Chem Notes

38 Pages of stuff to memoriseIt could be worse

Good Luck~ Kieran~

An Introduction Tuesday, 29 September 2015 6:05 AM

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LIST OF DEFINITIONS IN CHEMISTRY

ELEMENTS, MIXTURES AND COMPOUNDS

1. Boiling point is the temperature at which a pure liquid boils under atmospheric pressure.2. Crystallization is a process involving the recovery of a crystalline solid from its saturated solution by allowing the hot solution to cool to its saturation point.3. Distillation is a process used to recover a solvent from a solution where the solution is boiled and the vapour of the solvent condensed by cooling.4. Fractional distillation is a process used to separate a mixture of miscible liquids using a fractionating column.5. Melting point is the temperature at which a pure solid changes to the liquid state under atmospheric pressure.6. Paper chromatography is a process used to separate a mixture of solutes by their different rates of movement in a solvent over paper.7. A saturated solution is a solution that contains the maximum amount of solute that can dissolve in it at a particular temperature.

ATOMIC STRUCTURE AND BONDING

1. Atomic number (proton number) is the number of protons in an atom.2. A covalent bond is the force of attraction between two atoms of non-metals as a result of their sharing a pair of electrons.3. A covalent compound consists of molecules of non-metals.4. A molecule containing two atoms is called a diatomic molecule.5. An ionic bond is the electrostatic force of attraction between positive and negative ions formed by a transfer of valence electrons from a metal to a non-metal.6. Isotopes are atoms of an element with the same atomic number (number of protons) but different mass number (different number of neutrons).7. A macromolecule is a giant molecule made by large number of atoms covalently bonded together.8. Mass number (nucleon number) is the sum of the number of protons and neutrons in an atom of an element.9. A metallic bond is the electrostatic force of attraction between positive metal ions and negative electrons.10. Radioactive is a state in which an unstable atom decays to form a more stable atom and at the same time releases energy in the form of radiation such as alpha, beta and gamma rays.11. A refractory material is a substance with very high melting point and is heat resistant.12. Relative atomic mass of an element is the number of times one atom of that element is heavier than one-twelfth of an atom of carbon-12.13. The relative atomic mass of an isotope is the weighted average of the accurate masses of all its isotopes.14. Relative molecular mass of an element or compound is the number of times one molecule of it is heavier than one-twelfth of an atom of carbon-12.ACIDS, BASES AND SALTS

1. An acid is a substance that produces hydrogen ions as the only positive ions in water.2. An alkali is a basic hydroxide which is soluble in water.

Keywords Crash CourseSunday, 20 September 2015 3:57 PM

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2. An alkali is a basic hydroxide which is soluble in water.3. A basic oxide is a metallic oxide which reacts with an acid to produce a salt and water only.4. A base is a metal oxide or hydroxide which reacts with an acid to form a salt and water only.5. An acidic oxide is a non-metallic oxide which reacts with water to produce an acid.6. pH is a measure of the concentration of hydrogen ions in solution (elaborate).7. A precipitate is a solid which forms when two solutions are mixed.8. Precipitation is a chemical reaction between solutes in solution, during which an insoluble product is formed.9. A salt is a substance made up of positive metallic ion and negative non-metallic ion from the acid and is formed when the hydrogen ions of an acid are partly or completely replaced by metallic (or ammonium) ions.10. A strong acid is one that is completely or almost completely ionized to give hydrogen ions in aqueous solution.11. A strong alkali is one that is completely or almost completely ionized in aqueous solution.12. Titration is the process whereby a solution from a graduated vessel is added to a known volume of a second solution until the chemical reaction between the two is just completed.13. A weak acid is one that is only partially ionized to give hydrogen ions in aqueous solution.14. A weak alkali is one that is only partially ionized in aqueous solution.

HYDROGEN

1. The Reactivity Series is a list of metals in decreasing order of their tendency to become positive ions.

OXYGEN

1. Corrosion is the gradual destruction of any metal due to reaction with air, water or other chemicals.2. Rusting is the slow corrosion of iron, in the presence of atmospheric oxygen and water, to form brown hydrated iron(III) oxide.3. Pollutants are substances added which has an adverse effect on living things and the environment.

THE PERIODIC TABLE

1. A Group is a vertical column of elements in the Periodic Table where the elements in the same group have the same number of valence electrons.2. A metal is an element which ionizes by electron loss.3. A non-metal is an element which ionizes by electron gain.4. A Period is a horizontal row of elements in the Periodic Table where the elements in the same row have the same number of electron shells occupied by electrons.

STOICHIOMETRY AND THE MOLE CONCEPT

1. Empirical or simplest formula is the formula which shows the simplest whole number ratio of atoms in the formula.2. Molar solution is a solution containing one mole of solute in 1 dm3 of the solution.3. A mole is the amount of substance which contains the same number of particles

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3. A mole is the amount of substance which contains the same number of particles eg molecules, atoms, ions, electrons etc as there are in 12 g of carbon.4. Molecular formula is the formula which shows the actual number of atoms in the formula.

ELECTRICITY AND CHEMISTRY

1. Electrolysis is the chemical decomposition of a compound brought about by a flow of direct current through a solution of the compound or the molten compound.2. An electrolyte is a compound in solution or molten state which will conduct electricity with the decomposition at the electrodes as it does so.3. Electroplating is the electrical precipitation of one metal on another to secure an improved appearance or a greater resistance to corrosion.4. A non-electrolyte is a solution or molten compound which does not conduct electricity or be decomposed by it.

ENERGY CHANGES

1. An endothermic reaction is one in which heat is absorbed from the surrounding.2. An exothermic reaction is one in which heat is liberated to the surrounding.3. Photosynthesis is a process whereby green plants manufacture carbohydrates (glucose) from carbon dioxide and water in the presence of sunlight and chlorophyll as catalyst.

SPEED OF CHEMICAL REACTIONS

1. A catalyst is a substance that speeds up a reaction, but remains chemically and quantitatively unchanged at the end of the reaction.2. An enzyme is an organic catalyst which controls the rate of a biochemical reaction in living organism.3. The rate of a chemical reaction is the amount of reactants or products formed per unit time.

REDOX REACTION

Oxidation is the removal of hydrogen from a substance.Oxidation is the removal of electrons from a substance.Oxidation is an increase in oxidation number.

1. Oxidation is the addition of oxygen to a substance.

2. The oxidation number or state of an element in any molecule or ion is defined as the electrical charge it appears to have as determined by a set of rules.3. A redox reaction is a chemical reaction in which the oxidation state of any element changes (elaborate).

Reduction is the addition of hydrogen to a substance.Reduction is the additional of electrons to a substance.Reduction is a decrease in oxidation number.

4. Reduction is the removal of oxygen from a substance.

NITROGEN

1. Fertilizers are compounds containing plant nutrients which are essential for plant growth.2. Thermal decomposition is an irreversible reaction in which a compound is split up by heat into simpler substances, which do not recombine on cooling.3. Thermal dissociation is a reversible reaction which involves the splitting up of a compound by heat into simpler substances which recombine on cooling.

SULFUR

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SULFUR

1. Allotropes are the different forms of the same element in the same physical state.

METALS

1. An alloy is a uniform mixture prepared by adding other metals or non-metals to a basic metal, so as to obtain desirable qualities.3. Electrolytic reduction is the conversion of metal ion into metal atom by the gain of electron(s) at the cathode, brought about by passing direct current through an electrolyte.2. Recycling means recovering the materials using chemical process for reuse.

ORGANIC CHEMISTRY1. An addition reaction is one which occurs between an unsaturated compound and an attacking reagent, to form a single product.2. An alkane is a saturated hydrocarbon with the general molecular formula CnH2n+

2, where n is an integer.3. An alcohol is an organic compound containing a hydroxyl group with the general molecular formula CnH2n+2O, where n is an integer.4. An alkene is an unsaturated hydrocarbon containing carbon-carbon double bond, with the general formula CnH2n, where n is an integer.5. A carboxylic acid is an organic compound containing a carboxylic acid group with the general molecular formula CnH2nO2, where n is an integer.6. Hydration is the addition of molecules of water to a compound to form a single product.7. A hydrocarbon is a compound containing only carbon and hydrogen.8. Hydrolysis is a decomposition caused by the chemical action of water.9. A saturated compound is a compound which contains only single covalent bond and does not allow other atoms to add onto it.10. The structural formula of a compound is a formula which shows how the atoms are joined in the molecules.11. A substitution reaction involves the direct displacement of an atom or group of atoms e.g. hydrogen from a substance, which is then replaced by other atom of group of atoms of the attacking reagent.12. An unsaturated compound is a compound which contains a double or triple bond between pairs of carbon atoms, allowing other atoms to add across the bond.

