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Chemistry ATAR 11 > CHEMISTRY SEMESTER2 EXAM > Flashcards

CHEMISTRY SEMESTER2 EXAM Flashcards

1
Q

Relative Atomic Mass

A

Ar = (%a × Ar(a) +%b × Ar(b))/100

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2
Q

Structural formulas

A 
Meth………1 carbon atom
Eth…………2 carbon atoms
Prop……….3 carbon atoms
But…………4 carbon atoms
Pent………..5 carbon atoms
Hex…………6 carbon atoms
Hept……….7 carbon atoms
Oct………….8 carbon atoms
Non…………9 carbon atoms
Dec…………10 carbon atoms
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3
Q

Polarity/Non-Polarity

A

− Polar molecules: those with slight positive and slight negative charged ends. HAVE NET DIPOLE
− Occurs in all molecules with a single bond dipole
− Non-polar molecules: those where the bond dipoles are of equal size and act in symmetrically opposing directions where the dipoles cancel each other out. HAVE ZERO NET DIPOLE

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4
Q

Mass/Volume/Concentration

A

c=n/v

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5
Q

Properties of water

A

− Consists of small molecules with weak dispersion forces
− Strong polarity and extensive capacity of hydrogen bonding dictate its physical properties
− HIGH MELTING POINT (compared to other molecules of similar size); due to extensive hydrogen bonding with neighbouring molecules

− DENSITY IN THE SOLID PHASE IS LESS THAN IN LIQUID PHASE which is why ice floats; undergoes expansion
− Ice floating on water acts as an insulator
− Occurs because each water molecule forms four hydrogen bonds with four neighbouring water molecules

− HIGH SURFACE TENSION (liquids tendency to resist any increase in its surface area); results from strong intermolecular forces
− At water surface there is an imbalance of these forces causing molecules to be pulled in towards the bulk material
− Surface is trying to contract and achieve a minimum area

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6
Q

Trends In The Periodic Table

ATOMIC RADIUS

A

− Decreases from left to right (due to increasing nuclear charge and electrons are pulled closer to the nucleus)
− Increases down any group (as you go down a group the number of electron shells increases)

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7
Q

Trends In The Periodic Table

FIRST IONISATION ENERGY

A

a measure of how strongly it an atom holds onto its electrons
− Increases from left to right (due to the atom wanting to hold onto its electrons)
− Decreases down any group (as atomic radius gets bigger there is a further distance between electron and nucleus, thus weaker attraction)

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8
Q

Trends In The Periodic Table

ELECTRONEGATIVITY

A

ability of an atom in a molecule to attract electrons to itself
− Increases from left to right (due to the atom wanting to gain electrons)
− Decreases down any group

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9
Q

Addition reactions of hydrocarbons

A

− ONLY ALKENES UNDERGO ADDITION
− Double bond is replaced by bonds to other atoms such as H, F, Cl, Br or I
− Reagents used are H2, Cl2, Br2, HCl and HI
EXAMPLE
CH2=CHCH2CH2CH2CH3 + Br2 -> CH2BrCHBrCH2CH2CH2CH3

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10
Q

Substitution reactions of hydrocarbons

A

− Occur when alkane or benzene is combined with Cl2 or Br2
− Replacement of one of more H atoms with Cl or Br atoms
EXAMPLE
CH3CH3 + Br2 -> CH3CH2Br + HBr

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11
Q

Saturated

A

− An SATURATED solution contains as much solute as it can normally dissolve

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12
Q

Unsaturated

A

− An UNSATURATED solution contains less solute than it is able to dissolve

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13
Q

Supersaturated

A

− An SUPERSATURATED solution contains more solute than is can normally dissolve

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14
Q

KINETIC THEORY

A

− Particles of a gas are in rapid continuous random motion
− Attraction and repulsion between particles in a gas is negligible
− Total volume of all particles is negligible compared to volume the gas occupies
− Have kinetic energy
− Average kinetic energy of particles of a gas is proportional to its temperature
− Particle collisions are elastic and do not lose speed or slow down

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15
Q

KINETIC THEORY APPLIED TO REAL GASES

A

− Ideal gas has negligible volume, real gases do occupy space and volume is significant
− Ideal gas has negligible attraction, real gases do have forces of attraction and are significant

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16
Q

UNIVERSAL INDICATOR

A 
−	pH 0-3 = red
−	pH 4-6 = orange/yellow
−	pH 7 = green
−	pH 8-10 = blue
−	pH 11-14 = purple
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17
Q

