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CHEMISTRY FOUNDATIONS

Chemistry is the study of matter and its transformations. More specifically, it looks at how some materials interact with other materials.

Matter can be operationally defined as a substance that has:

• Mass

• Occupies space (i.e., it has volume)

Mass is a measure of the physical amount of a substance. Mass is not to be confused with weight although the two terms are sometimes used interchangeably in everyday life. Weight is the action of the acceleration due to gravity on mass; as such, it is a force. For example, a man with a mass of 50 kg on earth will still have a mass of 50 kg on the moon, but will weigh about 1/6 less on the moon than on the earth. This is because the moon’s gravity is about one-sixth that of the earth.

Matter can be found in three basic forms (referred to as the states of matter):

1. Solid – the material has a definite size and shape

2. Liquid – the material has a definite size, but takes on the shape of the container

3. Gas – the material has no definite size or shape; it expands (or contracts) to fill the container

1. PHYSICAL AND CHEMICAL PROPERTIES AND CHANGES The physical properties of a substance are the physical characteristics by which a substance may be recognised. These properties include for example, colour, odour, state (solid, liquid, gas etc.), volume, density, melting point and boiling point (substances have many other physical characteristics).

The chemical properties of a substance tell us how the substance combines with other substances to form new substances (having different physical properties). For example, a chemical property of hydrogen is that under certain well-defined conditions, it combines with oxygen to from water. Water has completely different properties from the substances of which it is formed (hydrogen and oxygen).

Exercise: Classify the following properties as physical or chemical.

1. Gold is yellow.

2. Gold resists a change in identity.

3. Iron rusts when exposed to air and water.

4. A 1 kilogram block of lead occupies less space than 1 kilogram of feathers.

5. Ammonia smells differently from hydrogen sulphide.

Physical changes occur when an object is subjected to a relatively mild physical stress (e.g., heating a block of ice (solid water) may result in the formation of liquid water. Physical changes may be recognised because they are often reversible. For example, when ice is warmed, it eventually melts to form liquid water. If the liquid water is sufficiently cooled, it becomes ice once more.

Chemical changes also occur when an object is subjected to physical stress. The difference between physical changes and chemical changes is that in chemical changes, new substances are formed; the new substances do not have the same properties as the original substances. Unlike physical changes, chemical changes are not often easily reversible. For example, if a mixture of hydrogen gas and oxygen

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gas is heated strongly, water will be produced. However cooling back the water will not regenerate the original hydrogen gas and oxygen gas. When two substances are mixed, one or more of the following observations may signal that a chemical change has occurred:

1. A solid is formed on mixing two liquids.

2. A gas is formed on mixing two substances.

3. A colour change has occurred on mixing two substances.

4. The temperature of the mixture changes (the mixture gets hotter or colder).

2. ATOMS All matter is composed of atoms. Atoms may vary in mass, size and in the way they interact with other atoms but they all share some common characteristics.

For the purposes of chemistry, all atoms are made up of three fundamental particles.

1. The electron (e–) is a subatomic particle (i.e., smaller than an atom) that has a negative charge.

2. The proton (p+) is a subatomic particle that has a positive charge equal in magnitude but opposite in sign to the electron.

3. The neutron (n0) is a subatomic particle that has no charge.

Each atom has a specific number of protons, electrons and neutrons.

The structure of all atoms is as follows.

All of the protons and neutrons exist in a tightly packed region at the centre of the atom called the nucleus. The electrons are in constant motion outside the nucleus; they move about the nucleus at random but do occupy preferred regions in space.

The atom is mostly empty space. For example, the nucleus of the smallest atom, hydrogen, is about 10–15 m. To understand the scale of an atom, suppose the hydrogen atom nucleus is expanded to the size of a tennis ball and centered at Dawson College; then the electron of the hydrogen atom will be, on average, about half a kilometre further west than Decarie Boulevard. The point is that the electron is rather far from its nucleus.