PARTICULATE NATURE OF MATTER

1. Diffusion describes the movement of solute particles through a medium, along a diffusion gradient, from a region of higher concentration to a region of lower concentration.2. Sublimation is the process in which a solid changes, upon heating, directly into a gas without going through the liquid state.

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HWA CHONG INSTITUTIONSEC 3 CHEMISTRY NOTE

LIST OF COMMON IONS

Ions is an tom of group of atoms which have become electrically charged due to the gain or loss of electrons.

Positive ions is known as cationNegative ions is known as anion

LIST OF CATIONS

+ 2+ 3+

Lithium Li+ Magnesium Mg2+ Aluminium Al3+

Sodium Na+ Iron(II) Fe2+ Iron(III) Fe3+

Potassium K+ Barium Ba2+ Chromium Cr3+

Copper(I) Cu+ Copper(II) Cu2+

Silver Ag+ Calcium Ca2+

Hydrogen H+ Zinc Zn2+

Ammonium NH4+ Lead Pb2+

LIST OF ANIONS

- 2- 3-

Fluoride F- Oxide O2- Phosphate PO43-

Chloride Cl- Carbonate CO32- Nitride N3-

Bromide Br- Sulfide S2-

Iodide I- Sulfite SO32-

Nitrate NO3- Sulfate SO4

2-

Nitrite NO2- Dichromate Cr2O7

2-

Hydroxide OH-

Hydrogencarbonate(bicarbonate)

HCO3-

Manganate(VII)(permanganate)

MnO4-

IonsFriday, 25 September 2015 1:42 AM

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1 Experimental Chemistry

1.1 Experimental design Candidates should be able to: name appropriate apparatus for the measurement of time, temperature, mass and volume, including burettes, pipettes, measuring cylinders and gas syringes

(a)

suggest suitable apparatus, given relevant information, for a variety of simple experiments, including collection of gases and measurement of rates of reaction

(b)

1.2 Methods of purification and analysis Candidates should be able to:

solid-solid (i)solid-liquid (ii)liquid-liquid (miscible and immiscible) (iii)

use of a suitable solvent, filtration and crystallization or evaporation (i)sublimation (ii)distillation and fractional distillation (iii)use of a separating funnel (iv)paper chromatography (v)

Techniques to be covered for separation and purification include:

describe methods of separation and purification for the components of the following types of mixtures:

(a)

describe paper chromatography and interpret chromatograms including comparison with ‘known’ samples and the use of Rf values

(b)

explain the need to use locating agents in the chromatography of colourless compounds

(c)

deduce from the given melting point and boiling point the identities of substances and their purity

(d)

explain that the measurement of purity in substances used in everyday life, e.g. foodstuffs and drugs, is important

(e)

1.3 Identification of ions and gases Candidates should be able to: describe the use of aqueous sodium hydroxide and aqueous ammonia to identify the following aqueous cations: aluminium, ammonium, calcium, copper(II), iron(II), iron(III), lead(II) and zinc (formulae of complex ions are not required)

(a)

describe tests to identify the following anions: carbonate (by the addition of dilute acid and subsequent use of limewater); chloride (by reaction of an aqueous solution with nitric acid and aqueous silver nitrate); iodide (by reaction of an aqueous solution with nitric acid and aqueous lead(II) nitrate); nitrate (by reduction with aluminium and aqueous sodium hydroxide to ammonia and subsequent use of litmus paper) and sulfate (by reaction of an aqueous solution with nitric acid and aqueous barium nitrate)

(b)

describe tests to identify the following gases: ammonia (using damp red litmus paper); carbon dioxide (using limewater); chlorine (using damp litmus paper); hydrogen (using a burning splint); oxygen (using a glowing splint) and sulfur dioxide (using acidified potassium manganate(VII))

(c)

Stuff to KnowFriday, 25 September 2015 1:43 AM

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2. ATOMIC STRUCTURE AND BONDING

2.1 Atomic Structure Candidates should be able to: identify and describe protons, neutrons and electrons in terms of their relative charges and relative masses

(a)

deduce the behaviour of beams of protons, neutrons and electrons in both electric and magnetic fields

(b)

describe the distribution of mass and charges within an atom (c)define proton (atomic) number and nucleon (mass) number (d)interpret and use such symbols as 12

6C (e)

deduce the numbers of protons, neutrons and electrons present in both atoms and ions given proton and nucleon numbers (and charge)

(f)

(i) describe the contribution of protons and neutrons to atomic nuclei in terms of proton number and nucleon number

(g)

(ii) define the term isotopes and distinguish between isotopes on the basis of different numbers of neutrons present

(writing of electronic configuration will be limited to the first 20 elements)

describe the number and relative energies of the s, p and d orbitals for the principal quantum numbers 1, 2 and 3 and also the 4s and 4p orbitals

(h)

describe the shapes of s and p orbitals (i)state the electronic configuration of atoms and ions given the proton number (and charge)

(j)

2.2 Chemical Bonding Candidates should be able to: describe the formation of ions by electron loss/gain in order to obtain the electronic structure of a noble gas

(a)

describe the formation of ionic (electrovalent) bonds between metals and non-metals as the electrostatic force which holds two oppositely charged ions together., e.g. NaCl, MgCl2 and MgO, including the use of ‘dot and cross’ diagrams

(b)

state that ionic materials contain a giant lattice in which the ions are held by electrostatic attraction, e.g. NaCl (candidates will not be required to draw diagrams of ionic lattices or explain the effect of ionic charge and ionic radii on the numerical magnitude of a lattice energy)

(c)

deduce the formulae of other ionic compounds from diagrams of their lattice structures, limited to binary compounds

(d)

relate the physical properties (including electrical property) of ionic compounds to their lattice structure

(e)

describe the formation of a covalent bond by the sharing of a pair of electrons in order to gain the electronic configuration of a noble gas

(f)

describe, including the use of ‘dot and cross’ diagrams, the formation of covalent bonds between non-metallic elements e.g. H2; O2; N2; Cl2; HCl; CO2; CH4; C2H4 (no knowledge of co-ordinate/dative bonds or hybridisation of orbitals will be expected)

(g)

deduce the arrangement of electrons in other covalent molecules (h)describe covalent bonding in terms of orbital overlap, giving σ and π bonds (the concept of bond angles in molecules analogous to those specified in (g) is notnecessary)

(i)

describe hydrogen bonding, using ammonia and water as examples of molecules containing –NH and –OH groups

(j)

outline the importance of hydrogen bonding to the physical properties including ice and water

(k)

explain the terms bond energy, bond length and bond polarity and use them to compare the reactivities of covalent bonds

(l)

describe intermolecular forces (van der Waals’ forces) (the concepts of permanent and induced dipoles are not required)

(m)

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and induced dipoles are not required) relate the physical properties (including electrical property) of covalent substances to their structure and bonding

(n)

describe metallic bonding in terms of a lattice of positive ions surrounded by mobile electrons (‘sea of electrons’)

(o)

relate the electrical conductivity of metals to the mobility of the electrons in the structure

(p)

describe, interpret and/or predict the effect of different types of bonding (ionic bonding; covalent bonding and metallic bonding) on the chemical and physical properties of substances

(q)

show understanding of chemical reactions in terms of energy transfer associated with the breaking and making of chemical bonds

(r)

compare the structure of simple molecular substances, e.g. methane, iodine, with those of giant molecular substances, e.g. poly(ethene); sand (silicon dioxide); diamond; graphite in order to deduce their properties

(s)

(t)) compare the bonding and structures of diamond and graphite in order to deduce their properties such as electrical conductivity, lubricating or cutting action (candidates will not be required to draw the structures) suggest from quoted physical data the type of structure and bonding present in a substance

(a)

deduce the physical and chemical properties of substances from their structures and bonding

(b)

3. ACIDS, BASES AND SALTS

3.1 The characteristic properties of acids and bases Candidates should be able to: state the properties of acids in reactions with metals (the ability of metals to react with acids should be linked to their position in the reactivity series), bases and carbonates

(a)

state the uses of sulfuric acid as in the manufacture of detergents and fertilisers; and as a battery acid

(b)

describe the characteristic properties of bases in reactions with acids and with ammonium salts

(c)

describe importance of water for acidity, i.e., water causes acid molecules to ionise and form hydrogen ions

(d)

define acid as a substance that produces hydrogen ions as the only positive ions in water

(e)

explain basicity of common acids and relate to concentration of hydrogen ions (f)describe qualitatively the difference between strong and weak acids in terms of extent of ionisation of acid in water

(g)

describe an alkali as a basic hydroxide which is soluble in water to produce hydroxide ions

(h)

describe the reaction between hydrogen ions and hydroxide ions to produce water, H+ + OH- → H2O as neutralisation

