Chemistry Comprehensive Block Test Revision Guide

Topic 1: States of Matter

Concise Notes & Key Formulas

• States of Matter:

  • Solids: Packed closely together in a fixed, regular lattice structure; vibrate about fixed positions; low kinetic energy, strong attractive forces; fixed shape and volume, incompressible.

  • Liquids: Close together but randomly arranged; move and slide around each other; moderate kinetic energy, weaker forces; fixed volume, takes the shape of its container.

  • Gases: Widely separated, random arrangement; rapid, random movement in all directions; high kinetic energy, negligible forces; no fixed shape or volume, highly compressible.

• Interconversion: Melting (solid to liquid), Freezing (liquid to solid), Boiling/Evaporation (liquid to gas), Condensation (gas to liquid), Sublimation (solid to gas), Deposition (gas to solid).

• Diffusion: Net movement of particles from higher concentration to lower concentration down a concentration gradient. Lighter particles with lower relative molecular mass ($M<em>r$) diffuse faster (e.g., $NH</em>3$ with $M<em>r = 17.0$ diffuses faster than $HCl$ with $M</em>r = 36.5$).

• Solubility Formula:
  $\text{Solubility} = \frac{\text{Mass of Solute}\,(g)}{\text{Mass of Solvent}\,(g)} \times 100$

Key Definitions

• Solute: The dissolved substance in a solution.

• Solvent: The liquid medium in which the solute dissolves.

• Solution: A homogeneous mixture formed when a solute dissolves completely in a solvent.

• Saturated Solution: A solution containing the maximum mass of dissolved solute possible at a given temperature in the presence of undissolved solute.

Common Mistakes

• Confusing boiling (occurs throughout liquid at boiling point) with evaporation (occurs only at the surface below boiling point).

• Forgetting that gas solubility decreases as temperature increases, whereas solid solubility usually increases.

• Forgetting to convert solvent mass to grams of solvent when applying the solubility formula.

Short Quiz

1. What is the process of a solid turning directly into a gas called?

2. Why does ammonia gas ($NH_3$) diffuse faster than hydrogen chloride gas ($HCl$)?

3. What state of matter has a fixed volume but no fixed shape?

Exam-Style Question

A student dissolves $15.0\,g$ of a salt in $50.0\,g$ of water at $25\,^{\circ}\text{C}$ to form a saturated solution. Calculate the solubility of the salt in $g$ per $100\,g$ of water.

Topic 2: Elements, Compounds, and Mixtures

Concise Notes & Key Formulas

• Elements: Pure substances containing only one type of atom; cannot be broken down by chemical means.

• Compounds: Pure substances formed when two or more different elements are chemically bonded in fixed proportions.

• Mixtures: Two or more substances physically mixed without chemical bonding; components retain individual properties.

• Separation Techniques:

  • Filtration: Separates an insoluble solid from a liquid (residue on paper, filtrate passes through).

  • Crystallization: Obtains a dissolved soluble solute from solution by evaporating solvent to saturation point and cooling.

  • Simple Distillation: Separates liquid solvent from a solution using evaporation and condensation.

  • Fractional Distillation: Separates miscible liquids with different boiling points using a fractionating column.

  • Paper Chromatography: Separates soluble components in a mixture based on solubility in mobile phase and adsorption to stationary phase.

• Retention Factor Formula:
  $R_f = \frac{\text{Distance travelled by solute}}{\text{Distance travelled by solvent front}}$

Key Definitions

• Residue: The insoluble solid remaining on the filter paper after filtration.

• Filtrate: The liquid that passes through the filter paper during filtration.

• Retention Factor ($R_f$): The ratio of the distance moved by a solute spot to the distance moved by the solvent front.

Common Mistakes

• Drawing the baseline in ink instead of pencil during chromatography (ink dissolves and interferes with results).

• Allowing the solvent level in chromatography to start above the baseline spot (it must start below).

• Confusing simple distillation (separates solvent from solute) with fractional distillation (separates liquid mixtures with different boiling points).

Short Quiz

1. Which separation technique is used to obtain pure water from sea water?

2. Why must the baseline in paper chromatography be drawn in pencil?

3. What is the maximum possible value for an $R_f$ factor?

Exam-Style Question

In a chromatography experiment, a dye spot travels $4.2\,cm$ while the solvent front travels $6.0\,cm$. Calculate the $R_f$ value of the dye.

Topic 3: Atomic Structure and the Periodic Table

Concise Notes & Key Formulas

• Subatomic Particles:

  • Proton: Relative Mass = $1$, Relative Charge = $+1$

  • Neutron: Relative Mass = $1$, Relative Charge = $0$

  • Electron: Relative Mass = $\frac{1}{1836}$, Relative Charge = $-1$

• Atomic Numbers & Mass:

  • Atomic Number ($Z$): Number of protons in nucleus.