In a free atom the number of protons equals the number of electrons.

The diverse character of different atoms is due to the number of protons that are packed into the small dense nucleus together with and as well, to the number and organization of the electrons in the volume outside the nucleus.

The nucleus consists of all of the protons and neutrons

The electrons are moving around in the space outside of the nucleus.

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2.1. Atomic Mass Electrons are so tiny that they weigh almost nothing (9.109×10–31 kg); protons and neutrons are nearly 2000-times heavier). Almost all of the mass of an atom therefor comes from the protons and neutrons, which reside in the nucleus. However, even protons and neutrons are too small to be conveniently weighed in kilograms (or pounds for that matter). Therefore scientists have defined a unit called the unified atomic mass unit (symbol, u), also known as the Dalton (symbol, Da). This 1 u = 1 Da. The unified atomic mass unit is approximately the mass of one proton (as well as one neutron). Thus , if we know the total number of protons and neutrons, we know the mass of the atom in units of Da.

Example: Calculate the mass of an atom that has 6 protons and 6 neutrons.

Since the mass of each proton is 1 Da and the mass of each neutron is 1 Da, the total mass 1 1 6 6 12 total

Da Dam protons neutrons Daproton neutrons

= + =

2.2. Isotopes We know that the chemical properties and many of the physical properties of an atom depend on the energy, number, and arrangement of its electrons. We also know that the number of protons must equal the number of electrons in a free atom. Thus the character of an atom is defined through its protons and electrons.

Unlike electrons and protons, neutrons do not contribute to the chemical or most of the physical properties of the atoms. Atoms of a given type can have variable number of number of neutrons. Thus, for example, a sample of pure carbon atoms can contain carbon atoms having 6, 7, 8 or more neutrons but must always contain only 6 protons and 6 electrons; if one of the atoms has a different number of protons, it cannot be a carbon atom.

Atoms of a given element always have the same number of protons but may have different numbers of neutrons. Such atoms are known as isotopes.

For example, the element hydrogen has three isotopes.

Name Symbol Atomic Number

Number of Protons

Number of electrons

Number of neutrons

Mass Number

Hydrogen-1 (Protium)

11 H 1 1 1 0 1

Hydrogen-2 (Deuterium)

11 H 1 1 1 1 2

Hydrogen-3 (Tritium)

31 H 1 1 1 2 3

Note that all of these atoms are equally “hydrogen” (although the hydrogen atom having 0 neutrons is by far the most abundant). These different types of atoms of hydrogen, having different numbers of neutrons are called isotopes. The conventional notation for writing isotopes is as follows:

ZASy

Here, Sy = the one to two letter symbol for the atom; Z = the atomic number of the atom (the number of protons, also equalling the number of

electrons in the free atom); and A = the mass number of the isotope.

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3. ELEMENTS Elements are composed of chemical combinations of isotopes. Thus, for example, the element oxygen is commonly found as a chemical combination of two oxygen atoms (two identical isotopes of oxygen or two different isotopes of oxygen). The element oxygen is written as O2 to indicate this.

There are more than a hundred known elements of which about 90 occur naturally on earth. The remainder of the elements have been produced artificially. Each element is composed of isotopes of a given atom.

The number of protons present in the nucleus identifies the atom uniquely as being the atom of a specific element.

Atoms of the same element:

• Always have exactly the same number of protons and electrons. This number of protons defines the element uniquely and is known as the atomic number (symbol: Z = the number of protons in the atom).

• May have different numbers of neutrons (i.e., may have different isotopes). Each element is made up of different isotopes in a fixed relative abundance.

Atoms of different elements always have a different number of protons (hence different values of Z).

Compounds (molecules) are chemical combination of atoms of different elements. A compound always consists of a chemical association of atoms of different elements. The physical and chemical properties of a compound are different from those of its building block elements. A compound can be broken down into its constituent elements through chemical changes, although this will often require a lot of energy.