(i)

describe aqueous ammonia as a weak alkali (j)identify aqueous sodium hydroxide and potassium hydroxide as strong alkalis (k)describe how to test hydrogen ion concentration and hence, relative acidity using Universal indicator and the pH scale

(l)

describe the importance of controlling pH in soil and how excess acidity can be treated using calcium hydroxide

(m)

classify oxides as either acidic, basic, amphoteric or neutral based on metallic / non-metallic character

(n)

classify sulfur dioxide as an acidic oxide and state its uses as a bleach, in the manufacture of wood pulp for paper and as a food preservative (by killing bacteria)

(o)

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3.2 Preparation of salts Candidates should be able to: define a salt as a substance formed when the hydrogen ions of an acid are partly or completely replaced by metallic or ammonium ions

(a)

describe the general rules of solubility for common salts to include nitrates, chlorides (including silver and lead), sulfates (including barium, calcium and lead), carbonates, hydroxides, Group I cations and ammonium salts

(b)

(methods for preparation should include precipitation and titration together with reactions of acids with metals, insoluble bases and insoluble carbonates)

describe the techniques used in the preparation, separation and purification of salts as examples of the techniques specified in section 1.2(a)

(c)

suggest a method of preparing a given salt from suitable starting materials, given appropriate information

(d)

4. PERIODIC TRENDS

4.1 Periodic Trends Candidates should be able to: describe the Periodic Table as an arrangement of the elements in the order of increasing proton (atomic) number

(a)

describe how the position of an element in the Periodic Table is related to proton number and electronic structure

(b)

describe the relationship between Group number and the ionic charge of an element (c)explain the similarities between the elements in the same Group of the Periodic Table in terms of their electronic structure

(d)

describe the change from metallic to non-metallic character from left to right across a Period of the Periodic Table

(e)

predict the properties of elements in Group I, VII and the Transition elements using the Periodic Table

(f)

describe the relationship between Group number, number of valence electrons and metallic/non-metallic character.

(g)

4.2 Group properties Candidates should be able to:

melting point and in their reaction with water

describe lithium, sodium and potassium in Group I (the alkali metals) as a collection of relatively soft, low density metals showing a trend in

(a)

describe chlorine, bromine and iodine in Group VII (the halogens) as a collection of diatomic non-metals showing a trend in colour, state and their displacement reactions with solutions of other halide ions

(b)

describe the elements in Group 0 (the noble gases) as a collection of monatomic elements that are chemically unreactive and hence important in providing an inert atmosphere, e.g. argon and neon in light bulbs; helium in balloons; argon in the manufacture of steel

(c)

describe the lack of reactivity of the noble gases in terms of their electronic structure (d)

4.3 Transition elements

describe the central block of elements (transition metals) as metals having high melting points, high densities, variable oxidation state and forming coloured compounds

(a)

e.g. iron in the Haber process; vanadium(V) oxide in the Contact process; nickel in the hydrogenation of alkenes

state the uses of these elements and/or their compounds as catalysts, (b)

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5. MOLE CONCEPT / STOICHIOMETRY

Candidates should be able to

state the symbols of the elements and formulae of the compounds mentioned in the syllabus

(a)

deduce the formulae of simple compounds from the relative numbers of atoms present and vice versa

(b)

determine the formulae of ionic compounds from the charges on the ions present and vice versa

(c)

interpret chemical equations with state symbols (d)write and/or construct chemical equations, with state symbols, including ionic equations

(e)

define the terms relative atomic, isotopic, molecular and formula masses, based on the 12C scale

(f)

define the term mole in terms of the Avogadro constant (g)analyse mass spectra in terms of isotopic abundances and molecular fragments [knowledge of the working of the mass spectrometer is not required]

(h)

calculate the relative atomic mass of an element given the relative abundances of its isotopes, or its mass spectrum

(i)

calculate the percentage mass of an element in a compound when given appropriate information

(j)

define the terms empirical and molecular formulae(k)calculate empirical formulae and molecular formulae from relevant data (using combustion data or composition by mass)

(l)

calculate stoichiometric reacting mass (from formulae and equations) and volumes of gases (e.g. in the burning of hydrocarbons) (one mole occupies 24 dm3 at room temperature and pressure); calculations involving the idea of limiting reactants may be set (The gas laws and the calculations of gaseous volumes at different temperatures and pressures are not required)

(m)

deduce stoichiometric relationships from calculations such as those in (m) (n)apply the concept of solution concentrations (in expressed in mol/dm3 or g/dm3) to process the results of volumetric experiments and to solve simple problems (appropriate guidance will be provided where unfamiliar reactions are involved)

(o)

calculate % yield and % purity (p)

6. ATMOSPHERE AND ENVIRONMENT

Candidates should be able to:

describe the percentage composition of clean air in terms of 79% nitrogen, 20% oxygen, and the remainder being noble gases (with argon as main constituent) and carbon dioxide.

(a)

name some common atmospheric pollutants (carbon monoxide; methane; nitrogen oxides (NO and NO2); ozone; sulfur dioxide; unburnt hydrocarbons)

(b)

carbon monoxide from the incomplete combustion of carboncontaining substances

(i)

nitrogen oxides from lightning activity and internal combustion engines. (ii)sulfur dioxide from volcanoes and combustion of fossil fuels. (iii)

state sources of these pollutants as (c)

redox reactions in catalytic converters to remove combustion pollutants, (i)the use of calcium carbonate to reduce the effects of 'acid rain and in flue gas desulfurisation.

(ii)

describe the reactions used in possible solutions to the problems arising from some of the pollutants in (b)

(d)

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desulfurisation.

the poisonous nature of carbon monoxide (i)the role of nitrogen dioxide and sulfur dioxide in the formation of acid rain and its effects on respiration, buildings and plants.

(ii)

the role of nitrogen dioxide, methane and unburnt hydrocarbons in the formation of photochemical smog.

(iii)

discuss some of the effects of these pollutants on health and on the environment (e)

describe the importance of the ozone layer and the problems involved with the depletion of ozone by reaction with chlorine containing compounds, chlorofluorocarbons (CFCs).

(f)

the process of combustion, respiration and photosynthesis (i)how carbon cycle regulates the amount of carbon dioxide in the atmosphere. (ii)

describe the carbon cycle in simple terms to include (g)

state that carbon dioxide and methane are greenhouse gases and may contribute to global warming, give the sources of these gases and discuss the possible consequences of an increase in global warming.

(h)

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Periodic TableFriday, 25 September 2015 6:12 AM

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Pure Substances is a single substance not mixed with anything else. A mixture consist of 2 or more substances that are not chemically combined together.Pure Substances have fixed melting and boiling points. The melting and boiling points of substances are also unique, and a pure substance can be identified via its melting and boiling points.

The substance has a lower melting point. The greater the amount of impurity, the lower the melting point.1)The substance melts over a range of temperatures.2)The substance has a higher boiling point. The greater the amount of impurities, the higher the boiling point.3)The liquid boils over a range of temperatures.4)

Impurities in substances can be identified in several ways:

Decanting1)A pure substance can be obtained through the use of various separation techniques.

Filtration2)Decanting separates an insoluble solid from a liquid by pouring off that liquid from the container.

Evaporation3)

Filtration is used to separate an insoluble solid from a liquid. The mixture is poured into a filter paper, A liquid passes through the small holes in the filter paper. The insoluble solid cannot pass through and is trapped in the filter paper. The solid is the residue and the liquid is the filtrate.

Crystallisation4)

Evaporation is used to separate dissolved salts from a solution. During evaporation, the solution is heated and water changes into steam. The salt left behind is the residue

Sublimation5)

Crystallisation separates a dissolved solid from a solution, forming pure crystals of the substance. Crystallisation is commonly done by heating a solution to evaporate off the solvent until a saturate solution is obtained. As it cools, pure crystals of the dissolved solid form. Impurities remain in the solution. This occurs because the solubility of solutes decreases as the temperature decreases. As a hot solution cools, it eventually becomes saturated, and can it can hold no more solute. The extra solute that cannot be dissolved separates as pure crystals. Impurities remain in the solution. Solids can be melted and allowed to cool to obtain pure crystals.

Simple Distillation6)

Sublimation separates a mixture of solids, one of which sublimes. Some substance can undergo sublimation, as they change directly from a solid to a vapour upon heating. The vapour changes back to solid directly on a cold surface.

Fractional Distillation7)

Simple Distillation separates a pure liquid from a solution. The liquid is changed into a vapour by boiling, and the vapour is pure as other substances are left behind. The vapour is then cooled and condenses to a pure liquid which is called the distillate.

Fractional Distillation separates mixtures of miscible liquids with widely different boiling points. Miscible liquids are completely soluble in each other. Fractional Distillation enlists the help of a fractionating column to help separate the liquids. Liquids with a lower boiling point boils first and is distilled, thus the temperature would remain constant when the liquid is being distilled. The temperature only rises when the liquid is mostly distilled, and then remains constant when it reaches the boiling point of the next liquid to be distilled.

