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Chemistry > IGCSE Chemistry > Flashcards

IGCSE Chemistry Flashcards

1
Q

Step 6 of extracting iron from bauxite (2)

A

molten slag removed - used for cement + roads

pure iron obtained

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2
Q

Step 5 of extracting iron from bauxite (4)

A

Limestone added to remove impurities

impurities react with limestone to produce molten slag

CaCO3 –> CaO + CO2

CaO + SiO2 –> CaSiO3 (slag)

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3
Q

Step 4 of extracting iron from haematite (2)

A

carbon reduces iron ore in hottest part of furnace

Fe2O3 (s) + 3C(s) –> 2FE(l) + 3CO (g)

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4
Q

Step 3 of extracting iron from bauxite (2)

A

carbon monoxide reduces iron oxide to iron by combining with oxygen in iron oxide

Fe2O3 + 3CO –> 2Fe + 3CO2

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5
Q

Step 2 of Extracting iron from haematite (2)

A

carbon monoxide produced in reduction reaction between carbon + carbon dioxide due to high temp.

CO2 + C –> 2CO

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6
Q

Step 1 of extracting Iron from haematite (3)

A

carbon combines with oxygen to from carbon dioxide in exothermic reaction

causes high temperature in furnace

C + O2 –> CO2

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7
Q

Main iron ore

A

Haematite

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8
Q

Extracting metals with heat and carbon (2)

A

slightly more reactive metals

ore dropped into fire –> heat + carbon caused metal to displace from compound –> carbon joins with oxygen

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9
Q

Extracting metals by heat

A

least reactive metals can be extracted using just heat

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10
Q

Define ore

A

naturally occuring metal compounds (usually with oxygen)

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11
Q

Reactivity Series (13)

A

Potassium
Sodium
Lithium
Calcium
Magnesium
Aluminium
Carbon
Zinc
Iron
Hydrogen
Copper
Silver
Gold

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12
Q

Copper reaction with water

A

no reaction

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13
Q

Magnesium reaction with water

A

reacts slowly with cold water

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14
Q

Calcium reaction with water

A

bubbles of gas immediately form

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15
Q

Alkali metal word equation with water

A

Alkali metal + water –> alkali metal hydroxide + hydrogen

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16
Q

Reaction of Potassium with water (2)

A

lilac flame produced

hydrogen gas produced

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17
Q

Reaction of Sodium with water (2)

A

fizzes around water

gives off hydrogen gas

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18
Q

Structure of Alloys (2)

A

alloys have different sized atoms

prevents layers sliding over each other –> stronger + less flexible

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19
Q

Reaction of Lithium with water

A

bubbles of hydrogen gas given off

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20
Q

Solder Alloy (2)

A

Tin + Lead

Lower melting point

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21
Q

Stainless Steel (2)

A

Iron + Carbon + Chromium + Nickel

resistant to corrosion

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22
Q

Cupronickel (2)

A

Copper + Nickel

cheaper than silver –> used for coins

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23
Q

Bronze Alloy (2)

A

Copper + Tin

Hard

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24
Q

Brass Alloy (2)

A

Copper + Zinc

Easier to shape

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25
Q

Hard Steel/High Carbon Steel Alloy (2)

A

Iron + 2.5% of carbon

hard

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26
Q

Mild Steel Alloy (2)

A

Iron + 0.25% of Carbon

can be easily shaped

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27
Q

Why are metals malleable and ductile (3)

A

metallic bonds not as strong

so layers of ions able to slide over each other

ions can move into different positions when hammered

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28
Q

Define metallic bonding (3)

A

metal atoms will give up their electrons to form cations

free electrons form “sea” + surround cations –> positive + negative attract holding bond together

strong attraction between closely packed cations and a “sea” of delocalised electrons

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29
Q

Properties of transition metals (6)

A

most have high density

high melting points

not very reactive

forms coloured compounds

used as catalysts

more than 1 valence electron

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30
Q

Properties of noble gases (2)

A

inert/unreactive

monoatomic

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31
Q

Define displacement

A

more reactive element takes less reactive element’s place

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32
Q

Halogen properties (4)

A

diatomic

colour becomes darker as you go down group

density increases as you go down group

melting point increases as you go down group

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33
Q

Colour of iodine solution (reaction with water)

A

yellow/brown

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34
Q

Colour of bromine solution (reaction with water)

A

orange

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35
Q

Colour of chlorine solution (reaction with water)

A

pale yellow

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36
Q

Halogen colours (4)

A

fluorine - pale-yellow gas

chlorine - pale-green gas

bromine - red-brown liquid

iodine - black, shiny solid

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37
Q

How alkali metals react with water (4)

A

react vigorously

produce effervescence

float on surface, moving rapidly

make water alkaline

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38
Q

Alkali metal reaction with halogen

A

metal + halogen –> metal halogen(ide)

