Chemistry Grade 9 - Vocabulary Flashcards (Unit 1-5)
Unit 1: Structure of the Atom
Unit outcomes
- Comprehend Dalton’s atomic theory and modern atomic theory
- Understand discovery of electron and nucleus; know terms: atomic number (Z), mass number (A), atomic mass, isotope, energy level, valence electrons, electron configuration
- Learn Dalton, Thomson, Rutherford, Bohr and quantum mechanical atomic models
- Develop skills: calculate protons/electrons/neutrons from Z and A; calculate atomic masses for isotopes; write ground-state electron configurations; draw diagrammatic representations of atoms
- Demonstrate scientific inquiry: observe, compare/contrast, communicate, question, apply concepts
1.1 Atomic theory
- Dalton’s postulates (the basic ideas behind his atomic theory):
1) All elements are made of small particles called atoms.
2) Atoms are indivisible and indestructible (later revised by subatomic discoveries).
3) All atoms of a given element are identical in mass and properties; atoms of different elements differ.
4) Atoms are neither created nor destroyed in chemical reactions.
5) Compounds form when atoms of more than one element combine.
6) In a given compound, the relative numbers/types of atoms are constant. - Modern atomic theory (summarized):
1) Atoms are the smallest particles that participate in chemical reactions; they are divisible into subatomic particles (electrons, protons, neutrons).
2) Atoms of the same element may have different masses due to isotopes; atoms are not all identical in mass.
3) Atoms of a given element have the same chemical properties.
4) Atoms of different elements have different chemical properties.
5) Atoms combine in simple whole-number ratios to form compounds. - Activity prompts: compare Dalton vs modern theory; discuss postulates that were revised.
- Dalton’s postulates (the basic ideas behind his atomic theory):
1.1.1 Dalton’s Atomic Theory: early postulates and experiments
- Origins: Dalton proposed atom-based explanation from chemical laws (conservation of mass, definite proportions, multiple proportions).
- 1804–1809: postulates summarized, including “atoms are indivisible” and “atoms of a given element are identical.”
- The three laws Dalton used as foundation:
- Law of conservation of mass: mass of reactants equals mass of products.
- Law of definite proportions: a pure compound has fixed elemental composition by mass (e.g., H2O = 11.19% H, 88.81% O by mass).
- Law of multiple proportions: when two elements form more than one compound, the masses of one element in the compounds are in small whole-number ratios.
- Limitations noted: some postulates (e.g., indivisibility) were later revised, but Dalton’s framework explained many observations.
- Postulates (summary):
1) All elements are composed of tiny particles called atoms.
2) Atoms are indivisible and indestructible (later revised).
3) All atoms of a given element have identical mass/properties.
4) Atoms are not created/destroyed in chemical reactions.
5) Compounds are formed when atoms of different elements combine.
6) In a given compound, the relative numbers/types of atoms are constant.
1.1.2 Modern Atomic Theory
- In the late 19th century, discoveries revealed subatomic particles, isotopes, and the atomic nucleus.
- Modern atomic theory (key points):
1) Atoms are the smallest particles that can take part in chemical reactions; they participate in reactions as whole units.
2) Atoms are divisible into electrons, protons, and neutrons; atoms are indestructible in ordinary chemical reactions, but can be split in nuclear reactions.
3) Atoms of the same element may not be identical in mass because of isotopes.
4) Atoms of the same element have identical chemical properties.
5) Atoms of different elements have different chemical properties.
6) Atoms of two or more elements combine in simple whole-number ratios to form compounds. - Exercises: compare Dalton’s theory with modern theory; discuss implications of isotopes.
1.2 Discoveries of fundamental subatomic particles and the atomic nucleus
- Electron discovery (Cathode rays):
- 1879: William Crookes studied cathode rays in discharge tubes; rays travel from cathode to anode in straight line.
- 1897: J. J. Thomson observed deflection by electric and magnetic fields, proving cathode rays are negatively charged particles; named electrons.
- Properties of cathode rays supporting particle nature: straight-line travel, paddle wheel deflection, deflection by electric field, charge-to-mass ratio e/m independent of gas/cathode material.
- Determined e/m = 1.76 × 10^8 C g^−1; later Millikan measured e− = 1.60 × 10^−19 C and mass m_e ≈ 9.11 × 10^−31 kg.
- Discovery of the atomic nucleus (radioactivity and Rutherford’s gold foil):
- Radioactivity defined by Becquerel/Maria Curie; spontaneous emission of radiation types: alpha (α), beta (β), gamma (γ).
- 1911: Rutherford, Geiger, Marsden bombarded gold foil with α-particles; most pass through, some deflected, few reflected back.
