As Chemistry Intro Resource Booklet

Name : .....................................

AS Chemistry Start module 2015

Resource material booklet

This booklet contains much useful information.

[The majority of this information has been covered in theYear 10 and 11 courses]

The more of it you know the better your results will able. In Chemistry knowledge is the key.

If you are not good at remembering dont lose this booklet, keep it handy!

It will prove very useful throughout the year.

AS intro resource booklet 2015 page 1

Index.

page 3 General ideas on chemical formulapage 5 Table of ionspage 6 Writing ionic formulapage 8 Molecular formulae you need to know.page 9 Common names of chemical substancespage 11 Colours of chemical substancespage 13 The Bohr atompage 15 Electron configurationspage 21 Types bondpage 23 Structurepage 28 Types of chemical reactionpage 29 Salt dissolving


Names and formulae of chemical compounds

There are millions of known chemical compounds.

These compounds have specific structures and are made up of simple particles, atoms, ions and molecules.

The name and formula give an idea of the composition and properties of a compound. So it is important to understand these names and formulae.

In Chemistry a compound often has a common name and a systematic name,

These compounds will also have a simple formula, but there can be other types of formula which can help to explain the properties of a particular compound.

Two examples are shown below to give you an idea of the importance of naming a substance, and writing the appropriate formula.

1. Compound 1

common name - common salt systematic name - sodium chloride emprical formula - NaCl

structural formula -


2. Compound 2.

common name - glucose

systematic name - 2,3,4,5,6-pentahydroxyhexanal

emprical formula - CHO molecular formula - CHO

structural formula

glucose powder or CHOHCH(OH)CH(OH)CH(OH)CH(OH)CHO

or

Properties of these two substances

i. Both these compounds are white crystalline solids.ii. The melting point of salt is far higher than glucose.iii. Both are soluble in water.iv. A salt solution conducts an electric current, a glucose solution does not v. Sugar dissolves in alcohol, salt is insoluble in alcohol. vi. Sugar contains covalent bonds, salt contains ionic bonds

All these facts are due to the structure of these two substances.


CationsCationsCations1+ ions 2+ ions 3+ ions

hydrogen Hlithium Lisodium Napotassium Ksilver Agcopper(I) Cuammonium NH

hydronium HOor hydroxonium

magnesium Mg2calcium Ca2iron(II) Fe2 pale greencopper(II) Cu2 bluezinc Zn2barium Ba2tin(II) Sn2lead(II) Pb2manganese(II) Mn2pale pinknickel Ni2 greencobalt(II) Co2pink

aluminium Alscandium Sciron(III) Fe yellow/brown

chromium(III) Cr green or blue

AnionsAnionsAnions1- ions 2- ions 3- ions

hydroxide OHhydrogencarbonate HCOhydrogensulfate HSOnitrate NOnitrite NOfluoride Fchloride Clbromide Briodide I

permanganate MnO purple

dihydrogenphosphate HPO

hypochlorite ClO or OClchlorate ClObromate BrOiodate IOethanoate CHCOOaluminate Al(OH) or AlO

oxide O2 various colours

carbonate CO2sulfate SO2

sulfite SO2

sulfide S2various colours

manganate MnO2 green

monohydrogenphosphate

HPO2

dichromate CrO2orange

chromate CrO2yellow

zincate Zn(OH)2 or ZnO2

nitride N

phosphide P

phosphate PO


READ THE IDEAS ON THIS PAGE - THEN - WORK THROUGH QUESTIONS

Ionic substances are made from positive ions - cations negative ions - anions

These solids are made up and an ordered arrangements of these ions.These ordered structures are called crystals. Sometimes these are large and characteristic in shape other times the crystals are very small and appear as a powder.

These compounds have a neutral charge. The charges of the positive ions are countered by the charges of the negative ions.

An ionic formula represents the combination of these ions.

The symbol for the positive is written first followed by the symbol for the negative ion. The charges of the ions are omitted.

Thus copper oxide, a black insoluble solid, Cu O CuO( charges equal and opposite thus one ion of each required )

lead chloride, a white insoluble solid, Pb Cl PbCl( two Cl ions are needed to oppose the Pb ion the formulaof the solid show one Pb and two Cl ions )

aluminium sulphide, a white solid, Al S AlS( two Al ions, 6+, and three S ions, 6-, are needed to form a formula with neutral charge. )

iron(III) hydroxide, a red insoluble solid, Fe OH Fe(OH)( one Fe is required to oppose three OH ions.A bracket ( ) is put around the hydroxide ion because the ion comprises the atoms of oxygen and hydrogen and it needs three of each to balance )