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39
Q

Alkali metal reaction with water

A

alkali metal + water –> metal hydroxide + hydrogen

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40
Q

Alkali metal reaction with air

A

metal + oxygen –> metal oxide

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41
Q

colour of potassium flame

A

lilac

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42
Q

colour of sodium flame

A

orange/yellow

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43
Q

Colour of lithium flame

A

red

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44
Q

Density trend of alkali metals

A

density increases as you go down

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45
Q

Melting point trend of alkali metals

A

decreases as you go down

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46
Q

Properties of group 1 metals (5)

A

soft to cut

shiny but react with air quickly

relatively low melting/boiling points

relatively low density

reactive

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47
Q

Volume which a mole of gas occupies in room temperature/pressure

A

24dm^3

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48
Q

Formula for volume of gas under room pressure/temperature

A

moles x 24dm^3

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49
Q

Formula for number of moles (2)

A

mass(g)/Mr (relative atomic mass)

concentration(mol/dm^3) x volume (dm^3)

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50
Q

Avogadro’s constant

A

6.02 x 10^23 particles

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51
Q

Define a mole (2)

A

amount of substance

6.02 x 10^23 particles

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52
Q

Why non-metals less reactive as they go down a group (4)

A

atom bigger as go down

electrons farther from nucleus

weaker electrostatic force of attraction

harder for atom to accept electrons

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53
Q

What do elements in the same period have the same of

A

same number of electron shells

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54
Q

Properties of non-metals (5)

A

smaller atomic radius

no conduction

low melting/boiling point

brittle

dull

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55
Q

Properties of metals (6)

A

large atomic radius

conduct heat/electricity

shiny

malleable

shiny

high melting/boiling point

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56
Q

Define metalloids

A

elements with characteristics of metals and non-metals

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57
Q

Why metals become more reactive going down a group (4)

A

atom bigger as go down

electrons farther from nucleus

electrostatic force of attraction weaker

easier for valence electron to be removed

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58
Q

Define groups in a periodic table

A

vertical columns

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59
Q

Periods in a periodic table

A

horizontal row

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60
Q

Define periodicity

A

trends in periodic table

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61
Q

Why do giant covalent bonds have high melting and boiling points (2)

A

every strong covalent bond must be broken

hence require lots of energy

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62
Q

Structure of giant covalent structures (2)

A

have strong covalent bonds - no weak

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63
Q

Melting and boiling points of giant covalent structures

A

high melting and boiling points

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64
Q

Are covalent substances insoluble or soluble in water

A

insoluble

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65
Q

Do covalent substances conduct electricity why? (3)

A

do not conduct electricity as no free electrons to carry charge

do not have ions

exceptions - some covalent substances will form ions when dissolved in water

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66
Q

What are covalent bonds said to be

A

volatile since they have low melting and boiling points

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67
Q

Define volatility

A

how easily substances vaporise

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68
Q

Why do covalent bonds have low melting and boiling points (2)

A

covalent bonds (intramolecular forces) not broken when substance melted or boiled

much weaker forces of attraction (intermolecular forces) broken

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69
Q

Melting and boiling point of covalent bonds

A

low melting and boiling points

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70
Q

Number of covalent bonds for silicon (IV) oxide

A

4

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71
Q

Examples of macromolecules (3)

A

silicon (IV) oxide

diamond

graphite

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72
Q

Why is graphite able to conduct electricity (2)

A

has delocalised electrons (electron left from not having 4th bond)

electron is free to move and allows conduction

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73
Q

Can graphite conduct electricity

A

yes

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74
Q

Can diamond conduct electricity

A

no

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75
Q

Properties of graphite (4)

A

strong bonds between atoms

weak intermolecular forces between layers - layers can slide over each other

flaky and can be used as lubricant

conducts electricity

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76
Q

Properties of diamond (3)

A

strong covalent bonds

high melting point - lots of energy needed to break strong bonds between atoms

hard substance

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77
Q

Graphite structure (3)

A

carbon atoms in hexagonal shape in layers

each carbon atom bonds to 3 other carbon atoms

layers held together by weak intermolecular forces of attraction

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78
Q

Diamond structure (3)

A

carbon atoms arranged in tetrahedral structure

carbon bonded to 4 other carbon atoms

held together by strong covalent bonds

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79
Q

2 giant covalent structures of carbon

A

diamond

graphite

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80
Q

Define allotropes

A

different forms of same element

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81
Q

Define macromolecules

A

molecules that contain large number of atoms

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82
Q

How is a covalent bond held together (3)

A

positively charged protons in the nucleus

negatively charged electrons in shared area

electrons and protons attracted as opposites attract

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83
Q

Group 8 valency

A

0

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84
Q

Group 7 valency

A

1

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85
Q

Group 6 valency

A

2

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86
Q

Group 5 valency

A

3

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87
Q

Group 4 valency

A

4

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88
Q

Group 3 valency

A

3

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89
Q

Group 2 valency

A

2

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90
Q

Group 1 valency

A

1

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91
Q

Define valency

A

number of bonds an element needs to form

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92
Q

Example of triatomic molecule

A

H2O

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93
Q

Define triatomic molecule

A

molecule made up of 3 atoms bonded together

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94
Q

Define a diatomic molecule

A

molecule made up of 2 atoms bonded together

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95
Q

Define a covalent bond

A

pair of electrons shared by two atoms

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96
Q

When will a covalent bond form with

A

non-metals

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97
Q

What determines strength of ionic bond

A

greater charge greater strength

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98
Q

Features of ionic substances (5)