METHYL VIOLET

A

− Acid colour = Yellow

− Base colour = Blue

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18
Q

BROMOCRESOL PURPLE

A

− Acid colour = Yellow

− Base colour = Violet

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19
Q

THYMOLPHTHALEIN

A

− Acid colour = Colourless

− Base colour = Blue

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20
Q

Chromatography

A

− A multi-step process used to separate and/or analyse the different components or solutes of a mixture or substance, whether it be liquid or gas, without causing any molecular changes to the chemicals involved

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21
Q

STATIONARY PHASE

A

− Solid or thick liquid

− Remains in fixed position

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22
Q

MOBILE PHASE

A

− Liquid or gas

− Carries components through or across stationary phase

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23
Q

Absorption

A

− Fluid is dissolved by a liquid or sold
− Endothermic
− Not affected by temperature
− Concentration uniform

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24
Q

ADSORPTION

A

− Atoms, ions or molecules of a substance adhere to the surface of the absorbent
− Exothermic
− Preferred at low temperature
− Concentration on surface is different from inside

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25
Paper chromatography
− Used to separate coloured chemicals or substances − STATIONARY PHASE: Paper − MOBILE PHASE: aqueous liquid or non-aqueous solvent − Ink is added to the bottom of the paper and placed in the solvent which uses capillary action to pull it up the paper and separate each component depending on polarity − How far the components are moved up the paper depends on the ability to be absorbed into the paper vs ability to be carried along − More polar components absorb more quickly as paper is polar
26
Thin-layer Chromatography
− Analysing mixtures − STATIONARY PHASE: thin glass plate coated in silica or alumina − MOBILE PHASE: solvent chosen on polarity − Small amount of solution added to the end of TLC place which is then placed in the mobile phase which uses capillary action to pull it up the glass and separates each component depending on polarity − Substances that aren’t visible are coated and shone under UV light or placed in iodine − SHARPER AND MORE DEFINED THAT PAPER CHROMATOGRAPHY − SIMPLE AND INEXPENSIVE − ONLY SMALL AMOUNT OF SAMPLE NEEDED
27
Gas Chromatography
− Ability to identify individual substances of a mixture and their concentrations − STATIONARY PHASE: high boiling point non-volatile viscous liquid, which has been coated in silica − MOBILE PHASE: inert carrier gas − Sample is dissolved and vaporised in a vaporisation chamber via syringe − Flows along carrier gas stream into chromatography column − Components with high volatility stay in mobile phase and leave quicker and have lower retention time − At the end a detector records a chromatograph
28
High Performance Liquid Chromatography
− Best suited for substances that decompose when heated or cannot be vaporised − STATIONARY PHASE: tightly packed column of fine particles such as silicon dioxide − MOBILE PHASE: liquid solvent − Sample is injected into high-pressure column containing the solvent where it travels through a chromatography column where the mixture is separated − Detector created chromatograph
29
NORMAL PHASE HPLC
− Filled with silica particles and non-polar solvent (polar components will attach to silica molecules longer that non-polar)
30
REVERSE PHASE HPLC
− Silica is non-polar due to attachment of hydrocarbon and polar solvent
31
Percentage mass
% element in compound= (mass of element in sample)/(total mass of sample)×100
32
Electron Configuration
1st shell = 2 electrons 2nd shell = 8 electrons 3rd shell = 8 electrons 4th shell = 2 electrons
33
JOHN DALTION
− Atomic theory:  Elements are composed of extremely small particles  All atoms of a given element are identical  Atoms are not created nor destroyed or changed  Chemical reaction involves only separation, combination, rearrangement of atoms  Compounds are formed when atoms of more than one element combine in a specific ration
34
J.J. THOMSON
− Plum pudding model (electrons embedded in the nucleus) − Discovered the electron − Electrons have a mass 1/1000th of the smallest atom known
35
ERNEST RUTHERFORD
− Gold foil experiment − Beam of alpha particles were targeted at a very thin sheet of gold foil − Alpha particles were deflected − Proposed that the atom consisted of mostly empty space occupied by electrons
36
SIR JAMES CHADWICK
− Discovered the neutron | − Accounted for all of the atoms mass
37
NIELS BOHR
− Quantum theory − Line spectra − Proposed that electrons moved about the central nucleus in circular orbits
38
SHAPES OF MOLECULES
