1
Q
What is a mixture?
A
A mixture is a physical combination of two or more substances not chemically bonded; components retain their properties and can be separated by physical methods.
2
Q
How does filtration separate mixtures?
A
Filtration separates an insoluble solid from a liquid using a filter; dissolved substances pass through (exception: very fine particles may require vacuum filtration).
2
Q
What is the difference between a homogeneous and heterogeneous mixture?
A
Homogeneous mixtures have uniform composition throughout (solutions); heterogeneous mixtures have visibly different components or phases.
2
Q
How does simple distillation work?
A
It separates liquids (or a solvent from a solution) based on boiling point differences; effective when boiling points differ by ~25 °C or more.
3
Q
When is fractional distillation used instead of simple distillation?
A
Fractional distillation is used when liquids have similar boiling points; a fractionating column allows repeated vaporization–condensation cycles.
4
Q
How does chromatography separate components?
A
Components partition differently between a stationary phase and a mobile phase, causing different movement rates; more soluble in mobile phase → travels further.
5
Q
What is an exothermic process?
A
An exothermic process releases energy (usually heat) to the surroundings; ΔH is negative and products have lower enthalpy than reactants.
6
Q
What is an endothermic process?
A
An endothermic process absorbs energy from the surroundings; ΔH is positive and products have higher enthalpy than reactants.
7
Q
How does temperature of surroundings change in exothermic vs endothermic reactions?
A
Exothermic reactions increase surrounding temperature; endothermic reactions decrease surrounding temperature.
8
Q
Are bond-breaking and bond-forming endothermic or exothermic?
A
Bond-breaking is endothermic; bond-forming is exothermic (overall ΔH depends on which dominates).
9
Q
What determines the identity of an element?
A
The number of protons (atomic number); changing protons changes the element.
10
Q
How do you find the number of neutrons in an atom?
A
Neutrons = mass number − atomic number.
11
Q
How does the number of electrons compare to protons in a neutral atom?
A
In a neutral atom, electrons = protons.
12
Q
How do ions differ in electron number?
A
Cations have fewer electrons than protons; anions have more electrons than protons.
13
Q
What does the x-axis represent in a mass spectrum?
A
Mass-to-charge ratio (m/z), usually with charge +1 in IB Chemistry.
13
Q
What does the y-axis represent in a mass spectrum?
A
Relative abundance (percentage compared to the most intense peak, which is set to 100%).
14
Q
What information can isotopic peaks provide?
A
They show isotopes of an element and their relative abundances, allowing calculation of average atomic mass.
15
Q
Why is the tallest peak not always the heaviest isotope?
A
The tallest peak represents the most abundant isotope, not the one with the greatest mass.
16
Q
What is the molecular ion peak (M⁺)?
A
The peak corresponding to the whole molecule after losing one electron; its m/z gives the relative molecular mass.
16
Q
What causes fragmentation in mass spectrometry?
A
High-energy electrons break bonds in the molecular ion, producing smaller fragment ions.
17
Q
Why are certain fragment peaks more intense?
A
More stable carbocations (e.g. tertiary > secondary > primary) form more frequently, giving higher peaks.
18
Q
How does fragmentation help identify structure?
A
Fragment patterns indicate functional groups and carbon skeletons, helping deduce molecular structure.
19
Q
What is the shape of an s orbital?
A
S orbitals are spherical and centered on the nucleus.
20
Q
What is the shape of a p orbital?
A
P orbitals are dumbbell-shaped with two lobes and a nodal plane at the nucleus.
21
How many p orbitals are in a p sublevel and how are they oriented?
Three, oriented at right angles along the x, y, and z axes.
22
What is a key feature of d orbitals (HL)?
D orbitals generally have more complex shapes (clover-like), except d_z² which has a dumbbell with a torus.
23
What is the maximum number of electrons in one orbital?
Two electrons, with opposite spins (Pauli exclusion principle).
23
What is the maximum number of electrons in each sublevel?
s = 2, p = 6, d = 10, f = 14.
24
What is the maximum number of electrons in an energy level n?
Maximum electrons = 2n².
24
Why must electrons in the same orbital have opposite spins?
Because no two electrons can have the same set of four quantum numbers.
25
What is meant by successive ionization energies?
Successive ionization energies are the energies required to remove electrons one at a time from the same atom in the gaseous state.
26
Why do successive ionization energies always increase?
After each electron is removed, remaining electrons experience greater nuclear attraction due to reduced shielding.