A

hard crystals - strong attraction between positive and negative ions

high melting points - lot of energy needed to overcome forces of attraction between ions

dissolves in water - water forms electrostatic attractions with ions and “pluck” them off lattice

does not conduct when solid - electrons held tightly in lattice and cannot move

conducts when molten/solution - lattice breaks and ions free to move

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99
Q

Can non-metals and covalent substances (except graphite) conduct electricity

A

no

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100
Q

At which states do metals conduct (2)

A

solid

liquid

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101
Q

Define an electrolyte

A

molten substance or solution that conducts

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102
Q

In metals and graphite which part of the atom moves and carries an electric current

A

outer electrons of atoms (delocalised electrons) which can move and carry current

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103
Q

Define electricity

A

flow of charged particles (ions or electrons)

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104
Q

Define reduction

A

when an atom/ion gains electrons

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105
Q

Define oxidation

A

when an atom/ion loses electrons

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106
Q

Ionic lattice structure

A

ion surrounded by ions of opposite charge

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107
Q

Name given to pattern which ionic compounds are organised in

A

ionic lattice

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108
Q

Define an ionic bond

A

electrostatic force of attraction between oppositely charged ions

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109
Q

How do ions form (3)

A

metal atoms give electrons to non-metal atoms

both atoms are stable charged particles

atoms attracted as they are oppositely charged

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110
Q

Which substances will an ionic bond form between (2)

A

metal

non-metal

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111
Q

Define anions

A

negative ions

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112
Q

Define cations

A

positive ions

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113
Q

What is the charge of an atom which has gained electrons

A

negative charge

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114
Q

Which atoms tend to lose electrons

A

metals

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115
Q

What is the charge of an atom which has lost electrons

A

positive charge

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116
Q

Properties of Isotopes

A

have same chemical properties as they have same number of valence electrons

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117
Q

Define an isotope

A

atoms of same element with same atomic number but different mass number due to different neutron number

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118
Q

What happens to the number of electrons as we move along a period

A

number of electrons in outer shell increases

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119
Q

What do elements in the same group have and why (2)

A

similar chemical properties

due to same number of valence electrons

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120
Q

Name of outer electrons

A

valence electrons

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121
Q

Features of noble gases (2)

A

full outer electron shell

are stable as electron shell full

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122
Q

How are electrons arranged (4)

A

arranged in shells

first shell holds 2 electrons

2nd shell holds 8 electrons

3rd shell holds 8 electrons

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123
Q

What is most of an atom

A

empty space

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124
Q

What does nucleon/mass number show

A

number of protons and neutrons an atom has

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125
Q

Define periodicity

A

trend that can be seen in periodic table

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126
Q

What happens to proton/atomic number when you move from left to right across a period

A

proton/atomic number increases

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127
Q

What does proton/atomic number show

A

number of protons an atom has

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128
Q

What does nuclide notation show (3)

A

element symbol

proton/atomic number

nucleon (mass) number

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129
Q

What is the charge of atoms (2)

A

atoms have no charge (neutral)

number of protons equal to number of electrons

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130
Q

Features of neutrons (3)

A

found in nucleus of atom

have no charge and are neutral (0)

have relative mass of 1 amu

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131
Q

Features of electrons (3)

A

found in electron shells orbiting nucleus

have negative charge (-1)

have relative mass of 1/2000

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132
Q

Features of protons (3)

A

found in nucleus

have a positive charge (+1)

have relative mass of 1 atomic mass unit (amu)

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133
Q

Name of centre of atom

A

nucleus

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134
Q

What 3 sub-atomic particles are in an atom (3)

A

proton

neutron

electron

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135
Q

Define concentration of a solution (2)

A

amount of solute dissolved in certain volume

can be expressed as number of moles of solute in one decimetre cubed of solution

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136
Q

Define solvent

A

substance which does the dissolving

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137
Q

Define solute

A

substance that is dissolved

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138
Q

Define a solution

A

mixture formed when one substance dissolves fully into another

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139
Q

Examples of alloys (3)

A

steel

bronze

brass

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140
Q

Define alloys

A

substances formed when metals mix with other elements

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141
Q

Define metalloids

A

elements that exhibit characters of both non-metals and metals

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142
Q

Exceptions to metal and non-metal properties (4)

A

carbon/silicon = high melting points

carbon conducts electricityas graphite

mercury = low melting point

alkali metals = soft, have low melting points, not sonorous

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143
Q

Properties of metals (7)

A

shiny/lustrous

good conductors (heat and electricity)

high melting and boiling points

high tensile strength

sonorous (rings when struck)

ductile (can be stretched into wire)

malleable (can be beaten into shape)