− VSEPR: valence shell electron pair repulsion theory − Molecular shape results from repulsion between groups of electrons (bonded and non-bonded) in the valence shell of the central atom − Valence electron groups around central atom have greatest possible angle of separation
39
Linear
1 or 2 groups of electrons gives linear shape (1800 bond angle)
40
Triangular Planar
3 groups of electrons and no lone pairs gives triangular planar shape (1200 bond angle)
41
Tetrahedral
4 groups of electrons and no lone pairs gives tetrahedral shape (109.50 bond angle)
42
Pyramidal
4 groups of electrons (one being a lone pair) gives pyramidal shape (few degrees less than 109.50)
43
Bent or V-Shape
4 groups of electrons (two of these being lone pairs) give bent or V shape (several degrees less than 109.50)
44
Concentration and reaction rate
− Raising concentration increases reaction rate as a higher concentration of reacting particles cause and increase in the rate of collisions between the reacting particles − Doubling concentration, doubles the reaction rate
45
Gas Pressure and reaction rate
− Raising pressure (by reducing volume or adding more gas to the same container) creates greater concentration, causing an increase in rate of collision between reacting particles − Doubling pressure, doubles the reaction rate
46
Temperature and reaction rate
− Raising temperature means particles have a greater kinetic energy; higher percentages of collisions have energy equal to or greater than the activation energy. − Greater percentage of collisions are successful
47
State of subdivision and reaction rate
− Heterogeneous reactions involve reactants that are in two separate phases (solid/solid, solid/liquid, liquid/gas etc) − Particles can only collide at the surface boundary where separate phases make contact − Increasing SA exposes greater amount of reaction particles to the possibility of collision − Increases rate of collisions between particles, hence increase in rate of collision
48
Catalysts and reaction rate
− Catalyst: substance that have the ability to speed up chemical reactions while they remain chemically unchanged at the end of the reaction − Many transition metals such as Mn, Pt, Pd, Au, Rh are catalytic − Increase rate of reaction as they PROVIDE A REACTION PATHWAY WITH A LOWER ACTIVATION ENERGY − Greater percentage of collisions will have equal to or greater than activation energy; greater percentage of collisions are successful
49
Enzymes
− Biological catalysts | − Very specific in reactions they catalyse and fast acting
50
Alkanes
− CnH2n+2 − Single bonds only − Saturated hydrocarbons − Suffix –ane
51
Alkanes: structural isomers
− Isomers are compounds that have the same molecular formula but different structural formula − Similar chemical and physical properties but not identical
52
Alkenes
− CnH2n − Contain a double bond − Unsaturated hydrocarbons − Suffix –ene
53
Alkenes: geometric isomers
− Same molecular formula and structural formula, ie same sequence of bonding between their atoms but DIFFERENT GEOMETRY − This is a result of the inability of double bonded carbon atoms to rotate − Trans: means across (diagonal from each other) − Cis: same side
54
Alkyne
− CnHn − Contain a triple bond − Unsaturated − Suffix –yne
55
Cycloalkanes
− CnH2n − 3 or more carbon atoms arranged into a ring − Every carbon atom is bonded to 2 hydrogen atoms and 2 carbon atoms − Prefix cyclo- is added to stem name − Can have side chains
56
Benzene
− Molecular formula of C6H6 − Carbon atoms are alternately double bonded − Flat hexagonal ring − Occurs naturally in coal and crude oil and is produced in burning of natural materials − Improves octane rating of fuels
57
Alkyl function groups
``` Alkyl functional groups Methyl………..CH3 Ethyl…………..CH3CH2 Propyl………...CH3CH2CH2 Butyl…………..CH3CH2CH2CH2 Pentyl…………CH3CH2CH2CH2CH2 ```
58
IUPAC nomenclature for hydrocarbons
1. Find longest continuous chain of hydrocarbons to get stem name ie. Prop, meth, eth 2. Number the carbon atoms of the chain sequentially according to the principal functional group Priority goes: type of bond (if double bond is present), halogen (F, Cl, Br, I) then alkyl groups (methyl, ethyl) 3. The principle functional group determines the suffix of the stem name 4. Prefixes are ordered alphabetically and numbered according to the carbon atom to which they are attached. Prefixes di, tri, tetra are used if multiples of a group are present (DO NOT AFFECT ALPHABETICAL ORDER). Numbers separated from letters with a dash and a comma is used to separate numbers
59
Endothermic
− gain heat to the surroundings − positive value for ∆H ie enthalpy increases during reaction − produce freezing temperatures (feel cold) ie sport cool packs
60
Exothermic
− lose heat to the surroundings − negative value for ∆H ie enthalpy decreases during reaction − Potential chemical energy stored in bonds is converted to particle kinetic energy − Produce heating effect (feel warm) ie sports heat packs
61
Energy effects making and breaking bonds