2. Other examples to show you how to do this work

a) sodium nitrate, Na NO NaNOb) calcium sulphate, Ca SO CaSOc) ammonium carbonate, NH CO (NH)COd) aluminium oxide, Al O AlO

1


3. Now try these, USE YOUR TABLE OF IONS and perhaps a scrap piece ofpaper to get the correct formula.

a.Zinc carbonate ................. b. Potassium fluoride .................

c. Magnesium sulphate ................. d. Copper oxide .................

e.Lead hydroxide ................. e. Sodium oxide .................

f. Aluminium chloride ................. g. Silver chloride .................

h. Potassium sulphide ................. i. Copper nitrate .................

j. Ammonium sulphate ................. k. Calcium iodide .................

l. Sodium phosphate ................. m. Iron (III) nitrate .................

n. Barium nitrate ................. o. Zinc nitride .................

p. Potassium chromate ................. q. Iron(II) sulphate .................

r. Calcium hydrogen carbonate .................

s. Aluminium sulphide .................

t. Magnesium ethanoate .................

u. Sodium zincate .................

v. Ammonium dichromate .................

w. Calcium dihydrogenphospate .................

x. Manganese (IV) oxide .................

2


Molecular formulae

a) Molecular substances contain only non metal atoms. These are substances where the atoms are held together by covalent bonds.b) Sometimes the name gives a clue as to the formula. c) Other formulae need to be learnt. LEARN AS MANY AS POSSIBLE !!

Name Formula Properties1. Water HO a colourless liquid at room temp 2. Hydrogen peroxide HO a colourless liquid used as a bleach3. Carbon monoxide CO a colourless poisonous gas 4. Carbon dioxide CO a colourless gas turns limewater milky5. Methane CH natural gas used as a fuel gas 6. Ethane CH a gas found in LPG7. Ethene CH a gas used to make plastics 8. Ethyne CH a gas mixed with O to weld iron9. Sulphur dioxide SO a pollutant gas formed when sulfur burns 10. Ammonia NH3 strong smelling gas, used as a cleaner

11. Nitrogen dioxide NO brown pungent poisonous gas, air pollutant 12. Nitrogen N 80% of air

13. Hydrogen H explosive colourless gas, 96% of the universe

14. Oxygen O a colourless gas supports life & combustion

15. Ozone O a reactive poisonous form of oxygen

16. Chlorine Cl a yellow/green poisonous gas

17. Glucose CHO white solid, simplest of the sugars

18. Iodine I grey/silver solid, a mild antiseptic

19. Sulphur S yellow solid

20. Phosphorus P white unstable flammable solid

21. Hydrochloric acid HCl colourless liquid, stomach acid.22. Sulphuric acid HSO colourless liquid, battery acid.23. Nitric acid HNO colourless liquid.24. Ethanoic acid CHCOOH colourless liquid, found in vinegar.25. Carbon C a black solid, also formula for diamond, graphite or charcoal.

26. Glucose CHO white solid, the simplest of the sugars27. Ethanol CHCHOH a colourless liquid, alcohol


Common names of substances in Chemistry

Alcohol ethanol CHCHOHAlum aluminium potassium sulphate 24-hydrate Al(SO).KSO.24HOAlumina aluminium oxide AlOBaking soda sodium hydrogen carbonate NaHCOBauxite ore containing aluminium oxide AlO ( FeO + SiO )Bleaching powder calcium chloride hypochlorite CaCl(OCl)Bluestone copper sulphate pentahydrate CuSO.5HOBrine sodium chloride solution/seawater NaClCarborundum silicon carbide SiCCaustic potash potassium hydroxide KOHCaustic soda sodium hydroxide NaOHChalk calcium carbonate CaCOCharcoal graphite, amorphous carbon CChile saltpeter sodium nitrate NaNOCommon salt sodium chloride NaClCryolite sodium aluminium fluoride NaAlFDiamond carbon CDry ice carbon dioxide solid COEpsom salts magnesium sulphate heptahydrate MgSO.7H0Ether diethyl ether CHCHOCHCHFlowers of sulphur sublimed sulphur SFools Gold iron sulphide FeSGlaubers salt sodium sulphate decahydrate NaSO.10HOGlycerin glycerol (from soap manufacture) CH(OH)CH(OH)CHOHGraphite crystalline carbon sheets CGypsum calcium sulphate dihydrate CaSO.2HOLaughing gas nitrous oxide NOLimewater calcium hydroxide solution Ca(OH)Litharge lead monoxide PbOMagnesia magnesium oxide MgOMagnetite triferric tetroxide, mixed iron oxide FeOMarble calcium carbonate CaCOMilk of magnesia magnesium hydroxide suspension Mg(OH)/HOMoth balls naphthalene CHNiter potassium nitrate KNONitroglycerine glyceryl trinitrate CH(NO)CH(NO)CHNOOil of vitriol sulphuric acid HSOPlaster of Paris calcium sulphate hemihydrate CaSO./HOQuartz silicon dioxide SiOQuicklime calcium oxide CaOQuicksilver mercury HgRed lead triplumbic tetroxide, mixed lead oxide PbORock salt sodium chloride NaClSal ammoniac ammonium chloride NHClSal volatile ammonium carbonate (NH)COSaltpeter potassium nitrate KNOSilica silicon dioxide SiOSlaked lime calcium hydroxide solid Ca(OH)Soda ash sodium carbonate NaCOSugar sucrose CHOTNT trinitrotoluene CH(NO)Washing soda sodium carbonate decahydrate NaCO.10HOWood alcohol methanol CHOH