- Conclusions: atom is mostly empty space; nucleus is very small, dense, positively charged; electrons orbit around nucleus (like planets around the sun), nucleus contains protons and neutrons.
- Rutherford’s nuclear model vs Thomson’s plum-pudding model.
- Discovery of neutrons (James Chadwick, 1932):
- Bombarded beryllium with α-particles; produced highly penetrating neutral particles with mass ~proton mass; neutrons (n⁰) with no charge and mass ~1.675 × 10^−24 g.
1.3 Composition of an atom and isotopes
- Subatomic particles: proton (p⁺, +1 charge, mass ≈ 1.673 × 10^−27 kg), neutron (n⁰, 0 charge, mass ≈ 1.675 × 10^−27 kg), electron (e⁻, −1 charge, mass ≈ 9.109 × 10^−31 kg).
- Relative and absolute masses and charges (Table 1.1):
- Electron: symbol e⁻; Relative charge −1; Relative mass 0.0005486; Actual mass ≈ 9.109 × 10^−31 kg.
- Proton: p⁺; Relative charge +1; Relative mass 1.007276; Actual mass ≈ 1.673 × 10^−27 kg.
- Neutron: n⁰; Relative charge 0; Relative mass 1.008665; Actual mass ≈ 1.675 × 10^−27 kg.
- Atomic number (Z): number of protons; also equals number of electrons in a neutral atom (Z = p⁺ = e⁻).
- Mass number (A): A = p⁺ + n⁰; nucleons total in nucleus.
- Isotopes: atoms of the same element with the same Z but different A (different numbers of neutrons). Isotopes have the same chemical properties (same Z/electrons) but different masses.
- Nuclear symbols and notations:
- Hyphen notation: Element name–A (e.g., Hydrogen-3).
- Nuclear notation: ^AZ X (e.g., ^31H, ^14_7N).
- Atomic mass concept: relative atomic mass is a weighted average of isotopic masses; carbon-12 is defined as exactly 12 amu; 1 amu = 1/12 mass of carbon-12 = 1.660 × 10^−24 g.
- Isotope designations (examples): Hydrogen isotopes: protium (¹H) with 1p/1e, deuterium (²H) with 1p/1e/1n, tritium (³H) with 1p/1e/2n.
1.3.1 Atomic Number and Mass Number
- Z = p⁺ = e⁻ (neutral atom).
- A = Z + n⁰; number of neutrons n⁰ = A − Z.
- Example: For neutral Br-80, Z = 35, A = 80; neutrons = 80 − 35 = 45.
- The concept of an electron cloud: region outside nucleus where electrons reside; atomic radius ≈ 10^−10 m (1 Å = 100 pm).
- Relative sizes and shielding: electrons are held by attraction to positively charged nucleus; inner electrons shield outer electrons; apparent nuclear charge felt by outer electrons is Zeff (effective nuclear charge).
1.3.2 Isotopes and Atomic Mass
- Isotopes have same Z, different A; examples: Hydrogen isotopes (protium, deuterium, tritium); Helium isotopes (³He, ⁴He).
- Relative atomic masses vs. average atomic mass: isotopic masses weighted by natural abundance give the average atomic mass.
- Calculating average atomic mass (weighted average):
- For chlorine:
ext{Average} = rac{(75.77oldsymbol{ imes}35) + (24.23oldsymbol{ imes}37)}{100} = 35.45 ext{ amu} - For boron with isotopes
ar{A} = rac{(10.0134oldsymbol{ imes}19.70) + (11.0094oldsymbol{ imes}80.30}{100} ext{ amu}
ightarrow 10.813 ext{ amu (example)} - Table 1.3 lists abundances and isotopic masses for several elements.
- Atomic mass unit scale and carbon-12 standard; 1 amu ≈ 1.660 × 10^−24 g.
1.3.3 Isotopes and Atomic Mass (designation and calculation)
- Isotopes are designated either by hyphen notation or nuclear symbol notation (A, Z, X).
- The number of neutrons is N = A − Z.
- If the element is neutral: p⁺ = e⁻ = Z.
- Activity: calculate Z, A, N for given isotopes; interpret neutral vs ionized species.
1.4 Atomic models
- Dalton’s model: solid, indivisible spheres (the billiard-ball model).
- Thomson’s model: plum-pudding model; electrons embedded in a positively charged sphere.
- Rutherford’s model: nucleus with dense positive core; electrons orbit around; atoms mostly empty space; planetary model; limitations on electron arrangement.