  • Mass Number ($A$): Protons + Neutrons.

  • Isotopes: Atoms of same element with same atomic number but different mass numbers.

• Relative Atomic Mass Formula:
  $A_r = \frac{\sum (\text{Isotope Mass} \times \text{Isotope Abundance})}{\text{Total Abundance}}$

• Periodic Table:

  • Groups (columns) = number of outer-shell electrons.

  • Periods (rows) = number of occupied electron shells.

  • Shell capacities: 1st shell max $2$, 2nd shell max $8$, 3rd shell max $8$ (for first 20 elements).

Key Definitions

• Atomic Number ($Z$): The number of protons in the nucleus of an atom.

• Mass Number ($A$): The total number of protons and neutrons in the nucleus of an atom.

• Isotopes: Atoms of the same element with the same number of protons but different numbers of neutrons.

Common Mistakes

• Confusing atomic number with mass number.

• Forgetting that neutral atoms have equal numbers of protons and electrons.

• Writing electronic configurations with incorrect shell capacities (e.g., putting 9 electrons in the second shell).

Short Quiz

1. What subatomic particle determines the atomic number of an element?

2. What is the electronic configuration of Sodium ($Na$, atomic number $11$)?

3. True or False: Isotopes have different chemical properties.

Exam-Style Question

A sample of chlorine contains $75\%$ of $^{35}\text{Cl}$ and $25\%$ of $^{37}\text{Cl}$. Calculate the relative atomic mass ($A_r$) of chlorine.

Topic 4: Chemical Formulae, Equations, and Calculations

Concise Notes & Key Formulas

• State Symbols: $(s)$ solid, $(l)$ liquid, $(g)$ gas, $(aq)$ aqueous solution.

• Conservation of Mass: Mass of reactants equals mass of products; chemical equations must be balanced.

• Relative Formula Mass ($M<em>r$): Sum of relative atomic masses ($A</em>r$) of all atoms in a formula.

  • Example for $H<em>2SO</em>4$:
    $M_r = (2 \times 1.0) + (1 \times 32.1) + (4 \times 16.0) = 98.1$

• Percentage Composition Formula:
  $\text{Percentage Mass of Element} = \frac{n \times A<em>r}{M</em>r} \times 100$
  where $n$ is the number of atoms of that element in the formula.

Key Definitions

• Law of Conservation of Mass: Total mass of products equals total mass of reactants in a chemical reaction.

• Relative Formula Mass ($M_r$): The sum of the relative atomic masses of all atoms present in a compound's formula.

Common Mistakes

• Changing chemical subscripts instead of coefficients when balancing equations.

• Forgetting to multiply $A_r$ by the number of atoms present in percentage composition calculations.

• Using $(l)$ instead of $(aq)$ for substances dissolved in water.

Short Quiz

1. What state symbol is used for a substance dissolved in water?

2. Balance the equation: $Fe + O<em>2 \rightarrow Fe</em>2O_3$.

3. What is the $M<em>r$ of water ($H</em>2O$), given $A<em>r(H) = 1.0$ and $A</em>r(O) = 16.0$?

Exam-Style Question

Calculate the percentage composition by mass of oxygen in calcium carbonate ($CaCO<em>3$). (Given: $A</em>r(Ca) = 40.1$, $A<em>r(C) = 12.0$, $A</em>r(O) = 16.0$).

Topic 5: Bonding

Concise Notes & Key Formulas

• Ionic Bonding:

  • Metal loses electrons to form positive cations; non-metal gains electrons to form negative anions.

  • Electrostatic attraction between oppositely charged ions in a giant ionic lattice.

  • High melting points; conduct electricity when molten or aqueous (mobile ions), but NOT when solid.

  • Polyatomic Ions: $NH<em>4^+$, $OH^-$, $NO</em>3^-$, $SO<em>4^{2-}$, $CO</em>3^{2-}$.

• Covalent Bonding:

  • Electrostatic attraction between shared pair of electrons and positively charged nuclei.

  • Simple Molecular: Low melting/boiling points (weak intermolecular forces), non-conductors.

  • Giant Covalent:

  • Diamond: Each C bonded to 4 C atoms; extremely hard, high melting point, non-conductor.

  • Graphite: Each C bonded to 3 C atoms in layers; slippery, high melting point, conducts electricity due to delocalised electrons.

  • Silicon Dioxide ($SiO_2$): High melting point, hard, non-conductor.

• Metallic Bonding:

  • Electrostatic attraction between positive metal cations in regular layers and a sea of delocalised electrons.

  • High melting points, high electrical/thermal conductivity, malleable and ductile.