A compound always has the same chemical combination of atoms of its building block elements. For example, the compound carbon dioxide always consists of one carbon isotope chemically combined with two oxygen isotopes. It is written as CO2. All samples of carbon dioxide, regardless of their origin, consist of exactly one atom of carbon and two atoms of oxygen.

4. PURE SUBSTANCES AND MIXTURES

4.1. Pure Substances A pure substance contains only one type of material. There are only two classes of pure substances:

• elements • compounds

Pure substances are rarely found in nature.

4.2. Mixtures Mixtures are by far the most common way that chemicals are found. They are a physical combination of different compounds or elements (i.e., physically combining different pure substances). Mixtures may be either homogeneous or heterogeneous. Homogeneous mixtures (more commonly referred to as solutions) are uniform in composition throughout. For example, when table salt (sodium chloride) is dissolved in pure water, the resulting salt water has a uniform appearance. No part of this mixture differs from any other part, so salt water is a homogeneous mixture.

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Heterogeneous mixtures do not have a uniform appearance. Parts of such mixtures physically appear different. For example, if sand is poured into water, it settles to the bottom. The mixture of sand and water is therefore a heterogeneous mixture. Schematically, the relationship between these forms of matter is depicted in the figure below.

Schematic Description of Substances Exercise: Classify the following substances as a pure substance, a homogeneous mixture or a

heterogeneous mixture:

1. Wine

2. Air

3. A salt-sand mixture

4. Zinc

5. Oxygen

6. Tap water

7. Steel

8. Wood

9. Blood

10. A page in the textbook

5. THE PERIODIC TABLE OF ELEMENTS The Periodic Table of Elements is an organisation of atoms according to their properties. It was devised in 1867 by the Russian chemist, Dimitry Mendeleev and independently that same year by the German, Lothar Meyer. Mendeleev is considered to be the “Father of Modern Chemistry” because Meyer did not make any further use of his table.

Mendeleev (and Meyer) noted that many of the known elements have very similar physical and chemical properties. Those atoms with such similar properties were put in columns. Thus, for example, atoms such as Na (sodium) and K (potassium) were soft, silvery, greyish metals which reacted explosively with water to form extremely strong basic solutions. Elements with such similar properties were organized into columns, called groups in the Periodic Table.

Fundamental Particles e–, p+, n0

Atoms Pure Substances Mixtures

Elements Compounds Heterogeneous mixtures

Homogeneous mixtures

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For example, the group consisting of Hydrogen (H), Lithium (Li), Sodium (Na), Potassium (K), etc., is known as the alkali metals. Other metals such as Mg and Ca found in the earth also formed alkaline solutions with water, but not in nearly the same explosive fashion as Na and K. These metals were known as the alkaline-earth metals.

This arrangement of atoms according to their properties soon proved its value. Gaps in Mendeelev’s Table soon began to appear, and he was able to predict the existence of as yet unknown elements. Chemists began to search for these “missing” elements and found several of them.

The modern Periodic Table has three basic characteristics that provide information about the atom. The example of the carbon atom is shown below.

The elements may be classified according to whether they are metals, nonmetals or metalloids (or semimetals).

A Periodic Table of Elements is provided on the following page. You are expected to know the names and symbols of some common elements (these have been highlighted in the Periodic Table on the following page).

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C

12.011

Atomic Number, Z, (number of protons and electrons present in the free atom)

Mass Number, u, (the average mass of a carbon atom in atomic mass units)

Symbol of the atom

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PERIODIC TABLE OF THE ELEMENTS 1 1.0079

