End of year Flashcards
(206 cards)
1
Q
What forms when atoms of different elements combine in fixed ratios?
A
Compounds, which have different properties from their component elements.
2
Q
What is a mixture?
A
A combination of two or more substances that are not chemically bonded and retain their individual properties.
3
Q
What are the two types of mixtures?
A
Homogeneous (uniform composition) and heterogeneous (non-uniform composition).
4
Q
What do chemical equations represent?
A
They describe what happens during a chemical reaction, showing reactants → products.
5
Q
What do the state symbols (s), (l), (g), and (aq) represent?
A
Solid, liquid, gas, and aqueous (dissolved in water), respectively.
6
Q
What happens to temperature during a state change?
A
Temperature remains constant during a state change.
7
Q
What’s the difference between a physical and chemical change?
A
Physical change doesn’t form new substances; chemical change forms new substances by rearranging atoms.
8
Q
What is a mole (mol)?
A
The amount of substance containing as many particles as atoms in 12g of carbon-12.
9
Q
What is Avogadro’s constant?
A
6.02 × 10²³ particles per mole.
10
Q
How do you calculate the number of particles?
A
N=n×L
11
Q
What is the formula for finding moles from mass and molar mass?
A
n=m/M
12
Q
Define relative atomic mass (Ar).
A
The weighted average mass of all isotopes of an element compared to 1/12 the mass of a carbon-12 atom.
13
Q
Define relative molecular mass (Mr).
A
The sum of the relative atomic masses of atoms in a molecule.
14
Q
What is molar mass (M) and its units?
A
Mass of one mole of a substance, in g mol⁻¹.
15
Q
What is empirical formula?
A
The simplest whole-number ratio of atoms in a compound.
16
Q
What is molecular formula?
A
The actual number of atoms of each element in a molecule.
17
Q
How do you calculate percentage composition by mass?
A
(Mass of element / Molar mass of compound) × 100
18
Q
What is atom economy?
A
(Mass of desired products / Total mass of reactants) × 100
19
Q
What is the limiting reactant?
A
The reactant that is completely used up in a chemical reaction.
20
Q
What is percentage yield?
A
(Actual yield / Theoretical yield) × 100
21
Q
What are the assumptions of the kinetic molecular theory for ideal gases?
A
Gases move randomly in straight lines
Have elastic collisions
Negligible volume
No intermolecular forces
and kinetic energy proportional to temperature.
22
Q
What is the ideal gas equation?
A
PV=nRT
23
Q
What is the value of the gas constant R?
A
8.31 J mol⁻¹ K⁻¹
24
Q
What are the standard conditions for STP?
A
273 K and 100 kPa
25
What is the molar volume of a gas at STP?
22.7 dm³ mol⁻¹
26
What is Avogadro’s Law?
Equal volumes of gases at the same temperature and pressure contain the same number of particles.
27
What is the formula for concentration?
C=n/V
28
Define solute, solvent, and solution.
Solute: dissolved substance; Solvent: does the dissolving; Solution: solute + solvent mixture.
29
What is a standard solution?
A solution with a known concentration.
30
What particles make up the nucleus of an atom?
Protons and neutrons (nucleons).
31
What occupies the space outside the nucleus of an atom?
Negatively charged electrons.
32
What is the role of a mass spectrometer?
It determines the relative atomic mass from the isotopic composition.
33
What does nuclear notation
A
X
Z
represent?
A = mass number (protons + neutrons), Z = atomic number (protons), X = element symbol.
34
How do you calculate the number of neutrons in an atom?
Mass number - Atomic number.
35
What are isotopes?
Atoms of the same element with the same number of protons but different numbers of neutrons.
36
Do isotopes have the same chemical properties? Why?
Yes, because they have the same number of electrons.
37
Do isotopes have the same physical properties?
No, because of different masses and nuclear composition.
38
What was the key finding of Rutherford’s gold foil experiment?
Atoms have a dense, positively charged nucleus and are mostly empty space.
39
Why were some alpha particles deflected backwards in Rutherford’s experiment?
They hit the dense, positively charged nucleus.
40
What is the relative mass and charge of a proton?
Mass = 1, Charge = +1
41
What is the relative mass and charge of a neutron?
Mass = 1, Charge = 0
42
What is the relative mass and charge of an electron?
Mass ≈ 1/1836, Charge = -1
43
What happens when an electron absorbs energy?
It jumps to a higher energy level (excited state).
44
What happens when an electron falls back to a lower energy level?
It emits a photon.
45
What evidence supports Bohr’s idea of discrete energy levels?