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144
Q

Define mixture

A

contains elements and compounds not chemically joined together

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145
Q

Chemical equation of Hydrated cobalt(II) chloride <-> Anhydrous cobalt(II) chloride + Water

A

CoCl2 6H2O(s) -> CoCl2(s) + 6H2O(l)

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146
Q

Word equation of Hydrated copper(II) sulfate crystals <-> Anhydrous copper(II) sulfate + Water

A

CuSO4 5H2O(s) -> CuSO4(s) + 5H2O(l)

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147
Q

Define an element

A

a substance which contains only one type of atom

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148
Q

Colour change from hydrated cobalt chloride crystals to anhydrous cobalt chloride

A

pink crystals to blue

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149
Q

Colour change from hydrated copper sulfate crystals to anhydrous copper sulfate

A

blue crystals to white powder

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150
Q

Name reversible reactions (2)

A

Hydrated copper(II) sulfate crystals <-> Anhydrous copper(II) sulfate + Water

Hydrated cobalt(II) chloride <-> Anhydrous cobalt(II) chloride + Water

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151
Q

Name indicators of chemical reaction (4)

A

colour change

gas produced

solid formed

temperature change

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152
Q

Define a physical change (3)

A

involve a change of state

no new substance is made

easily reversible

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153
Q

Define a chemical reaction (2)

A

reaction which occurs when 2 or more substances react to form a new substance

difficult to reverse/irreversible

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154
Q

Determinants of rate of diffusion (2)

A

molecular mass - larger molecular mass is slower

vacuum - faster as no other particles to collide with

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155
Q

Define diffusion (2)

A

movement of particles through a liquid or gas

movement of particles from an area of high concentration to an area of low concentration

156
Q

Define an ion

A

charged particles formed by atom gaining or losing electron

157
Q

Define molecules

A

particles of atoms joined together by covalent bonds

158
Q

How process of evaporation works (4)

A

liquid particles have some particles which move faster

faster particles have more kinetic energy

hence they have enough energy to escape liquid surface as gas

not all particles evaporate as boiling point not reached

159
Q

Process of change from liquid to solid

A

freezing

160
Q

Process of change from gas to liquid

A

condensing

161
Q

Process of change from liquid to gas

A

boiling

162
Q

Process of change from solid to liquid

A

melting

163
Q

Define boiling point

A

temperature of substance when it changes from liquid to a gas

164
Q

Define melting point

A

temperature of substance when it changes from solid to liquid

165
Q

Which state can be compressed and why (2)

A

gas

particles are spread out

166
Q

Which states cannot be compressed and why (3)

A

solids

liquids

particles already packed together

167
Q

Particle arrangement of gas (3)

A

spaced far apart

particles move about freely

spread to fill container they are in

168
Q

Particle arrangement of liquid (4)

A

packed tightly together

particles have strong forces of attraction

particles can move over one another

liquids take shape of their container

169
Q

Particle arrangement of solid (3)

A

packed tightly together

has strong forces of attraction holding particles together

particles only vibrate on the spot

170
Q

Which state of matter flows (2)

A

liquid

gas

171
Q

Which state of matter has a fixed volume (2)

A

solid

liquid

172
Q

Which state of matter has a fixed shape

A

solid

173
Q

Define a property

A

characteristic of something

174
Q

Define compounds

A

two or more types of atoms chemically joined together

175
Q

Define atoms

A

simplest particles that make up everything

176
Q

Define a particle

A

simplest and smallest unit of matter

177
Q

State symbol of gas

A

g

178
Q

State symbol of liquid

A

l

179
Q

State symbol of solid

A

s

180
Q

Name the states of matter (3)

A

solids

liquids

gases

181
Q

Filtration function

A

separates insoluble solid from solution

182
Q

SI measurement for temperature

A

kelvin (K)

183
Q

SI measurement for electric current

A

ampere (A)

184
Q

SI measurement for amount of substance

A

mole (mol)

185
Q

SI measurement for time

A

second

186
Q

SI measurement for length

A

metre

187
Q

SI measurement for mass

A

kilogram

188
Q

What does 1 L equal to in cm cubed

A

1000

189
Q

What does 1 L equal to in decimetres cubed

A

1

190
Q

What does 0K equal to in degrees celsius

A

-273 degrees celsius

191
Q

Purpose of paper chromatography (2)

A

used to test for impurities in substance

impure substance has more spots than expected

192
Q

Conditions for keeping Rf values the same (3)

A

same type of chromatography

same conditions

same solvent used

193
Q

Rf formula

A

distance moved by spot from baseline/distance moved by solvent front from baseline

194
Q

Purpose of Rf values (2)

A

substances have unique Rf values

identify unknown Rf value to known Rf values to find substance

195
Q

What is responsible for how far spots travel (2)

A

mass of spots - smaller mass greater distance

solubility of spots - more solubility greater distance

196
Q

Setting up chromatography procedure (4)