H hydrogen hydrogène

2 4.003

He helium hélium

3 6.941

Li lithium lithium

4 9.012

Be beryllium béryllium

atomic number atomic mass

Symbol English name French name*

5 10.811

B boron bore

6 12.011

C carbon carbone

7 14.007

N nitrogen azote

8 15.9994

O oxygen oxygène

9 18.998

F fluorine fluor

10 20.18

Ne neon néon

11 22.99

Na sodium sodium

12 24.31

Mg magnesium magnésium

*all are masculine 13 26.98

Al aluminum aluminium

14 28.086

Si silicon silicium

15 30.974

P phosphorus phosphore

16 32.07

S sulfur soufre

17 35.453

Cl chlorine chlore

18 39.95

Ar argon argon

19 39.098

K potassium potassium

20 40.08

Ca calcium calcium

21 44.96

Sc scandium scandium

22 47.87

Ti titanium titane

23 50.94

V vanadium vanadium

24 51.996

Cr chromium chrome

25 54.94

Mn manganese manganèse

26 55.85

Fe iron fer

27 58.93

Co cobalt cobalt

28 58.69

Ni nickel nickel

29 63.546

Cu copper cuivre

30 65.39

Zn zinc zinc

31 69.72

Ga gallium gallium

32 72.61

Ge germanium germanium

33 74.922

As arsenic arsenic

34 78.96

Se selenium sélénium

35 79.904

Br bromine brome

36 83.80

Kr krypton krypton

37 85.468

Rb rubidium rubidium

38 87.62

Sr strontium strontium

39 88.91

Y yttrium yttrium

40 91.22

Zr zirconium zirconium

41 92.91

Nb niobium niobium

42 95.94

Mo molybdenum molybdène

43 98

Tc technetium technétium

44 101.07

Ru ruthenium ruthénium

45 102.91

Rh rhodium rhodium

46 106.42

Pd palladium palladium

47 107.87

Ag silver argent

48 112.411

Cd cadmium cadmium

49 114.82

In indium indium

50 118.71

Sn tin étain

51 121.76

Sb antimony antimoine

52 127.60

Te tellurium tellure

53 126.904

I iodine iode

54 131.29

Xe xenon xénon

55 132.91

Cs cesium césium

56 137.33

Ba barium baryum

57 138.91

La lanthanum lanthane

72 178.49

Hf hafnium hafnium

73 180.95

Ta tantalum tantale

74 183.84

W tungsten tungstène

75 186.21

Re rhenium rhénium

76 190.21

Os osmium osmium

77 192.22

Ir iridium iridium

78 195.08

Pt platinum platine

79 196.97

Au gold or

80 200.59

Hg mercury mercure

81 204.38

Tl thallium thallium

82 207.2

Pb lead plomb

83 208.98

Bi bismuth bismuth

84 209

Po polonium polonium

85 210

At astatine astate

86 222

Rn radon radon

87 223

Fr francium francium

88 226.03

Ra radium radium

89 227.03

Ac actinium actinium

104 261

Rf rutherfordium rutherfordium

105 [268]

Db dubnium

106 [271]

Sg seaborgium

107 [270]

Bh bohrium

108 [277]

Hs hassium

109 [276]

Mt meitnerium

110 [281]

Ds darmstsdtium

111 [280]

Rg Roentgenium

112 [285]

Cn copernicium

113 [284]

Uut Ununtrium

114 [289]

Uuq Ununquadium

115 [288]

Uup Ununpentium

116 [293]

Uuh Ununhexium

117 [294]

Uus Ununseptium

118 [294]