2 The following are of interest , but are not as important as those listed above.

Aqua fortis nitric acid HNOAqua regia nitric acid and hydrochloric acid HNO/HClBorax sodium tetraborate decahydrate NaBO.10HOBrimstone liquid sulphur SCinnabar mercury(II) sulphide HgSCorrosive sublimate mercury(II)chloride HgClCream of tartar potassium hydrogen tartrate KHCHODolomite calcium and magnesium carbonate CaCO.MgCOGreen vitriol ferrous sulphate heptahydrate FeSO.7HOHalite sodium chloride NaClHypo sodium thiosulphate pentahydrate NaSO.5HOLye sodium hydroxide NaOHMarsh gas methane CHMuriatic acid hydrochloric acid HClPrussic acid hydrocyanic acid HCNWhite vitriol zinc sulphate heptahydrate ZnSO.7HO


Colours in ChemistryYear 11 level

SOLIDSSOLIDS1. White solids All Na, K compounds (except NaO, KO)

All Al, Mg , Ca , NH compounds (except some alums) (unless anion All Pb (except PbO, PbO, PbO, PbI, PbBr ) is coloured) All Zn (except hot ZnO)

All Ag (except AgBr cream, AgI pale yellow , AgO brown ) SiO pure sandP sometimes called yellowTiOCuSO anhydrous

2. White Precipitates Carbonates of Pb , Mg , Zn , Ca (Al(CO) does not exist)PbClPbSO, BaSO, CaSOHydroxides of Mg , Ca , Al, Pb , Zn , Sn

3. White precipitates Al, Zn , Pb , Sn in excess NaOH that dissolve AgCl, Zn(OH) in ammonia

all carbonates (any colour) in acidsall hydroxides (any colour) in acids

4. Yellow solids Sulphur, Au, AgI, FeS (iron pyrites)PbI, PbO, PbCrO and most CrO salts. ZnO when hot

5. Black solids Charcoal / graphiteCuO, CuS, PbSFeO, FeOMnO (or brown)

6. Brown solids PbO, AgOFeCl anhydrous is dark brown crystalsFeCl.6HO is yellow brown crystalsFeO is red brown, rusty

7. Red solids PbO , CuO, HgO , P, AgCrO

8. Grey solids CaC, FeS, I crystals ( vapour purple )

9. Orange solids (NH)CrO, KCrO, NaCrO


10. Purple solids KMnO

11. Green solids CuCO, (really a basic carbonate, CuCO.Cu(OH) )Fe compounds (except FeO black, FeS yellow)Ni compoundsCrO

12. Blue solids CuSO.5HO, Cu(NO).3HOanhydrous CoCl

13. Pink crystals CoCl.6HOMn salts are very pale pink

14. Green precipitates CuCO.Cu(OH) from Cu + CO FeCO and Fe(OH) (both darken on standing in air)

15. Blue precipitates Cu(OH) dissolves in NH to dark blue solution Cu(NH)

16. Brown precipitates Fe(OH) , AgO from Ag + OH

LIQUIDSLIQUIDS17. Colourless Aqueous solutions of most soluble salts, acids and alkalis

HO, HO, Organic liquids, e.g. alkanes, alcohols, esters

18. Red Br

19. Coloured solut ions ( i) Green Fe , Ni , Cr, MnO ( ii) Blue Cu (sometimes green/blue) ( iii Purple MnO, I in non-polar organic solvents eg. hexane ( iv) Orange CrO (v) Brown Br solutions, I in polar solvents eg. propanone , I (I + I) (vi) Yellow Fe, CrO

GASESGASES

20.Colourless Most gases, Some fume in moist air e.g. HCl

21.Coloured Cl green/yellow F pale yellow NO brown

22.Coloured vapours [i.e. having B.Pts. above room temperature]Br orange/brown,I purple, Sulphur dark red


nucleuscontains neutrons and protons

1234

electron shell, n

2 8 18 32

maximum number of electrons in shell, 2n

5

Bohr model of the atom

50

Atoms 1 - 20

Period 1

H H H He

Period 2

Li Be B C N flO F Ne

Period 3

Na Mg Al Si P


Period 3

S fiCl Ar

Period 4 and then transition metals

K Ca fiSc


Electron configurations of selected elements.