A

solvent used must start below baseline - otherwise spots travel into solvent

baseline done in pencil

paper sits vertically and is upright by a wooden splint

spots travel up stationary phase at different rates causing them to separate

197
Q

Define the mobile phase in chromatography (2)

A

part that moves

e.g solvent which moves spots up paper

198
Q

Define stationary phase in chromatography (2)

A

part that doesn’t move

e.g filter paper

199
Q

Chromatography purpose

A

used to separate small quantities of substance from a solution

200
Q

Where is fractional distillation used (2)

A

crude oil industry to separate crude oil

separated into petrol, diesel and kerosene

201
Q

How fractional distillation works (5)

A

same set up with distillation

components of mixture have similar boiling points and vaporise at same time

components travel up flask and meet cool glass beads

glass beads condense liquid with higher boiling point and cause it to drop back in round bottom flask

liquid with lower boiling point travels along condenser

202
Q

Purpose of fractional distillation (2)

A

separate more complex liquid mixtures

used when difference in boiling point of separate components is roughly 25 degrees C

203
Q

How distillation works (3)

A

solvent boils in solution, turning into a gas

travels into condenser where it meets cool atmosphere and condenses into liquid

solute remains in flask while solvent collected as distillate

204
Q

Function of distillation (3)

A

separate liquid from soluble solid

obtain liquid solvent

separates mixture of liquids with different boiling points (however temperature difference between mixtures must be 50 degrees C or above)

205
Q

How does crystallisation work (3)

A

crystals form when temperature of solution decreases

as a result, solid particles have less space in the liquid

solid particles are pushed out of solution and grow as crystals

206
Q

Function of crystallisation

A

allows pure crystals of solid to be obtained from solution

207
Q

Are crystals usually soluble

A

yes

208
Q

Examples of crystals (2)

A

salt

copper sulfate

209
Q

Define crystals

A

small, regularly shaped solids which reflect light

210
Q

Why is purity important (2)

A

important for food and pharmaceutical industries

impurities can be harmful and need to be checked

211
Q

What happens during melting of an impure substance (3)

A

substance begins to melt before melting point reached

other substances get in between original substance, breaking intermolecular forces

substance begins turning into a liquid as intermolecular forces begin to break

212
Q

What happens during melting of a pure substance (3)

A

particles arranged in regular pattern held together by intermolecular forces

intermolecular forces begin to break as substance begins to melt

molecules aren’t held as tightly as before which disorders arrangement of atoms

213
Q

Impure solid melting and boiling point compared to pure solid

A

impure solids will have a lower melting point

214
Q

Impure substance molecular structure (2)

A

contains more than one type of particle

original arrangement of pure substance altered

215
Q

Pure substance molecular structure

A

specific arrangement of atoms/molecules

216
Q

What is the melting/boiling point for an impure substance (2)

A

differing melting/boiling point

separate components melt at different temperatures

217
Q

Melting/boiling point for a pure substance (2)

A

sharp melting/boiling point

set temperature at which IMFs break (as they are all the same)

218
Q

Define an impure substance

A

substance which has more than one type of particle

219
Q

Define a pure substance

A

substance having only one type of particle

220
Q

Function of evaporation (2)

A

separates soluble solid from liquid

liquid evaporated to leave soluble solid behind

221
Q

Define filtrate

A

liquid seperated

222
Q

Define residue

A

insoluble solid left after filtration

223
Q

How does filtration work (2)

A

insoluble solid get stuck in filter paper

liquid easily passes through

224
Q

Define accurate results

A

results close to actual value

225
Q

Define precise results

A

set of results close to each other

226
Q

Burettes function (4)

A

used in titration reactions

allow volumes of liquid to be added to another

addition of liquid controlled by tap

scientist able to measure how much liquid needed in reaction

227
Q

Pipette function

A

transfers volume of liquid from one vessel to another

228
Q

Measuring cylinder function

A

measure volume of liquid to be used in reaction

229
Q

Define a fair investigation

A

investigation in which only one variable is changed (independent variable)

230
Q

Define a control variable

A

factors controlled during investigation

231
Q

Define the dependent variable (4)

A

variable measured during investigation

is effect

is on y-axis

is on right column of table

232
Q

Define the independent variable (4)

A

variable deliberately changed during investigation

is cause

is on x-axis

is on left column for table

233
Q

Percentage by mass formula

A

m/Mr x 100

234
Q

Purpose of percentage of mass calculations

A

determine percentage of certain element in compound

235
Q

Define relative molecular mass

A

overall mass of compound

236
Q

Define relative atomic mass (3)

A

measured in relation to carbon-12 atom : because carbon-12 mass is exactly 12

mass number on periodic table

average mass of naturally occurring isotopes of element

237
Q

Unit given to mass of atom (3)

A

called atomic mass unit

symbol : u

can also be called Dalton (Da)

238
Q

Define spectator ions (2)

A

ions present in reaction but don’t take part

appear unchanged on both sides of chemical equation

239
Q

Why are chemical equations balanced

A

law of conservation of mass - mass cannot be created or destroyed

240
Q

State symbol for solution (aqueous)