- Bohr model (1913): electrons move in circular orbits around nucleus; quantized energy levels; energy levels designated by principal quantum number n; maximum electrons per shell determined by 2n^2:
- n = 1 → 2 electrons (K-shell)
- n = 2 → 8 electrons (L-shell)
- n = 3 → 18 electrons (M-shell)
- n = 4 → 32 electrons (N-shell)
- Bohr postulates include energy absorption/emission when moving between levels; shells (K, L, M, N) correspond to principal quantum numbers.
- Quantum Mechanical Model (1920s): electrons are associated with orbitals, not fixed paths; orbitals are regions of space with high probability of finding an electron; sublevels: s, p, d, f; filling follows the Aufbau principle and diagonal rule; electrons fill lowest-energy sublevels first; noble-gas core representation (e.g., [Ne] 3s^2 for Mg).
- Electronic configurations and the periodic table: the arrangement of electrons in shells and subshells explains periodicity and chemical properties; how to write ground-state configurations e.g. Mg: 1s^2 2s^2 2p^6 3s^2 = [Ne]3s^2; copper example: ext{Cu} = [Ar] 4s^1 3d^{10} due to energy-level order.
1.4.1 Atomic models: quick recap
- Dalton: solid sphere; no internal structure.
- Thomson: pudding with embedded electrons.
- Rutherford: nucleus and empty space; electrons around the nucleus but no arrangement mechanism.
- Bohr: quantized orbits and energy levels; 2n^2 rule for shell capacities.
- Quantum Mechanical: probabilistic electron positions; orbitals; four sublevels per energy level (s, p, d, f).
1.5 Electron configurations and the Aufbau principle
- Aufbau principle: electrons occupy the lowest-energy orbital available (build-up).
- Diagonal rule (Madelung rule): order of filling sublevels often follows a diagonal path due to energy considerations; example sequence: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f, etc.
- Orbital capacities: s-sublevel (1 orbital) holds 2 e^−; p-sublevel (3 orbitals) holds 6 e^−; d-sublevel (5 orbitals) holds 10 e^−; f-sublevel (7 orbitals) holds 14 e^−.
- Ground-state electron configurations are written as a main-level notation (e.g., 1s^2 2s^2 2p^6 3s^2 for Mg) or with noble-gas cores: [Ne] 3s^2 for Mg; the first 14 elements table demonstrates their configurations.
1.5.1 Examples and practice
- Examples: Li: 1s^2 2s^1; Na: 1s^2 2s^2 2p^6 3s^1; Mg: [Ne] 3s^2; Cu: [Ar] 4s^1 3d^10.
- Practice exercises: identify correct electron configurations; fill configurations for the first 14 elements; determine valence electrons and group placement from outer-shell configurations.
1.6 Summary of Unit 1 concepts
- The Greek Democritus concept vs modern atomic theory; how the modern view evolved to include subatomic particles and nucleus.
- The nucleus concentrates most of the atom’s mass; electrons form the electron cloud surrounding the nucleus.
- The atomic number Z identifies the element; the mass number A identifies total nucleons; isotopes differ by neutron count.
- The periodic table arises from electron configurations and the arrangement of electrons in shells and subshells.
Checklists and Review exercises (highlights)
- True/False type questions about protons, neutrons, electrons, s- and d-sublevels, alpha particles, and isotopes.
- Problems asking for atomic numbers, mass numbers, and neutrons from given isotopes.
- Exercises on isotopic symbols, nuclear notation, and neutron calculations.
Key formulas and numbers to remember
- Mass number: A = p^+ + n^0
- Atomic number: Z = p^+ = e^- in neutral atoms
- Neutron count: n^0 = A - Z
- Electron mass: m_e = 9.11 imes 10^{-31} ext{ kg}; Electron charge: e = 1.60 imes 10^{-19} ext{ C}
- Proton mass: mp
oughly 1.673 imes 10^{-27} ext{ kg}; Neutron mass: mn
oughly 1.675 imes 10^{-27} ext{ kg} - Electron-to-mass ratio: rac{e}{m} = 1.76 imes 10^{8} ext{ C g}^{-1}
- 1 amu = 1.660 imes 10^{-24} ext{ g}
- Nobel-gas core notation: e.g., ext{Mg} = [ ext{Ne}] 3s^2
- Maximum electrons per shell (Bohr): 2n^2 ext{ (n = 1,2,3,4) } o 2, 8, 18, 32
Connections to previous lectures
- Builds on Dalton’s laws and chemical laws (conservation of mass, definite proportions) by showing atomic-scale justification.
- Links to periodic trends via electron configurations and nuclear charge shielding.
Real-world relevance and ethical/philosophical implications
- Understanding isotopes underpins medical isotopes, tracing, and radiometric dating; responsible use of radioactive materials is essential.