Key Definitions

• Ionic Bond: Strong electrostatic attraction between oppositely charged ions.

• Covalent Bond: Strong electrostatic attraction between a shared pair of electrons and the positively charged nuclei of the bonding atoms.

• Metallic Bond: Strong electrostatic attraction between positive metal ions and delocalised electrons.

Common Mistakes

• Stating that covalent bonds break when simple molecular structures melt (only weak intermolecular forces break).

• Stating that ionic solids conduct electricity because electrons move (it is mobile ions in liquid/aqueous states that carry charge).

Short Quiz

1. Why does graphite conduct electricity while diamond does not?

2. What type of bonding involves a sea of delocalised electrons?

3. What is the chemical formula of ammonium sulfate?

Exam-Style Question

Explain, in terms of structure and bonding, why sodium chloride ($NaCl$) has a high melting point and conducts electricity when dissolved in water, but not when solid.

Topic 6: Gases in the Atmosphere

Concise Notes & Key Formulas

• Composition of Clean, Dry Air:

  • Nitrogen ($N_2$): $\approx 78\%$

  • Oxygen ($O_2$): $\approx 21\%$

  • Argon ($Ar$): $\approx 0.9\%$

  • Carbon Dioxide ($CO_2$): $\approx 0.04\%$

• Determining Percentage of Oxygen:

  • Copper Turnings: $2Cu(s) + O_2(g) \rightarrow 2CuO(s)$ (volume drops by $\approx 21\%$).

  • Iron Rusting: Wet iron filings in inverted tube over water (water level rises by $\approx 21\%$).

• Combustion Reactions:

  • Magnesium: $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$ (bright white flame, white powder).

  • Hydrogen: $2H<em>2(g) + O</em>2(g) \rightarrow 2H_2O(l)$ (pale blue flame).

  • Sulfur: $S(s) + O<em>2(g) \rightarrow SO</em>2(g)$ (blue flame, colorless pungent gas).

• Thermal Decomposition:

  • Metal carbonate $\rightarrow$ Metal oxide + Carbon dioxide.

  • Example: $CuCO<em>3(s) \rightarrow CuO(s) + CO</em>2(g)$ (green powder turns black; gas turns limewater cloudy).

Key Definitions

• Thermal Decomposition: The breaking down of a substance into two or more simpler substances using heat.

• Greenhouse Gas: A gas (such as $CO_2$) that absorbs infrared radiation emitted from Earth's surface and re-radiates it, trapping heat in the atmosphere.

Common Mistakes

• Misidentifying the color change in copper(II) carbonate decomposition (it is green solid to black solid).

• Forgetting that argon is present in larger amounts than carbon dioxide in dry air.

Short Quiz

1. What percentage of clean, dry air is made up of nitrogen?

2. What observation indicates that carbon dioxide gas is produced when limewater is used?

3. What is the color of the flame when magnesium burns in oxygen?

Exam-Style Question

Describe an experiment using copper turnings to determine the percentage of oxygen in air, including the equation for the reaction and expected percentage result.

Topic 7: Group 1 Alkali Metals

Concise Notes & Key Formulas

• Group Overview:

  • Lithium ($Li$), Sodium ($Na$), Potassium ($K$), Rubidium ($Rb$), Caesium ($Cs$), Francium ($Fr$).

  • Soft, low density, low melting points compared to transition metals.

• Reactivity Trend:

  • Reactivity INCREASES down Group 1 ($Li < Na < K < Rb < Cs$).

  • Explanation: Group 1 metals react by losing 1 outer electron to form $+1$ cations. Down the group, atomic radius increases and electron shielding increases, weakening attraction between nucleus and outer electron.

• Reaction with Water:

  • $2M(s) + 2H<em>2O(l) \rightarrow 2MOH(aq) + H</em>2(g)$

  • Lithium: Effervesces steadily, floats, shrinks, disappears.

  • Sodium: Effervesces rapidly, melts into shiny sphere, floats and moves quickly, disappears.

  • Potassium: Effervesces violently, melts into sphere, burns with lilac flame, disappears with a pop.

Key Definitions

• Alkali Metals: Elements in Group 1 of the periodic table, characterized by having one electron in their outer shell and forming alkaline solutions when reacting with water.

Common Mistakes

• Stating that alkali metals gain electrons when they react (they lose 1 outer electron).

• Forgetting that the solution formed with water is alkaline (contains $OH^-$ ions).

Short Quiz

1. Why do Group 1 metals become more reactive down the group?

2. What flame color is observed when potassium reacts with water?

3. What gas is produced when sodium reacts with water?

Exam-Style Question

Write a balanced chemical symbol equation, including state symbols, for the reaction between sodium metal and water. Describe two physical observations made during this reaction.