Module 2 – Foundations In Chemistry Flashcards
1
Q
What was stated in Dalton’s atomic theory? (4)
A
● Atoms are tiny particles made of elements
● Atoms cannot be divided
● All the atoms in an element are the same
● Atoms of one element are different to those of other elements
2
Q
What did Thompson discover about electrons? (3)
A
● They have a negative charge
● They can be deflected by magnetic and electric field
● They have very small mass
3
Q
Explain the plum pudding model
A
Atoms are made up of negative electrons moving around in a sea of positive charge
4
Q
What were Rutherford’s proposal after the gold leaf experiment? (4)
A
● Most of the mass and positive charge of the atom are in the nucleus
● Electrons orbit the nucleus
● Most of atom’s volume is the space between the nucleus and the electrons
● Overall positive and negative charges must balance
5
Q
Explain the current model of the atom
A
● Protons and neutrons are found in the nucleus
● Electrons orbit in shells
● The nucleus is tiny compared to the total volume of an atom
● Most of the atom’s mass is in the nucleus
● Most of the atom is empty space between the nucleus and the electrons
6
Q
What is the charge of a proton?
A
1+
7
Q
What is the charge of an electron?
A
1-
8
Q
Which particle has the same mass as proton?
A
Neutron
9
Q
Which two particles make up most of atom’s mass?
A
Protons and neutrons
10
Q
Which letter is used to represent the atomic number of an atom?
A
Z
11
Q
What does the atomic number tell about an element?
A
Atomic number = number of protons in an atom
12
Q
Which letter represents the mass number?
A
A
13
Q
How is mass number calculated?
A
Mass number = number of protons + number of neutrons
14
Q
How to calculate the number of neutrons?
A
Number of neutrons = mass number - atomic number
15
Q
Define isotope
A
Atoms of the same element with different number of neutrons
16
Q
Why does different isotopes of the same element react in the same way?
A
● Neutrons have no impact on the chemical reactivity
● Reactions involve electrons, isotopes have the same number of electrons in the same arrangement
17
Q
What are ions?
A
Charged particles that is formed when an atom loses or gains electrons
18
Q
What is the charge of the ion when electrons are gained?
A
Negative
There is a positive charge when electrons are lost
E.g. 3+ ion has lost 3 electrons
19
Q
What is the unit used to measure atomic masses called?
A
Unified atomic mass unit, u
20
Q
Define relative atomic mass
A
The weighted mean mass of an atom of an element compared with one twelfth of the mass of an atom of carbon -12
21
Q
What is the unit of relative atomic mass?
A
No units
22
Q
Define relative isotopic mass
A
The mass of an atom of an isotope compared with one twelfth of the mass of an atom of carbon-12
23
Q
The relative isotopic mass is same as which number?
A
Mass number
24
Q
What two assumptions are made when calculating mass number?
A
- The contribution of the electron is neglected
- The mass of both proton and neutron is taken as 1.0 u
25
How to calculate the relative molecular mass and relative formula mass?
Both can be calculated by adding the relative atomic masses of each of the atom making up the molecule or the formula
26
What are the uses of mass spectrometry?
● Identify unknown compounds
● Find relative abundance of each isotope of an element
● Determine structural information
27
How does a mass spectrometer work?
● The sample is made into positive ions.
● They pass through the apparatus and are separated according to mass-to-charge ratio.
● A computer analyses the data and produces a mass spectrum.
28
How is the group number related to the number of electrons?
Group number = number of electrons in the outer shell
29
Does the group number indicates horizontal or vertical column in the periodic table?
Vertical column
30
Do metals usually gain or lose electrons?
Lose electrons (non metals generally gain electrons)
31
Which are the 4 elements that don’t tend to form ions and why?
The elements are beryllium, boron, carbon and silicon
Requires a lot of energy to transfer outer shell electrons
32
What are molecular ions?
Covalently bonded atoms that lose or gain electrons
33
What is the charge of an ammonium ion?
+1 → NH4+
34
What is the charge of a hydroxide ion?
-1 → OH-
35
What is the charge of a nitrate ion?
-1 → NO3-
36
What is the charge of a carbonate ion?
-2 → CO32-
37
What is the charge of a sulphate ion?
-2 → SO42-
38
What is an empirical formula?
Simplest whole number ratio of atoms of each element present in a compound
39
How to calculate empirical formula
● Divide the amount of each element by its molar mass
● Divide the answers by the smallest value obtained
● If there is a decimal, divide by a suitable number to make it into a whole number
40
What is the symbol for amount of substance?
n
41
What is the unit used to measure amount of substance?
Mole
42
What does the Avogadro constant represent?
The number of atoms per mole of the carbon-12 isotope
43
How to calculate the mass of 1 mole of the element?