- The concept of models (Dalton, Thomson, Rutherford, Bohr, quantum mechanical) illustrates how scientific knowledge evolves with new evidence.
Quick recall prompts
- What is the difference between mass number and atomic number?
- How does the modern atomic theory differ from Dalton’s theory?
- How many electrons can the 4th shell hold? 32 electrons.
Unit 2: Periodic Classification of the Elements
Unit outcomes
- Understand periodic classification and its importance in chemistry
- Correlate electron configuration with periodic properties; predict trends in periodic properties
- Appreciate the role of periodic classification in real-world contexts; develop scientific inquiry skills
2.1 Introduction
- Historical basis for classifying elements; periodicity concepts in nature
2.2 The Modern Periodic Table
- Periodic Law (modern form): properties are a periodic function of atomic number Z (not atomic mass)
- Mosley’s contribution: atomic number as fundamental property; “No two elements have the same atomic number.”
- Structure of the modern table: 7 periods, 18 groups; long-form table; division into blocks: s, p, d, f
- Groups and periods: vertical groups share similar outer-electron configurations; horizontal periods indicate same number of shells
- Blocks correspond to the outer-sublevel being filled (s-block, p-block, d-block, f-block)
- Representative elements (s- and p-block, main-group elements, groups IA–VIIIA plus He in VIIIA), Transition elements (d-block, groups IB–VIIIB), Inner-transition elements (f-block: Lanthanides and Actinides)
- Periodic table relationships: diagonal relationships (e.g., Li–Mg, Be–Al, B–Si)
2.2.1 The Periodic Law (history and predictions)
- Mendeleev’s law (1871): properties recur periodically with increasing atomic mass; he left gaps for undiscovered elements (eka-aluminium, eka-silicon) and predicted properties
- Eka-silicon (Ge) predictions vs. actuals (Ge, 1886) show close agreement in several properties (atomic mass, density, oxide/cloride formulas, boiling points, etc.)
- Modern periodic law (Mosley): atomic number governs periodicity; elements in same group display similar chemical properties
- Problems with Mendeleev’s table: isotopes not placed distinctly; some elements out of order by atomic mass before isotopes were known
2.2.2 Characteristics of Groups and Periods
- Periods: 7 total; periods 1–3 are short; 4–6 are long; 7 is incomplete/radioactive
- Block concept: outer electron configuration determines block (s, p, d, f)
- Groups: 18 now; Group IA–VIIIA (main groups); IB–VIIIB (transition blocks)
- Group location guides properties: e.g., Group VIIA halogens have seven valence electrons; Group VIIIA noble gases have full outer shells
- Group names: IA Alkali metals; IIA Alkaline earth metals; IIIA Boron family; IVA Carbon family; VA Nitrogen family; VIA Oxygen family; VIIA Halogens; VIIIA Noble gases
- Representative elements (main-group): ns^1 or ns^2 for Group IA/IIA; ns^2np^1–np^6 for Groups IIIA–VIIIA
- Transition elements (d-block): outer electrons fill d-orbitals; characterized as metals; located between s-block and p-block
- Rare-earth elements (f-block): inner-transition metals; Lanthanide series (period 6, 4f orbitals) and Actinide series (period 7, 5f orbitals)
2.2.3 Classification of the elements
- Three broad categories based on outer electron configurations: Representative (s- and p-block), Transition (d-block), Rare-earth (f-block)
- Table 2.5 lists valence electron configurations and blocks for representative elements (IA to VIIIA)
2.2.4 Periodic table organization and examples
- How to determine group/period from electron configuration; how to identify the block from the last electron sublevel (s, p, d, f)
- Example: Cl (Z=17) has config 1s^2 2s^2 2p^6 3s^2 3p^5; outer electrons = 7; belongs to Group VIIA; block p
- Table 2.8 (illustrative) shows sample elements with given atomic numbers and radii; students build a model to reflect trends
2.3 Periodic properties within a group and across a period
- Group trends (e.g., down a group): valence electron count constant, but atomic size increases; ionization energy generally decreases; electronegativity generally decreases; metallic character increases down a group
- Period trends: across a period, atomic size decreases (Z increases, effective nuclear charge increases); ionization energy and electronegativity increase; electron affinity generally increases across a period
- Shielding effect: inner electrons screen the nuclear charge; Zeff = Z − S
- For Na: Zeff ≈ +1; Na has 3 shells; [Ne] 3s^1 example
- Probes of periodic trends: atomic radius, ionization energy, electron affinity, electronegativity, metallic character
2.4 Advantages of periodic classification
- Predicts formulas of compounds by similar valence electron configurations (e.g., Na2O, Li2O, K2O all alkali metal oxides)
- Predicts oxide behavior and basicity vs acidity; trends in oxides across a period (Na2O, MgO basic; Al2O3 amphoteric; SiO2 acidic)
- Summaries and checklists included (Unit Summary + Review Exercises)
Key formulas and concepts to remember
- Periodic law (modern): properties are periodic functions of atomic number Z
- Zeff = Z − S (shielding/effective nuclear charge)
- Max electrons per shell (Bohr): 2n^2; shell capacities: 2, 8, 18, 32
- Electron configuration writing: Aufbau principle; diagonal rule
- Noble gas core notation: [Ne] 3s^2, [Ar] 4s^2 3d^10, etc.