The line emission spectrum of hydrogen.
46
What does each line in an emission spectrum represent?
A photon emitted when an electron transitions between energy levels.
47
What is the Balmer series?
Transitions to n=2 in hydrogen (visible spectrum).
48
What part of the spectrum corresponds to transitions to n=1?
Ultraviolet (Lyman series).
49
What part of the spectrum corresponds to transitions to n=3?
Infrared (Paschen series).
50
How are wavelength and frequency related?
Inversely – as wavelength decreases, frequency increases.
51
What is the order of the electromagnetic spectrum?
Radio, Microwave, Infrared, Visible, UV, X-ray, Gamma.
52
Which has more energy: red light or violet light?
Violet light.
53
What is an orbital?
A region where there’s a high probability (≥95%) of finding an electron.
54
How many electrons can one orbital hold?
2, with opposite spins.
55
What shapes do s and p orbitals have?
s = spherical, p = dumbbell-shaped.
56
How many orbitals and electrons are in each sub-level?
s: 1 orbital (2e), p: 3 orbitals (6e), d: 5 orbitals (10e), f: 7 orbitals (14e)
57
What are the 3 main rules for filling orbitals?
Pauli Exclusion, Aufbau Principle, Hund’s Rule.
58
What is the electron configuration of Chromium (Z = 24)?
[Ar] 3d⁵ 4s¹
59
What is the electron configuration of Copper (Z = 29)?
[Ar] 3d¹⁰ 4s¹
60
Which orbital is filled first, 4s or 3d?
4s is filled first, but 4s electrons are removed first in ionization.
61
What is first ionization energy?
Energy required to remove one mole of electrons from one mole of gaseous atoms.
62
What equation represents first ionization energy?
X(g) → X⁺(g) + e⁻
63
What factors affect ionization energy?
Nuclear charge, atomic radius, electron shielding.
64
How does ionization energy change across a period?
It increases due to greater nuclear charge.
65
How does ionization energy change down a group?
It decreases due to increased atomic radius and shielding.
66
Why is there a drop in ionization energy from Be to B?
Electron in B is in a higher energy 2p orbital.
67
Why do successive ionization energies increase?
Fewer electrons, stronger attraction to the nucleus.
68
What do large jumps in ionization energies indicate?
Removal of an electron from a new inner shell.
69
What does the convergence limit in an emission spectrum represent?
The ionization energy.
70
What equations relate frequency, wavelength, and energy?
E = hv
c = λv
(Where h = Planck’s constant, c = speed of light)
71
What forms a cation?
A cation is formed when a metal loses valence electrons.
72
What forms an anion?
An anion is formed when a non-metal gains electrons.
73
What determines how many electrons an atom gains or loses to form an ion?
The atom’s electron configuration.
74
What is an ionic bond?
An ionic bond is the electrostatic attraction between oppositely charged ions.
75
What is the physical state and structure of ionic compounds under normal conditions?
They are usually solid and form lattice structures.
76
How can you deduce the formula of an ionic compound?
By balancing the charges of the component ions, including polyatomic ions.
77
Why do ionic compounds have high melting points?
Due to strong electrostatic forces between ions that require a lot of energy to break.
78
Why don’t solid ionic compounds conduct electricity?
Because the ions are fixed in place and cannot move.
79
Why do molten or aqueous ionic compounds conduct electricity?
The ions are free to move and carry current.
80
Why are ionic compounds soluble in polar solvents like water?
The polar water molecules can separate the positive and negative ions from the lattice.
81
What is a covalent bond?
A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
82
Between which types of elements do covalent bonds typically form?
Between non-metal atoms.
83
How many electrons are shared in single, double, and triple covalent bonds?
2, 4, and 6 electrons respectively.
84
What happens to bond length and strength as more electrons are shared?
Bond length decreases and bond strength increases.
85
What causes bond polarity?
Differences in electronegativity between bonded atoms.
86
What kind of bond results when electrons are shared unequally?
A polar covalent bond.
87
What does a Lewis structure show?
All the valence electrons in a covalently bonded molecule or ion.
88
What is the octet rule?
The tendency of atoms to have 8 electrons in their valence shell.
89
Which elements commonly form incomplete octets?
Beryllium (Be) and Boron (B).
90
What are resonance structures?
Different Lewis structures for the same molecule showing multiple positions for double bonds.
91
What theory is used to predict molecular shape?
VSEPR theory (Valence Shell Electron Pair Repulsion).
92
What is the molecular geometry of a molecule with 4 bonding pairs and 0 lone pairs?
Tetrahedral (109.5°).
