unit of mass, provided we use it both in the numerator and inthe denominator. The result has no units, it is expressed in % mass.

3.3.2 Percent by volume (Textbook, p.67)

Percent by volume of a solution is the volume of solute in 100 units ofvolume of solution:

% vol. =solute vol.solution vol.

· 100

We can use any unit of volume, provided we use it in the whole expression. Theresult has no units, it is expressed in % vol.


3.3.3 Grams of solute per litre (Textbook, p.67)

Concentration in grams per litre of a solution is the mass of solute in1 litre of solution:

c(g/L) =grams of solutelitres of solution

3.4 Solubility (Textbook, p.68)

Solubility of a solute in a solvent is the maximum amount of solute thatcan be solved either in a given amount of solvent or in a given amount ofsolution. Solubility depends on temperature.

The solubility of a solute can be given as the concentration of a saturatedsolution of the solute on a given solvent. Solubility of solids normally increasewith temperature. Solubility of gases decrease with temperature.

3.4.1 Solubility curve (Textbook, p.69)

We can draw the solubility of some salts in water (see textbook, page 69).

3.4.2 Solubility of gases in liquids (Textbook, p.69)

As temperature increases, the solubility of a gas will decrease.As pressure increases, the solubility of a gas will increase.

3.5 Osmosis (Textbook, p.70)

3.6 Separation of mixtures (Textbook, p.72)

3.6.1 Separation of heterogeneous mixtures (Textbook, p.72-73)

3.6.2 Separation of homogeneous (Textbook, p.74-75)

4. Atoms

4.1 Dalton’s atomic theory (Textbook p.90)

Between 1803 and 1805, British chemist John Dalton thought about some chem-ical laws and came up with the conclusion that matter is made of atoms.The word atom comes from the ancient greek and means indivisible.

4.1.1 The model (Textbook p.90)

1. Elements are made of atoms, which are tiny, indivisible andindestructible.

2. All atoms of a given element are identical in mass and properties.3. Different elements have different types of atoms.4. Compounds are formed by a combination of two or more different

kinds of atoms.

4.2 Thomsom’s atomic theory (Textbook p.94)

J. J. Thomsom was a British scientist who discovered the electron and thatelectrons are subatomic particles.

4.2.1 Cathode rays (Textbook p.92)

In a famous experiment with vacuum tubes he discovered the cathode rayswhich are a beam of particles that have mass and negative charge. Theseparticles where called electrons. In short, he found out that atoms had elec-trons inside them.

4.2.2 The plum cake model (Textbook p.90)

An atom consists of a positively charged sphere in which the electrons areembedded like plums or raisins in a pudding.

25

26 CHAPTER 4. ATOMS

4.3 Rutherford’s atomic theory (Textbook p.96)

A few years after the Thomsom’s atomic model, a new particle was discovered:the proton.

A proton has one positive charge (the exact opposite as the electron, which hasa negative one) and its mass is far larger than the mass of the electron.

It soon became clear that atoms had both electrons and protons. Rutherford,along with Geiger and Marsden carried out an experiment that led to his atomictheory.

4.3.1 The gold foil experiment (Textbook p.96)

They fired a beam of alpha particles coming from a radioactive material ata thin sheet of gold (gold foil). They made these observations:

1. Most of the alpha particles passed right through the gold foil.2. Some particles suffered a small scattering due to something with a

positive charge.3. Very few alpha particles bounced backwards.

4.3.2 Rutherford’s model (Textbook p.96)

This model is called the planetary model because it resembles a small solarsystem.

An atom has two different parts:

1. A nucleus, which is a small, positively charged blob of matter in thecenter of the atom. Almost all the mass of an atom is in the nucleus.It is composed of protons and neutronsa.

2. A shell, which is a big space surrounding the nucleus. It is almostempty. Electrons are moving in circles around the nucleus, so theshell has a negative charge.

aNeutrons where discovered later, in 1932. The mass of neutrons is almost thesame as the mass of the protons. Neutrons have no electric charge.

4.3. RUTHERFORD’S ATOMIC THEORY 27

4.3.3 Subatomic particles

There are three subatomic particlesa:

• Electrons: They are particles with a very tiny mass and one negativecharge. Electrons are in the atom’s shell.

• Protons: They are more massive than electrons and have onepositive charge. Protons are in the nucleus of the atom.

• Neutrons: They have approximately the same mass as protons. Theyhave no charge and they are in the nucleus.

aWe now know that protons and neutrons are not elementary particles, they aremade of quarks

4.3.4 Atomic number (Textbook p.98)

Atomic number, Z, of an atom is the number of protons in this atom.

4.3.5 Mass number (Textbook p.98)

Mass number, A, of an atom is the number of protons and neutrons inthis atom.

Almost all the mass of an atom is due to the protons and neutrons (nucleons),so this number gives an idea of its mass (hence its name).

4.3.6 Chemical element

A chemical element is a pure substance that is composed of atoms withthe same atomic number.

The Periodic Table of the element consists of all known elements ordered bytheir atomic number.

4.3.7 Isotopes of an element (Textbook p.99)

Isotopes of an element are atoms with the same atomic number but differ-ent mass number.

28 CHAPTER 4. ATOMS

4.3.8 Ions

Atoms are electrically neutral because they have the same number of electronsand protons.

If the number of electrons is not equal to the number of protons then theatom has an electric charge and it is called an ion.

It is very difficult to remove or add protons to an atom because they aredeep inside it and they are closed together due to very big attraction forces. Itis easier to add or remove electrons from atoms because they are in the shell(the outer part).

A ion is an atom with an electric charge.If the atom has a negative charge (more electrons than protons), then it iscalled anion.If the atom has a positive charge (less electrons than protons), then it iscalled cation.

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