- Group/period names and block designations (s, p, d, f)
- Diagonal relationships: Li–Mg, Be–Al, B–Si
Unit 3: Chemical Bonding and Intermolecular Forces
Unit outcomes
- Understand ionic, covalent, and metallic bonds; properties of substances with these bonds
- Draw Lewis structures (electron-dot) for simple ions and molecules
- Understand polarity within molecules; origin and significance of intermolecular forces
- Appreciate the role of intermolecular forces in biology and chemistry; distinguish between intra- and inter-molecular forces
- Demonstrate scientific inquiry skills: predicting, modeling, observing, communicating
3.1 Chemical Bonding
- Chemical bond: the attractive force that binds atoms together in a molecule or crystal lattice
- Lewis theory (G. N. Lewis): atoms tend to achieve stable electron configurations (octet) via sharing or transfer of electrons; Lewis dot structures (bonding/pair electrons vs. lone pairs)
- Bond types and bonding tendencies summarized:
- Ionic bonds: transfer of electrons from metal to non-metal; formation of cations and anions; crystalline solids with high melting points; conducts electricity when molten or in solution; typically ionic compounds like NaCl
- Covalent bonds: sharing of electrons between atoms; variety of bond orders (single, double, triple); gases, liquids or low-melting solids; generally poor conductors in all states
- Metallic bonds: delocalized electrons in a lattice of positive ions; sea of electrons; high electrical and thermal conductivity; malleability and ductility
3.2 Ionic Bonding
- Ion formation: metals lose electrons to form cations; non-metals gain electrons to form anions
- Example: Group IA metals form +1 cations; Halogens (Group VIIA) form -1 anions
- Formula-based ionic compounds: MX, M2X, M3X2, etc., depending on valence charges
- Lewis symbols for ions: Na^+ ; Cl^−; NaCl as an ionic compound
- Practice tasks: predict charges and formulae for given elements
3.3 Covalent Bonding
- Covalent bond formation: sharing of electron pairs (e.g., H2, Cl2, HCl)
- Lewis structures for covalent molecules; example H2: H–H; HCl: H–Cl with octet satisfied
- Bond types: single, double, triple bonds; examples CO2 (double bonds), C2H4 (double), N2 (triple)
- Polar vs non-polar covalent bonds: polarity arises from differences in electronegativity; partial charges δ+ and δ− assigned to bonded atoms
- Coordinate (dative) covalent bonds: both electrons supplied by one atom (e.g., NH3 donating to H+ in NH4^+; CO also has a coordinate bond)
- 3.3.1 Polarity in Covalent Molecules: dipole moments and electronegativity differences
3.3.2 Coordinate Covalent Bond
- Donor-acceptor model: donor supplies both electrons; acceptor has an empty orbital; examples include NH3 + H^+ → NH4^+; H2O with H^+; CO with C═O where O donates to C
- Properties of covalent compounds: generally liquids/gases; lower melting/boiling points than ionic compounds; do not conduct electricity when molten/aqueous; insoluble in polar solvents but soluble in non-polar solvents
3.4 Metallic Bonding
- Description: metal cations in a lattice embedded in a “sea” of delocalized electrons; electrons move freely; explains electrical and thermal conductivity, malleability, ductility
- Demonstrations: metal crystal models and conductivity experiments
3.5 Intermolecular Forces
- Intermolecular forces (Van der Waals forces): dipole-dipole interactions (polar molecules), London dispersion forces (present in all molecules; stronger with more electrons), hydrogen bonding (special dipole-dipole between H and N/O/F)
- Hydrogen bonding explains high boiling points for HF, H2O, NH3; represented by dotted lines between H and electronegative lone pairs
- Distinctions: intra-molecular bonds (covalent/ionic) vs. inter-molecular forces (dipole-dipole, London, hydrogen bonding)
3.6 Practical applications and activities
- Lewis structures, electron-dot notation, and predicting bond types
- Exercises on identifying polar vs non-polar molecules; determining bond types in given species
- Experimental activities include properties of ionic and covalent compounds (melting points, solubility, conductivity)
3.7 Summary and practice