Uuo Ununoctium

58

140.12

Ce cerium cérium

59 140.91

Pr praseodymium praséodyme

60 144.24

Nd neodymium néodyme

61 145

Pm promethium prométhium

62 150.36

Sm samarium samarium

63 151.96

Eu europium europium

64 157.25

Gd gadolinium gadolinium

65 158.93

Tb terbium terbium

66 162.50

Dy dysprosium dysprosium

67 164.93

Ho holmium holmium

68 167.26

Er erbium erbium

69 168.93

Tm thulium thulium

70 173.04

Yb ytterbium ytterbium

71 174.97

Lu lutetium lutécium

90 232.038

Th thorium thorium

91 231.036

Pa protactinium protactinium

92 238.029

U uranium uranium

93 237.048

Np neptunium neptunium

94 244

Pu plutonium plutonium

95 243

Am americium américium

96 247

Cm curium curium

97 247

Bk berkelium berkélium

98 251

Cf californium californium

99 252

Es einsteinium einsteinium

100 257

Fm fermium fermium

101 258

Md mendelevium mendélévium

102 259

No nobelium nobélium

103 262

Lr lawrencium lawrencium

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1.1 Classification of the Elements Atoms of the elements are organized in the Periodic Table according to property in the sequence of their atomic numbers, Z (the number of protons in the atom). There are several ways in which the elements are classified and it is necessary to know them.

Metals and Nonmetals 1.1.1The vast majority of elements are metals.

Alka

li m

etal

s

Alka

line-

Eart

h m

etal

s

Pnic

toge

ns

Chal

coge

ns

Halo

gens

B

Aℓ Si

Ge As

Sb Te

Po At

Non-metals are to the right of the stepladder structure (bold lines).

Metalloids (semi-metals) are at the boundary of the stepladder and are indicated on the Periodic Table above by the lighter shade of grey. They have many properties typical of the metals but are generally semi-conductors rather than the excellent conductors of electricity and heat that metals are.

MONATOMIC AND DIATOMIC ELEMENTS 1.1.2The elements written in outline font in the Periodic Table above are diatomic elements. That is, they are found in nature under normal conditions as a chemically bonded pair of atoms. Thus, for example, the elements hydrogen, nitrogen, oxygen and iodine (as well as fluorine and bromine) are found respectively as:

H2, N2, O2, Cℓ2 and I2.

The first four diatomic elements in the list above (H2, N2, O2and Cℓ2) are gases; of the last element I2, is a solid.

Apart from this group of diatomic elements, all other elements are treated as if they are monatomic; that is, the atom is the same as the element. This is strictly speaking not true but is useful for the purposes of simplicity.)

Transition metals

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6. FORMATION OF COMPOUNDS In nature, all systems tend to move to the state of greatest stability. Thus, for example, atoms combine with each other to form compounds only so that they can become more stable.

Recall that an atom consists of protons and neutrons densely packed into a nucleus at the centre of the atom. The electrons are moving rapidly outside of, and far away, from the nucleus. Because of this separation between the electron and its nucleus, the electron can “forget” to which nucleus it belongs and attach itself to the nucleus of another atom that may be closer. Therefore the electrons of can become “lost”. Thus, atoms are capable of losing and gaining electrons. This movement of electrons between nuclei results in the formation of chemical bonds – either ionic bonds or covalent bonds.

1.2 Ionic Bonds If an atom gains an electron, it acquires an overall negative charge because each electron carries a negative charge.

e.g., Cℓ + e– → Cℓ–

Similarly, if an atom loses an electron, it acquires a positive charge (i.e., it loses a negative charge).

e.g., Na → Na+ + e–

Atoms acquire a charge only by gaining or losing electrons NOT by gaining or losing protons. Note that the total number of electrons is conserved. The electrons are simply jumping from one nucleus to another.

Substances may lose or gain more than one electron.

e.g. Fe → Fe3+ + 3 e–

O + 2 e– → O2–

Anions are chemical substances which have negative charges (e.g., N3–, O2–, I–)

Cations are chemical substances which have positive charges (e.g., H+, Na+, Ca2+)

+

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Typically:

Metals, to the left of the Periodic Table, tend to lose electrons thereby forming cations.

Non-metals tend to gain electrons thereby forming anions.