27
What does a large jump in successive ionization energies indicate?
A large jump shows removal of an electron from a new, inner shell closer to the nucleus.
28
How can successive ionization energies be used to determine group number?
The number of electrons removed before a large jump corresponds to the number of valence electrons.
29
What causes the hydrogen emission spectrum?
Electrons fall from higher to lower energy levels, emitting photons with specific energies.
29
What is the Lyman series?
Transitions ending at n = 1; all lines are in the ultraviolet region.
30
What is the Balmer series?
Transitions ending at n = 2; includes visible lines and some UV.
31
What is the Paschen series?
Transitions ending at n = 3; lines are in the infrared region.
32
What is the general trend of first ionization energy down a group?
It decreases down a group because of increased atomic radius and greater electron shielding.
33
What is the general trend of first ionization energy across a period?
It generally increases across a period due to increasing nuclear charge and similar shielding.
34
Why is the first ionization energy of oxygen lower than nitrogen?
Oxygen has paired electrons in a p orbital, increasing repulsion and making removal easier.
35
Why is the first ionization energy of boron lower than beryllium?
Boron’s outer electron is in a higher-energy p orbital, which is easier to remove than Be’s s electron.
36
What is percent composition?
Percent composition is the percentage by mass of each element in a compound.
37
How is percent composition calculated?
(Mass of element ÷ molar mass of compound) × 100%.
38
What is an empirical formula?
The simplest whole-number ratio of atoms in a compound.
39
How is percent composition related to empirical formula?
Percent composition is used to find mole ratios, which are simplified to determine the empirical formula.
40
What is stoichiometry?
Stoichiometry is the quantitative relationship between reactants and products based on a balanced chemical equation.
41
What is a limiting reactant?
The reactant that is completely consumed first and limits the amount of product formed.
42
How does percent yield relate to stoichiometry?
Percent yield compares actual yield to theoretical yield: (actual ÷ theoretical) × 100%.
43
What assumptions are made for an ideal gas?
Gas particles have negligible volume and no intermolecular forces, and collisions are perfectly elastic.
43
Why do real gases deviate from ideal behavior?
Real gases have particle volume and intermolecular attractions, especially at high pressure and low temperature.
44
Under what conditions do real gases behave most ideally?
At low pressure and high temperature, where particles are far apart and moving rapidly.
45
What is the ideal gas law equation?
PV = nRT, relating pressure, volume, amount of gas, and temperature.
45
What units must be used in the ideal gas law?
P in atm or Pa, V in L or m³, T in kelvin, and R consistent with chosen units.
46
What is the combined gas law?
(P₁V₁) ÷ T₁ = (P₂V₂) ÷ T₂ for a fixed amount of gas.
46
When should the combined gas law be used instead of the ideal gas law?
When the amount of gas is constant and conditions change between two states.
47
What common mistake must be avoided when using the combined gas law?
Temperature must always be converted to kelvin before substitution.
48
What factors affect the strength of ionic bonds?
Ionic bond strength depends on charge of ions and ionic radius; higher charges and smaller ions give stronger electrostatic attraction.
48
Why do compounds with 2+ and 2− ions have stronger ionic bonds?
Greater ionic charge increases Coulombic attraction, raising lattice enthalpy and bond strength.
49
How does ionic radius affect ionic bond strength?
Smaller ionic radii reduce distance between ions, increasing attraction and bond strength.
50
Why do some ionic compounds dissolve easily in water despite strong bonds?
Hydration enthalpy can overcome lattice enthalpy, allowing ions to separate.
50
What are typical properties of ionic substances?
High melting points, conduct electricity when molten or dissolved, often soluble in water, and brittle solids.
51
What are typical properties of simple covalent substances?
Low melting points, poor electrical conductivity, often gases or liquids at room temperature.
52
Why do giant covalent substances have high melting points?
Extensive networks of strong covalent bonds require large amounts of energy to break.
53
What causes metallic substances to conduct electricity?
Delocalized electrons move freely through the lattice, carrying charge.
54
What is lattice enthalpy?
The enthalpy change when one mole of an ionic solid forms from its gaseous ions.
55
What does a more negative lattice enthalpy indicate?
Stronger ionic bonding and a more stable ionic lattice.
56
How do charge and ionic radius affect lattice enthalpy?
Higher ionic charge and smaller ionic radius make lattice enthalpy more negative.