Pure SubstancesSunday, 20 September 2015 3:57 PM

Experimental Chemistry Page 14



Chromatography is a method of separating and identifying mixtures.

To separate and identify mixtures of coloured substances found in foods and in dyes-

To separate and identify substances in drugs and in blood-

To separate and identify substances in urine to find out if athletes have been using drugs-

Chromatography has a variety of uses, such as:

Get a filter paper and draw a line with pencil near the bottom of the filter paper. Pen is not used because the ink would separate during chromatography. Hang the filter paper on something.

1)

A solution of whatever you're testing is obtained, and a drop of the solution is placed on the pencil line near the bottom of a strip of filter paper.

2)

The paper is dipped into a suitable solvent, with the solvent level below the spot. This is important as if the solution is in the solvent, it may dissolve directly into the solvent and taint it, affecting the chromatography results.

3)

The solvent travels up the paper. The solutions on the pencil line dissolve in the solvent and travel up the paper at different speeds, and hence the dyes are separated. The result is called a chromatogram.

4)

The process of chromatography goes like so:

The number of dots above the pencil line tell you how many different substances where in the solution you tested. If two solutions are tested and they have dots that stop at the same distance above the pencil line, they are probably the same.When the solution is colourless, the chromatogram is sprayed with a locating agent. A locating agent is a substance that reacts with the substances on the paper to produce a coloured product.

Substances on a chromatogram have a Rf value, which is derived by:

The distance moved by the substance divided by the distance moved by the solvent

The Rf value differs with the solvent and the temperature.

ChromatographyFriday, 25 September 2015 2:51 AM


Cations and Anions and GasesFriday, 25 September 2015 4:09 AM



An element is a substance that cannot be broken into simpler substances by chemical methods. If it can be broken down, it is a compound. Heat or electricity can help break down a compound. This is called decomposition. These elements have a name and a chemical symbol, and can be found on the periodic table. 118 elements have been discovered at the moment.

Elements can be classified in 2 ways. The first is by state. There are 92 naturally occurring elements that can be found on earth. 11 are gases, 2 are liquids, and 79 are solids. They can also be classified by how well they conduct electricity, and can be sorted into 3 groups: Metals, Metalloids and Non-metals.

Metals1)Good Conductors of Electricity-

Most known elements are metals-

Metalloids2)Properties between those of metals and non-metals-

Non-metals3)Poor conductors of Electricity-

Only 22 of the known elements are non-metals-

Some are Noble Gases-

All Noble Gases are colourless and unreactive-

An atom is the smallest unit of an element, having the properties of that element. A molecule is a group of two or more atoms chemically combined together. Most non-metals are made up of molecules. If all the atoms in the molecule are the same kind, it is an element. Symbols are used to represent atoms, and a chemical formula is used to represent a molecule. The chemical formula of a molecule shows the number and kinds of atoms contained in it. There are different types of molecules with different names, based on the number of atoms in it.

1 Atom: Monatomic2 Atoms: Diatomic3 Atoms: Triatomic4 or More Atoms: Polyatomic

A compound is a substance containing two or more elements chemically combined together. Compounds formed have different properties from their constituent elements. Heat is usually given out when elements combine to form compounds. There are some general rules that are followed when naming a compound.

If a compound contains both a metal and non-metal, the metal usually comes first in the formula

-

A compound with only two elements often ends with -ide-

A compound that contains OH is a hydroxide-

The name of a compound with an ion containing oxygen usually ends in -ate-

Compounds are made of 2 different particles: Molecules and Ions.We already know about Molecules, so let's talk about ions.An atom or a group of atoms that has an electrical charge is called an ion. Compounds made of ions are called ionic compounds. Ions that lose an electron have a positive charge, and are called cations. Ions that gain an electron have a negative charge, and are called anions.

Compounds and mixtures are notably different. A mixture consists of 2 or more substances that are not chemically combined together. Here are some properties of Compounds and Mixtures.

Elements, Compounds and MixturesSunday, 20 September 2015 3:57 PM

Atomic Structure and Bonding Page 20

not chemically combined together. Here are some properties of Compounds and Mixtures.

Has a fixed composition by mass-

Has fixed melting and boiling points-

Has its own physical and chemical properties which are different from its elements-

Cannot be separated into two or more substances by physical means; a chemical reaction is needed to separate the elements.

-

Compound

Has a variable composition by mass-

Has variable melting and boiling points-

Does not have its own properties, and has the same properties as its components-

Can be easily separated into its components by physical means without a chemical reaction-

Mixture


Protons carry a positive electric charge, and have a relative mass of 1-

Neutrons also have a relative mass of 1, but does not have a relative charge-

Electrons carry a negative electric charge and a negligible mass-

An atom consists of the following: Protons, Neutrons and Electrons.

A cloud of electrons that move at nearly the speed of light surrounds a nucleus containing protons and neutrons. Almost all mass of an atom is contained in its nucleus. Most of the atom is made of empty space. In normal atoms, the number of protons is equal to the number of electrons, so the atom is electronically neutral. As we move down the periodic table, that number of protons and electrons are increased by 1. However, there is no pattern in the number of neutrons in an atom. The number of protons in an atom is called the proton number. It is also called the atomic number. Nucleon number is the number of protons and neutrons in the nucleus of an atom. It is also called the mass number.

Number of Protons = Proton Number/Atomic NumberNumber of Neutrons= Nucleon Number - Proton NumberNumber of Electrons= Number of Protons

Isotopes are atoms with the same number of protons but different number of neutrons. Atoms of an Element always have the same number of protons. Isotopes have the same chemical properties and form compounds with the same formula. However, there are small differences in their physical properties, such as differing masses.

Electrons in an atom are arranged in orbits around the nucleus of the atom. The orbits are also called shells. These shells are numbered 1,2,3,4 and so on, outwards from the nucleus. Each shell can hold a certain maximum number of electrons. The first shell is closest to the nucleus, can hold a maximum of 2 electrons and is always filled first. The second shell can hold a maximum of 8 electrons, and is usually filled before the third shell fills, which can also hold a maximum of 8 electrons. The shell which is farthest from the nucleus is called the outer shell. It is also called the valence shell. The electrons in this shell are known as outer electrons or valence electrons. Most of the time, only the outer electron is drawn in the electronic structure, and this is called the outer electronic structure.

Noble Gas Atoms have the most stable electronic structures, thus atoms of other elements try to achieve this electronic structure, also known as a noble gas configuration. They can either form ions or share outer electrons to achieve this. When atoms share electrons, molecules are formed.

In an atom, the number of protons is the same as the number of electrons, and thus the atom is electronically neutral. When electrons are added or taken away from an atom, the atom is no longer electrically neutral, and becomes a charged particle, also known as an ion.

Metal Ions have few outer shell electrons, and lose electrons in their outer shell to obtain a noble gas configuration. They form positive ions, because they lose electrons. Non-metals have more outer shell electrons then metals, and gain electrons to obtain a noble gas configuration, The form negative ion, because they gain electrons.

Atomic StructureFriday, 25 September 2015 5:42 AM


Pd s block 3d/4d blocks of Transition Metals (Periods 4/5), the 1st/10th are NOT true transition elements, they have

no partially filled d shell in an ion.

p block elements

Gp1 Gp2 Gp3/13 Gp4/14 Gp5/15 Gp6/16 Gp7/17 Gp0/18

1 1H 1s12He 1s2

2 3Li[He]2s1

4Be[He]2s2

The electronic structure of Elements 1 to 56, ZSymbol, Z = atomic or proton number = total

electrons in neutral atom, [He] = 1s2, [Ne] = 1s22s22p6, [Ar] = 1s22s22p63s23p6, [Kr] =

1s22s22p63s23p63d104s24p6

Between Groups 2 and 3 (13) are the d–blocks and f–blocks where the quantum energy level rules permit

their inclusion and electron filling. Periods 4 and 5 have 18 elements each, including the 3d and 4d

blocks of elements respectively (Groups 3 to 12 –new notation).