Mass of 1 mole = relative atomic mass in grams
44
How to calculate moles when mass and molar mass are given?
Moles (mol) = mass (g) / molar mass (g mol-1)
45
State Avogadro’s law
Under the same temperature and pressure, one mole of any gas would occupy the same volume.
46
How much volume does a gas occupy, at room temperature and pressure?
24 dm3 or 24000 cm3
47
Define molar gas volume
The volume per mole of gas molecules
48
Why do different gas particles occupy the same volume?
The gas particles are very spread out, hence individual differences has no effect.
49
How to calculate moles when gas volume is given?
Moles (mol) = volume (dm3) / 24 Moles (mol) = volume (cm3) / 24000
50
What are the ideal ways in which gases behave?
● They are in continuous motion
● No intermolecular forces experienced
● Exert pressure when they collide with each other or container
● No kinetic energy is lost in the collisions
● When temperature increases, the kinetic energy of gases also increase
51
Write down the ideal gas equation (in words and symbols, including the units)
pV = nRT
Pressure (Pa) x volume (m3) = number of moles (mol) x gas constant x temperature (K)
52
1 atmospheric pressure is equal to how many pascal?
1 atm = 101325 Pa (101KPa)
53
0°C is equal to how many kelvin?
273 K
54
What does concentration of a solution mean?
It is the amount of solute dissolved in 1 dm3 of solvent
55
How do you calculate moles when concentration and volume are given?
Moles (mol) = concentration (mol dm-3) x volume (dm3)
56
What is a standard solution?
A solution of known concentration
57
Write down the steps to prepare a standard solution
1. Weigh the solute using the weigh by difference method
2. In a beaker dissolve the solute using the solvent
3. Pour the solution into a volumetric flask
4. Rinse the beaker using the solution and add it to the flask
5. Add solvent to the flask carefully until it reaches the graduation line
6. Mix the solution thoroughly to ensure complete mixing
58
What does the terms concentrated and dilute mean?
Concentrated - large amount of solute per dm3 of solvent
Dilute - small amount of solute per dm3 of solvent
59
What is a species in a chemical reaction?
Any particle that takes part in a reaction
60
What are the four common state symbols?
1. Solid (s)
2. Liquid (l)
3. Gaseous (g)
4. Aqueous (aq)
61
What does percentage yield mean?
The efficiency of which reactants are converted into products
62
What are the reasons for not obtaining 100% yield?
● Reaction may be at equilibrium
● The reactants may be impure
● Side reactions could happen
● Reactants or products may be left behind while transferring
● Loss of products during separation and purification
63
How is percentage yield calculated?
Percentage yield = (actual amount of product (mol) / theoretical amount of product (mol)) x 100
64
What does atom economy tell us about?
The proportion of desired products compared with all the products formed in the reaction
65
How is atom economy calculated?
Atom economy = (molecular mass of desired product / sum of the molecular masses of all products) x 100
66
Does 100% yield mean 100% atom economy?
No, even if all the reactants are converted into products, not all products of the reaction will be the required products
67
Which type of reaction has 100% atom economy?
Addition reactions (two or more reactants are combined to form a product)
68
When an acid is added to water what ion is released into the solution?
Hydrogen ion , H+
69
Define acid
Proton donor
70
Describe the dissociation of a strong acid
Full dissociation
71
Define base
Proton acceptor
72
What are alkalis?
Bases that can dissolve in water to form aqueous hydroxide ions
73
Which base is used to treat acid ingestion?
Magnesium hydroxide
74
What are amphoteric substances?
Substances that can act as acids and bases
75
What is formed when acids react with carbonate?
Salt, carbon dioxide and water
76
What is a salt?
A compound that is formed when H+ of an ion is replaced by a metal ion or positive ion
77
What is formed when acids react with metal oxide?
Salt and water
78
What is formed when acid reacts with alkali?
Salt and water
79
What is formed when acids react with metal?
Salt and hydrogen
80
Why are the products same when acids react with alkali or metal oxides?
Both alkali and metal oxides are types of bases
81
How are ammonium salts formed?
When acid reacts with aqueous ammonia
82
What are hydrated crystals?
A crystalline structure containing water
83
What does anhydrous crystals mean?
Crystalline form that contains no water
84
What does a dot formula indicate?
The amount of water present in a crystalline structure
85
Write down the methods to carry out a titration
1. Using a pipette, measure the volume of a solution
2. Add the solution into a conical flask and add an indicator into it
3. Add the other solution into a burette and record the volume
4. Slowly add the solution in the burette into the conical flask
5. Swirl the mixture continuously until the end point is reached
6. Repeat until concordant results are obtained
86
What is the colour of methyl orange in an acid, base and at end point?
● Acid - red
● Base - yellow
● End point - orange
87
What is an oxidation number?