57
Why is lattice enthalpy difficult to measure directly?
It involves gaseous ions, so it is calculated using Born–Haber cycles.
57
What is meant by allotropy?
Allotropy is the existence of an element in different structural forms in the same physical state.
58
Why is diamond very hard?
Each carbon forms four covalent bonds in a tetrahedral giant covalent lattice.
58
Why does graphite conduct electricity?
Each carbon bonds to three others, leaving one delocalized electron per atom.
59
Why is graphene stronger than steel by mass?
A single layer of hexagonally bonded carbon atoms forms extremely strong covalent bonds.
60
What determines bond angles in molecules?
Bond angles are determined by repulsion between electron pairs around a central atom (VSEPR theory).
61
What is paper chromatography used for?
It separates and identifies components of a mixture based on solubility and attraction to the stationary phase.
61
What is the stationary phase in paper chromatography?
Water trapped in the cellulose fibers of the paper.
61
What is the mobile phase?
The solvent that moves up the paper carrying dissolved substances.
62
How is the retardation factor (Rf) calculated?
Rf = distance traveled by substance ÷ distance traveled by solvent front.
62
What are intermolecular forces?
Weak attractions between molecules, much weaker than covalent or ionic bonds.
63
What are the three main types of IMFs?
London dispersion forces, dipole–dipole interactions, and hydrogen bonding.
64
When does hydrogen bonding occur?
When hydrogen is bonded to nitrogen, oxygen, or fluorine.
65
How do IMFs affect boiling point?
Stronger IMFs increase boiling point by requiring more energy to separate molecules.
65
What is an electron domain?
An electron domain is a region of electron density around a central atom, including bonding pairs and lone pairs.
65
How does electron-domain geometry differ from molecular geometry?
Electron-domain geometry considers all electron domains; molecular geometry considers only the positions of atoms.
66
Why do lone pairs affect molecular shape more than bonding pairs?
Lone pairs repel more strongly because they are closer to the nucleus and occupy more space.
66
What is formal charge?
Formal charge is the hypothetical charge on an atom assuming equal sharing of bonding electrons.
67
How is formal charge calculated?
Formal charge = valence electrons − nonbonding electrons − (bonding electrons ÷ 2).
67
How does formal charge relate to molecular stability?
More stable structures minimize formal charges and place negative charges on more electronegative atoms.
67
Why can molecules with nonzero formal charges still be stable?
Resonance structures can delocalize charge, increasing overall stability.
68
What defines a transition metal in IB Chemistry?
A transition metal is an element that forms at least one ion with a partially filled d subshell.
68
Why do transition metals show variable oxidation states?
The 4s and 3d electrons are similar in energy, so different numbers can be lost.
69
Why do many transition metal compounds are colored?
d–d electron transitions absorb specific wavelengths of visible light.
70
Why do transition metals act as catalysts?
They provide alternative reaction pathways by forming temporary complexes with reactants.
70
What causes metallic bonding?
Attraction between positive metal ions and a sea of delocalized electrons.
71
What factors increase the strength of metallic bonding?
Higher charge on metal ions and greater number of delocalized electrons.
71
Why do transition metals have stronger metallic bonding than Group 1 metals?
They contribute more delocalized electrons, increasing electrostatic attraction.
72
How does metallic bond strength affect melting point?
Stronger metallic bonding leads to higher melting points.
72
What is a monomer?
A small molecule that can join with others to form a polymer.
73
What is the difference between addition and condensation polymerization?
Addition polymerization forms polymers without by-products; condensation polymerization releases small molecules like water.
74
Why are many polymers flexible?
Weak intermolecular forces between long chains allow them to slide past each other.
74
What type of bonding holds monomers together in polymers?
Strong covalent bonds form the polymer backbone.
74
What principle governs electron filling of orbitals?
The Aufbau principle: electrons fill lowest-energy orbitals first.
75
Why are electrons removed from the 4s orbital before the 3d in transition metals?
After filling, 3d orbitals become lower in energy than 4s.
76
How does electron configuration change when forming cations?
Electrons are removed from the highest-energy occupied orbitals first.
76
Why do some ions have particularly stable configurations?
They achieve noble gas or half-filled/full subshell configurations.
76
What is oxidation state?
The hypothetical charge an atom would have if all bonds were ionic.
77
What are the common oxidation states of oxygen and hydrogen?
Oxygen is usually −2 (exception: peroxides −1); hydrogen is usually +1 (exception: metal hydrides −1).