5B[He]2s2

2p16C[He]2s2

2p27N[He]2s2

2p38O[He]2s2

2p49F[He]2s2

2p510Ne[He]2

s22p6

3 11Na[Ne]3s1

12Mg[Ne]3s2

13Al[Ne]3s23p1

14Si[Ne]3s23p2

15P[Ne]3s23p3

16S[Ne]3s23p4

17Cl[Ne]3s23p5

18Ar[Ne]3s23p6

4 19K[Ar]4s1

20Ca[Ar]4s2

21Sc[Ar] 3d14s222Ti [Ar] 3d24s2

23V [Ar] 3d34s2

24Cr [Ar] 3d54s1

25Mn[Ar] 3d54s2

26Fe [Ar] 3d64s2

27Co[Ar] 3d74s2

28Ni [Ar] 3d84s2

29Cu [Ar] 3d104s1

30Zn [Ar] 3d104s2

31Ga [Ar] 3d104s24p

1

32Ge [Ar] 3d104s24p

2

33As [Ar] 3d104s24p

3

34Se [Ar] 3d104s24p

4

35Br [Ar] 3d104s24p

5

36Kr [Ar] 3d104s24p

6

5 37Rb[Kr]5s1

38Sr[Kr]5s2

39Y [Kr] 4d15s240Zr [Kr] 4d25s2

41Nb[Kr] 4d45s1

42Mo[Kr] 4d55s1

43Tc [Kr] 4d55s2

44Ru[Kr] 4d75s1

45Rh[Kr] 4d85s1

46Pd[Kr] 4d10

47Ag [Kr] 4d105s1

48Cd [Kr] 4d105s2

49In [Kr] 4d105s25p

1

50Sn [Kr] 4d105s25p

2

51Sb [Kr] 4d105s25p

3

52Te [Kr] 4d105s25p

4

53I [Kr] 4d105s25p

5

54Xe [Kr] 4d105s25p

6

6 55Cs[Xe]6s1

56Ba[Xe]6s2

4f–block (14) and 5d–block (10) 32 elements in period 6 including the Lanthanide Series of

Metals.

81Tl [Xe] 5d106s26p

1

82Pb [Xe] 5d106s26p

2

83Bi [Xe] 5d106s26p

3

84Po [Xe] 5d106s26p

4

85At [Xe] 5d106s26p

5

86Rn [Xe] 5d106s26p

6

7 87Fr[Rn]7s1

88Ra[Rn]7s2

5f–block and 6d–block including the Actinide Series of Metals.

Energy level filling order up to Z = 56 is 1s 2s 2p 3s 3p 4s 3d 4p (for Z = 1 to 36) then 5s 4d 5p 6s 4f/5d (for Z = 37 to 56)However, when writing out the electron configuration you must write them out in order of strict principal quantum with the accompanying s, p, d, f notation i.e. the order 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 6s (upto Z = 58)The s–block consists of Groups 1 and 2 where the only outer electrons are in an s sub –energy level orbital (no outer p electrons, 2 per period).•The p–block corresponds to Groups 3 to 0 (old notation) where the three p sub –energy level orbitals are being filled (6 per period).•

Therefore Sc to Zn form the head elements of Groups 3 to 12 using the 'new' group number notation.○

Similarly on period 5 there is a 4d block where the 4d sub–shell level is filled.○

So 10 d block elements per period are now permitted\under the quantum number rules.○

Starting with period 4, where the first of the d sub–shells is low enough in energy to be filled, there are ten elements 'inserted' between groups 2 and 3, the so –called d blocks of ten elements (the 1st block, the 3d block Sc–Zn is on Period 4).

•

Starting with cerium (Z=58, period 6), see in full table below, there is a further insertion of fourteen elements where the s even f–orbital sub–shell is being filled after the first of the d–block metals and similarly with thorium (Z=90) in period 7 and these are known as the f blocks (14 per period where permitted).

•

There are differing number of atomic orbitals in the various subshells. You need 2 electrons to fill up 1 atomic orbital. The re is 1 S Orbital, 3 P Orbitals, 5 D orbitals and 7 F Orbitals. Every orbital is first filled with 1 electron, before it is f illed up. Once it is filled up, then the next subshell can be filled with electrons. Useful Site: http://www.docbrown.info/page07/ASA2ptable2a.htm

SPDFFriday, 25 September 2015 6:17 AM


http://www.docbrown.info/page07/ASA2ptable2a.htm

Ionic Bonding, which results when electrons are transferred from one atom to another, forming positive and negative ions.

1.

Covalent Bonding, which results when atoms are joined together by sharing electrons, forming molecules.

2.

Metallic Bonding, which is only found in metals.3.

There are 3 main ways of forming chemical bonds between atoms.

Ionic Bonding occurs when atoms gain or lose electrons to form ions. Atoms of metal lose electrons to form positive ions, and atoms of non-metals gain electrons to form negative ions. An ionic bond is the force of attraction between oppositely charged ions in a compound. When drawing a dot and cross diagram, you generally only draw the outermost shell.

In an ionic compound, there are no definitive pairs, and instead, there is a repetition of positive and negative ions throughout the whole structure. Due to the large number of ions in its structure, it is called a giant ionic structure or a giant lattice structure, as it forms a giant ionic crystal lattice. All structures formed formed from metal and non-metals have giant lattice structures. The metal or ammonium ion is always written first in the name of the compound and in the formula. If a compound has more than one polyatomic atom, the ions is placed in brackets. Polyatomic Ions consists of more than one atom. The total charge is the total number of protons minus the number of electrons. To balance out an ionic bond, simply cross multiply the charges. For example, Calcium ions have a charge of 2+, and Phosphate ions have a charge of 3-, so you would need 3 calcium ions and 2 phosphate ions to balance out the equation.

An Ionic Compound forms an Ionic Crystal Lattice that contains a network of ions arranged in a regular pattern. The ions are held together by strong electrostatic forces of attraction. Ionic compound are hard, crystalline solids that have flat sides and regular shapes, and are hard due to the strong ionic bonds. They have high melting and boiling points, To melt an ionic compound, a large amount of heat is needed to break the strong electrostatic forces holding the ions together. Therefore, ionic compounds have high melting points. To boil an ionic compound, a large amount of heat is needed to overcome the strong electrostatic forces of attraction holding the ions in the liquid. An ionic compound cannot evaporate easily due to the strong forces holding the ions together. Solid ionic compounds do not conduct electricity as ions cannot move in the solid state as they are held in place, so the current cannot flow. However, when in molten state or dissolved in water, the ion can move and can act as electrical charge carriers and carry the electric current. Many ionic compounds are soluble in water, as there is an attraction between the ions and the water molecules, causing the ions to separate and go into the solution. However, most ionic compounds are insoluble in non-aqueous solvent.

A covalent bond is formed by sharing a pair of electrons. Covalent bonds are formed between atoms of non-metals. This bond can be formed between atoms of the same elements, or between atoms of different elements. Covalent bonds occur in substances consisting of molecules, and are called covalent compounds. Dot and cross diagrams of covalent compounds generally only require you to draw the outermost shells. Valence electrons are electrons in the outer most shell which are used to form bonds between atoms. Valency is the number of electrons an atom uses to form bonds between atoms.

Chemical BondingFriday, 25 September 2015 6:41 AM


Elements and compounds consisting of small molecules have simple molecular structures. In the molecules, the atoms are held together by strong covalent bonds However, between the molecules, there are only weak forces of attraction called intermolecular forces. The intermolecular forces are weak at liquid state. These forces are broken and reformed when the molecules slide past one another. When the liquid is heated, heat energy is converted to kinetic energy, allowing the molecules to overcome the intermolecular forces holding them together, which cause the liquid to turn into gas.

A sigma bond is a bond where electron density is concentrated n he region between the nuclei. It is a bond that lies along the internuclear axis. All single bonds in molecules are sigma bonds. When two atomic orbitals, each containing a single unpaired electron, approach each other and overlap, the two electron can occupy a molecular orbital which is at a lower energy. When the electrons exist between the two atoms, the electron is attracted by both nuclei which is a much more stable situation than when the atoms existed separately. When two s-orbitals in H atom, each with a single electron, approach and overlap, a new molecular orbital called the sigma orbital is formed, and the bond is a sigma bond.

A pi bond is a result of side to side overlap of P orbitals. Electron density lies above and below the internuclear axis. Before a pi bond is formed, a sigma bond must be present. All molecules with a double bond will have one sigma bond and one pi bond, and all molecules with a triple bond will


double bond will have one sigma bond and one pi bond, and all molecules with a triple bond will have one sigma bond and 2 pi bonds.

Most substances with simple molecular structures are liquids or gases at room temperature. As the forces between the molecules are weak, the molecules are not held together tightly like particles in a solid, but are free to move. Substances with a simple molecular structure usually have low melting and boiling points. Only the weak intermolecular forces between the molecules are overcome when the substance boils. The covalent bonds that hold the atoms together are strong, the van der waals forces of attraction between molecules are weak and easy to overcome. Covalent compounds are volatile and evaporate easily. Substances with a simple molecular structure do not conduct electricity. To conduct electricity, substances must contain either ions or electrons that are free to move. Molecules do not contain either. Most molecular substances are insoluble in water. Instead, they dissolve in organic solvents.