A

aq

241
Q

State symbol for gas

A

g

242
Q

State symbol for liquid

A

l

243
Q

State symbol for solid

A

s

244
Q

Name the Diatomic elements (7)

A

Hydrogen

Nitrogen

Fluorine

Oxygen

Iodine

Chlorine

Bromine

245
Q

Define a chemical equation

A

names of chemicals within word equations replaced with chemical formulae

246
Q

Order of word equations (3)

A

reactants start

products end

seperated by arrow

247
Q

Definition of products

A

new substances formed during reaction

248
Q

Definition of reactants

A

chemicals at the start of the reaction

249
Q

Phosphate radical

A

PO 3- 4

250
Q

Sulfate radical

A

SO 2- 4

251
Q

Carbonate radical

A

CO 2- 3

252
Q

Ammonium radical

A

NH + 4

253
Q

Nitrate radical

A

NO - 3

254
Q

Hydrogen carbonate radical

A

HCO - 3

255
Q

Hydroxide radical

A

OH -

256
Q

What is the overall charge of an ion

A

0

257
Q

Name the radicals (7)

A

Hydroxide

Hydrogen Carbonate

Nitrate

Ammonium

Carbonate

Sulfate

Phosphate

258
Q

What is a radical (2)

A

group of elements bonded together

exist as ions

259
Q

When is radical method used

A

when compound has more than 2 elements

260
Q

What does a roman numeral signify of an element

A

the valency of the element

261
Q

When is valency method used

A

when compound has no prefix

262
Q

prefix for 6

A

hexa

263
Q

prefix for 5

A

penta

264
Q

prefix for 4

A

tetra

265
Q

prefix for 3

A

tri

266
Q

prefix for 2

A

di

267
Q

Prefix for one

A

mono

268
Q

How does a prefix help in a chemical formula

A

prefix helps to show the number of atoms for an element

269
Q

Cations and Anions order in compound name (2)

A

cations come first

anions come second

270
Q

Exception to rules (Ammonium) (2)

A

Ammonium exists as a compound with 3 elements

however it ends with -ide

271
Q

Exception to rules (Hydroxide) (2)

A

hydroxide usually exists as a compound with 3 elements

however it ends with -ide

272
Q

How will the compound end if there are only 2 elements

A

second element will end in -ide

273
Q

Order of two elements in compound

A

element further to left on periodic table comes first in name

274
Q

How will oxygen be represented as if there are 3 elements in a compound and one of them is oxygen

A

compound name will end in ate or ide

275
Q

Uses of limestone, lime, slaked lime (6)

A

treat waste products

neutralise acidic soil

road building

cement

neutralise lake acidity

extract iron in blast furnace

276
Q

pH of limestone, lime, slaked lime

A

basic

277
Q

How is slaked lime produced

A

Calcium oxide (CaO) + Water (H2O) –> Calcium hydroxide Ca(OH)2

278
Q

Slaked lime chemical compound/formula (2)

A

calcium hydroxide

Ca(OH)2

279
Q

Define thermal decomposition

A

when one reactant is heated and 2 or more products are formed

280
Q

Word/chemical equation of calcium carbonate thermal decomoposition

A

Calcium carbonate (CaCO3) –> Calciunm oxide (CaO) + carbon dioxide (CO2)

281
Q

How is lime produced

A

thermal decomposition of limestone

282
Q

Lime chemical formula and compound (2)

A

calcium oxide

CaO

283
Q

Define limestone (2)

A

sedimentary rock

made of calcium carbonate (CaCO3)

284
Q

Conditions for turning sulfur dioxide into sulfur trioxide (4)

A

High temp. of catalyst (450 Celsius)

Excess air so sulfur dioxide converted into sulfur trioxide

Exothermic reaction - faster reaction (lower temp. produce more but reaction too slow)

Low pressure (less than 2 atm) - high pressure is wasteful

285
Q

Describe the contact process to manufacture sulfuric acid (5)

A

sulfur burned in oxygen to produce sulfur dioxide

sulfur dioxide mixed with excess air

then passed over vanadium oxide catalyst to form sulfur trioxide

sulfur trioxide dissolved in concentrated sulfuric acid to form oleum

water added to oleum to make concentrated sulfuric acid

286
Q

Uses of sulfur (4)

A

gunpowder

vulcanization (hardening) of rubber

fungicide

making of sulfuric acid

287
Q

Which compounds is sulfur found in (2)

A

sulfates

sulfides

288
Q

Characteristic of sulfur (2)

A

yellow substance

abundant non-metallic element

289
Q

Gas test for oxygen (O2)

A

relights glowing splint

290
Q

Gas test for Hydrogen (H2)

A

pops with lighted splint

291
Q

Gas test for Chlorine (Cl2)

A

Bleaches litmus paper

292
Q

Gas test for CO2

A

limewater turns milky

293
Q

Gas Test for Ammonia (NH3)

A

red litmus paper blue

294
Q

Positive result for iodide anion (I^-)