In general, substances will not give up electrons unless there are other atoms ready to receive those electrons. Likewise, a substance cannot accept electrons without other substances ready to supply those electrons. Thus, for example, sodium can lose an electron:

Na → Na+ + e–

but will not do so unless it is in the presence of another atom willing to receive that electron, such as chlorine:

Cℓ + e– → Cℓ–

Combining these equations

Na + Cℓ + e– → Na+ + e– + Cℓ–

Cancelling the electrons on either side of the equations yields the net result:

Na + Cℓ → Na+ + Cℓ–

Note that there are no net free electrons remaining at the end. All of the electrons lost by Na atoms have been absorbed by the Cℓ atoms. At this point the system consists of a cation (Na +) and an anion (Cℓ–). These two oppositely charged ions now attract each other to form an ionic bond.

Na+ + Cℓ– → Na+ Cℓ– = NaCℓ

The attraction of positive to negative results in the formation of an ionic bond. The resulting neutrally charged compound is known as an ionic compound. Ionic compounds are crystalline solids.

1.3 Covalent Bonds Nonmetals, to the right of the step-ladder in the Periodic Table) tend to gain electrons. In ionic compounds, a metal atom is present to supply the needed electron(s). If, however, no such supplier is present (i.e., there is no metal), another method of acquiring the needed electrons must be found.

Consider the case of a sample consisting only of 17Cℓ atoms. Each chlorine atom in the sample needs one electron which it can only obtain by stealing it from a nearby chlorine atom. However none of the chlorine atoms want to give up their electrons. Each chlorine atom seeks to seize one electron from the other chlorine atom while retaining all of its own electrons.

Cl Cl

The result is that the two atoms are drawn together with the two electrons, one from each fluorine atom, located between the two nuclei.

Cℓ Cℓ

This type of bond is known as a covalent bond. In covalent bonds, the electrons are shared between two atoms.

Covalent bonds are formed between two nonmetal atoms.

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7. BOND POLARITY Covalent bonds are formed when two nonmetal atoms combine chemically by sharing electrons. However the electrons are not always shared equally. Consider the bond between a carbon and oxygen atom. Both atoms are nonmetals so the two atoms share a pair of electrons. However the carbon atom has 6 protons, whereas the oxygen atom has 8 protons. The shared electrons (negative) will therefore be pulled more strongly toward the more positive oxygen nucleus:

The result is that one half of the bond (in this case on the oxygen side) appears to have a negative charge while the other half of the bond appears to be positive (the carbon side). The negative side is referred to as having a partial negative charge (δ–); remember this is a covalent bond. In an ionic bond this would be a full negative charge (–). Similarly, the positive side is referred to as having a partial positive charge (δ+).

A bond polar is sometimes simply represented by an arrow pointing from the positive end to the negative end of the bond.

The unequal sharing of electrons in a covalent bond thus results in one end of the bond being partially negative while the other end is partially positive. Thus the bond is said to be a polar covalent bond (or just simply a polar bond). Most covalent bonds are polar. The common bond types in biology and their relative polarity are provided below.

H—C C—N C—Cℓ N—O H—N H—Cℓ C—O H—O

Least Polar Most Polar

6C 8O

Midpoint between the two atoms

Electron pair is pulled toward the oxygen

6C 8O

This half of the bond appears to be negative because the electrons are on this side

This half of the bond appears to be positive because the electrons are on the other side

δ– Partial negative charge

δ+ Partial positive charge

A δ+

B δ–

Equal bond polarities Equal bond polarities

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8. Hydrogen Bonding and Water There are several very important consequences of bond polarity. These relate to the shape of molecules, and consequently, the overall polarity of the molecule and some physical properties of the molecule.

Consider the case of the molecule carbon dioxide, CO2. Carbon dioxide is a linear molecule with carbon in between two oxygen atoms.

O C O

180°

The double lines represent double bonds, namely two electron pairs linking the atoms together, but these do not change the discussion. From the discussion in Section 7, the C—O is quite polar.

C—O

There are two polar bonds present in the structure of CO2, one for each (C—O) bond.

O C O

Since the O —C—O angle is 180°, the “pull” of electrons to the left is precisely balanced by the “pull” of the electrons to the right. The net effect is that there is no net “pull” so the molecule overall is nonpolar.