Metal atoms are packed together closely in regular three-dimensional patterns to form giant lattice structures. A metallic bond is the attractive force between positively charged ions and negatively charged, free or mobile. Each atom gives up its outer shell electrons to become a positive ion. These electrons move and occupy the spaces between the ions. The moving electrons are called mobile electrons or delocalised electrons, and the metal is often described as consisting of positive ions in a sea of electrons.

Metals are malleable, which means they can be bent. Metals are also ductile, which means they can be stretched. This is because layers of atoms in a metal can slide over each other easily when a force is applied. The metal does not break as the sea of electrons hold the atoms in the metal together. The metallic bonding in most metals is strong and hence a lot of energy is needed to separate the atoms, resulting in a high melting and boiling point. Electricity conducts electricity, due to the movement of the free or mobile electrons through the metal, and acts as electrical charge carriers. Metal conducts electricity both when solid and when molten. Metals conduct heat, and they allow heat to pass through them easily. When one end of a metal is heated, the delocalised electrons get more energy. They move faster, colliding with neighbouring electrons. Heat is transferred in the collisions, causing the whole piece of metal thus becomes hot.

Some substances have giant structures consisting of atoms joined together by covalent bonds. These are called giant molecular structures.Diamond is a form of the element carbon. In diamond, each carbon atom forms covalent bonds with four other carbon atoms. The structure is a giant network of carbon atoms held together by covalent bonds. The strong covalent bonds throughout diamond make it the hardest natural substance. Because of its hardness, diamond is used in cutting other hard solids. Hardness is measured on a scale called the Mohs scale. The scale ranges from 1 to 10, diamond is a 10.


Silicon Dioxide is a compound of silicon and oxygen, and has a giant molecular structure like diamond. Each silicon atom is joint by covalent bonds to four oxygen atoms. Each oxygen atoms is joined to two silicon atoms. In the structure, there are no separate molecules.

All substances with a giant molecular structure are solids at room temperature. The solid structure has many strong covalent bonds between the atoms, which make it quite hard. Substances with a giant molecular structure have high melting and boiling points. This is because strong covalent bonds must be broken to melt and boil them. Diamond melts at about 3500 degrees Celsius. Substances with a giant molecular structure do not conduct electricity, with the exception of graphite. For electricity to be conducted, there must be movement of free mobile electrons. All the outer shell electrons in the atoms of covalent structures are used to form covalent bonds. Thus, there are no free mobile electrons and conduction of electricity does not occur. Substances with a giant molecular structure are insoluble in any solvent.

Graphite is another form of the element carbon and also has a giant molecular structure. However, this structure is different from that of diamond or silicon dioxide. Graphite has a macromolecular structure. In graphite, the carbon atoms are arranged in flat layers. The carbon atoms in each layer are arranged in hexagonal rings of six atoms, which each atom joined to others by strong covalent bonds. The covalent bonds between the carbon atoms are strong. However, the intermolecular forces between the layers are weak. For every three covalent carbon bonds, one electron is delocalised over the whole sheet of atoms in one layer, and is free to move between these layers of carbon. The mobile electron can be used as an electrical charge carrier, and thus allows graphite to conduct electricity. Graphite is also soft, as the forces between the layers of carbon atoms are weak and so they can easily slide past each other.


In a molecule, atoms with the higher electronegativity has more electron density around it, and thus has a partial negative charge. The other atom has less electron density, and thus has a partial negative charge. The centres of positive charge and negative charge do not coincide. This charge separation is called dipole. Such bonds between atoms with differing electronegativity are polar. Bond polarity increases as the electronegativity difference increases. A dipole moment is the moment that forces a set of separate charges to rotate in an electric field. The dipole moment is represented by an arrow with the tail at the positive centre and the head at the negative centre. A molecule with an electric dipole is said to be polar and to possess a dipole. A bond with a dipole moment gives rise to bond polarity. A polar molecule has positive and negative poles, areas where there is a partial charge along the molecule. Polar molecules orient themselves in an electric field, with the negative ends towards the positive end of the field and vice versa. A molecule is polar if there is a net charge separation between two ends of the molecule or net dipole. If there are no polar bonds and the bond dipoles cancels, the molecule is non -polar. If the bond dipole does not cancel, the molecule is polar and has a partial charge. The balance of the attractive and repulsive forces occurs when the nuclei are separated by a distance. The distance is known as the bond length. Energy is released as the two atoms come together to form the bond, and the same amount of energy required to break the bond. This is the bond energy.

Bond strength increases as bond multiplicity increases. 1)Bond strength increases as orbital overlap increases. 2)Bond strength increases as differences in electronegativity increases. 3)

Bond PolarityMonday, 28 September 2015 2:38 AM


Electrostatic attraction between a hydrogen bonded to nitrogen, oxygen or fluorine and the lone pair of a neighbouring nitrogen, oxygen or fluorine atom is called a hydrogen bond. The hydrogen is attached directly to one of the most electronegative elements (N,O,F), causing the hydrogen to acquire a significant amount of positive charge. Each of the elements to which the hydrogen is attaches also has at least one "active" lone pair. H-X has a very polar bond due to the large electronegativity difference. When X is bonded to H, they pull the bonding pair well away from H, making H positive and poorly shielded. Lone pairs have the electrons contained in a relatively small volume of space which therefore has a high density of negative charge. Hydrogen bonded molecules have higher than expected melting points and boiling points compared with the hydrides of other elements in their respective Group.

Hydrogen BondingMonday, 28 September 2015 3:09 AM


Acids have a sour taste1)Acids are hazardous2)Acids turn blue litmus paper red3)Acids react with metals to produce hydrogen gas and a salt4)Acids react with carbonates and hydrogencarbonates to produce carbon dioxide and water and a salt5)Acids react with metal oxides and hydroxides to form water and a salt6)

Properties of Acid

Hydrogen ions need to be present in acid solutions for the acids to have acidic properties. Pure acids without water consists of small covalent molecules. In the presence of water, `the acid molecules form ions, and the acid is ionised. The properties and reactions of acids are due to the hydrogen ions. Therefore, acids only behave as acids when they are dissolved in water. Thus, and acid is a substance that produces hydrogen ions when dissolved in water.

The basicity of the acid is the maximum number of hydrogen ions produced by a molecule of an acid. In a solution of a strong acid, all the acid molecules become ions in the water. In a solution of a weak acid, most of the acid molecules remain unchanged in the water, and few acid molecules are ionised to become hydrogen ions. A strong acid is one that completely ionises in water. A weak acid is one that partially ionises in water. Strong acids react more vigorously than weak acids. Strong acids and concentrated acids are not the same. Similarly, weak acids are not dilute acids. Strong and weak refer to the extent of ionisation, while concentration refer to the amount of solute in a solution.

Bases are oxides or hydroxides of metals. Soluble bases are called alkalis. An alkali is a substance that produces hydroxide ions in water. All alkalis are bases, but not all bases are alkalis.

Alkalis feel slippery1)Alkalis are caustic and hazardous2)Alkalis change red litmus paper blue3)Alkalis react with acids to form a salt and water4)Alkalis react with ammonium compounds to form ammonia gas, a salt and water5)Alkalis react with solutions of metal ions to form a precipitate6)

Properties of Bases

Alkalis can also be weak or strong. When strong alkalis are added to water, they form hydroxide ions in solution. An alkali that only has a small fraction of its molecules react with water to form hydroxide ions are known as weak alkalis.

Acids and BasesSunday, 20 September 2015 3:57 PM

Acid, Bases and Salts Page 30

An indicator is a substance that has different colours in acidic and alkaline solutions.

pH is a measure of the acidity or alkalinity of a solution. All aqueous solutions contain both hydrogen ions and hydroxide ions. The concentration of these ions differs in solutions of different pH. The pH is a measure of the concentration of the hydrogen ions in a solution. A solution with pH of less than 7 is acidic. Acids have a greater concentration of hydrogen ions. The lower the pH level, the greater the concentration of hydrogen ions and the more acidic the solution. A pH of 7 is neutral. This is the pH of pure water. The concentration of hydrogen ions and hydroxide ions is the same. A solution with pH greater than 7 is alkaline. Alkalis have a greater concentration of hydroxide ions. The higher the pH, the greater the concentration of hydroxide ions and the more alkaline the solution.

An ionic equation is an equation involving ions in an aqueous solution, only the ions formed or changed during the reaction are included. In ionic equations, formulae of ions that change are included. Ions that do not change are spectator ions and are not included. Formulae of solids, liquids and gases are written in full.

Indicators and pHMonday, 28 September 2015 4:16 AM


The reaction of an acid and an alkali is called neutralisation. During neutralisation, the properties of the acid and the base are taken away or destroyed, thus an acid and a base cancel out each other's properties. An acid-alkali neutralisation produces a salt and water. Neutralisation is the reaction between an acid and a base to form a salt and water only.