A

yellow precipitate

295
Q

Test for iodide anion (I^-) (2)

A

add dilute nitric acid

then add aqueous silver nitrate

296
Q

Positive result for Sulfate anion (SO4^2-) test

A

white precipitate

297
Q

Test for Sulfate (SO4 ^2-) (2)

A

add dilute HCl acid

add aqueous barium nitrate/chloride

298
Q

Positive result for nitrate anion (NO3 ^-) (2)

A

ammonia gas produced

red litmus paper turned blue

299
Q

Test for Nitrate anion (NO3^-) (3)

A

add aqueous sodium hydroxide

add aluminum foil (or powder)

warm

300
Q

Positive result for Bromine anion

A

cream precipitate

301
Q

Test for Bromide anion (Br^-) (2)

A

add dilute nitric acid

then add aqueous silver nitrate

302
Q

Positive result for Chloride anion test

A

white precipitate

303
Q

Test for Chloride anion Cl^- (2)

A

add dilute nitric acid

then add aqueous silver nitrate

304
Q

Positive result for carbonate anion (CO3 ^2-)

A

effervescence/carbon dioxide produced

305
Q

Test for carbonate anion (CO3 ^2-)

A

add dilute acid

306
Q

Flame test colour of copper cation

A

blue-green

307
Q

Flame test colour of potassium cation

A

lilac

308
Q

Flame test colour of sodium cation

A

yellow

309
Q

Flame test colour of lithium cation

A

red

310
Q

Zinc (II) reaction with aqueous ammonia cation test result (2)

A

white precipitate

soluble in excess - gives colorless solution

311
Q

Zinc (II) reaction with aqueous sodium hydroxide cation test result (2)

A

white precipitate

soluble in excess - gives colorless solution

312
Q

Iron (III) reaction with aqueous ammonia cation test result (2)

A

red-brown precipitate

insoluble in excess

313
Q

Iron (III) reaction with aqueous sodium hydroxide cation test result (2)

A

red-brown precipitate

insoluble in excess

314
Q

Iron (II) reaction with aqueous ammonia cation test result (2)

A

green precipitate

insoluble in excess

315
Q

Iron(II) reaction with aqueous sodium hydroxide cation test result (2)

A

green precipitate

insoluble in excess

316
Q

Copper(II) reaction with aqueous ammonia cation test result (2)

A

light blue precipitate

soluble in excess will give dark blue solution

317
Q

Calcium (Ca2+) reaction with aqueous ammonia cation test result

A

no precipitate or slight white precipitate)

318
Q

Copper(II) (Cu2+) reaction with aqueous sodium hydroxide cation test result (2)

A

light blue precipitate

insoluble in excess

319
Q

Calcium (Ca2+) reaction with aqueous sodium hydroxide cation test result (2)

A

white precipitate

insoluble in excess

320
Q

Ammonium (NH4+) reaction with aqueous sodium hydroxide cation test result (2)

A

ammonia gas produced when heated

turns red litmus blue

321
Q

Preparing soluble salt without acid (5)

A

if metal oxide/carbonate soluble in water titrate with acid

evaporate some water and leave to cool

if metal oxide/carbonate not soluble in water add excess metal oxide/carbonate to acid

filter to remove excess solid

evaporate some water and leave to cool

322
Q

Preparing soluble salt with dilute acid (4)

A

add excess solute to acid

excess solute removed by filtration

solution evaporated

evaporate some water and leave to cool

323
Q

How to prepare insoluble salts (2)

A

mix 2 solutions of cations and anions

filter, wash, dry

324
Q

Reactions to prepare soluble salt (4)

A

acid + metal

acid + metal oxide

acid + metal carbonate

acid + metal hydroxide

325
Q

Insoluble salts (5)

A

silver chloride/bromide/iodide

lead chloride

barium/calcium/lead sulfate

most carbonates

most hydroxides

326
Q

Soluble salts (6)

A

sodium/potassium/ammonium salts

all nitrate salts

most chloride/bromide/iodide salts

most sulfates

sodium/potassium/ammonium carbonate salts

sodium/potassium/ammonium hydroxide salts

327
Q

Titration calculation - mole ratios method (3)

A

write balanced equation

write mole relationship (ratio of n)

calculate number of moles (n = c x v)

328
Q

Titration calculations - power method (3)

A

write balanced equation

assign values (n of moles according to number above compound)

substitute into equation

329
Q

Titration equation (4)

A

(c1V1)/n1 = (c2V2)/n2

c = concentration

v = volume

n = number of moles

330
Q

Method of titrations (9)

A

pour solution 1 into small beaker

pipette set volume of solution 1 into conical flask

add indicator to flask

pour solution 2 into other beaker

use filter funnel to pour solution 2 in burette

place conical flask under burette

add solution2 to conical flask until change observed

record volume of solution 2 used (rough titre)

repeat

331
Q

Titration accuracy points (4)

A

reading from bottom of meniscus

indicators to determine end point

rough titres not used in average volume

volume measured using pipette and conical flask

332
Q

Equipment used in a titration (2)