93
What conditions must be met for a molecule to be polar?
It must have polar bonds and an asymmetric shape.
94
What are the two types of covalent structures?
Simple covalent and giant covalent structures.
95
Why do simple covalent compounds have low melting points?
Due to weak intermolecular forces between molecules.
96
Why is diamond hard and non-conductive?
Each carbon atom forms 4 strong covalent bonds and there are no free electrons.
97
Why is graphite conductive?
It has delocalized electrons that can move between layers.
98
What are the three main types of intermolecular forces?
London dispersion forces, dipole-dipole forces, and hydrogen bonds.
99
Which is the weakest intermolecular force?
London dispersion forces.
100
What causes dipole-dipole attractions?
The attraction between the positive end of one polar molecule and the negative end of another.
101
What makes hydrogen bonds especially strong?
The large electronegativity difference between hydrogen and N, O, or F atoms.
102
What is the order of strength of intermolecular forces (weakest to strongest)?
London dispersion < dipole-dipole < hydrogen bonding.
103
What is a metallic bond?
An electrostatic attraction between a lattice of positive metal ions and delocalized electrons.
104
What affects the strength of a metallic bond?
The charge and size of the metal ions.
105
Why do metals conduct electricity?
Because they have mobile delocalized electrons.
106
Why are metals malleable and ductile?
Because the layers of metal ions can slide over each other without breaking the bond.
107
What is an alloy?
A mixture of metals or a metal with another element with enhanced properties.
108
Why are alloys often stronger than pure metals?
Different-sized atoms disrupt the regular metal lattice, making it harder for layers to slide.
109
What is heat in terms of energy?
Heat is a form of energy associated with the total kinetic energy of particles in a substance.
110
What does temperature measure in a substance?
Temperature measures the average kinetic energy of particles in a substance.
111
What is the law of conservation of energy?
Energy cannot be created or destroyed, only transferred or transformed; total energy is conserved.
112
What does ΔH represent in thermodynamics?
ΔH represents the enthalpy change, measured in kJ mol⁻¹.
113
Under what conditions is standard enthalpy change measured?
Standard conditions: 298 K (25°C), 101.3 kPa, 1 mol dm⁻³ concentration, substances in standard states.
114
What equation is used to calculate heat change in a substance?
q = mcΔT
115
What does each symbol in q = mcΔT represent?
q = heat energy (J), m = mass (g), c = specific heat capacity (J/g·K), ΔT = temperature change (K or °C)
116
What is a calorimetry experiment used for?
Measuring the amount of heat released or absorbed in a chemical reaction.
117
What is an exothermic reaction?
A reaction that releases heat to the surroundings; ΔH is negative.
118
What is an endothermic reaction?
A reaction that absorbs heat from the surroundings; ΔH is positive.
119
Give an example of an exothermic reaction.
Combustion or condensation (e.g., rain formation).
120
Give an example of an endothermic process.
Photosynthesis or melting ice.
121
Define standard enthalpy change of formation (ΔHf°).
The energy change when one mole of a compound is formed from its elements in their standard states.
122
What is the standard enthalpy change of combustion (ΔHcomb°)?
The energy released when one mole of a compound is completely burned in excess oxygen under standard conditions.
123
What is the standard enthalpy change of neutralization (ΔHneut°)?
The enthalpy change when one mole of water is formed from a neutralization reaction between a strong acid and base.
124
How do you calculate enthalpy change from heat energy?
ΔH = q / n, where q = heat (kJ), n = moles.
125
What is the specific heat capacity?
The energy needed to raise the temperature of 1 gram of a substance by 1 K.
126
State Hess’s Law.
The total enthalpy change of a reaction is the same, regardless of the route taken, provided the initial and final conditions are the same.
127
How do you apply Hess’s Law using equations?
Combine known reactions to derive the desired equation and sum their ΔH values.
128
What is bond enthalpy?
The energy required to break one mole of bonds in the gaseous state.
129
What is the formula for calculating ΔH using bond enthalpies?
ΔH = ΣE(bonds broken) − ΣE(bonds formed)
130
What does a positive ΔH from bond enthalpies indicate?
The reaction is endothermic (more energy required to break bonds than released forming them).
131
Why is ozone important in the atmosphere?
It absorbs harmful UV radiation, protecting living organisms.
132
Why does O₂ require higher energy UV to break compared to O₃?
O₂ has stronger bonds (double bonds) compared to the weaker bonds in O₃.
133
What is the energy-wavelength formula?
E = hc/λ, where h = Planck’s constant, c = speed of light, λ = wavelength.