Contrast the CO2 molecule with water, H2O. Water is has a bent structure with oxygen in between two hydrogen atoms.

HO

H104.5°

Oxygen attracts electrons much better than hydrogen, so each H—O is highly polar. However, the fact that the water molecule is bent means that all of the “pull” of electrons do not cancel off as in CO2.

HO

H In fact, water has a net overall polarity, that is, it has an overall positive end and a negative end. Therefore the whole water molecule can be represented by an arrow as for individual polar bonds.

If two water molecules are brought in contact with each other, since they each have positive and negative ends, the positive end of one water molecule will naturally align with the negative end of the other water molecule. The attraction between the positive end of one molecule with the negative end of another molecule in polar molecules have consequences such as high boiling temperatures. In the examples discussed above

Molecule Polarity Boiling Temperature at normal pressure

O C O Nonpolar –57°C

HO

H Highly Polar +100°C

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The interaction between different molecules is known as intermolecular forces (i.e., forces between molecules).

All polar molecules will have relatively high boiling points but the bond polarities of the N—H and O—H bonds are particularly strong. Because of this, polar interactions between molecules that have N—H or O—H bonds are so strong (much stronger than average) that they are given the special term hydrogen-bonding (H-bonding).

9. SOLUTIONS, ELECTROLYTES AND pH When an ionic material is dissolved in water, something similar to the interaction between two water molecules occurs. The positive end of a water molecule will align with the anion (negative ion) of the ionic compound. Similarly, negative end of a water molecule will align with the cation (positive ion) of the ionic compound. These attractions can pull the ionic compound apart resulting in the freed ions; the ionic compound is said to have dissociated (dis + associated) in water. The freed ions can conduct electricity, thus any substance that releases ions in water will be an electrical conductor and is therefore called an electrolyte. Strong electrolytes dissociate completely or almost completely in water. For weak electrolytes, only some of the ionic compound will dissociate so there will not be many ions in solution.

Of particular interest in chemical and biological systems are acids and bases, so the following definitions are important.

9.1. Acids Acids can be defined as substances that release H+ in water. Acids can be strong (strong electrolytes) or weak (weak electrolytes).

Examples of common acids

Hydrochloric acid (HCℓ) Strong

Carbonic acid (H2CO3) Weak

Citric acid (C6H8O7) Weak

Phosphoric acid (H3PO4) Weak

Sulfuric acid (H2SO3) Strong

Acetic acid* (CH3COOH) Weak *Acetic acid is the important component of vinegar.

Do not confuse the terms strong and weak with more dangerous or less dangerous. The terms “strong” and “weak” in chemistry simply refer to the degree of dissociation. For example, one of the most dangerous acids is hydrofluoric acid which is classified as a weak acid.

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9.2. Bases Bases can be defined as substances that can bind H+. As with acids, bases can be either strong or weak.

Examples of common bases

Ammonia (NH3) Weak

Sodium bicarbonate (NaHCO3) Weak

Sodium carbonate(Na2CO3) Weak

Sodium hydroxide (NaOH) Strong

Sodium hypochlorite (NaOCl) Weak *Sodium hypochlorite is the important component of bleach.

9.3. pH The pH scale is used as a measure of acidity or basicity in water solutions. Acids release H+ in water, so the greater the number of H+ ions present in the solution the more acidic will be the solution. Bases bind H+, therefore the smaller the number of H+ ions present in the solution the more basic will be the solution. In water, the pH scale has a range of 0–14 (although of concentrated solutions it is possible for the range to extend outside of these limits).

Pure water has equal amounts of H+ and OH– ions. The OH– ion is an acceptor of H+ so it is a base:

H+ + OH– → H2O

The middle of the pH range is 7. Therefore a neutral solution has a pH = 7. Acidic solutions will have pH < 7 while basic solutions will have pH > 7.

pH = 7

pH < 7 Acid Stronger Acid pH > 7 Base Stronger Base