Noble Gases Group VIII or 0

98

818321882Rn

8181882!Xe

81882Kr

882Ar

82Ne

2He

72503218822n

helium gas in blimp excited neon atoms krypton filled bulb



Alkali Metals - Group I

all have 1 valence [outer] electron

98

1818321882Fr

18181882!!Cs

181882Rb

1882K

182Na

12Li

72503218822n



Halogens - Group VII

all have 7 valence [outer] electrons

98

78321882!At

7181882!I

71882!Br

782Cl

72F

72503218822n

chlorine gas bromine liquid iodine solid



First row of Transition metals Sc Zn

11882Cu

21682Ni

21382!Mn

11382Cr

21182V

21082Ti

2

28

98

31882Ga

21882Zn

21582Co

21482Fe

982Sc

82Ca

72503218822n

Scandium metal Vanadium metal Manganese metal


Other elements

2

29

118

118

118

98

11832321882Rg

825321882!Am

21321882U

32882Au

1882Ag

82Cu

72503218822n

Biggest known nugget Americium-241 cubes of Uraniumof Gold in a smoke alarm


The only atoms that can exist on there own at room temperature are the Noble Gases.

These exist as monatomic gases at room temperature.

At room temperature all other atoms have to form links or bonds with each other.

Four types of chemical bond

1. Ionic bond - between cation, +ve and anion, -ve

2. Covalent bond - between non metal atoms,[ sigma and pi , non-polar, polar, strongly polar & coordinate bonds ]

3. Metallic bond - between metals, metal cations and their valence electrons

4. Intermolecular - between molecules or atoms of Noble Gasesbond [ H - bonds, permanent dipoles, temporary or induced dipoles ]

[ Understand and identify these in a substance and you will begin

to make sense of this subject. ] ! !


Cl Cle

e

a non-polar bond

a pair of electrons shared between two non-metal atoms with the same electronegativity ens Cl = 3.0

HCle

e

a polar bond

a pair of electrons sharedunequally between two non-metal atoms with different electronegativitiesens Cl = 3.0, H = 2.1

HOe

e

a strongly polar bond

a pair of electrons sharedunequally between two non-metal atoms with very different electronegativities, O-H, N-H, F-Hens O = 3.5, H = 2.1

Cl Na an ionic bond

a pair of electrons not sharedbetween anon-metal and a metalens Cl = 3.0, Na = 0.9

NaeNa a metallic bond

an attraction between metal cationsand their mobile valence electronsens Na = 0.9

types of bond


Structure

Types of structure that you should be familar with

i.e. you should know about

a. particles holding structure together

b. types of force holding particles together

c. the relative strengths of these forces

d. the physical properties involved with these structures.


TYPES OF SOLIDS

Forces of Attraction in Crystall ine Solids Crystalline solids have an ordered array of atoms or molecules. This array is called a lattice. The atoms or molecules in solids are held together in four different ways and these are: Molecular Solids: Molecules are held together by relatively weak forces (van der Waals, hydrogen bonding) so molecular solids have low melting points and in many cases only form solids at low temperatures. e.g. ice, solid CO. Most of the non-metal elements e.g. chlorine and neon, form solids of this type. The forces between the molecules are weak BUT within the molecule e.g. CO there are strong forces holding the atoms together.


Ionic Solids: The lattice is made up of positive and negative ions. The force of attraction between the oppositely charged ions is strong compared to molecular solids. They have high melting points. Each ion is surrounded by neighbours of opposite charge and there are no separate molecules. Ions are not easily moved from their lattice sites, so typical ionic solids are hard and rigid, though brittle [If an ion is displaced, it moves closer to an ion of like charge which repels it, thus breaking the lattice]. They shatter easily and cannot be shaped by hammering and bending. The electrons are localized on the ions and ionic solids only conduct electricity when they are molten and the ions are free to move. Examples are NaCl [sodium chloride] and KI [potassium iodide].