134
What is lattice enthalpy (ΔHlat°)?
The energy released when one mole of an ionic compound is formed from gaseous ions.
135
What is ionization enthalpy (ΔHi°)?
The energy needed to remove one mole of electrons from one mole of gaseous atoms.
136
What is electron affinity (ΔHe°)?
The energy change when one mole of electrons is added to one mole of gaseous atoms.
137
What is the Born-Haber cycle used for?
Calculating lattice enthalpy using known enthalpy changes and Hess’s Law.
138
What three conditions must be met for a reaction to occur according to collision theory?
Particles must collide, collide with proper orientation, and collide with sufficient energy.
139
What is activation energy (Ea)?
The minimum amount of energy required for a successful collision that leads to a reaction.
140
How does a catalyst affect the rate of a chemical reaction?
It lowers the activation energy, increasing the rate without being permanently changed.
141
How does increasing temperature affect the rate of reaction?
It increases particle speed and the proportion of particles with energy ≥ Ea, resulting in more successful collisions.
142
How does increasing concentration or pressure affect the rate of reaction?
It increases the frequency of collisions, leading to more successful collisions per unit time.
143
How does particle size (surface area) affect the reaction rate?
Smaller particles have a larger surface area, allowing more frequent collisions and a faster rate.
144
What does the Maxwell-Boltzmann distribution curve represent?
The distribution of kinetic energy among particles in a sample.
145
How does the curve change with temperature increase?
It flattens and shifts right, increasing the number of particles with energy ≥ Ea.
146
What is the effect of a catalyst on the Maxwell-Boltzmann distribution curve?
The activation energy line shifts left, showing more particles now have enough energy to react.
147
How is the rate of reaction defined?
As the change in concentration of reactants or products per unit time.
148
What is the equation for reaction rate in terms of concentration?
Rate = Δ[Product]/Δt or Rate = -Δ[Reactant]/Δt
149
How can concentration changes be monitored during a reaction?
By measuring gas volume, mass change, pressure, titration, color change, or conductivity.
150
What determines the overall rate of a reaction with multiple steps?
The rate-determining step, which is the slowest step.
151
What is the rate law for a reaction A + B → C?
rate = k[A]^x[B]^y, where x and y are the reaction orders.
152
How is the overall order of a reaction determined?
By summing the powers of the concentration terms: x + y.
153
Can reaction orders be inferred from a balanced chemical equation?
No, they must be determined experimentally.
154
What is the unit of the rate constant for a first-order reaction?
s⁻¹
155
How does the rate change for a second-order reaction if concentration is doubled?
The rate increases by a factor of 4.
156
What is the molecularity of an elementary step?
The number of reactant particles involved in that step.
157
How do catalysts affect reaction mechanisms?
They introduce a new pathway with lower activation energy.
158
What is the Arrhenius equation?
k = Ae^(−Ea/RT)
159
What does the term "A" in the Arrhenius equation represent?
The frequency factor – how often molecules collide with correct orientation.
160
What is the linear form of the Arrhenius equation?
ln(k) = -Ea/R × (1/T) + ln(A)
161
What does the slope of the plot of ln(k) vs 1/T represent?
-Ea/R
162
How does temperature affect the rate constant (k)?
As temperature increases, k increases because more particles have sufficient energy.
163
What does a graph of 1/T vs ln(k) allow you to determine?
The activation energy (Ea) and the frequency factor (A).
164
What is chemical equilibrium?
A state in a closed system where the rates of the forward and reverse reactions are equal.
165
What does the equilibrium constant 𝐾𝑐 represent?
The ratio of the concentrations of products to reactants at equilibrium for a given reaction.
166
How does the magnitude of 𝐾𝑐 indicate the position of equilibrium?
If 𝐾𝑐>1, equilibrium favors products; if 𝐾𝑐<1, equilibrium favors reactants.
167
What is the reaction quotient (Q)?
The ratio of product and reactant concentrations at any point in time (not necessarily at equilibrium).
168
How can comparing 𝑄 and 𝐾𝑐 predict the direction of a reaction?
If 𝑄<𝐾𝑐, the reaction proceeds forward; if 𝑄>𝐾𝑐, the reaction goes backward; if 𝑄=𝐾𝑐, the system is at equilibrium.
169
Does a catalyst affect the position of equilibrium or 𝐾𝑐?
No, a catalyst speeds up both forward and reverse reactions equally but does not change 𝐾𝑐 or the position of equilibrium.
170
What are the characteristics of dynamic equilibrium?
Forward and reverse reactions occur at the same rate; concentrations of reactants and products remain constant; occurs in a closed system.