The number of electrons an atom uses to bond with any other atom
88
What is the oxidation number of an uncombined element such as C, H, O2?
0
89
What is the oxidation number of combined oxygen such as in H2O?
-2
90
What is the oxidation number of oxygen in peroxides?
-1
91
What is the oxidation number of combined hydrogen such as in NH3, H2S?
+1
92
What is the oxidation number of hydrogen in metal hydrides such as LiH?
-1
93
What is the oxidation number of a simple ion?
Charge on the ion E.g Na+→ +1 ; Cl-→ -1
94
What is the oxidation number of combined fluorine such as in NaF, CaF2?
-1
95
When an element has more than one stable oxidation number how is it indicated?
Written as a Roman numeral
96
What is the oxidation number of Fe in iron (III) chloride?
+3
97
What are oxyanions?
Negative ions that have an element along with oxygen
98
What is the oxidation number of S in SO42-?
+6 Because combined oxygen has an oxidation number of -2.
4 x -2 = -8.
The charge on the compound is -2.
The sum of oxidation numbers must equal -2 So, -2 -(-8) = +6
99
Define oxidation in terms of electron transfer and oxidation number
Oxidation is
● Loss of electrons
● An increase in oxidation number
100
Define reduction in terms of electron transfer and oxidation number
Reduction is
● Gain of electrons
● A decrease in oxidation number
101
What is a redox reaction?
A reaction in which both oxidation and reduction takes place
102
What is the oxidation number of a metal?
0, because it is an uncombined element
103
What does the principal quantum number indicate?
The shell occupied by the electrons
104
What is a shell?
A group of orbitals with the same principal quantum number
105
How many electrons can the 1st shell hold?
2
106
How many electrons can the 2nd shell hold?
8
107
How many electrons can the 3rd shell hold?
18
108
How many electrons can the 4th shell hold?
32
109
What is an orbital?
A region around the nucleus that can hold up to two electrons with opposite spins
110
How many electrons can an orbital hold?
2
111
What are the 4 types of orbitals?
● s orbital
● p orbital
● d orbital
● f orbital
112
What is the shape of a s-orbital?
Spherical
113
What is the shape of a p-orbital?
Dumb-bell shape
114
How many orbitals are found in a S subshell?
1
115
How many electrons can be held in a S subshell?
2
116
How many orbitals does P subshell have?
3
117
How many electrons can be held in a P subshell?
6
118
How many orbitals are present in a D subshell?
5
119
How many electrons can be held in a d-sub shell?
10
120
How many orbitals are found in a F subshell?
7
121
How many electrons can fill F subshell?
14
122
When using ‘electrons in box’ representation, what shape is used to represent the electrons?
Arrows
123
What letter is used to represent the shell number?
n
124
From which shell onwards is S orbital present?
n = 1
125
From which shell onwards is P orbital present?
n = 2
126
From which shell onwards is D orbital present?
n = 3
127
From which shell onwards is F orbital present?
n = 4
128
What are the rules by which electrons are arranged in a shell?
● Electrons are added one at a time
● The lowest available energy level is filled first
● Each energy level must be filled before the next one can fill
● Each orbital is filled singly before pairing
● 4s is filled before 3d
129
Why does 4s orbital fill before 3d orbital?
4s orbital has a lower energy than 3d before it is filled
130
Which electrons are lost when an atom becomes a positive ion?
Electrons in the highest energy levels
131
What are the 3 main types of chemical bonds?
● Ionic
● Covalent
● Metallic
132
Define ionic bonding
The electrostatic attraction between positive and negative ions
133
Given an example of an ionically bonded substance
NaCl (Sodium Chloride - salt)
134
Define covalent bonding
Electrostatic attraction between a shared pair of electrons and the nuclei
135
Define metallic bonding
Electrostatic attraction between the positive metal ions and the sea of delocalised electrons
136
Electrons in which shell are represented in a dot and cross diagram?
The outer shell
137
Why does giant ionic lattices conduct electricity when liquid but not when solid?
In solid state the ions are in fixed positions and thus cannot move. When they are in liquid state the ions are mobile and thus can freely carry the charge
138
Giant ionic lattices have high or low melting and boiling point? Explain your answer
They have high melting and boiling point because a large amount of energy is required to overcome the electrostatic bonds
139
In what type of solvents do ionic lattices dissolve?
Polar solvents E.g water
140
Why are ionic compounds soluble in water?
Water has a polar bond. Hydrogen atoms have a delta+ charge and oxygen atoms have a delta - charge. These charges are able to attract charged ions
141
What is it called when atoms are bonded by a single pair of shared electrons?
Single bond
142
How many covalent bonds does carbon form?
4
143
How many covalent bonds does oxygen form?
2
144
What is a lone pair?
Electrons in the outer shell that are not involved in the bonding
145
What is formed when atoms share two pairs of electrons?