- Key terms: bond types, Lewis structures, octet rule, polarity, coordinate covalent bonds, metallic bonding, intermolecular forces, hydrogen bonding, dipole-dipole, London forces
- Common exercises: drawing Lewis structures; identifying bond types; predicting polarity; comparing ionic and covalent compounds
Key formulas and concepts to remember
- Octet rule: atoms tend to have eight electrons in the valence shell
- Electronegativity and polarity: electronegativity differences create dipoles; polar covalent bonds yield partial charges
- Hydrogen bonding strength and implications for boiling points
- Ionic bond formation and lattice structure; covalent bond sharing and bond order
- Ionic vs covalent properties: melting points, electrical conductivity (ionic melts/solutions conduct; covalent generally do not), solubility in polar vs non-polar solvents
Unit 4: Chemical Reactions and Stoichiometry
Unit outcomes
- Understand fundamental laws of chemical reactions; write and balance chemical equations; understand energy changes in reactions
- Types of chemical reactions; stoichiometry (mass-mass, volume-volume, mass-volume problems); limiting and excess reagents; theoretical vs actual yields; percentage yield
- Oxidation-reduction (redox) reactions; identifying oxidizing/reducing agents and agents; oxidation numbers
- Rate of reaction and chemical equilibrium; Le Châtelier’s principle; factors affecting rates and equilibrium
- Demonstrate scientific inquiry skills: observing, predicting, modeling, communicating, measuring, problem solving
4.1 Introduction
- A chemical reaction is a process where reactants are transformed into products; energy and property changes accompany most reactions
- Examples: combustion, fermentation, metabolism, etc.
- Notation convention: Reactants → Products; signs of energy changes may be included (e.g., heat)
4.2 Fundamental laws of chemical reactions
- Law of Conservation of Mass: mass of reactants equals mass of products; mass is conserved in all reactions
- Law of Definite Proportions: a compound has fixed elemental composition by mass
- Law of Multiple Proportions: when two elements form more than one compound, the masses of one element combined with a fixed mass of the other are in simple whole-number ratios
- Demonstrations and experiments: Mass conservation using reactions in sealed systems
4.3 Chemical Equations
- Distinguish between chemical equation and chemical reaction
- Steps to write chemical equations: word equation → symbols/formulas → balance
- Balancing methods: Inspection method; Least Common Multiple (LCM) method
- Substances exist in diatomic form for certain elements (H2, O2, N2, etc.) in equations
- Qualitative and quantitative meanings: stoichiometry expresses mole/mass relationships
4.4 Energy changes in chemical reactions
- Endothermic reactions: absorb heat; ΔH > 0; products have greater energy than reactants
- Exothermic reactions: release heat; ΔH < 0; products have lower energy than reactants
- ΔH is enthalpy change of reaction: ΔH = Hproducts − Hreactants
- Energy diagrams illustrate energy changes during a reaction
4.5 Types of chemical reactions
- Combination (synthesis): two or more substances form a single product (A + B → AB)
- Decomposition: one compound breaks down into simpler substances (AB → A + B)
- Single displacement (replacement): one element displaces another from a compound (A + BC → B + AC)
- Double displacement (metathesis): exchange of ions between two compounds (AB + CD → AD + CB)
- Examples and experiments illustrate each type
4.6 Stoichiometry
- Molar ratios from balanced equations; mole-mass, volume-volume, and mass-volume relationships
- Mass-mass method: convert given mass to moles, apply mole ratios, convert to desired mass
- Mole-ratio method: use molar ratios directly from balanced equation
- Gas volumes at STP: 1 mole of any gas occupies 22.4 L at STP; relate volume to moles via n = V/22.4 L
- Mass-volume and volume-volume problems explained with worked examples
- Limiting and excess reagents: identify limiting reactant; calculate theoretical yield; determine actual yield; compute percent yield
- Theoretical vs actual yield definitions and calculations
- Example problem sets include Ca + 2HCl → CaCl2 + H2, decomposition of KClO3, etc.