77
Why do transition metals have multiple oxidation states?
Similar energies of s and d electrons allow variable electron loss.
78
How is oxidation state used in redox reactions?
Oxidation is an increase in oxidation state; reduction is a decrease.
78
What information does a structural formula show in organic chemistry?
It shows how atoms are connected, functional groups present, and bonding, which determines properties and reactions.
79
What rules are used to name organic compounds in IB Chemistry?
IUPAC rules: longest carbon chain, lowest possible numbers for functional groups, correct suffix/prefix.
79
What are isomers?
Compounds with the same molecular formula but different arrangements of atoms.
80
What is the difference between structural and stereoisomers?
Structural isomers differ in connectivity; stereoisomers have same connectivity but different spatial arrangement.
81
Why do isomers often have different physical properties?
Different shapes or functional group positions affect intermolecular forces and polarity.
81
What does IR spectroscopy measure?
Absorption of infrared radiation causing bond vibrations in molecules.
82
What is the purpose of the fingerprint region?
It uniquely identifies compounds, even if functional groups are similar.
82
Which bonds give strong IR absorptions?
Polar bonds (e.g. O–H, C=O); nonpolar bonds absorb weakly.
83
What causes cis–trans isomerism?
Restricted rotation around a C=C bond with two different groups on each carbon.
83
What is the difference between cis and trans isomers?
Cis: similar groups on same side; trans: similar groups on opposite sides.
84
Why do cis and trans isomers have different boiling points?
Cis isomers are often more polar, giving stronger intermolecular forces.
84
What is the system in thermochemistry?
The part of the universe being studied (usually the reacting chemicals).
84
What are the surroundings?
Everything outside the system that can exchange energy with it.
84
How are system and surroundings related in exothermic reactions?
Energy is transferred from system to surroundings.
84
What equation is used in calorimetry?
q = mcΔT, where q is heat energy transferred.
84
What does a reaction coordinate diagram show?
Energy changes from reactants to products during a reaction.
85
What does Hess’s law state?
The total enthalpy change is independent of the reaction pathway.
86
How is ΔH calculated using enthalpies of formation?
ΔH = ΣΔHf(products) − ΣΔHf(reactants).
86
What is activation energy (Ea)?
The minimum energy needed for a reaction to occur.
87
How does a catalyst affect a reaction coordinate diagram?
It lowers activation energy without changing ΔH.
88
When are bond enthalpies used to calculate ΔH?
When formation or combustion data are unavailable; values are averages.
89
Why are bond enthalpy calculations less accurate?
Bond enthalpies are averaged over many compounds.
90
What are typical properties of metal oxides?
Basic oxides that react with acids to form salts and water.
90
What are typical properties of nonmetal oxides?
Acidic oxides that react with bases to form salts and water.
91
What are amphoteric oxides?
Oxides (e.g. Al₂O₃) that react with both acids and bases.
92
When does entropy increase in a reaction?
When number of gas molecules increases or solids form liquids/gases.
92
Why do gases have higher entropy than solids?
Gas particles have greater freedom of movement.s
92
What is entropy?
A measure of disorder or randomness in a system.
93
What does a negative ΔH indicate?
An exothermic reaction.
94
What does a positive ΔS indicate?
An increase in disorder.
95
When is ΔG negative?
When a reaction is spontaneous under given conditions.
95
What is the relationship between ΔG and equilibrium constant K?
ΔG = −RT ln K.
96
What does a negative ΔG imply about K?
K > 1, equilibrium favors products.
96
What does ΔG = 0 indicate?
The system is at equilibrium.
97
What is the Pauli exclusion principle?
No two electrons in an atom can have the same set of four quantum numbers.
98
How do IMFs affect physical properties?
Stronger IMFs increase boiling point and viscosity and decrease volatility.
99
What is Hund’s rule?
Electrons occupy degenerate orbitals singly with parallel spins before pairing.
99
What is the Aufbau principle?
Electrons fill orbitals in order of increasing energy.
99
What are four allotropes of carbon?
Diamond, graphite, graphene, and fullerenes (e.g. C₆₀).
100
What is ionic bonding (IB definition)?
Ionic bonding is the electrostatic attraction between oppositely charged ions formed by electron transfer.
101
What is covalent bonding (IB definition)?
Covalent bonding is the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
102
What is metallic bonding (IB definition)?
Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and delocalized electrons.
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IB Chemistry SL
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