171
Write the expression for the equilibrium constant 𝐾𝑐 for the reaction 𝑎𝐴+𝑏𝐵⇌𝑐𝐶+𝑑𝐷
Kc=[C]c[D]d/[A]a[B]b
172
How does increasing concentration affect equilibrium according to Le Châtelier’s Principle?
The equilibrium shifts to the side that consumes the added substance (to the side with fewer moles of solute if concentration increases).
173
How does increasing pressure affect equilibrium in reactions involving gases?
Equilibrium shifts toward the side with fewer moles of gas.
174
How does temperature change affect equilibrium and 𝐾𝑐?
Temperature changes shift the equilibrium position and change the value of 𝐾𝑐; increasing temperature favors the endothermic direction.
175
How do you manipulate 𝐾𝑐 for the reverse reaction?
Kc′=1/Kc
176
How do you find 𝐾𝑐 for a reaction that is a multiple of another reaction?
Raise the original 𝐾𝑐 to the power of the multiple (e.g. 𝐾𝑐^𝑥).
177
How do you combine 𝐾𝑐 values for two reactions added together?
Multiply the individual 𝐾𝑐 values:
𝐾𝑐^𝑡𝑜𝑡𝑎𝑙=𝐾𝑐^1×𝐾𝑐^2
178
What is an ICE table used for in equilibrium problems?
To track Initial concentrations, Changes, and Equilibrium concentrations for reactants and products.
179
What is the relationship between Gibbs free energy change Δ𝐺 and the equilibrium constant 𝐾?
ΔG=−RTlnK
180
What condition must be met for a system to be in dynamic equilibrium?
It must be a closed system where neither reactants nor products can escape.
181
In an exothermic reaction, how does increasing temperature affect equilibrium?
The equilibrium shifts towards the reactants (the endothermic side).
182
In an endothermic reaction, how does decreasing temperature affect equilibrium?
The equilibrium shifts towards the reactants (the exothermic side).
183
What are three ways to define oxidation and reduction?
Oxidation: oxygen gain, hydrogen loss, or electron loss
Reduction: hydrogen gain, oxygen loss, or electron gain
184
What does OIL RIG stand for?
Oxidation Is Loss (of electrons); Reduction Is Gain (of electrons)
185
What happens to an oxidizing agent during a redox reaction?
It is reduced and gains electrons.
186
What happens to a reducing agent during a redox reaction?
It is oxidized and loses electrons.
187
What is the oxidation number of an element in its elemental form?
0
188
What does an increase in oxidation number indicate?
Oxidation (loss of electrons)
189
What does a decrease in oxidation number indicate?
Reduction (gain of electrons)
190
What are common oxidizing agents?
Halogens, ozone, permanganate ions (MnO₄⁻), hydrogen peroxide (H₂O₂)
191
What are common reducing agents?
Hydrogen, carbon, carbon monoxide, sulfur dioxide (SO₂)
192
What is the oxidation state of chromium in K₂Cr₂O₇?
+6
193
What is the purpose of half-equations in redox?
To separate and balance the oxidation and reduction parts of the redox reaction.
194
What are the steps to write redox half-equations in acidic solution?
1. Balance all atoms except O and H
2. Balance O using H₂O
3. Balance H using H⁺
4. Balance charge using e⁻
5. Add the two half-reactions
195
What is the balanced redox equation for MnO₄⁻ + Fe²⁺ → Mn²⁺ + Fe³⁺?
MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O
196
What does the activity series show?
The relative reactivity of metals and their ease of oxidation.
197
How does the activity series predict redox feasibility?
A metal can only displace another from a compound if it is higher in the activity series.
198
What is redox titration used for?
To determine the concentration of an oxidizing or reducing agent.
199
What type of energy conversion occurs in a voltaic cell?
Chemical energy → Electrical energy
200
Where does oxidation occur in a voltaic cell?
At the anode (negative electrode)
201
Where does reduction occur in a voltaic cell?
At the cathode (positive electrode)
202
What is the purpose of a salt bridge in voltaic cells?
Maintains electrical neutrality by allowing ion flow.
203
In a voltaic cell, what is the cell diagram notation?
Anode (oxidation) on the left || cathode (reduction) on the right
Example: Zn | Zn²⁺ || Cu²⁺ | Cu
204
What is the standard electrode potential of the standard hydrogen electrode (SHE)?
0 V
205
What is the equation relating E° and Gibbs free energy?
ΔG° = −nFE°
206
What does a positive E° value indicate about a reaction?
It is spontaneous (ΔG° is negative)