4.7 Oxidation and Reduction (Redox Reactions)
- Oxidation: loss of electrons; increase in oxidation number
- Reduction: gain of electrons; decrease in oxidation number
- Oxidizing agent: substance that causes oxidation (is reduced)
- Reducing agent: substance that causes reduction (is oxidized)
- Rules for assigning oxidation numbers (including exceptions for O, H, and group trends)
- Examples illustrating oxidation state changes and agent roles
- Non-redox reactions: many double displacement and acid-base reactions are not redox
4.8 Rates of Chemical Reactions and Chemical Equilibrium
- Rate of reaction: change in concentration of reactants/products over time; rate = Δ[C]/Δt (or related concentration measures)
- Collision theory: reactions require effective collisions with sufficient energy (activation energy) and proper orientation
- Pre-conditions for reaction: molecule collisions, activation energy, proper orientation
- Factors affecting rate: Nature of reactants, Temperature, Concentration, Surface area, Catalysts
- Activation energy and catalysts: catalysts provide alternate pathways with lower activation energy, increasing rate
- Reversible vs irreversible reactions; chemical equilibrium (forward rate = reverse rate) and the equilibrium constant K_eq
- Le Châtelier’s principle: stress (temperature, concentration, pressure) shifts equilibrium to counteract the stress; applications to industrial processes (Haber, Contact)
- Expressions for Kf and Kr and K_eq; interpretation in practical terms
4.8.2 Chemical Equilibrium
- Dynamic equilibrium: forward and reverse reactions continue; net concentrations do not change
- Equilibrium constant expressions; example: N2 + 3H2 ⇌ 2NH3; K_eq = ([NH3]^2)/([N2][H2]^3)
- Graphical depiction of forward vs reverse rates approaching equilibrium
4.8.3 Review and exercises
- Review exercises on identifying reaction types, balancing, and applying Le Chatelier’s principle
Key formulas and concepts to remember
- Balancing methods: balancing coefficients, not changing subscripts
- Enthalpy change: riangle H = Hp - Hr
- Reaction rate: ext{Rate} = rac{ ext{Change in concentration}}{ ext{Time}}
- Ideal gas law: PV = nRT and related constants (R ≈ 0.0821 L·atm·K^−1·mol^−1, or R ≈ 8.314 J·mol^−1·K^−1)
- 22.4 L mol^−1 at STP for gases
- Stoichiometric relationships: mole ratios from balanced equations; mass/volume relationships depending on the problem type
Unit 5: Physical States of Matter
Unit outcomes
- Understand the kinetic molecular theory and properties of solid, liquid and gas
- Explain gas behavior via V, T, P, and n; introduce ideal gas concepts, diffusion, evaporation, boiling, vapor pressure, and phase changes
- Solve problems with gas laws; perform activities to illustrate gas laws; measure boiling/melting points and phase changes
- Demonstrate inquiry skills: observing, predicting, measuring, drawing conclusions
5.1 Introduction
- Matter exists as solid, liquid, gas; plasma at very high temperatures
- Basic definitions for solid, liquid, gas in terms of shape, volume, compressibility, density, and fluidity
- Phase changes overview: fusion/melting, freezing, evaporation/boiling, condensation, sublimation, deposition
5.2 Kinetic theory and properties of matter
- Kinetic molecular theory assumptions:
1) Matter is composed of particles in constant motion; particles possess kinetic energy and potential energy
2) The differences among solids, liquids, gases arise from energy and motion differences - The three states differ in energy content and particle motion
- Properties of matter: density, compressibility, fluidity
- Kinetic molecular theory assumptions:
5.3 The Gaseous State (gases)
- Kinetic molecular theory: gas particles are widely spaced; negligible particle volume; no interparticle attraction under ideal conditions; continuous, random motion; average kinetic energy proportional to Kelvin temperature
- Gas laws: Boyle’s law (P vs V at constant T and n), Charles’ law (V vs T at constant P and n), and the Combined Gas Law (PV/T = constant)
- Avogadro’s law: equal volumes of gases at same T and P contain equal numbers of molecules; V ∝ n
- Ideal gas equation: PV = nRT; molar volume at STP: 22.4 L/mol; ideal gas constant values
- Graham’s law of diffusion: rate ∝ 1/√M; lighter gases diffuse faster
- Diffusion experiments comparing gases (NH3 vs HCl) illustrate rates and distances
5.4 The Liquid State
- Evaporation vs boiling; vapor pressure; factors affecting evaporation: temperature, intermolecular forces, surface area
- Volatile vs non-volatile liquids; vapour pressure increases with temperature; weaker intermolecular forces yield higher vapor pressure
- Boiling point: temperature at which vapor pressure equals external pressure; at STP, water boils at 100°C; the notion of absolute zero; the heating curve for phase changes