Acidic Oxides 1)Acidic Oxides are the oxides of non-metals. They are called acidic oxides because they react with water to produce acids. Acidic oxides react with alkalis to form salts and water.Basic Oxides2)Basic Oxides are the oxides of metals. These react with acids to form salts and water.Amphoteric Oxides3)Some metallic oxides are amphoteric. This means they can behave as am acidic oxide or as a basic oxide. Amphoteric oxides react with acids and alkalis to form salts and water.Neutral Oxides4)Neutral Oxides do not react with either acids or bases. Therefore, they do not form salts.

Oxides can be classified into four types. These are acidic, basic, amphoteric and neutral.

A salt is produced when an acid reacts with a base. The method used to prepare a salt depends on the chemicals being reacted together to make the salt and on the solubility of the salt in water.

An Insoluble Salt is prepared by precipitation.

Soluble Salts of Sodium, Potassium and Ammonium are prepared via the titration method. (1) A known volume of acid is pipetted into a conical flask and universal indicator added. The acid is titrated with the alkali from the burette.(2) The acid is added until the indicator turns green, pH 7 neutral. This means all the acid has been neutralised to form the salt. I've illustrated the method using universal indicator BUT it isn't that accurate an indicator for titrations.You should use a more precise indicator like phenolphthalein or methyl orange. I didn't repeat all the titration details here again, I've just kept to the basic ideas and description, but there lots of detailed examples on the page(3) The volume of alkali needed for neutralisation is then noted, this is called the endpoint volume. (1)-(3) are repeated with both known volumes mixed together BUT without the contaminating indicator, such as phenolphthalein or methyl orange.(4) The solution is transferred to an evaporating dish and heated to partially evaporate the water causing crystallisation or can be left to very slowly evaporate - which tends to give bigger and better crystals.(5) The residual liquid can be decanted away and the crystals can be carefully collected and dried by 'dabbing' with a filter paper OR the crystals can be collected by filtration and dried (as above).

Other Soluble Salts are prepared by adding excess metal or insoluble carbonate or insoluble oxide to a suitable dilute acid.(1) The required volume of acid is measured out into the beaker with a measuring cylinder. The excess of insoluble metal, oxide, hydroxide or carbonate is weighed out and the solid added in small portions to the acid in the beaker with stirring. Doing the weighing will minimise trial and error especially if the reaction is slow, as long as you know how to do the theoretical calculation!(2) The mixture may be heated to speed up the reaction. When no more of the solid dissolves it means ALL the acid is neutralised and there should be a little excess solid. You should see a residue of the solid (oxide, hydroxide, carbonate) left at the bottom of the beaker.(3) The hot solution (with care!) is filtered to remove the excess solid metal/oxide/carbonate, into an evaporating dish.(4) You may need to carefully heat the solution to evaporate some of the water. Then hot solution is left to cool and crystallise. After crystallisation, you collect and dry the crystals with a filter paper.

Neutralisation and SaltsMonday, 28 September 2015 5:22 PM


Shiny, silvery solids1)All soft and easily cut with a knife2)Low densities and low melting points, densities increase down the group while melting point decreases.

3)

The Group I elements are known as alkali metals. They are very similar and they are the most reactive metals in the Periodic Table. All of them have one outer shell electron in their atoms, which is what makes them similar. They share the following physical properties:

Tarnish easily as they react easily with air, thus they are kept in oil1)Group I elements are called the alkali metals because they react with water to give alkaline solutions, turning red litmus blue

2)

Elements become more reactive down the group3)Metals reacts to form ionic compounds, and the ions have a charge of 1+4)

They also share the following chemical properties:

The elements consist of small molecules, which each contain 2 atoms, thus they are diatomic1)Fluorine and Chlorine are gases, Bromine is a liquid and Iodine is a solid2)Chlorine is greenish-yellow, Bromine is reddish-brown, Iodine is Shiny Black, and the Elements become darker down a group

3)

They have low melting and boiling points, which increases down the group.4)

Group VII elements are known as halogens. They are very similar and are very reactive. All of them have seven outer shell electrons in their atoms, which is what makes them similar. The share the following physical properties:

Halogens react with most metals to form compounds called halides1)Halogens become less reactive down the group2)A more reactive halogen will displace a less reactive halogen from an aqueous solution of its ions, and this reaction is called a displacement reaction

3)


The elements are all colourless gases, and consist of single atoms, thus they are monoatomic1)All the noble gases have stable electronic arrangements, and with the exception of helium, all have 8 electrons in the outer shells

2)

They have very low melting and boiling points3)

Group 0 elements are called the noble gases, and are the least reactive elements in the Periodic Table. They share the following physical properties:

They are all unreactive and usually do not form bonds.

They are Shiny and Silvery1)Good Conductors of Electricity2)They are hard and strong, thanks to the strong metallic bonds that are stronger than those in Group I and II metals

3)

They have high densities and high melting points, like most other metals4)Many solid compounds of the transitional metals and their aqueous solutions are coloured5)

The transition elements are a block of elements between Group II and III, occurring in Periods 4, 5 and 6. These elements are all metals, and are known as transition metals. They share the following physical properties:

Transition Metals have variable valency in compounds1)Several transition metals and their compounds are good catalysts, and can speed up chemical reactions

2)


Group Properties and Transition ElementsSunday, 20 September 2015 3:57 PM

Periodic Trends Page 33

I swear the rest of this chapter is useless the whole thing can be condensed into this

Formulas Monday, 28 September 2015 7:56 PM

Mole Concept Page 34

The Earth is surrounded by a layer of air about eight kilometres thick called the atmosphere. Air is important for living things and without it, there can be no life. Air is a mixture of gases. It consists of 78% Nitrogen, 21% Oxygen, 0.03% Carbon Dioxide and almost 1% Noble Gases. The amount of water vapour in air varies from day to day and from place to place. Air pollutants are also present in small amounts. Oxygen, nitrogen and the noble gases are useful gases that are obtained from the air. The air is first compressed and cooled to change into a liquid. The liquid air mixture is then separated by fractional distillation. In the process, nitrogen, which has the lowest boiling point, distils over first at -196 degrees Celsius, followed by argon at -186 and oxygen at -183.

This apparatus can be used to find the percentage of oxygen in the air. The copper powder is heated strong and air in the syringes passed back and forth over the hot copper until there is no further change in volume. The oxygen in the air combines with the hot copper to form black copper (II) oxide. The volume of gas in the syringes decreases as the oxygen reacts with the copper and is removed.Volume of Oxygen in Air = Initial Volume of air in the Syringes - Final Volume of Gas in the Syringes

Carbon Monoxide1)Caron Monoxide is a colourless, odourless gas, with the formula CO. Some of the carbon monoxide in the air is natural, resulting from forest fires. However, most of it comes from their incomplete combustion of fuel in motor engines. Carbon in the fuel burns to form carbon monoxide instead of carbon dioxide. Carbon monoxide is poisonous, and it is especially dangerous because it is colourless and odourless. It can cause breathing problems and reduces the ability of blood to carry oxygen, as it reacts with haemoglobin to produce carboxyhaemoglobin, causing it to no longer combine with oxygen. Nitrogen Oxides2)Oxides of nitrogen include nitrogen monoxide and nitrogen dioxide. At high temperatures, the nitrogen and oxygen in the air combine to form nitrogen monoxide. The nitrogen monoxide combines with more oxygen to become nitrogen dioxide. These reactions occur naturally in lightning and forest fires. The two oxides of nitrogen are often described simply as NO x. Most NO x is produced inside vehicle engines, although some come from burners in power stations, factories and incinerators. The gas irritates and damage the lungs. They are also a cause of acid rain. The gases also react with sunlight and other pollutants to form ozone. Sulfur Dioxide3)Coal and petroleum are important fuels. Large amounts are burnt around the world in power stations to generate electricity and in industries to provide energy. Both fuels contain sulfur as an impurity, although there is more of it in coal. Diesel fuel, used in vehicles, also contain a little sulfur. When the fuels are burnt, the sulfur is oxidised into sulfur dioxide. Sulfur dioxide is also produced in large quantities during volcanic eruptions. It irritates the eyes and causes breathing problems, as well as affecting plant growth as it enters the leaves. It is the main cause of acid rain.Other Pollutants4)Unburnt Hydrocarbons come mainly from hydrocarbons in fuel that have not been burnt in -

The presence of substances in the atmosphere that are harmful to living things and the environment contributes to air pollution. Pure air is colourless, odourless and safe to breathe. However, harmful substances are often present in the air. The most common air pollutants are carbon monoxide, nitrogen oxide and sulfur dioxide.

The AtmosphereFriday, 25 September 2015 1:55 AM

Air and Environment Page 35

Unburnt Hydrocarbons come mainly from hydrocarbons in fuel that have not been burnt in vehicle engines. Some can cause cancer, and they react with sunlight and other pollutants to form ozone.