Double bond
146
What is formed when atoms share three pairs of electrons?
Triple bond
147
What is average bond enthalpy?
Measure of average energy needed to break the bond
148
What is a dative covalent bond?
A bond where both of the shared electrons are supplied by one atom
149
How are oxonium ions formed?
Formed when acid is added to water, H3O+
150
What does expansion of the octet mean?
When a bonded atom has more than 8 electrons in the outer shell
151
What are the types of covalent structure?
● Simple molecular lattice
● Giant covalent lattice
152
Describe the bonding in simple molecular structures?
Atoms within the same molecule are held by strong covalent bonds and different molecules are held by weak intermolecular forces
153
Why do simple molecular structures have low melting and boiling point?
Small amount of energy is enough to overcome the intermolecular forces
154
Can simple molecular structures conduct electricity?
No, they are non conductors.
155
Why do simple molecular structures not conduct electricity?
They have no free charged particles to move around
156
Simple molecular structures dissolve in what type of solvent?
Non polar solvents
157
Give examples of giant covalent structures
● Diamond
● Graphite
● Silicon dioxide, SiO2
158
List some properties of giant covalent structures
● High melting and boiling point
● Non-conductors of electricity, except graphite
● Insoluble in polar and non-polar solvents
159
How does graphite conduct electricity?
Delocalised electrons present between the layers are able to move freely carrying the charge
160
Why do giant covalent structures have high melting and boiling point?
Strong covalent bonds within the molecules need to be broken which requires a lot of energy
161
Draw and describe the structure of a diamond
3D tetrahedral structure of C atoms, with each C atom bonded to four others
162
What does the shape of a molecule depend on?
Number of electron pairs in the outer shell Number of these electrons which are bonded and lone pairs
163
What is the shape, diagram and bond angle in a shape with 2 bonded pairs and 0 lone pairs?
Linear 180°
164
What is the shape, diagram and bond angle in a shape with 3 bonding pairs and 0 lone pairs?
Trigonal planar 120°
165
What is the shape, diagram and bond angle in a shape with 4 bonded pairs and 0 lone pairs?
Tetrahedral 109.5°
166
What is the shape, diagram and bond angle in a shape with 5 bonded pairs and 0 lone pairs?
Trigonal bipyramid 90° and 120°
167
What is the shape, diagram and bond angle in a shape with 6 bonded pairs and 0 lone pairs?
Octahedral 90°
168
What is the shape, diagram and bond angle in a shape with 3 bonded pairs and 1 lone pairs?
Pyramidal 107°
169
What is the shape, diagram and bond angle in a shape with 2 bonded pairs and 2 lone pairs?
Non linear 104.5°
170
By how many degrees does each lone pair reduce the bond angle?
2.5°
171
Define electronegativity
The ability of an atom to attract the pair of electrons (the electron density) in a covalent bond
172
In which direction of the periodic table does electronegativity increase?
Top right, towards fluorine
173
What does it mean when the bond is non-polar?
The electrons in the bond are evenly distributed
174
What is the most electronegative element?
Fluorine
175
How is a polar bond formed?
Bonding atoms have different electronegativities
176
Why is H2O polar, whereas CO2 is non polar?
CO2 is a symmetrical molecule, so there is no overall dipole
177
What is meant by intermolecular force?
Attractive force between neighbouring molecules
178
What are the 3 types of intermolecular forces?
● Hydrogen bonding
● Permanent Dipole-Dipole interactions
● Induced Dipole-Dipole interactions (London Forces)
179
What is the strongest type of intermolecular force?
Hydrogen bonding
180
Describe permanent dipole-dipole interactions
Some molecules with polar bonds have permanent dipoles → forces of attraction between those dipoles and those of neighbouring molecules
181
Describe London forces (induced dipole-dipole interactions)
● London forces are caused by random movements of electrons
● This leads to instantaneous dipoles
● Instantaneous dipole induces a dipole in nearby molecules
● Induced dipoles attract one another
182
Are London forces greater in smaller or larger molecules?
Larger due to more electrons
183
Does boiling point increase or decrease down the noble gas group? Why?
Boiling point increases because the number of electrons increases and hence the strength of London forces also increases
184
What conditions are needed for hydrogen bonding to occur?
O-H, N-H or F-H bond, lone pair of electrons on O, F, N
Because O, N and F are highly electronegative, H nucleus is left exposed
A strong force of attraction between H nucleus and lone pair of electrons on O, N, F
185
Why is ice less dense than liquid water?
● In ice, the water molecules are arranged in a orderly pattern. It has an open lattice with hydrogen bonds.
● In water, the lattice is collapsed and the molecules are closer together.
186
Why does water have a melting/ boiling point higher than expected?
Hydrogen bonds are stronger than other intermolecular forces so extra strength is required to overcome the forces