Network Covalent Solids: Atoms are covalently bonded throughout the entire solid, thus the force of attraction is a covalent bond. The strong attractive force gives them very high melting points. The electrons are localized and are not free to move making them insulators. Examples are diamond, quartz (SiO2) and the silicate minerals. In diamond each carbon atom is bonded tetrahedrally to four other carbon atoms. Quartz has a lattice of SiO4 tetrahedra where every oxygen is bonded to two silicon atoms. The other two allotropes of carbon (different forms of the same element in the same physical state) are also network solids and are graphite and the fullerenes. Both conduct heat and electricity as the electrons can move freely through the solid. Graphite consists of sheets which can move over each other due to weak forces holding the sheets together. The fullerenes consist of balls of between 30 - 72 Carbon atoms.

diamond graphite


Metall ic Solids: The metal atoms occupy regular lattice sites, however the valence electrons can move throughout the entire solid. These delocalized electrons readily carry heat and electricity throughout the solid. The mobile electrons produce the typical sheen or metallic lustre characteristic of metals. Since the force of attraction is not localised, the metal atoms can slide from one lattice site to another, making metals malleable (shaped by hammering) and ductile (drawn into wire). The strength of the metallic bond is quite variable as indicated by the range of melting points from tungsten, 3422 0C to mercury, !39 0C.


This table summarises some of the facts about the four types of solids.

TYPE of SOLID

METALLIC

IONIC

MOLECULAR

COVALENT

Entity

atoms

ions

molecules

atoms/molecules

Force of Attraction

metallic

bond

ionic bond

Van der Waals

covalent bond

Strength of the Force of

Attraction

variable

strong

weak

strong

Texture

soft

hard / brittle

soft

very hard

Conductor of

Heat or Electricity

conductors

insulators [conduct

when molten]

insulators

insulators

(except graphite)

Melting Point

variable mp

-39 to 4000C

high mp 600 to 3000C

low mp

-272 to 400C

very high mp

1200 to 4000C

Solubility

insoluble in

water/organic

soluble in

water

soluble in organic solvents

insoluble

Examples

Fe Pb Sn

NaCl MgO

Na2SO4

sugar CH4 CO2

diamond graphite quartz

Amorphous Solids An amorphous solid is not crystall ine as the atoms or molecules are not arranged in a regular lattice and is formed when the shape of the molecule is too complex, the molecules are frozen too rapidly or when impuriites prevent a regular lattice. They include glass, rubber and plastics. Amorphous solids can be considered to be supercooled liquids of high viscosity.


Types of chemical reaction

1. Combustion A reaction where a substance burns, normally in oxygen, to produce energy in the form heat, light and sometimes sound.

Example : calcium + oxygen calcium oxide

2. Exothermic A reaction where energy is released, in the form of heat or light or sound. The enthalpy [stored energy] of the products is less than that of the reactants.

Example : any combustion reaction, most spontaneous reactions.

3. Endothermic A reaction where energy is absorbed. Normally the reaction mixture gets cooler. Its enthalpy increases.

Example : ammonium salts dissolving.

4. Dehydration A reaction where a substance loses water to form a new substance. [it is not dried]

Example : sucrose with conc sulfuric acid carbon + water 5. Hydrolysis A reaction with water to form a new compound. [not dissolving something in water]

Example : sulfur dioxide + water sulfurous acid

6. Displacement A reaction where a reactive element displaces a less reactive element Example : zinc + copper sulphate copper + zinc sulphate

7. Effervescence A reaction where bubbles of gas are vigorously produced.

Example : acid + metal salt + hydrogen

8. Decomposition A reaction where a substance breaks down into simpler substances. Can be achieved by heat, a catalyst and by electrolysis.

A B + C

Example : magnesium hydroxide magnesium oxide + water


9. Precipitation A reaction where two solutions are mixed together to produce a suspension which contains an insoluble product known as a precipitate.

Example : sodium chloride + silver nitrate sodium nitrate + silver chloride

a white ppt.

Solubility of ionic solids

------------all common ones



Pb,Ba, Ca and (Ag)the rest

the rest Ag,Pb also Cu

the restNa,K,Ba react with

water

the restNa,K, NH,

Ca and Ba slightly

the restNa,K and NH

insolublesoluble

CHCOO

NO

HCO

SO

Cl, Br, I

O

OH

CO

10. Combination A reaction where two or more substances combine to form a more complicated product. D + E F

Example : magnesium + nitrogen magnesium nitride

11. Neutralisation A reaction where an acid reacts with a base to form a salt and water (and CO when the base is a carbonate or hydrogencarbonate)

Example : nitric acid + lead(II) oxide lead nitrate + water


What happens when salt dissolves? Why does it disappear? Why is it soluble?


As Chemistry Intro Resource Booklet

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