A

burette

pipette

333
Q

Steps in carrying out a titration (2)

A

measured volume of solution added from burette to known volume of another solution in conical flask

continue until indication that reaction has ended (e.g pH)

334
Q

When is titration used (2)

A

determining concentration of unknown solution from concentration of known solution

reacting acids and metal hydroxides/alkalis

335
Q

Define amphoteric oxides (2)

A

oxides able to act as acid or base

will react with both acids and bases to produce salt + water

336
Q

Examples of neutral oxides (3)

A

carbon monoxide (CO)

dinitrogen oxide (N2O)

nitrogen monoxide (NO)

337
Q

Define neutral oxides

A

oxides with neither basic nor acidic properties

338
Q

Groups which metals have soluble basic oxides (2)

A

group 1

group 2

339
Q

pH of non-metal oxides

A

acidic oxides

340
Q

pH of metal oxides

A

basic oxides

341
Q

When is an oxide formed

A

when element combines with oxygen

342
Q

Define bases

A

proton (hydrogen ion) acceptors

343
Q

Define acids

A

proton (hydrogen ion) donors

344
Q

How can acidic soil be neutralised (2)

A

slaked lime Ca(OH)2 –> quicklime (CaO) + water

slaked lime neutralises soil

345
Q

How can acidic rain be caused

A

burning fossil fuels

346
Q

Define salt

A

ionic compounds made from cations/anions

347
Q

Word equation acid + metal carbonate

A

acid + metal carbonate –> salt + water + carbon dioxide

348
Q

Word equation of acid + metal hydroxide

A

acid + hydroxide –> salt + water

349
Q

Word equation of acid + metal oxide

A

acid + metal oxide –> salt + water

350
Q

Reaction word equation of acid + metal

A

acid + metal –> salt + hydrogen

351
Q

Instrument which is a more accurate way of determining pH

A

pH meter

352
Q

Alkaline/basic pH

A

greater than 7

353
Q

Neutral pH

A

7

354
Q

Acidic pH

A

less than 7

355
Q

Define the pH scale

A

measure of concentration of hydrogen ions present

356
Q

Phenolphtalein colour with acid and base

A

acid - colourless

base - pink

357
Q

Universal indicator paper colour with acid and alkali (4)

A

strong acid - red

weak acid - orange/yellow

strong base - purple

weak base - blue

358
Q

Methyl orange colour with acid and base (2)

A

acid - red

basic - yellow

359
Q

Litmus colour with acid and basic (2)

A

acid - red

basic - blue

360
Q

Testing indicators used (4)

A

Litmus

Methyl orange

Universal Indicator paper

Phenolphtalein

361
Q

Purpose of indicators

A

determine if substance is acidic or basic

362
Q

Where are bases used

A

soap

363
Q

Characteristics of bases (2)

A

corrosive

soapy

364
Q

Examples of common acids (5)

A

hydrochloric acid (HCL)

sulfuric acid (H2SO4)

nitric acid (HNO3)

ethanoic acid (CH3COOH)

citric acid (C6H8O7)

365
Q

Characteristics of acids (2)

A

sour tasting

corrosive

366
Q

Define an alkali

A

soluble base

367
Q

Example of a base

A

Ammonia (NH3)

368
Q

What ion do metal hydroxides contain

A

OH^- ions

369
Q

What ion do metal oxides contain

A

O^2- ion

370
Q

What are bases either (2)

A

metal oxides

metal hydroxides

371
Q

What ion do acids contain

A

hydrogen ion (H+)

372
Q

What happens if particles collide with energy less than the activation energy (2)

A

successful collision will not occur

particles will not react

373
Q

Enthalpy change for endothermic reaction

A

positive

374
Q

Endothermic reaction pathway diagram (2)

A

reactants start with potential energy

products take energy from surroundings

375
Q

Activation energy meaning

A

minimum amount of energy for chemical particles to collide

376
Q

Exothermic reaction pathway diagram (3)

A

reactants start with potential energy

activation energy required to complete reaction

products have less energy than reactants

377
Q

Enthalpy change for exothermic reaction

A

negative

378
Q

Define enthalpy (2)

A

measure of potential energy of a chemical

symbol H

379
Q

How do exothermic reactions release energy

A

during reaction, energy stored in reactant bond is released to surroundings

380
Q

Define surroundings

A

matter surrounding site of chemical reaction

381
Q

Define system

A

consists of particles involved in chemical reactions

382
Q

Chemical equation for butane bubble reaction

A

2C4H10 + 13O2 –> 8CO2 + 10H2O

383
Q

Word equation for butane bubbles reaction

A

Butane + oxygen –> carbon dioxide + water

384
Q

Define an endothermic reaction (3)

A

reaction which takes in thermal energy from surroundings

decreases temperature of surroundings

energy to break bonds > energy to form bonds

385
Q

Define an exothermic reaction (3)

A

reaction which releases thermal energy

increases temperature of surroundings

energy in making bonds > energy to break the bonds

Chemistry (5 decks)

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