- Heat of vaporization (ΔHvap) and heat of condensation (ΔHcond); ΔHvap = −ΔHcond
5.5 The Solid State
- Phase changes from solid to liquid (melting) and liquid to gas (vaporization) and solid to gas (sublimation)
- Heat of fusion (ΔHfus) for melting; molar heat of fusion (ΔHfus) is energy to melt one mole; examples (ice ≈ 6.01 kJ/mol)
- Heat of solidification (ΔHcryst) equals −ΔHfus for solidification; sublimation energy DHsub = ΔHfus + ΔH_vap
- Heating curves illustrate phase transitions and temperature plateaus during phase changes
5.3–5.5 Useful formulas and concepts
- Molar volume at STP: V_m = 22.4 ext{ L mol}^{-1} for ideal gases
- Ideal gas equation: PV = nRT with appropriate units
- Charles’ Law: rac{V1}{T1} = rac{V2}{T2} at constant P and n (T in Kelvin)
- Boyle’s Law: PV = ext{constant} at constant T and n
- Combined gas law: rac{P1 V1}{T1} = rac{P2 V2}{T2}
- Graham’s law: rate ratio rac{r1}{r2} = rac{ ext{√}(M2)}{ ext{√}(M1)}
- Phase change definitions: fusion (melting), solidification, vaporization, condensation, sublimation, deposition
- Absolute zero: 0 K = −273.15°C; relates to gas behavior and the limitation of temperature scales
5.5.1 Experiments and activities (selected examples)
- Measuring vapour pressure of water via an apparatus with an open/closed system
- Determining boiling point and using porclain chips in heating experiments
- Observing melting/boiling curves with heating curves for pure substances
- Diffusion experiments (Graham’s law) using gas diffusion in a tube with NH3 and HCl
5.6 The solid state and phase behavior
- Iodine sublimation demonstration; ether and other volatiles; phase behavior under various conditions
Unit 5 unit summaries and review prompts
- Summary of key phase-change relationships; definitions of solid, liquid, gas; phase-change energy terms; the ideal gas model and limitations
Key formulas and concepts to remember
- Boyle’s Law: P imes V = ext{constant} (at constant T and n)
- Charles’ Law: rac{V}{T} = ext{constant} (at constant P and n)
- Combined gas law: rac{P1 V1}{T1} = rac{P2 V2}{T2}
- Ideal gas equation: PV = nRT
- 1 atm = 101325 Pa ≈ 101.3 kPa; 22.4 L per mole at STP
- Vaporization/condensation: ΔHvap and ΔHcond; ΔHvap = −ΔHcond
- Heat of fusion (ΔHfus) and heat of crystallization (ΔHcryst)
- Sublimation energy: ΔHsub = ΔHfus + ΔH_vap
Connections and overarching themes
- Matter exists in three main states with transitions driven by energy and intermolecular forces; kinetic theory connects microscopic motion to macroscopic properties (pressure, volume, temperature)
- The gas laws, kinetic theory, and Dalton’s and Schrodinger’s ideas about atoms connect to chemical bonding and periodic trends studied in Units 1–3
Practice prompts and exercises (highlights)
- Convert among units (mmHg, atm, Pa); apply PV = nRT to problems; calculate molar mass from gas data; determine volumes at STP; evaluate changes in temperature/volume/pressure relationships
- Investigate rate of diffusion using Graham’s law; determine diffusion speeds for gases with known molar masses; compare rates experimentally and via theoretical predictions
Quick recall prompts
- What is the difference between evaporation and boiling?
- How does vapor pressure depend on temperature and intermolecular forces?
- What is the significance of the heating curve during phase changes?
- How does the ideal gas law relate P, V, T, and n? What does it assume about real gases?
Overall cross-unit connections
- Atomic structure (Unit 1) informs chemical bonding (Unit 3) and periodic properties (Unit 2)
- Bonding types affect properties of substances (Unit 3), which in turn influence chemical reactions and stoichiometry (Unit 4)
- Gas laws (Unit 5) tie back to molecular interactions and kinetic theory discussed in Unit 5 and Unit 3 concepts (molecular interactions and phase changes)
Quick reference: key constants and conversions to memorize
1 mol of gas at STP: volume = 22.4 ext{ L}
Ideal gas constant: R egin{cases}0.0821 ext{ L·atm·mol}^{-1} ext{·K}^{-1}\ 8.314 ext{ J·mol}^{-1} ext{·K}^{-1} ext{ (SI)} ag*{} \ ext{(choose units consistently)}
Absolute zero: 0 ext{ K} = -273.15^ ext{o}C$$
Note on structure for exam prep
- Focus on understanding concepts, definitions, and how to apply formulas to solve problems (balancing equations, calculating moles, using gas laws, applying Le Chatelier’s principle)
- Practice: write balanced equations, determine limiting reagents, compute theoretical/actual yields, and predict reaction outcomes using the concepts from these units.
Final reminder
- Mastery comes from linking model ideas (atom, bonding, periodic trends) to problem-solving (stoichiometry, energetics, kinetics, equilibrium) and to real-world applications (industrial processes, material properties, environmental chemistry).