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Methane is a colourless, odourless gas produced when plant and animal matter decay. Cows and other farm animals produce methane, large amounts of methane are also produced in rice fields. It contributes to global warming.

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Ozone, O 3, is a form of oxygen and is a colourless gas. Sunlight acts on other air pollutants such as nitrogen dioxide, carbon monoxide and hydrocarbons to produce a mixture called photochemical smog, and ozone is part of this mixture. Ozone irritates the eyes, nose and throat, and can cause asthma attacks. It damages plant crops.

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Air pollutants are harmful to the environment, causing the formation of acid rain, which has a pH value of below 5. Regular rain is slightly acidic with a pH value of 5.6 due to carbon dioxide in the air. Acid rainfall is any rainfall that has an acidity level beyond what is expected in non-polluted rainfall. Acid rain is formed from two air pollutants, sulfur dioxide and nitrogen dioxide.

Sulfur dioxide in the air reacts with oxygen and water to form sulfuric acid. The sulfuric acid dissolves in rainwater, making it acidic. Nitrogen dioxide also undergoes a similar reaction in the air, forming nitric acid. The nitric acid also dissolves in rainwater, making it acidic. Acid rain with a pH of about 4 is 40 times more acidic than unpolluted rainwater. Some acid rain has been found to have a pH of less than 1.5, more acidic than lemon juice. Sometimes, wind carries air pollutants over long distances before they dissolve in rainwater. Therefore air pollutants from one country can produce acid rain in another country.

Acid rain makes soils more acidic, and many plants do not grow well in acidic soil.-

Fish cannot survive in acidic water, with acid rain killing fish in thousands in lakes in Scandinavia and North America.

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Acid rain corrodes buildings and objects made of calcium carbonate, such as limestone, marble or cement.

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Acid rain damages trees, and destroys whole forests throughout Europe.-

Acid rain attacks metals, causing galvanised iron sheets to corrode more quickly. -

Acid rains have harmful effects:

A layer of ozone surrounds the Earth about 40 km above the ground. The ozone absorbs UV radiation from the Sun, preventing it from reaching the ground. UV radiation is dangerous, as it causes skin cancer in animals and damages plants. The amount of ozone in the ozone layer has decreased since 1978, and the depletion is particularly great at the South Pole, where up to half the ozone over Antarctica has been destroyed. This formed the ozone hole and dangerous ultraviolet radiation is now streaming through this hole. In the tropics, there is much less depletion of Ozone. The ozone hole is caused by chlorofluorocarbons, or CFCs. They are compounds of carbon, chlorine and fluorine. CFCs have been widely used as propellants in aerosol cans and as coolant liquids in refrigerators and air conditioners. CFCs in the atmosphere are decomposed by sunlight to produce


refrigerators and air conditioners. CFCs in the atmosphere are decomposed by sunlight to produce chlorine atom. These chlorine atoms then react with ozone molecules and destroy them by converting them into oxygen molecules. In the last few decades, large amounts of CFCs were released into the atmosphere. CFCs are very stable and can remain in the atmosphere for a long time. Slowly they react with ozone, destroying the ozone layer. An international treaty called the Montreal Protocol was signed in 1987 to drastically reduce the use of CFCs. In 1992, major industrial nations agreed to completely phase out ozone-destroying chemicals by 2020. Although limited use of CFCs is allowed, other compounds are now used for aerosol propellants, refrigerators and producing plastics. These low boiling point hydrocarbons do not contain chlorine or are rapidly destroyed in the lower atmosphere to form small molecules that do not react with ozone. Singapore signed the Montreal Protocol in 1989 and implemented policies to comply with its obligations, and since 1996, import or manufacture of products containing CFCs has been banned. Unfortunately, a lot of chlorine from CFCs has accumulated in the atmosphere, though quantities are now beginning to fall. Still, ozone depletion is expected to be serious until the middle of this century.

Catalytic Converters attached to the exhaust system of vehicles contains a ceramic catalyst for the removal of carbon monoxide, hydrocarbons and nitrogen oxides helps to reduce air pollution in Singapore, and has been used since 1994. When the hot exhaust gases from the engine pass over the catalyst, the harmful pollutants are converted to harmless substances. Nitrogen Oxides react with Carbon Monoxide to from carbon dioxide and nitrogen and water. Unburnt Hydrocarbons such as Octane are oxidised to form carbon dioxide and water.

Oil fuels used in Singapore's factories contain no more than 2% sulfur, and the amount of sulfur in diesel fuel has decreased, from 0.5% in 1996 to 0.05% in 1999. Power stations are now burning more natural gas while buses and taxis increasingly use compressed natural gas instead of diesel. Natural gas is mainly methane and contains no sulfur, producing non-polluting products from combustion, carbon dioxide and water.

Flue Gas Desulfurisation is a process used to remove sulfur dioxide from waste gases produced by a coal or oil burning power station. The waste gases are called flue gases, and are made up of sulfur dioxide and small amounts of nitrogen oxides and are produced when fuels are burnt. To remove sulfur dioxide, powdered limestone is added to hot flue gases, and the heat decomposes the limestone to give calcium oxide. The calcium oxide then reacts with sulfur dioxide to form calcium sulfite. This step removes sulfur dioxide from the flue gases, and the remaining gases can now be safely released. In many instances, calcium sulfite is further oxidised to form calcium sulfate.

The atmosphere is about 0.03% carbon dioxide. However, this small percentage represents many billion tonnes of the gas. Carbon Dioxide in the atmosphere is continuously being removed and returned to the atmosphere in a variety of processes, and is called the carbon cycle. Photosynthesis, combustion and respiration are three processes involved in the carbon cycle.


Plants remove carbon dioxide from the atmosphere by photosynthesis which occurs in the green leaves on the plants. In the presence of sunlight, chlorophyll in the leaves converts carbon dioxide and water into glucose and oxygen. The glucose is converted into starch and stored in the plants. During photosynthesis, solar energy is converted into chemical energy which is stored in the glucose and the starch.

Combustion1)Fuels such as coal, petrol and natural gas contain hydrocarbons, which are compounds containing carbon and hydrogen. When the fuels burn in excess oxygen, carbon dioxide and water are produced. The combustion of methane produces carbon dioxide and water. Respiration2)All living things carry out respiration. During this process, carbon and hydrogen in foods are changed into carbon dioxide and water. Energy is also released during respiration. The equation for the reaction that takes place during respiration in living cells has food and oxygen react to form carbon dioxide and water.

Carbon dioxide is returned into the air via combustion and respiration.

The carbon dioxide produced during combustion and respiration enters the atmosphere. Over the thousands of years, these processes have kept the amount of carbon dioxide in the atmosphere fairly consistent.

Solar energy from the Sun keeps the Earth warm, and without this energy, the earth would be a cold, lifeless planet. The Earth is just the right temperature for life to exist. However, scientist believe that the Earth is becoming warmer. The gases in the atmosphere that trap heat are called greenhouse gases. Carbon dioxide is a greenhouse gas. Methane, nitrogen dioxide and CFCs are other greenhouse gases. The gradual increase of greenhouse gases in the atmosphere is causing Earth to become warmer. The phenomenon is called global warming, Solar energy includes visible radiation and infrared radiation. We detect infrared radiation as heat. Gases in the atmosphere allow this radiation to reach the Earth's surface and are able to trap infrared radiation emitted from the Earth. The warming effect due to the trapping of heat energy by gases in the atmosphere is called the greenhouse effect.


Stage 3 is where the greenhouse effect occurs. An increases in the amount of greenhouse gases present in the atmosphere causes more infrared radiation to be absorbed. This leads to an overall increase in the earth's temperature. It is 0.5 degrees Celsius warmer than it was 100 years ago, and by 2050, Singapore may be 10 degrees Celsius hotter than it is today. Greenhouse gases is the main cause of global warming. The more greenhouse gases there are in the atmosphere, the more heat from the sun that will be trapped. Over the past 200 years, since the beginning of the Industrial Revolution, the amount of carbon dioxide in the atmosphere has increased by about 25%. This is mainly due to the burning of fossil fuels in power stations, factories and motor vehicles. At the same time, an increase in agriculture is increasing the quantity of methane released into the atmosphere.

Land Ice, such as glaciers and the Antarctic ice cap will melt and cause sea levels to rise even further, with a rise of 20cm predicted by 2050. This will result in low-lying islands and coastal lands disappearing.

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There will be big changes in global climate. Equatorial countries such as Singapore will probably experience more rain, but countries in North America and Europe will probably suffer drought, resulting in food shortages.

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Hotter weather results in more suitable habitats for Mosquitoes. Mosquito-borne diseases such as malaria and dengue fever are already starting to spread. Warmer weather in the oceans will produce more powerful and destructive typhoons

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There are various consequences of Global Warming:


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