− Bond breaking: endothermic, as it requires an input of energy − Stronger the bonds, the more energy to be absorbed − Bond making: exothermic, as it releases energy − Stronger the bonds, the more energy is released when forming them
62
Reaction rate
reaction rate= (amount of substance used or produced)/(time taken)
63
Collision Theory
− describes chemical changes in terms of collisions between individual reacting particles − if reacting particles collide with sufficient energy and suitable orientation then they can form a transition state − Transition state: where original bonds break and new ones form  Individual particles of the reacting substances must collide  The collision energy must be equal to or greater than a certain minimum amount known as the activation energy, (Ea)  The reacting particles must collide with a suitable orientation
64
Intramolecular Forces
− Bonds that keep atoms clustered together within the molecule − Strong IONIC, COVALENT ETC
65
Intermolecular Forces
− Bonds that keep molecules bonded to one another | DISPERSON, DIPOLE DIPOLE
66
Dipole-Dipole
− POLAR MOLECULES are able to attract one another due to the weak electrostatic attraction between their dipoles − Dipole-dipole attraction occurs between the oppositely charged ends of polar molecules − They increase melting points and boiling points − Like dissolves in like
67
Dispersion
− ALL MOLECULES and are the only intermolecular force present in non-polar substances − Weak for molecules with few electrons and show increasing strength for molecules with a greater number of electrons (POWER IN NUMBERS YO) − Stronger in molecules whose shape allows maximum surface contact between molecules ie linear
68
Hydrogen Bonding
− Only occur in molecules having an H – F, H – O or H – N arrangement of atoms − Due to the electrostatic attraction between lone pair of electrons of a F, O or N atom and an H atom already covalently bonded to another F, O or N atom − Melting and boiling points are much higher for substances that can hydrogen bond than for those of similar molecular size that cannot hydrogen bond − Substances which experience hydrogen bonding between molecules are extremely soluble in other substances that can also hydrogen bond
69
Solutions and Ion formation
− STRONG ELECTOLYTES: when these dissolve in water they are entirely present as independent mobile ions − WEAK ELECTROLYTES: when these dissolve in water they are partly present as independent mobile ions but mostly as molecules − NON-ELECTROLYTES: when these dissolve in water they do not produce independent mobile ions. They are entirely present as molecules − Strong acids = strong electrolytes − Weak acids and bases = weak electrolytes
70
Acids
− Conduct electricity − Turn blue litmus red − Sour taste
71
Bases
− Conduct an electric current − Turn red litmus blue − Bitter taste − Soapy slippery feel
72
ACID + METAL
SALT + H2
73
ACID + METAL HYDROXIDE
SALT + H2O
74
ACID + METAL OXIDE
SALT + H2O
75
ACID + CARBONATE
SALT + H2O + CO2
76
ACID + HYDROGENCARBONATE
SALT + H2O + CO2
77
ACID + METAL SULFITE
SALT + H2O + SO2
78
BASE + AMMONIUM SALT
SALT + H2O + NH3
79
BASE + NON-METAL OXIDE
SALT + H2O
80
Non-metal oxides that form an acid in water
CO2 (carbon dioxide)…………..H2CO3 (carbonic acid) SO2 (sulphur dioxide)………….H2SO3 (sulphurous acid) SO3 (sulphur trioxide)…………H2SO4 (sulphuric acid) NO2 (nitrogen dioxide)………...HNO3 and HNO2 (nitric and nitrous acid) P4O10 (phosphorous oxide)…H3PO4 (phosphoric acid)
81
Ionisation of water
H2O(l) ⇄ H+(aq) + OH-(aq) | 2H2O(l) ⇄ H3O+(aq) + OH-(aq)
82
Dissociation
− BASES DISSOCIATE RELEASING OH- IONS | − When ions present in the ionic solid are released into water to form a solution of independent mobile ions
83
Ionisation
− ACIDS IONISE PRODUCING H+ IONS | − Molecules react with water to form ions not originally present within the substance
84
Arrhenius Theory
− Acid-base behaviour centred on the ability of certain substances that contain H or OH to produce H+ or OH- ions when dissolved in water
85
Common Acids
``` H2O water H2O2 hydrogen peroxide NH3 ammonia NO nitric oxide H2CO3 carbonic acid CH3COOH acetic acid HCl hydrochloric acid H2SO4 sulfuric acid H3PO4 phosphoric acid NHO3 nitric acid C2H2 acetylene ```
86
Mass Spectrometry
Step 1: Ionisation − Atom or molecule is ionised by knocking one or more electrons off to give a positive ion (even for elements that would usually form negative ions or never form ions Step 2: Acceleration − Ions are accelerated so they have the same kinetic energy Step 3: Deflection − Ions are deflected by a magnetic field according to their masses (lighter they are the more they are deflected). Also depends on the number of positively charged ions (more the ion is charger the more it is deflected) Step 4: Detection − Beam of ions passing through the machine is detected electrically

Chemistry ATAR 11 (7 decks)

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